Properties, Measurements and Units
1
CHAPTER 1
PROPERTIES, MEASUREMENTS AND UNITS
OBJECTIVES
After studying this chapter, you will be able to:
1. Describe the properties of substances
2. Explain the measurements and units
3. Define the international system of units
1.1. THE PROPERTIES OF SUBSTANCES
All matter has physical and chemical properties. Physical properties are
characteristics that scientists can measure without changing the composition of the
sample under study, such as mass, color, and volume (the amount of space occupied
by a sample). Chemical properties describe the characteristic ability of a substance
to react to form new substances; they include its flammability and susceptibility to
corrosion. All samples of a pure substance have the same chemical and physical
properties. For example, pure copper is always a reddish-brown solid (a physical
property) and always dissolves in dilute nitric acid to produce a blue solution and
a brown gas (a chemical property).
2
University Chemistry
Physical properties can be extensive or intensive. Extensive properties vary
with the amount of the substance and include mass, weight, and volume. Intensive
properties, in contrast, do not depend on the amount of the substance; they include
color, melting point, boiling point, electrical conductivity, and physical state at a
given temperature. For example, elemental sulfur is a yellow crystalline solid that
does not conduct electricity and has a melting point of 115.2 °C, no matter what
amount is examined (Figure 1). Scientists commonly measure intensive properties to
determine a substance’s identity, whereas extensive properties convey information
about the amount of the substance in a sample.
Figure.1: The Difference between Extensive and Intensive Properties of Matter. Because
they differ in size, the two samples of sulfur have different extensive properties, such
as mass and volume. In contrast, their intensive properties, including color, melting
point, and electrical conductivity, are identical.
Although mass and volume are both extensive properties, their ratio is an
important intensive property called density. Density is defined as mass per unit
volume and is usually expressed in grams per cubic centimeter (g/cm3). As mass
increases in a given volume, density also increases. For example, lead, with its
greater mass, has a far greater density than the same volume of air, just as a brick
has a greater density than the same volume of Styrofoam. At a given temperature
and pressure, the density of a pure substance is a constant:
densityρ=massvolume=mVdensity=massvolumeρ=mV
Pure water, for example, has a density of 0.998 g/cm3 at 25 °C. The average
densities of some common substances are in Table.1. Notice that corn oil has a lower
mass to volume ratio than water. This means that when added to water, corn oil will
“float” (Figure 2).
Table 1: Densities of Common Substances
Substance
Density at 25 °C (g/cm3)
Substance
Density at 25 °C (g/cm3)
blood
1.035
corn oil
0.922
body fat
0.918
mayonnaise
0.910
whole milk
1.030
honey
1.420
Properties, Measurements and Units
3
Figure 2: Water and oil. Since the oil has a lower density than water, it floats on top.
1.1.1 Physical Property and Change
Physical changes are changes in which no chemical bonds are broken or
formed. This means that the same types of compounds or elements that were
there at the beginning of the change are there at the end of the change. Because
the ending materials are the same as the beginning materials, the properties (such
as color, boiling point, etc) will also be the same. Physical changes involve moving
molecules around, but not changing them. Some types of physical changes include:
•
•
•
•
Changes of state (changes from a solid to a liquid or a gas and vice versa)
Separation of a mixture
Physical deformation (cutting, denting, stretching)
Making solutions (special kinds of mixtures) .
As an ice cube melts, its shape changes as it acquires the ability to flow.
However, its composition does not change. Melting is an example of a physical
change (Figure 3), since some properties of the material change, but the identity
of the matter does not. Physical changes can further be classified as reversible or
irreversible. The melted ice cube may be refrozen, so melting is a reversible physical
change. Physical changes that involve a change of state are all reversible. Other
changes of state include vaporization (liquid to gas), freezing (liquid to solid),
and condensation (gas to liquid). Dissolving is also a reversible physical change.
When salt is dissolved into water, the salt is said to have entered the aqueous state.
The salt may be regained by boiling off the water, leaving the salt behind.
4
University Chemistry
Figure 3: Ice Melting is a physical change. When solid water (H2O) as ice melts into
a liquid (water), it appears changed. However, this change is only physical as the the
composition of the constituent molecules is the same: 11.19% hydrogen and 88.81%
oxygen by mass.
1.1.2 Chemical Properties and Change
Chemical changes occur when bonds are broken and/or formed between
molecules or atoms. This means that one substance with a certain set of properties
(such as melting point, color, taste, etc) is turned into a different substance with
different properties. Chemical changes are frequently harder to reverse than
physical changes.
One good example of a chemical change is burning paper. In contrast to the
act of ripping paper, the act of burning paper actually results in the formation
of new chemicals (carbon dioxide and water, to be exact). Another example of
chemical change occurs when water is formed. Each molecule contains two atoms
of hydrogen and one atom of oxygen chemically bonded.
Another example of a chemical change is what occurs when natural gas
is burned in your furnace. This time, before the reaction we have a molecule of
methane, CH4, and two molecules of oxygen, O2, while after the reaction we have
two molecules of water, H2O, and one molecule of carbon dioxide, CO2. In this
case, not only has the appearance changed, but the structure of the molecules has
also changed. The new substances do not have the same chemical properties as the
original ones. Therefore, this is a chemical change.
The combustion of magnesium metal is also chemical change (Magnesium +
Oxygen → Magnesium Oxide):
2Mg+O2→2MgO2Mg+O2→2MgO
as is the rusting of iron (Iron + Oxygen → Iron Oxide/ Rust):
Properties, Measurements and Units
5
4Fe+3O2→2Fe2O34Fe+3O2→2Fe2O3
Using the components of composition and properties, we have the ability to
distinguish one sample of matter from the others.
1.1.3 Physical and Chemical Properties
The characteristics that enable us to distinguish one substance from another
are called properties. A physical property is a characteristic of matter that is not
associated with a change in its chemical composition. Familiar examples of physical
properties include density, color, hardness, melting and boiling points, and electrical
conductivity. We can observe some physical properties, such as density and
color, without changing the physical state of the matter observed. Other physical
properties, such as the melting temperature of iron or the freezing temperature of
water, can only be observed as matter undergoes a physical change. A physical
change is a change in the state or properties of matter without any accompanying
change in its chemical composition (the identities of the substances contained in
the matter). We observe a physical change when wax melts, when sugar dissolves
in coffee, and when steam condenses into liquid water (Figure 4). Other examples
of physical changes include magnetizing and demagnetizing metals (as is done
with common antitheft security tags) and grinding solids into powders (which can
sometimes yield noticeable changes in color). In each of these examples, there is a
change in the physical state, form, or properties of the substance, but no change in
its chemical composition.
Figure 4: (a) Wax undergoes a physical change when solid wax is heated and forms
liquid wax. (b) Steam condensing inside a cooking pot is a physical change, as water
vapor is changed into liquid water.
The change of one type of matter into another type (or the inability to change)
is a chemical property. Examples of chemical properties include flammability,
toxicity, acidity, reactivity (many types), and heat of combustion. Iron, for example,
combines with oxygen in the presence of water to form rust; chromium does not
oxidize (Figure 5). Nitroglycerin is very dangerous because it explodes easily; neon
poses almost no hazard because it is very unreactive.
6
University Chemistry
Figure 5: (a) One of the chemical properties of iron is that it rusts; (b) one of the chemical properties of chromium is that it does not.
To identify a chemical property, we look for a chemical change. A chemical
change always produces one or more types of matter that differ from the matter
sspresent before the change. The formation of rust is a chemical change because
rust is a different kind of matter than the iron, oxygen, and water present before the
rust formed. The explosion of nitroglycerin is a chemical change because the gases
produced are very different kinds of matter from the original substance. Other
examples of chemical changes include reactions that are performed in a lab (such
as copper reacting with nitric acid), all forms of combustion (burning), and food
being cooked, digested, or rotting (Figure 6).
Figure 6: (a) Copper and nitric acid undergo a chemical change to form copper nitrate
and brown, gaseous nitrogen dioxide. (b) During the combustion of a match, cellulose in the match and oxygen from the air undergo a chemical change to form carbon
dioxide and water vapor. (c) Cooking red meat causes a number of chemical changes,
including the oxidation of iron in myoglobin that results in the familiar red-to-brown
color change. (d) A banana turning brown is a chemical change as new, darker (and less
tasty) substances form.
Properties, Measurements and Units
7
Properties of matter fall into one of two categories. If the property depends on
the amount of matter present, it is an extensive property. The mass and volume of
a substance are examples of extensive properties; for instance, a gallon of milk has
a larger mass and volume than a cup of milk. The value of an extensive property
is directly proportional to the amount of matter in question. If the property of a
sample of matter does not depend on the amount of matter present, it is an intensive
property. Temperature is an example of an intensive property. If the gallon and
cup of milk are each at 20 °C (room temperature), when they are combined, the
temperature remains at 20 °C. As another example, consider the distinct but related
properties of heat and temperature. A drop of hot cooking oil spattered on your
arm causes brief, minor discomfort, whereas a pot of hot oil yields severe burns.
Both the drop and the pot of oil are at the same temperature (an intensive property),
but the pot clearly contains much more heat (extensive property).
While many elements differ dramatically in their chemical and physical
properties, some elements have similar properties. We can identify sets of elements
that exhibit common behaviors. For example, many elements conduct heat and
electricity well, whereas others are poor conductors. These properties can be used to
sort the elements into three classes: metals (elements that conduct well), nonmetals
(elements that conduct poorly), and metalloids (elements that have properties of
both metals and nonmetals).
The periodic table is a table of elements that places elements with similar
properties close together (Figure 7). You will learn more about the periodic table as
you continue your study of chemistry.
Figure 7: The periodic table shows how elements may be grouped according to certain
similar properties.
8
University Chemistry
1.1.4 Substances and Mixtures
In chemistry, a chemical substance is a form of matter that has constant chemical
composition and characteristic properties. It cannot be separated into components
without breaking chemical bonds. Chemical substances can be solids, liquids, gases,
or plasma. Changes in temperature or pressure can cause
substances to shift between the different phases of matter.
An element is a chemical substance that is made
up of a particular kind of atom and hence cannot be
broken down or transformed by a chemical reaction into
a different element. All atoms of an element have the
same number of protons, though they may have different
numbers of neutrons and electrons.
Note
The background
color denotes
whether an element
is a metal, metalloid, or nonmetal,
whereas the element
symbol color indicates whether it is a
solid, liquid, or gas.
A pure chemical compound is a chemical substance
that is composed of a particular set of molecules or ions that
are chemically bonded. Two or more elements combined
into one substance through a chemical reaction, such as
water, form a chemical compound. All compounds are
substances, but not all substances are compounds. A
chemical compound can be either atoms bonded together
in molecules or crystals in which atoms, molecules or ions form a crystalline lattice.
Compounds made primarily of carbon and hydrogen atoms are called organic
compounds, and all others are called inorganic compounds. Compounds containing
bonds between carbon and a metal are called organometallic compounds.
Chemical substances are often called ‘pure’ to set them apart from mixtures.
A common example of a chemical substance is pure water; it always has the same
properties and the same ratio of hydrogen to oxygen whether it is isolated from a
river or made in a laboratory. Other chemical substances commonly encountered
in pure form are diamond (carbon), gold, table salt (sodium chloride), and refined
sugar (sucrose). Simple or seemingly pure substances found in nature can in fact be
mixtures of chemical substances. For example, tap water may contain small amounts
of dissolved sodium chloride and compounds containing iron, calcium, and many
other chemical substances. Pure distilled water is a substance, but seawater, since it
contains ions and complex molecules, is a mixture.
Chemical Mixtures
A mixture is a material system made up of two or more different substances,
which are mixed but not combined chemically. A mixture refers to the physical
combination of two or more substances in which the identities of the individual
substances are retained. Mixtures take the form of alloys, solutions, suspensions,
and colloids.
Properties, Measurements and Units
9
Figure 8: Chemical Mixtures.
Naturally occurring sulfur crystals
Sulfur occurs naturally as elemental sulfur, sulfide, and sulfate minerals and
in hydrogen sulfide. This mineral deposit is composed of a mixture of substances.
Heterogeneous Mixtures
A heterogeneous mixture is a mixture of two or more chemical substances
(elements or compounds), where the different components can be visually
distinguished and easily separated by physical means. Examples include:
•
•
•
•
•
•
•
mixtures of sand and water
mixtures of sand and iron filings
a conglomerate rock
water and oil
a salad
trail mix
mixtures of gold powder and silver powder
Interactive: Oil and Water
Explore the interactions that cause water and oil to separate from a mixture.
Homogenous Mixtures
A homogeneous mixture is a mixture of two or more chemical substances
(elements or compounds), where the different components cannot be visually
distinguished. The composition of homogeneous mixtures is constant. Often
10
University Chemistry
separating the components of a homogeneous mixture is more challenging than
separating the components of a heterogeneous mixture.
Distinguishing between homogeneous and heterogeneous mixtures is a matter
of the scale of sampling. On a small enough scale, any mixture can be said to be
heterogeneous, because a sample could be as small as a single molecule. In practical
terms, if the property of interest is the same regardless of how much of the mixture
is taken, the mixture is homogeneous.
A mixture’s physical properties, such as its melting point, may differ from those
of its individual components. Some mixtures can be separated into their components
by physical (mechanical or thermal) means.
1.2. MEASUREMENTS AND UNITS
A measuring unit is a standard quantity used to express a physical quantity.
Let us learn about the physical quantities and some of the standard units used to
measure them.
The Physical quantities are further classified into 4 types:
•
•
•
•
Lengths of given objects
Weights of given objects
Money
Time
1.2.1 Lengths of given objects
Length describes how long a thing is from one end to the other. Length is used
to identify the size of the object or a distance from one point to another.
Metric Unit
Millimetre (mm): Used to measure very
short lengths or thicknesses.
Example: length of a pen tip.
Inch (in): Used to measure
the length of small objects.
Example: Length of a rod.
Centimetre (cm): Used to measure small
lengths.
Example: Length of a pen.
Foot (ft): Used to measure
short distances and heights.
Example: Heights of buildings.
Meter (m): Used to measure big lengths.
Example: Length of a classroom.
Yard (yd): It is bigger than
a foot.
Example: Length of a football field.
Kilometre (km): Used to measure very
long lengths or distance.
Example: Distance between two places.
Mile (mi): Used to measure
long distances.
Example: Distance between
two places.
US Standard Unit (English Unit or Customary
Unit)
11
Properties, Measurements and Units
The object can be differentiated or compared by its length.
Here’s an image showing 3cm length.
The above example is an illustration of how the object is compared by its length.
Here are few conversions of length:
1. 1 km = 1000m
2. 1m = 100 cm. 1cm = 10mm
1.2.2 Weights of given objects
US Standard Unit (English Unit
or Customary Unit)
Metric Unit
Millimetre (mm): Used to
measure very light things.
Example: Medicines
Ounce (oz): Used to measure small quantities.
Example: Bread.
Gram (g): Used to measure
small things.
Example: Potato
Pound (lb): Used to measure body weight, etc.
Example: Bread
12
University Chemistry
US Standard Unit (English Unit
or Customary Unit)
Metric Unit
Kilogram (kg): Used to
measure heavy things.
Example: Bodyweight.
Ton: Used to measure
much heavier things.
Example: Trucks and heavy
load
Here are few conversions of weight:
•
•
•
•
1 kg = 1000 grams.
1 gram = 1000 milligrams.
1 gram = (1/1000) kg.
1 milligram = (1/1000) gram.
The object can be compared with its weight.
The above example is an illustration of how an object is compared by its weight.
Note: The heavier object weighs down.
1.2.3 Money
Money can be defined as anything that people use to buy goods and services.
Money is a part of everyone’s life. Since the money/currency differs for each and
every country because each and every country use their own currencies. Each
country’s currency value differs from other countries on the basis of their economy.
13
Properties, Measurements and Units
US Standard Unit (English Unit or Customary Unit)
Country
India
Indian Rupees.
United States
United States Dollars.
United Kingdom.
Pound.
Japan
Japanese Yen.
Australia
Australian dollar.
Russia
Russian Rouble.
Here are few currency conversions:
One U.S Dollar = 74.74 rupee.
One Pound = 96.97 rupee.
One Rupee = 1.40 Japanese yen.
According to Indian currency:
•
•
1 = 100 paisa
1 paisa = (1/100) rupee
1.2.4. Time
The ongoing sequence of events is time.
The basic unit of time is the second.
There are also minutes, hours, days, weeks, months and years.
Metric Units and US Standard Unit are the same for the time
Second (s)
Minute (min)
Hour (hr)
Day
Week
Month
Yea
14
University Chemistry
Whereas,
•
•
•
•
•
•
1 year = 12 months
1 month = 4 weeks approx.
1 week = 7days.
1 day = 24 hours.
1 hour = 60 minutes
1 minute = 60 seconds.
Measuring units is a basic concept which we come across in our daily life.
Teaching kids about the measuring unit such as length, weight, money and time
give them a good idea and will let the kid explore more.
It gives a proper image to the kids in comparing the objects based on the units
because kids must understand the importance of measurement and be familiar
with their use in everyday life. It is an essential life skill.
1.3 THE INTERNATIONAL SYSTEM OF UNITS
In earlier time scientists of different countries were using different systems of
units for measurement. Three such systems, the CGS, the FPS (or British) system
and the MKS system were in use extensively till recently.
The base units for length, mass and time in these systems were as follows:
•
•
•
In CGS system they were centimetre, gram and second respectively.
In FPS system they were foot, pound and second respectively.
In MKS system they were metre, kilogram and second respectively.
The system of units which is at present internationally accepted for
measurement is the Système International d’ Unites (French for International
System of Units), abbreviated as SI. The SI, with standard
scheme of symbols, units and abbreviations, developed
by the Bureau International des Poids et measures (The
Note
International Bureau of Weights and Measures, BIPM) in
1971 were recently revised by the General Conference on
When mole is used,
Weights and Measures in November 2018. The scheme
the elementary entiis now for international usage in scientific, technical,
ties must be speciindustrial and commercial work. Because SI units used
fied. These entities
decimal system, conversions within the system are quite
may be atoms, molecules, ions, elecsimple and convenient. We shall follow the SI units in this
trons, other particles
book. In SI, there are seven base units as given in Table 2.
or specified groups
Besides the seven base units, there are two more units that
of such particles.
are defined for (a) plane angle dθ as the ratio of length of
Properties, Measurements and Units
15
arc ds to the radius r and (b) solid angle dΩ as the ratio of the intercepted area dA
of the spherical surface, described about the apex O as the centre, to the square of
its radius r, as shown in Fig. 9 (a) and (b) respectively. The unit for plane angle is
radian with the symbol rad and the unit for the solid angle is steradian with the
symbol sr. Both these are dimensionless quantities.
Figure 9: Description of (a) plane angle dθ and (b) solid angle dΩ
1.3.1 Extensive and Intensive Properties
Extensive Properties
Some properties of matter depend on the size of the sample, while some do
not. An extensive property is a property that depends on the amount of matter in
a sample. The mass of an object is a measure of the amount of matter that an object
contains. A small sample of a certain type of matter will have a small mass, while a
larger sample will have a greater mass. Another extensive property is volume. The
volume of an object is a measure of the space that is occupied by that object.
The figure below illustrates the extensive property of volume. The pitcher and
glass both contain milk. The pitcher holds approximately two quarts and the glass
will hold about 8 ounces of milk. The same milk is in each container. The only
difference is the amount of milk contained in the glass and in the pitcher
16
University Chemistry
Figure 10: Milk pitcher and glass.
Intensive Properties
The electrical conductivity of a substance is a property that depends only on
the type of substance. Silver, gold, and copper are excellent conductors of electricity,
while glass and plastic are poor conductors. A larger or smaller piece of glass will
not change this property. An intensive property is a property of matter that depends
only on the type of matter in a sample and not on the amount. Other intensive
properties include color, temperature, density, and solubility.
The copper wire shown in the picture below has a certain electrical conductivity.
You could cut off the small end sticking out and it would have the same conductivity
as the entire long roll of wire shown here. The conductivity is a property of the
copper metal itself, not of the length of the wire.
Figure 11: Copper wire.
1.3.2 Conversion Factors
A unit conversion expresses the same property as a different unit of measurement.
For instance, time can be expressed in minutes instead of hours, while distance
17
Properties, Measurements and Units
can be converted from miles to kilometers, or feet, or any
other measure of length. Often measurements are given
in one set of units, such as feet, but are needed in different
units, such as chains. A conversion factor is a numeric
expression that enables feet to be changed to chains as an
equal exchange.
A conversion factor is a number used to change one
set of units to another, by multiplying or dividing. When
a conversion is necessary, the appropriate conversion
factor to an equal value must be used. For example, to
convert inches to feet, the appropriate conversion value
is 12 inches equal 1 foot. To convert minutes to hours,
the appropriate conversion value is 60 minutes equal
1 hour. A unit cancellation table is developed by using
known units, conversion factors, and the fact that a unit
of measure ÷ the same unit of measure cancels out that
unit. The table is set up so all the units cancel except for
the unit desired. To cancel a unit, the same unit must
be in the numerator and in the denominator. When you
multiply across the table, the top number will be divided
by the bottom number, and the result will be the answer
in the desired units.
Note
The hour units on
the top and bottom
cancel along with
the minutes, leaving
seconds as the only
unit. Setting up a
unit cancellation table helps keep units
straight, even for the
most seasoned professional firefighter.
These tables are particularly important
when more than one
unit conversion is
necessary to obtain
the desired unit.
Answers should
always be presented
with the appropriate
number of significant digits.
Example 1 - Ralph wants to know how many seconds
are in 3 hours and 36 minutes.
Step 1. Change 3 hours and 36 minutes to the same units. This unit can be hours
or minutes. Using minutes is easier because the end time value will need to be in
seconds.
The appropriate conversion factor is: 1 hour = 60 minutes.
3 hours and 36 minutes = 180 minutes plus 36 minutes = 216 minutes
Step 2. Set up the cancellation table so all units will cancel, except the desired
unit, seconds.
The appropriate conversion factor is:
1 minute = 60 seconds.
There are 12,960 seconds in 3 hours 36 minutes.
18
University Chemistry
1.3.3 The Reliability of Measurements and Calculations
Counting is the only type of measurement that is free from uncertainty, provided
the number of objects being counted does not change while the counting process
is underway. The result of such a counting measurement is an example of an exact
number. If we count eggs in a carton, we know exactly how many eggs the carton
contains. The numbers of defined quantities are also exact. By definition, 1 foot
is exactly 12 inches, 1 inch is exactly 2.54 centimeters, and 1 gram is exactly 0.001
kilogram. Quantities derived from measurements other than counting, however,
are uncertain to varying extents due to practical limitations of the measurement
process used.
Significant Figures in Measurement
The numbers of measured quantities, unlike defined or directly counted
quantities, are not exact. To measure the volume of liquid in a graduated cylinder,
you should make a reading at the bottom of the meniscus, the lowest point on the
curved surface of the liquid.
Figure 12: To measure the volume of liquid in this graduated cylinder, you must mentally subdivide the distance between the 21 and 22 mL marks into tenths of a milliliter,
and then make a reading (estimate) at the bottom of the meniscus.
Refer to the illustration in Figure 12. The bottom of the meniscus in this
case clearly lies between the 21 and 22 markings, meaning the liquid volume
is certainly greater than 21 mL but less than 22 mL. The meniscus appears to be a
bit closer to the 22-mL mark than to the 21-mL mark, and so a reasonable estimate
of the liquid’s volume would be 21.6 mL. In the number 21.6, then, the digits 2 and
1 are certain, but the 6 is an estimate. Some people might estimate the meniscus
position to be equally distant from each of the markings and estimate the tenthplace digit as 5, while others may think it to be even closer to the 22-mL mark and
estimate this digit to be 7. Note that it would be pointless to attempt to estimate
a digit for the hundredths place, given that the tenths-place digit is uncertain. In
Properties, Measurements and Units
19
general, numerical scales such as the one on this graduated cylinder will permit
measurements to one-tenth of the smallest scale division. The scale in this case has
1-mL divisions, and so volumes may be measured to the nearest 0.1 mL.
This concept holds true for all measurements, even if you do not actively make
an estimate. If you place a quarter on a standard electronic balance, you may obtain
a reading of 6.72 g. The digits 6 and 7 are certain, and the 2 indicates that the mass
of the quarter is likely between 6.71 and 6.73 grams. The quarter weighs about 6.72
grams, with a nominal uncertainty in the measurement of ± 0.01 gram. If we weigh
the quarter on a more sensitive balance, we may find that its mass is 6.723 g. This
means its mass lies between 6.722 and 6.724 grams, an uncertainty of 0.001 gram.
Every measurement has some uncertainty, which depends on the device used (and
the user’s ability). All of the digits in a measurement, including the uncertain last
digit, are called significant figures or significant digits. Note that zero may be a
measured value; for example, if you stand on a scale that shows weight to the
nearest pound and it shows “120,” then the 1 (hundreds), 2 (tens) and 0 (ones) are
all significant (measured) values.
Whenever you make a measurement properly, all the digits in the result
are significant. But what if you were analyzing a reported value and trying to
determine what is significant and what is not? Well, for starters, all nonzero digits
are significant, and it is only zeros that require some thought. We will use the terms
“leading,” “trailing,” and “captive” for the zeros and will consider how to deal
with them.
Starting with the first nonzero digit on the left, count this digit and all remaining
digits to the right. This is the number of significant figures in the measurement
unless the last digit is a trailing zero lying to the left of the decimal point.
Captive zeros result from measurement and are therefore always significant.
Leading zeros, however, are never significant—they merely tell us where the
decimal point is located.
20
University Chemistry
The leading zeros in this example are not significant. We could use exponential
notation (as described in Appendix B) and express the number as 8.32407 ×× 10−3;
then the number 8.32407 contains all of the significant figures, and 10−3 locates the
decimal point.
The number of significant figures is uncertain in a number that ends with a
zero to the left of the decimal point location. The zeros in the measurement 1,300
grams could be significant or they could simply indicate where the decimal point
is located. The ambiguity can be resolved with the use of exponential notation:
1.3 ×× 103 (two significant figures), 1.30 ×× 103 (three significant figures, if the tens
place was measured), or 1.300 ×× 103 (four significant figures, if the ones place was
also measured). In cases where only the decimal-formatted number is available, it
is prudent to assume that all trailing zeros are not significant.
When determining significant figures, be sure to pay attention to reported
values and think about the measurement and significant figures in terms of what is
reasonable or likely—that is, whether the value makes sense. For example, the official
January 2014 census reported the resident population of the US as 317,297,725.
Significant Figures in Calculations
A second important principle of uncertainty is that results calculated from a
measurement are at least as uncertain as the measurement itself. We must take
the uncertainty in our measurements into account to avoid misrepresenting the
uncertainty in calculated results. One way to do this is to report the result of a
calculation with the correct number of significant figures, which is determined by
the following three rules for rounding numbers:
•
•
•
When we add or subtract numbers, we should round the result to the same
number of decimal places as the number with the least number of decimal
places (the least precise value in terms of addition and subtraction).
When we multiply or divide numbers, we should round the result to the
same number of digits as the number with the least number of significant
figures (the least precise value in terms of multiplication and division).
If the digit to be dropped (the one immediately to the right of the digit to
be retained) is less than 5, we “round down” and leave the retained digit
unchanged; if it is more than 5, we “round up” and increase the retained
21
Properties, Measurements and Units
digit by 1; if the dropped digit is 5, we round
up or down, whichever yields an even value
for the retained digit. (The last part of this
rule may strike you as a bit odd, but it’s based
on reliable statistics and is aimed at avoiding
any bias when dropping the digit “5,” since it
is equally close to both possible values of the
retained digit.)
The following examples illustrate the application of
this rule in rounding a few different numbers to three
significant figures:
•
•
•
•
0.028675 rounds “up” to 0.0287 (the dropped
digit, 7, is greater than 5)
18.3384 rounds “down” to 18.3 (the dropped
digit, 3, is less than 5)
6.8752 rounds “up” to 6.88 (the dropped digit is
5, and the retained digit is even)
92.85 rounds “down” to 92.8 (the dropped digit
is 5, and the retained digit is even)
Let’s work through these rules with a few examples.
Example 1.7.11.7.1: Rounding Numbers
Round the following to the indicated number of
significant figures:
•
•
•
•
31.57 (to two significant figures)
8.1649 (to three significant figures)
0.051065 (to four significant figures)
0.90275 (to four significant figures)
Important
Do you think
the US population was correctly
determined to the
reported nine significant figures, that is,
to the exact number
of people? People
are constantly being born, dying, or
moving into or out
of the country, and
assumptions are
made to account for
the large number of
people who are not
actually counted.
Because of these uncertainties, it might
be more reasonable
to expect that we
know the population to within
perhaps a million
or so, in which case
the population
should be reported
as 317 million, or
3.17×1083.17×108
people.
Solution
•
•
•
•
31.57 rounds “up” to 32 (the dropped digit is 5, and the retained digit is
even)
8.1649 rounds “down” to 8.16 (the dropped digit, 4, is less than 5)
0.051065 rounds “down” to 0.05106 (the dropped digit is 5, and the
retained digit is even)
0.90275 rounds “up” to 0.9028 (the dropped digit is 5, and the retained
digit is even)
22
University Chemistry
1.3.4 Significant Figures in Calculations
Before dealing with the specifics of the rules for determining the significant figures
in a calculated result, we need to be able to round numbers correctly. To round a
number, first decide how many significant figures the number should have. Once
you know that, round to that many digits, starting from the left. If the number
immediately to the right of the last significant digit is less than 5, it is dropped and
the value of the last significant digit remains the same. If the number immediately to
the right of the last significant digit is greater than or equal to 5, the last significant
digit is increased by 1. Consider the measurement 207.518m207.518m. Right now, the
measurement contains six significant figures. How would we successively round it to
fewer and fewer significant figures? Follow the process as outlined in Table.2.
Table 2: Rounding examples
Number of Significant Figures
Rounded Value
Reasoning
6
207.518
All digits are significant
5
207.52
8 rounds the 1 up to 2
4
207.5
2 is dropped
3
208
5 rounds the 7 up to 8
2
210
8 is replaced by a 0 and rounds the 0 up
to 1
1
200
1 is replaced by a 0
Notice that the more rounding that is done, the less reliable the figure is. An
approximate value may be sufficient for some purposes, but scientific work requires
a much higher level of detail.
It is important to be aware of significant figures when you are mathematically
manipulating numbers. For example, dividing 125 by 307 on a calculator gives
0.4071661238… to an infinite number of digits. But do the digits in this answer have
any practical meaning, especially when you are starting with numbers that have
only three significant figures each? When performing mathematical operations,
there are two rules for limiting the number of significant figures in an answer—one
rule is for addition and subtraction, and one rule is for multiplication and division.
In operations involving significant figures, the answer is reported in such a way that it
reflects the reliability of the least precise operation. An answer is no more precise than the
least precise number used to get the answer.
1.3.5 Multiplication and Division
For multiplication or division, the rule is to count the number of significant
figures in each number being multiplied or divided and then limit the significant
figures in the answer to the lowest count. An example is as follows:
23
Properties, Measurements and Units
The final answer, limited to four significant figures, is 4,094. The first digit
dropped is 1, so we do not round up.
Scientific notation provides a way of communicating significant figures without
ambiguity. You simply include all the significant figures in the leading number.
For example, the number 450 has two significant figures and would be written
in scientific notation as 4.5 × 102, whereas 450.0 has four significant figures and
would be written as 4.500 × 102. In scientific notation, all significant figures are listed
explicitly.
Example 1
Write the answer for each expression using scientific notation with the
appropriate number of significant figures.
•
•
23.096 × 90.300
125 × 9.000
Solution
a
Explanation
Answer
The calculator answer is 2,085.5688, 2.0856×1032.0856×103
but we need to round it to five significant figures. Because the first digit to be
dropped (in the tenths place) is greater
than 5, we round up to 2,085.6.
b
Explanation
Answer
The calculator gives 1,125 as the answer, but 1.13×1031.13×103
we limit it to three significant figures.
Addition and Subtraction
How are significant figures handled in calculations? It depends on what type
of calculation is being performed. If the calculation is an addition or a subtraction,
the rule is as follows: limit the reported answer to the rightmost column that all
numbers have significant figures in common. For example, if you were to add 1.2
and 4.71, we note that the first number stops its significant figures in the tenths
column, while the second number stops its significant figures in the hundredths
column. We therefore limit our answer to the tenths column.
24
University Chemistry
We drop the last digit—the 1—because it is not significant to the final answer.
The dropping of positions in sums and differences brings up the topic of
rounding. Although there are several conventions, in this text we will adopt the
following rule: the final answer should be rounded up if the first dropped digit is 5
or greater, and rounded down if the first dropped digit is less than 5.
Calculations Involving Multiplication/Division and Addition/Subtraction
In practice, chemists generally work with a calculator and carry all digits
forward through subsequent calculations. When working on paper, however,
we often want to minimize the number of digits we have to write out. Because
successive rounding can compound inaccuracies, intermediate rounding needs to
be handled correctly. When working on paper, always round an intermediate result
so as to retain at least one more digit than can be justified and carry this number
into the next step in the calculation. The final answer is then rounded to the correct
number of significant figures at the very end.
1.3.5. Mass Percentage Composition
Bicarbonate of soda (sodium hydrogen carbonate) is used in many commercial
preparations. Its formula is NaHCO3. Find the mass percentages (mass %) of Na, H,
C, and O in sodium hydrogen carbonate.
Solution
First, look up the atomic masses for the elements from the Periodic Table. The
atomic masses are found to be:
•
•
•
•
Na is 22.99
H is 1.01
C is 12.01
O is 16.00
Properties, Measurements and Units
25
Next, determine how many grams of each element are present in one mole of
NaHCO3:
•
•
•
•
22.99 g (1 mol) of Na
1.01 g (1 mol) of H
12.01 g (1 mol) of C
48.00 g (3 mole x 16.00 gram per mole) of O
The mass of one mole of NaHCO3 is:
22.99 g + 1.01 g + 12.01 g + 48.00 g = 84.01 g
And the mass percentages of the elements are
•
•
•
•
mass % Na = 22.99 g / 84.01 g x 100 = 27.36 %
mass % H = 1.01 g / 84.01 g x 100 = 1.20 %
mass % C = 12.01 g / 84.01 g x 100 = 14.30 %
mass % O = 48.00 g / 84.01 g x 100 = 57.14 %
Answer
•
•
•
•
mass % Na = 27.36 %
mass % H = 1.20 %
mass % C = 14.30 %
mass % O = 57.14 %
When doing mass percent calculations, it’s always a good idea to check to make
sure your mass percents add up to 100% (helps catch math errors):
27.36 + 14.30 + 1.20 + 57.14 = 100.00
Percent Composition of Water
Another simple example is finding the mass percent composition of the
elements in water, H2O.
First, find the molar mass of water by adding up the atomic masses of the
elements. Use values from the periodic table:
•
•
H is 1.01 grams per mole
O is 16.00 grams per mole
Get the molar mass by adding up all the masses of elements in the compound.
The subscript after the hydrogen (H) indicates there are two atoms of hydrogen.
There is no subscript after oxygen (O), which means only one atom is present.
•
•
molar mass = (2 x 1.01) + 16.00
molar mass = 18.02
Now, divide the mass of each element by the total mass to get the mass
percentages:
26
University Chemistry
•
•
•
•
mass % H = (2 x 1.01) / 18.02 x 100%
mass % H = 11.19%
mass % O = 16.00 / 18.02
mass % O = 88.81%
The mass percentages of hydrogen and oxygen add up to 100%.
Mass Percent of Carbon Dioxide
What are the mass percentages of carbon and oxygen in carbon dioxide, CO2?
Mass Percent Solution
Step 1: Find the mass of the individual atoms.
Look up the atomic masses for carbon and oxygen from the Periodic Table.
It›s a good idea at this point to settle on the number of significant figures you›ll be
using. The atomic masses are found to be:
•
•
C is 12.01 g/mol
O is 16.00 g/mol
Step 2: Find the number of grams of each component make up one mole of CO2.
One mole of CO2 contains 1 mole of carbon atoms and 2 moles of oxygen atoms.
•
•
12.01 g (1 mol) of C
32.00 g (2 mol x 16.00 gram per mole) of O
The mass of one mole of CO2 is:
•
12.01 g + 32.00 g = 44.01 g
Step 3: Find the mass percent of each atom.
mass % = (mass of component/mass of total) x 100
And the mass percentages of the elements are
For carbon:
•
•
•
mass % C = (mass of 1 mol of carbon/mass of 1 mol of CO2) x 100
mass % C = (12.01 g / 44.01 g) x 100
mass % C = 27.29 %
For oxygen:
•
•
•
mass % O = (mass of 1 mol of oxygen/mass of 1 mol of CO2) x 100
mass % O = (32.00 g / 44.01 g) x 100
mass % O = 72.71 %
Answer
•
mass % C = 27.29 %
Properties, Measurements and Units
•
27
mass % O = 72.71 %
Again, make sure your mass percents add up to 100%. This will help catch any
math errors.
•
27.29 + 72.71 = 100.00
The answers add up to 100%, which was expected.
Tips for Success Calculating Mass Percent
•
•
•
You won’t always be given the total mass of a mixture or solution. Often,
you’ll need to add up the masses. This might not be obvious. You may be
given mole fractions or moles and then need to convert to a mass unit.
Watch your significant figures.
Always make sure the sum of the mass percentages of all components
adds up to 100%. If it doesn’t, you need to go back and find your mistake.
EXERCISE
Answer the following questions
1. What are examples of pure substances and mixtures?
2. What are some examples of mixture substances?
3. What is the difference between a unit and a measurement?
4. Is energy an extensive or intensive property?
5. What are some conversion factors?
6. What determines the reliability of a measurement?
MULTIPLE CHOICE QUESTIONS
Tick the correct answer:
1.
Which of the following may not be a physical property?
a. Odor
b. Color
c. Density
d. Composition
2.
The observation of __________ properties needs a chemical change to occur.
a. Chemical
b. Physical
c. Extrinsic
d. Intrinsic
3.
Candela is the S.I. unit of _____
a. Luminous intensity
b. Thermodynamic temperature
28
4.
5.
6.
7.
8.
9.
10.
University Chemistry
c. Amount of substance
d. Electric current
How many scientific fundamental quantities are given S.I. units?
a. 5
b. 7
c. 3
d. 9
What is the symbol of the amount of substance’s S.I. unit?
a. K
b. s
c. mol
d. kg
What are the multiples for the prefixes yocto, atto respectively?
a. 10-24, 10-18
b. 10-9, 10-15
c. 10-15, 10-24
d. 10-24, 10-21
1Litre = _______ m3.
a. 1000
b. 0.001
c. 1
d. 10
What is the difference in units between Kelvin and centigrade scales of
temperature?
a. 212.15
b. 32
c. 298
d. 273.15
What is the human body temperature in Fahrenheit?
a. 212
b. 98.6
c. 273.15
d. 32
Convert 40°C to °F.
a. 104K
b. 313°F
29
Properties, Measurements and Units
c. 104°F
d. 313K
ANSWERS
1. (d)
2. (a)
3. (a)
4. (b)
5. (c)
6. (a)
7. (b)
8. (d)
9. (b)
10. (c)
REFERENCES
1.
2.
3.
4.
5.
6.
7.
8.
Designing Buildings Wiki, Standard Method of Measurement, accessed 1 July
2020
Dodd, Richard (2012). Using SI Units in Astronomy. Cambridge University
Press. p. 246. doi:10.1017/CBO9781139019798. ISBN 9780521769174.
Douglas Hubbard: “How to Measure Anything”, Wiley (2007), p. 2
Gill, Simeon; Parker, Christopher J. (2017). “Scan posture definition and
hip girth measurement: the impact on clothing design and body scanning”.
Ergonomics. 60 (8): 1123–1136. doi:10.1080/00140139.2016.1251621. PMID
27764997. S2CID 23758581.
Montévil, Maël (2019). “Measurement in biology is methodized by theory”.
Biology & Philosophy. 34 (3). doi:10.1007/s10539-019-9687-x. ISSN 0169-3867.
S2CID 96447209.
RICS, RICS standards and guidance - SMM7: Standard method of measurement
of building works, accessed 1 July 2020
Young, Hugh D; Freedman, Roger A. (2012). University Physics (13 ed.).
Pearson Education Inc. ISBN 978-0-321-69686-1.
Yunus A. Çengel; Michael A. Boles (2002). Thermodynamics: An Engineering
Approach (Eighth ed.). McGraw Hill. p. 996. ISBN 9780073398174.
31
The Composition of Matter
CHAPTER 2
THE COMPOSITION OF MATTER
OBJECTIVES
After studying this chapter, you will be able to:
1. Describe the names and symbols of the elements
2. Learn the atoms
3. Describe the types of compounds
INTRODUCTION
Matter is everything around us. Atoms and molecules are all composed of
matter. Matter is anything that has mass and takes up space. Matter consist of atoms
that are divisible and composed of protons, neutrons and electrons.
Anything that occupies space and has mass is known as Matter. So everything
that we see around us is known as Matter. Matter is basically made up of Atoms and
Molecules. Depending upon its physical state, the nature of matter falls into three
categories: Solids, Liquids and Gases.
•
Solids- These are the substances where the constituent particles (atoms
or molecules) are held together so tightly, that it is impossible for them
to move from there position. They have definite shape and fixed volume.
32
University Chemistry
•
•
Liquids - These are the substances where forces between the particles are
weak enough to allow their movement. They don’t have a specific shape
but they have specific volume.
Gases - These substances have very weak forces between their particles.
This causes the constituent particles to move freely. They have neither
fixed shape nor a definite volume. They tend to completely occupy the
container in which they are placed.
The nature of matter is also determined by its composition. If the matter is
composed of more than one type of particles then it is called as a mixture while if it
consists of a single type of particles then it is known as a pure substance.
2.1 ELEMENTS
An element is a pure substance consisting only of atoms that all have the same
numbers of protons in their nuclei. Unlike chemical compounds, chemical elements
cannot be broken down into simpler substances by any chemical reaction. The
number of protons in the nucleus is the defining property of an element, and is
referred to as its atomic number (represented by the symbol Z) – all atoms with the
same atomic number are atoms of the same element. All of the baryonic matter of
the universe is composed of chemical elements. When different elements undergo
chemical reactions, atoms are rearranged into new compounds held together by
chemical bonds. Only a minority of elements, such as silver and gold, are found
uncombined as relatively pure native element minerals. Nearly all other naturally
occurring elements occur in the Earth as compounds or mixtures. Air is primarily
a mixture of the elements nitrogen, oxygen, and argon, though it does contain
compounds including carbon dioxide and water.
The history of the discovery and use of the elements began with primitive
human societies that discovered native minerals like carbon, sulfur, copper and
gold (though the concept of a chemical element was not yet understood). Attempts
to classify materials such as these resulted in the concepts of classical elements,
alchemy, and various similar theories throughout human history. Much of the
modern understanding of elements developed from the work of Dmitri Mendeleev,
a Russian chemist who published the first recognizable periodic table in 1869. This
table organizes the elements by increasing atomic number into rows (“periods”) in
which the columns (“groups”) share recurring (“periodic”) physical and chemical
properties. The periodic table summarizes various properties of the elements,
allowing chemists to derive relationships between them and to make predictions
about compounds and potential new ones.
By November 2016, the International Union of Pure and Applied Chemistry
had recognized a total of 118 elements. The first 94 occur naturally on Earth, and
the remaining 24 are synthetic elements produced in nuclear reactions. Save for
unstable radioactive elements (radionuclides) which decay quickly, nearly all of the
The Composition of Matter
33
elements are available industrially in varying amounts. The discovery and synthesis
of further new elements is an ongoing area of scientific study.
2.1.1 The Names and Symbols of the Elements
Each element has a name. Many of these names are already familiar to you gold, silver, copper, chlorine, platinum, carbon, oxygen, and nitrogen. The names
themselves are interesting. Many refer to a property of the element. The Latin
name for gold is aurum, meaning “shining dawn.” The Latin name for mercury,
hydrargyrum, means “liquid silver.”
The practice of naming an element after one of its properties continues. Cesium
was discovered in 1860 by the German chemist Bunsen (the inventor of the Bunsen
burner). Because this element imparts a blue color to a flame, Bunsen named it
cesium from the Latin word caesius, meaning “sky blue.”
Other elements are named for people. Curium is named for Marie Curie (18671934), a pioneer in the study of radioactivity. Marie Curie, a French scientist of
Polish birth, was awarded the Nobel Prize in Physics in 1903 for her studies of
radioactivity. She was also awarded the Nobel Prize in Chemistry in 1911 for her
discovery of the elements polonium (named after Poland) and radium (Latin, radius,
“ray”).
Some elements are named for places. The small town of Ytterby in Sweden has
four elements named for it: terbium, yttrium, erbium, and ytterbium. Californium
is another example of an element named for the place where it was first observed.
This element does not occur in nature. It was first produced in 1950 in the Radiation
Laboratory at the University of California, Berkeley, by a team of scientists headed
by Glenn Seaborg. Seaborg was also the first to identify curium at the metallurgical
laboratory at the University of Chicago (now Argonne National Laboratory) in 1944.
Seaborg himself was named a Nobel laureate in 1951 in honor of his pioneering
work in the preparation of other unknown elements.
Each element has a symbol, one or two letters that represent the element much
as your initials represent you. The symbol of an element represents one atom of that
element. For 14 of the elements, the symbol consists of one letter. With the possible
exceptions of yttrium (Y) and vanadium (V), you are probably familiar with the
names of all elements having one-letter symbols. These elements are listed in Table
1. For 12 of these elements, the symbol is the first letter of the name.
Potassium was discovered in 1807 and named for potash, the substance from
which potassium was first isolated. Potassium’s symbol, K, comes from kalium,
the Latin word for potash. Tungsten, discovered in 1783, has the symbol W, for
wolframite, the mineral from which tungsten was first isolated.
34
University Chemistry
Table 1. Elements with one-letter symbols
Symbol
Element
B
boron
C
carbon
F
fluorine
H
hydrogen
I
iodine
N
nitrogen
O
oxygen
P
phosphorus
K
potassium
S
sulfur
W
tungsten
U
uranium
V
vanadium
Y
yttrium
Most other elements have two-letter symbols. In these two-letter symbols,
the first letter is always capitalized and the second is always lowercased. Eleven
elements have names (and symbols) beginning with the letter C. One of these,
carbon, has a one-letter symbol, C. The other ten have two-letter symbols (see Table
2).
Table 2. Elements whose name begins with the letter C
Symbol
Element
Cd
cadmium
Ca
calcium
Cf
californium
C
carbon
Ce
cerium
Cs
cesium
Cl
chlorine
Cr
chromium
Co
cobalt
Cu
copper
Cm
curium
2.1.2 The Periodic Table
The properties of the chemical elements are often summarized using the periodic
table, which powerfully and elegantly organizes the elements by increasing atomic
The Composition of Matter
35
number into rows (“periods”) in which the columns (“groups”) share recurring
(“periodic”) physical and chemical properties. The current standard table contains
118 confirmed elements as of 2021.
Although earlier precursors to this presentation exist, its invention is generally
credited to the Russian chemist Dmitri Mendeleev in 1869, who intended the table
to illustrate recurring trends in the properties of the elements. The layout of the table
has been refined and extended over time as new elements have been discovered
and new theoretical models have been developed to explain chemical behavior.
Use of the periodic table is now ubiquitous within the academic discipline of
chemistry, providing an extremely useful framework to classify, systematize and
compare all the many different forms of chemical behavior. The table has also
found wide application in physics, geology, biology, materials science, engineering,
agriculture, medicine, nutrition, environmental health, and astronomy. Its principles
are especially important in chemical engineering.
2.2 ATOMS
An atom is the smallest unit of ordinary matter that forms a chemical element.
Every solid, liquid, gas, and plasma is composed of neutral or ionized atoms. Atoms
are extremely small, typically around 100 picometers across. They are so small that
accurately predicting their behavior using classical physics—as if they were tennis
balls, for example—is not possible due to quantum effects.
Every atom is composed of a nucleus and one or more electrons bound to the
nucleus. The nucleus is made of one or more protons and a number of neutrons.
Only the most common variety of hydrogen has no neutrons. More than 99.94% of
an atom’s mass is in the nucleus. The protons have a positive electric charge, the
36
University Chemistry
electrons have a negative electric charge, and the neutrons have no electric charge. If
the number of protons and electrons are equal, then the atom is electrically neutral.
If an atom has more or fewer electrons than protons, then it has an overall negative
or positive charge, respectively – such atoms are called ions.
The electrons of an atom are attracted to the protons in an atomic nucleus by the
electromagnetic force. The protons and neutrons in the nucleus are attracted to each
other by the nuclear force. This force is usually stronger than the electromagnetic
force that repels the positively charged protons from one another. Under certain
circumstances, the repelling electromagnetic force becomes stronger than the
nuclear force. In this case, the nucleus splits and leaves behind different elements.
This is a form of nuclear decay.
The number of protons in the nucleus is the atomic number and it defines to which
chemical element the atom belongs. For example, any atom that contains 29 protons
is copper. The number of neutrons defines the isotope of the element. Atoms can
attach to one or more other atoms by chemical bonds to form chemical compounds
such as molecules or crystals. The ability of atoms to associate and dissociate is
responsible for most of the physical changes observed in nature. Chemistry is the
discipline that studies these changes.
2.2.1 The Nuclear Atom
An atom is defined as the smallest particle into which an element can be divided
without losing the chemical properties that characterize it.
It is composed by the atomic nucleus, a positively charged center where most
of the mass is concentrated, and by the electrons, a certain number of negatively
charged particles that make up the cortex.
The atomic nucleus is composed by protons and neutrons, thus denominated
nucleons, with a positive charge equal to the negative charge of the electrons, in
such a way that the total electric charge of the atom is neutral (protons have positive
charge and neutrons have no electric charge).
The Composition of Matter
37
The particles that constitute the atom, along with their mass and charge, are:
•
•
•
Electron: Mass (kg) = 9.1 x 10-31, Charge (C) = 1.602 x 10-19 (-)
Proton: Mass (kg) = 1.673 x 10-27, Charge (C) = 1.602 x 10-19 (+)
Neutron: Mass (kg) = 1.696 x 10-27, Charge (C) = 0
2.2.2 The Masses of Atoms
The atomic mass (ma or m) is the mass of an atom. Although the SI unit of
mass is the kilogram (symbol: kg), atomic mass is often expressed in the non-SI
unit atomic mass unit (amu) or unified mass (u) or dalton (symbol: Da), where 1
amu or 1 u or 1 Da is defined as 1⁄12 of the mass of a single carbon-12 atom, at rest.
The protons and neutrons of the nucleus account for nearly all of the total mass of
atoms, with the electrons and nuclear binding energy making minor contributions.
Thus, the numeric value of the atomic mass when expressed in daltons has nearly
the same value as the mass number. Conversion between mass in kilograms and
mass in daltons can be done using the atomic mass constant
mu =
m(12 C )
=1 Da.
12
The formula used for conversion is:
1 Da
= m=
u
M u M (12 C )
=
= 1.660 539 066 60(50) ×10−27 kg ,
NA
12 N A
where Mu is the molar mass constant, NA is the Avogadro constant, and M(12C)
is the experimentally determined molar mass of carbon-12.
The relative isotopic mass (see section below) can be obtained by dividing the
atomic mass ma of an isotope by the atomic mass constant mu yielding a dimensionless
value. Thus, the atomic mass of a carbon-12 atom is 12 Da by definition, but the
relative isotopic mass of a carbon-12 atom is simply 12. The sum of relative isotopic
masses of all atoms in a molecule is the relative molecular mass.
The atomic mass of an isotope and the relative isotopic mass refers to a certain
specific isotope of an element. Because substances are usually not isotopically pure,
it is convenient to use the elemental atomic mass which is the average (mean) atomic
mass of an element, weighted by the abundance of the isotopes. The dimensionless
(standard) atomic weight is the weighted mean relative isotopic mass of a (typical
naturally-occurring) mixture of isotopes.
The atomic mass of atoms, ions, or atomic nuclei is slightly less than the sum
of the masses of their constituent protons, neutrons, and electrons, due to binding
energy mass loss (per E = mc2).
38
University Chemistry
2.2.3 Moles and Molar Mass
In Dalton’s theory each chemical compound has a
Important
particular combination of atoms and that the ratios of the
One mole
numbers of atoms of the elements present are usually small
always has the same
whole numbers. We also described the law of multiple
number of objects:
proportions, which states that the ratios of the masses
6.022 × 1023.
of elements that form a series of compounds are small
whole numbers. The problem for Dalton and other early
chemists was to discover the quantitative relationship
between the number of atoms in a chemical substance and its mass. Because the
masses of individual atoms are so minuscule (on the order of 10−23 g/atom), chemists
do not measure the mass of individual atoms or molecules. In the laboratory, for
example, the masses of compounds and elements used by chemists typically range
from milligrams to grams, while in industry, chemicals are bought and sold in
kilograms and tons. To analyze the transformations that occur between individual
atoms or molecules in a chemical reactionit is therefore absolutely essential for
chemists to know how many atoms or molecules are contained in a measurable
quantity in the laboratory—a given mass of sample. The unit that provides this link
is the mole (mol). , from the Latin moles, meaning “pile” or “heap” (not from the
small subterranean animal!).
Many familiar items are sold in numerical quantities that have unusual names.
For example, cans of soda come in a six-pack, eggs are sold by the dozen (12), and
pencils often come in a gross (12 dozen, or 144). Sheets of printer paper are packaged
in reams of 500, a seemingly large number. Atoms are so small, however, that even
500 atoms are too small to see or measure by most common techniques. Any readily
measurable mass of an element or compound contains an extraordinarily large
number of atoms, molecules, or ions, so an extraordinarily large numerical unit is
needed to count them. The mole is used for this purpose.
A mole is defined as the amount of a substance that contains the number of
carbon atoms in exactly 12 g of isotopically pure carbon-12. According to the most
recent experimental measurements, this mass of carbon-12 contains 6.022142 ×
1023 atoms, but for most purposes 6.022 × 1023 provides an adequate number of
significant figures. Just as 1 mol of atoms contains 6.022 × 1023 atoms, 1 mol of eggs
contains 6.022 × 1023 eggs. The number in a mole is called Avogadro’s number, after
the 19th-century Italian scientist who first proposed how to measure the number
of molecules in a gas. Since the mass of the gas can also be measured on a sensitive
balance, knowing both the number of molecules and their total mass allows us to
simply determine the mass of a single molecule in grams.
The mole provides a bridge between the atomic world (amu) and the laboratory
(grams). It allows determination of the number of molecules or atoms by weighing
them. The numerical value of Avogadro’s number, usually written as No, is a
consequence of the arbitrary value of one kilogram, a block of Pt-Ir metal called the
The Composition of Matter
39
International Prototype Kilogram, and the choice of reference for the atomic mass
unit scale, one atom of carbon-12. A mole of C-12 by definition weighs exactly 12 g
and Avogadro’s number is determined by counting the number of atoms. It is not
so easy. Avogadro’s number is the fundamental constant that is least accurately
determined.
The definition of a mole—that is, the decision to base it on 12 g of carbon-12—is
arbitrary but one arrived at after some discussion between chemists and physicists
debating about whether to use naturally occurring carbon, a mixture of C-12 and
C-13, or hydrogen. The important point is that 1 mol of carbon—or of anything else,
whether atoms, compact discs, or houses—always has the same number of objects: 6.022 ×
1023.
To appreciate the magnitude of Avogadro’s number, consider a mole of pennies.
Stacked vertically, a mole of pennies would be 4.5 × 1017 mi high, or almost six times
the diameter of the Milky Way galaxy. If a mole of pennies were distributed equally
among the entire population on Earth, each person would get more than one trillion
dollars. Clearly, the mole is so large that it is useful only for measuring very small
objects, such as atoms.
The concept of the mole allows us to count a specific
number of individual atoms and molecules by weighing
Important
measurable quantities of elements and compounds. To
obtain 1 mol of carbon-12 atoms, we would weigh out 12
The molar
mass of any subg of isotopically pure carbon-12. Because each element
stance is its atomic
has a different atomic mass, however, a mole of each
mass, molecular
element has a different mass, even though it contains the
mass, or formula
23
same number of atoms (6.022 × 10 ). This is analogous to
mass in grams per
the fact that a dozen extra-large eggs weighs more than
mole.
a dozen small eggs, or that the total weight of 50 adult
humans is greater than the total weight of 50 children.
Because of the way in which the mole is defined, for every element the number
of grams in a mole is the same as the number of atomic mass units in the atomic
mass of the element. For example, the mass of 1 mol of magnesium (atomic mass
= 24.305 amu) is 24.305 g. Because the atomic mass of magnesium (24.305 amu) is
slightly more than twice that of a carbon-12 atom (12 amu), the mass of 1 mol of
magnesium atoms (24.305 g) is slightly more than twice that of 1 mol of carbon-12
(12 g). Similarly, the mass of 1 mol of helium (atomic mass = 4.002602 amu) is
4.002602 g, which is about one-third that of 1 mol of carbon-12. Using the concept
of the mole, we can now restate Dalton’s theory: 1 mol of a compound is formed by
combining elements in amounts whose mole ratios are small whole numbers. For example,
1 mol of water (H2O) has 2 mol of hydrogen atoms and 1 mol of oxygen atoms.
Molar Mass
The molar mass is defined as the mass in grams of 1 mol of that substance.
One mole of isotopically pure carbon-12 has a mass of 12 g. For an element, the
40
University Chemistry
molar mass is the mass of 1 mol of atoms of that element; for a covalent molecular
compound, it is the mass of 1 mol of molecules of that compound; for an ionic
compound, it is the mass of 1 mol of formula units. That is, the molar mass of a
substance is the mass (in grams per mole) of 6.022 × 1023 atoms, molecules, or formula
units of that substance. In each case, the number of grams in 1 mol is the same as the
number of atomic mass units that describe the atomic mass, the molecular mass, or
the formula mass, respectively.
The periodic table lists the atomic mass of carbon as 12.011 amu; the average
molar mass of carbon—the mass of 6.022 × 1023 carbon atoms—is therefore 12.011
g/mol:
Substance (formula)
Atomic, Molecular, or
Formula Mass (amu)
Molar Mass (g/
mol)
carbon (C)
12.011 (atomic mass)
12.011
ethanol (C2H5OH)
46.069 (molecular mass)
46.069
calcium phosphate
[Ca3(PO4)2]
310.177 (formula mass)
310.177
The molar mass of naturally occurring carbon is different from that of carbon-12
and is not an integer because carbon occurs as a mixture of carbon-12, carbon-13,
and carbon-14. One mole of carbon still has 6.022 × 1023 carbon atoms, but 98.89%
of those atoms are carbon-12, 1.11% are carbon-13, and a trace (about 1 atom in
1012) are carbon-14. Similarly, the molar mass of uranium is 238.03 g/mol, and the
molar mass of iodine is 126.90 g/mol. When we deal with elements such as iodine
and sulfur, which occur as a diatomic molecule (I2) and a polyatomic molecule (S8),
respectively, molar mass usually refers to the mass of 1 mol of atoms of the element—
in this case I and S, not to the mass of 1 mol of molecules of the element (I2 and S8).
The molar mass of ethanol is the mass of ethanol (C2H5OH) that contains 6.022
× 1023 ethanol molecules. As you calculated in Example 1, the molecular mass of
ethanol is 46.069 amu. Because 1 mol of ethanol contains 2 mol of carbon atoms (2
× 12.011 g), 6 mol of hydrogen atoms (6 × 1.0079 g), and 1 mol of oxygen atoms (1
× 15.9994 g), its molar mass is 46.069 g/mol. Similarly, the formula mass of calcium
phosphate [Ca3(PO4)2] is 310.177 amu, so its molar mass is 310.177 g/mol. This is the
mass of calcium phosphate that contains 6.022 × 1023 formula units.
The mole is the basis of quantitative chemistry. It provides chemists with a way
to convert easily between the mass of a substance and the number of individual
atoms, molecules, or formula units of that substance. Conversely, it enables chemists
to calculate the mass of a substance needed to obtain a desired number of atoms,
molecules, or formula units. For example, to convert moles of a substance to mass,
we use the relationship
(moles)(molarmass)→mass
or, more specifically,
The Composition of Matter
41
 grams 
moles 
 = grams
 mole 
mass
(
) → moles
molarmass


 mole 
grams
=
=
 grams

 moles
 grams / mole 
 grams 
Be sure to pay attention to the units when converting between mass and moles.
Figure 2 is a flowchart for converting between mass; the number of moles; and
the number of atoms, molecules, or formula units.
Figure 2. A Flowchart for Converting between Mass; the Number of Moles; and the
Number of Atoms, Molecules, or Formula Units.
2.3 COMPOUNDS
A chemical compound is a chemical substance composed of many identical
molecules (or molecular entities) composed of atoms from more than one element
held together by chemical bonds. A molecule consisting of atoms of only one
element is therefore not a compound.
There are four types of compounds, depending on how the constituent atoms
are held together:
•
•
•
•
molecules held together by covalent bonds
ionic compounds held together by ionic bonds
intermetallic compounds held together by metallic bonds
certain complexes held together by coordinate covalent bonds.
A chemical formula specifies the number of atoms of each element in a
compound molecule, using the standard abbreviations for the chemical elements
and numerical subscripts. For example, a water molecule has formula H2O indicating
two hydrogen atoms bonded to one oxygen atom. Many chemical compounds
have a unique CAS number identifier assigned by the Chemical Abstracts Service.
Globally, more than 350,000 chemical compounds (including mixtures of chemicals)
have been registered for production and use.
42
University Chemistry
A compound can be converted to a different chemical substance by interaction
with a second substance via a chemical reaction. In this process, bonds between
atoms may be broken in either or both of the interacting substances, and new bonds
formed.
2.3.1 Molecules and Molecular Compounds
A molecule is an electrically neutral group of two or
more atoms held together by chemical bonds. A molecule
may be homonuclear, that is, it consists of atoms of one
chemical element, as with two atoms in the oxygen
molecule (O2); or it may be heteronuclear, a chemical
compound composed of more than one element, as with
water (two hydrogen atoms and one oxygen atom; H2O).
Hints
Shades of
gray exist between
ionic and molecular
compounds.
Many compounds do not contain ions but instead
consist solely of discrete, neutral molecules. These
molecular compounds (covalent compounds) result when atoms share, rather than
transfer (gain or lose), electrons. Covalent bonding is an important and extensive
concept in chemistry, and it will be treated in considerable detail in a later module of
this text. We can often identify molecular compounds on the basis of their physical
properties. Under normal conditions, molecular compounds often exist as gases,
low-boiling liquids, and low-melting solids, although many important exceptions
exist.
Whereas ionic compounds are usually formed when a metal and a nonmetal
combine, covalent compounds are usually formed by a combination of nonmetals.
Thus, the periodic table can help us recognize many of the compounds that are
covalent. While we can use the positions of a compound’s elements in the periodic
table to predict whether it is ionic or covalent at this point in our study of chemistry,
you should be aware that this is a very simplistic approach that does not account for
a number of interesting exceptions.
2.3.2 Ions and Ionic Compounds
An ion is an atom or molecule with a net electrical charge.
The charge of an electron is considered to be negative by convention and this
charge is equal and opposite to the charge of a proton, which is considered to be
positive by convention. The net charge of an ion is not zero because its total number
of electrons is unequal to its total number of protons.
A cation is a positively charged ion with fewer electrons than protons while
an anion is a negatively charged ion with more electrons than protons. Opposite
electric charges are pulled towards one another by electrostatic force, so cations and
anions attract each other and readily form ionic compounds.
The Composition of Matter
43
Ions consisting of only a single atom are termed atomic or monatomic ions,
while two or more atoms form molecular ions or polyatomic ions. In the case of
physical ionization in a fluid (gas or liquid), “ion pairs” are created by spontaneous
molecule collisions, where each generated pair consists of a free electron and a
positive ion. Ions are also created by chemical interactions, such as the dissolution
of a salt in liquids, or by other means, such as passing a direct current through a
conducting solution, dissolving an anode via ionization.
An ionic compound is a chemical compound composed of ions held together
by electrostatic forces termed ionic bonding. The compound is neutral overall,
but consists of positively charged ions called cations and negatively charged ions
called anions. These can be simple ions such as the sodium (Na+) and chloride
(Cl−) in sodium chloride, or polyatomic species such as the ammonium (NH+4) and
carbonate (CO2−3) ions in ammonium carbonate. Individual ions within an ionic
compound usually have multiple nearest neighbours, so are not considered to be
part of molecules, but instead part of a continuous three-dimensional network,
usually in a crystalline structure.
Ionic compounds containing basic ions hydroxide (OH−) or oxide (O2−) are
classified as bases. Ionic compounds without these ions are also known as salts
and can be formed by acid–base reactions. Ionic compounds can also be produced
from their constituent ions by evaporation of their solvent, precipitation, freezing, a
solid-state reaction, or the electron transfer reaction of reactive metals with reactive
non-metals, such as halogen gases.
Ionic compounds typically have high melting and boiling points, and are hard
and brittle. As solids they are almost always electrically insulating, but when melted
or dissolved they become highly conductive, because the ions are mobilized.
2.3.3 Chemical Nomenclature
A chemical nomenclature is a set of rules to generate systematic names for
chemical compounds. The nomenclature used most frequently worldwide is the one
created and developed by the International Union of Pure and Applied Chemistry
(IUPAC).
The IUPAC’s rules for naming organic and inorganic compounds are contained
in two publications, known as the Blue Book and the Red Book, respectively. A third
publication, known as the Green Book, describes the recommendations for the use of
symbols for physical quantities (in association with the IUPAP), while a fourth, the
Gold Book, contains the definitions of many technical terms used in chemistry. Similar
compendia exist for biochemistry (the White Book, in association with the IUBMB),
analytical chemistry (the Orange Book), macromolecular chemistry (the Purple Book)
and clinical chemistry (the Silver Book). These “color books” are supplemented by
shorter recommendations for specific circumstances that are published periodically
in the journal Pure and Applied Chemistry.
44
University Chemistry
EXERCISE
Answer the following questions
1. What is matter?
2. What you mean by elements?
3. What is Latin name of gold?
4. What is atom?
5. What is the SI unit symbol for mass?
6. What is a chemical compound composed of?
7. List the four types of compounds.
MULTIPLE CHOICE QUESTIONS
Tick the correct answer:
1.
Substances formed from atoms of two or more elements are called
a. mixtures
b. compounds
c. colloids
d. elements
2.
A heterogeneous mixture in which visible particles settle is called a
a. colloid
b. solution
c. suspension
d. Tyndall effect
3.
According to the ______________, matter is neither created nor destroyed
during a chemical change.
a. Tyndall effect
b. law of conservation of mass
c. law of mass
d. chemical property law
4.
Compounds are made from the atoms of two or more ______.
a. elements
b. colloids
c. substances
d. solutions
5.
Which of the following is an example of a chemical change?
a. boiling
b. evaporation
45
The Composition of Matter
6.
7.
8.
9.
10.
c. burning
d. melting
What type of substance is gelatin?
a. colloid
b. compound
c. substance
d. suspension
The formation of water when hydrogen burns is an example of ______.
a. chemical change
b. chemical property
c. physical change
d. physical property
Which of these warning refers to a chemical property of the material?
a. fragile
b. flammable
c. handle with care
d. shake well
Which of the following is a substance?
a. colloid
b. mixture
c. element
d. solution
Copper is an example of _______.
a. an element
b. a compound
c. a colloid
d. a suspension
ANSWERS
1. (b)
2. (c)
3. (b)
4. (a)
5. (c)
6. (a)
7. (a)
8. (b)
9. (c)
10. (a)
REFERENCES
1.
2.
Ball, P. (2004). The Elements: A Very Short Introduction. Oxford University
Press. ISBN 978-0-19-284099-8.
Emsley, J. (2003). Nature’s Building Blocks: An A–Z Guide to the Elements.
Oxford University Press. ISBN 978-0-19-850340-8.
46
3.
University Chemistry
Gray, T. (2009). The Elements: A Visual Exploration of Every Known Atom in
the Universe. Black Dog & Leventhal Publishers Inc. ISBN 978-1-57912-814-2.
4. Iannone, A. Pablo (2001). Dictionary of World Philosophy. Routledge. ISBN
978-0-415-17995-9. OCLC 44541769.
5. Kean, Sam (2011). The Disappearing Spoon: And Other True Tales of Madness,
Love, and the History of the World from the Periodic Table of the Elements.
Back Bay Books.
6. McEvilley, Thomas (2002). The shape of ancient thought: comparative studies
in Greek and Indian philosophies. Allworth Press. ISBN 978-1-58115-203-6.
7. Robert Siegfried (2002), From elements to atoms: a history of chemical
composition, American Philosophical Society, ISBN 978-0-87169-924-4
8. Scerri, E.R. (2007). The Periodic Table, Its Story and Its Significance. Oxford
University Press. ISBN 978-0-19-530573-9.
9. Siegfried, Robert (2002). From Elements to Atoms: A History of Chemical
Composition. DIANE. ISBN 978-0-87169-924-4. OCLC 186607849.
10. Strathern, P. (2000). Mendeleyev’s Dream: The Quest for the Elements. Hamish
Hamilton Ltd. ISBN 978-0-241-14065-9.
11. Teresi, Dick (2003). Lost Discoveries: The Ancient Roots of Modern Science.
Simon & Schuster. pp. 213–214. ISBN 978-0-7432-4379-7.
47
Chemical Reactions
CHAPTER 3
CHEMICAL REACTIONS
OBJECTIVES
After studying this chapter, you will be able to:
1. Understand the chemical equations
2. Discuss on precipitation reactions
3. Explain acid-base reactions
4. Define redox reactions
INTRODUCTION
A chemical reaction is a process that leads to the chemical transformation of one
set of chemical substances to another. Classically, chemical reactions encompass
changes that only involve the positions of electrons in the forming and breaking
of chemical bonds between atoms, with no change to the nuclei (no change to the
elements present), and can often be described by a chemical equation. Nuclear
chemistry is a sub-discipline of chemistry that involves the chemical reactions of
unstable and radioactive elements where both electronic and nuclear changes can
occur.
The substance (or substances) initially involved in a chemical reaction are called
reactants or reagents. Chemical reactions are usually characterized by a chemical
48
University Chemistry
change, and they yield one or more products, which usually have properties
different from the reactants. Reactions often consist of a sequence of individual
sub-steps, the so-called elementary reactions, and the information on the precise
course of action is part of the reaction mechanism. Chemical reactions are described
with chemical equations, which symbolically present the starting materials, end
products, and sometimes intermediate products and reaction conditions.
Chemical reactions happen at a characteristic reaction rate at a given
temperature and chemical concentration. Typically, reaction rates increase with
increasing temperature because there is more thermal energy available to reach the
activation energy necessary for breaking bonds between atoms.
Reactions may proceed in the forward or reverse direction until they go to
completion or reach equilibrium. Reactions that proceed in the forward direction to
approach equilibrium are often described as spontaneous, requiring no input of free
energy to go forward. Non-spontaneous reactions require input of free energy to
go forward (examples include charging a battery by applying an external electrical
power source, or photosynthesis driven by absorption of electromagnetic radiation
in the form of sunlight).
3.1 CHEMICAL EQUATIONS
Chemical equations are symbolic representations of chemical reactions in which
the reactants and the products are expressed in terms of their respective chemical
formulae. They also make use of symbols to represent factors such as the direction
of the reaction and the physical states of the reacting entities.
Chemical reactions are represented on paper by chemical equations. For
example, hydrogen gas (H2) can react (burn) with oxygen gas (O2) to form water
(H2O). The chemical equation for this reaction is written as:
(1)
Chemical formulas and other symbols are used to indicate the starting materials,
or reactants, which by convention are written on the left side of the equation, and
the final compounds, or products, which are written on the right. An arrow points
from the reactant to the products.
49
Chemical Reactions
(2)
The arrow is read as “yields” or “reacts to form.” Equation 2 indicates that
ammonium dichromate (the reactant) yields chromium(III) oxide, nitrogen, and
water (the products). The equation for this reaction is even more informative when
written as follows:
(3)
Equation 3 is identical to Equation 2 except for the addition of abbreviations in
parentheses to indicate the physical state of each species. The abbreviations are (s)
for solid, (l) for liquid, (g) for gas, and (aq) for an aqueous solution, a solution of the
substance in water.
Consistent with the law of conservation of mass, the numbers of each type of
atom are the same on both sides of Equations 2 and 3. Each side of the reaction
has two chromium atoms, seven oxygen atoms, two nitrogen atoms, and eight
hydrogen atoms.
In a balanced chemical equation, both the numbers of each type of atom and
the total charge are the same on both sides. Equations 2 and 3 are balanced chemical
equations. What is different on each side of the equation is how the atoms are
arranged to make molecules or ions. A chemical reaction represents a change in the
distribution of atoms, but not in the number of atoms. In this reaction, and in most
chemical reactions, bonds are broken in the reactants (here, Cr–O and N–H bonds),
and new bonds are formed to create the products (here, O–H and N≡N bonds).
If the numbers of each type of atom are different on the two sides of a chemical
equation, then the equation is unbalanced, and it cannot correctly describe what
happens during the reaction. To proceed, the equation must first be balanced.
3.1.1 Symbolizing Reactions
Symbol Meaning
+
used to separate one reactant or product from another
used to separate the reactants from the products - it is pronounced “yields” or
“produces” when the equation is read
used when the reaction can proceed in both directions - this is called an
equilibrium arrow and will be used later in the course
(g)
indicates that the substance is in a gaseous state
an alternative way of representing a substance in a gaseous state
(s)
indicates that the substance is in a solid state
50
University Chemistry
an alternative way of representing a substance in a solid state
(aq)
indicates that the substance is dissolved in water - the aq comes from aqueous
indicates that heat is applied to make the reaction proceed
3.1.2 Balancing Equations
According to the law of conservation of mass, matter can neither be created
nor destroyed. Thus, mass of each element present in the products of a chemical
reaction must be equal to its mass present in the reactants. In other words, the
number of atoms of each element remains the same before and after a chemical
reaction. In a balanced chemical equation number of atoms of a particular element
present in the reactants and products must be equal. If not, equation is said to be
‘not balanced.
One balances a chemical equation by changing the scalar number for each
chemical formula. Simple chemical equations can be balanced by inspection, that is,
by trial and error. Another technique involves solving a system of linear equations.
Balanced equations are often written with smallest whole-number coefficients.
If there is no coefficient before a chemical formula, the coefficient is 1.
The method of inspection can be outlined as putting a coefficient of 1 in front
of the most complex chemical formula and putting the other coefficients before
everything else such that both sides of the arrows have the same number of each
atom. If any fractional coefficient exists, multiply every coefficient with the smallest
number required to make them whole, typically the denominator of the fractional
coefficient for a reaction with a single fractional coefficient.
As an example, seen in the above image, the burning of methane would be
balanced by putting a coefficient of 1 before the CH4:
1 CH4 + O2 → CO2 + H2O
Since there is one carbon on each side of the arrow, the first atom (carbon) is
balanced.
Looking at the next atom (hydrogen), the right-hand side has two atoms, while
the left-hand side has four. To balance the hydrogens, 2 goes in front of the H2O,
which yields:
51
Chemical Reactions
1 CH4 + O2 → CO2 + 2 H2O
Inspection of the last atom to be balanced (oxygen) shows that the right-hand
side has four atoms, while the left-hand side has two. It can be balanced by putting
a 2 before O2, giving the balanced equation:
CH4 + 2 O2 → CO2 + 2 H2O
This equation does not have any coefficients in front of CH4 and CO2, since a
coefficient of 1 is dropped.
3.2 PRECIPITATION REACTIONS
Precipitation refers to a chemical reaction that occurs in aqueous solution when
two ions bond together to form an insoluble salt, which is known as the precipitate.
A precipitation reaction can occur when two solutions containing different salts
are mixed, and a cation/anion pair in the resulting combined solution forms an
insoluble salt; this salt then precipitates out of solution.
The following is a common laboratory example of a precipitation reaction.
Aqueous silver nitrate (AgNO3) is added to a solution containing potassium chloride
(KCl), and the precipitation of a white solid, silver chloride (AgCl), is observed:
AgNO3 (aq) + KCl (aq) → AgCl (s) + KNO3(aq)
Note that the product silver chloride is the precipitate,
and it is designated as a solid. This reaction can be also be
written in terms of the individual dissociated ions in the
combined solution. This is known as the complete ionic
equation:
NO
Ag+ (aq) + NO3−(aq) + K+ (aq) + Cl−(aq) → AgCl
−
3 (aq)
(s)
+ K+ (aq) +
A final way to represent a precipitation reaction is
known as the net ionic equation. In this case, any spectator
ions (those that do not contribute to the precipitation
reaction) are left out of the formula completely. Without
the spectator ions, the reaction equation simplifies to the
following:
Ag+(aq) + Cl−(aq) → AgCl (s)
Hints
A precipitation reaction refers
to the formation of
an insoluble salt
when two solutions
containing soluble
salts are combined.
The insoluble salt
that falls out of solution is known as the
precipitate, hence
the reaction’s name.
Observing precipitation reactions can be useful in
the laboratory to determine the presence of various ions in solution. For instance, if
silver nitrate is added to a solution of an unknown salt and a precipitate is observed,
the unknown solution might contain chloride (Cl–).
52
University Chemistry
3.2.1 Net Ionic Equations
The net ionic equation is a chemical equation for a reaction that lists only those
species participating in the reaction. The net ionic equation is commonly used
in acid-base neutralization reactions, double displacement reactions, and redox
reactions. In other words, the net ionic equation applies to reactions that are strong
electrolytes in water.
We can write a molecular equation for the formation of silver chloride precipitate:
The corresponding ionic equation is:
If you look carefully at the ionic equation, you will notice that the sodium ion
and the nitrate ion appear unchanged on both sides of the equation. When the two
solutions are mixed, neither the Na+ nor the NO−3 ions participate in the reaction.
They can be eliminated from the reaction.
A spectator ion is an ion that does not take part in the
chemical reaction and is found in solution both before and
after the reaction. In the above reaction, the sodium ion
and the nitrate ion are both spectator ions. The equation
can now be written without the spectator ions:
The net ionic equation is the chemical equation that
shows only those elements, compounds, and ions that
are directly involved in the chemical reaction. Notice
that in writing the net ionic equation, the positivelycharged silver cation was written first on the reactant
side, followed by the negatively-charged chloride anion.
This is somewhat customary because that is the order
in which the ions must be written in the silver chloride
product. However, it is not absolutely necessary to order
the reactants in this way.
Net ionic equations must be balanced by both mass
and charge. Balancing by mass means ensuring that
there are equal masses of each element on the product
and reactant sides. Balancing by charge means making
sure that the overall charge is the same on both sides of
the equation. In the above equation, the overall charge
is zero, or neutral, on both sides of the equation. As a
Important
The key to
knowing which species dissociate into
ions and which form
solids (precipitates) is
to be able to recognize molecular and
ionic compounds,
know the strong
acids and bases, and
predict the solubility
of compounds. Molecular compounds,
like sucrose or sugar,
don’t dissociate in
water. Ionic compounds, like sodium
chloride, dissociate
according to solubility rules. Strong acids
and bases completely
dissociate into ions,
while weak acids and
bases only partially
dissociate.
Chemical Reactions
53
general rule, if you balance the molecular equation properly, the net ionic equation
will end up being balanced by both mass and charge.
3.2.2 Using Precipitation Reactions in Chemistry
Precipitates are insoluble ionic solid products of a reaction, formed when
certain cations and anions combine in an aqueous solution. The determining factors
of the formation of a precipitate can vary. Some reactions depend on temperature,
such as solutions used for buffers, whereas others are dependent only on solution
concentration. The solids produced in precipitate reactions are crystalline solids,
and can be suspended throughout the liquid or fall to the bottom of the solution.
The remaining fluid is called supernatant liquid. The two components of the mixture
(precipitate and supernate) can be separated by various methods, such as filtration,
centrifuging, or decanting.
The use of solubility rules require an understanding of the way that ions react.
Most precipitation reactions are single replacement reactions or double replacement
reactions. A double replacement reaction occurs when two ionic reactants dissociate
and bond with the respective anion or cation from the other reactant. The ions
replace each other based on their charges as either a cation or an anion. This can
be thought of as “switching partners”; that is, the two reactants each “lose” their
partner and form a bond with a different partner:
A double replacement reaction is specifically classified as a precipitation reaction
when the chemical equation in question occurs in aqueous solution and one of the
of the products formed is insoluble. An example of a precipitation reaction is given
below:
Both reactants are aqueous and one product is solid. Because the reactants are
ionic and aqueous, they dissociate and are therefore soluble. However, there are six
54
University Chemistry
solubility guidelines used to predict which molecules are insoluble in water. These
molecules form a solid precipitate in solution.
3.3 ACID-BASE REACTIONS
Acid–base reaction, a type of chemical process
typified by the exchange of one or more hydrogen ions,
H+, between species that may be neutral (molecules, such
as water, H2O; or acetic acid, CH3CO2H) or electrically
charged (ions, such as ammonium, NH4+; hydroxide,
OH−; or carbonate, CO32−). It also includes analogous
behaviour of molecules and ions that are acidic but do
not donate hydrogen ions (aluminum chloride, AlCl3,
and the silver ion AG+).
Acids are chemical compounds that show, in water
solution, a sharp taste, a corrosive action on metals, and
the ability to turn certain blue vegetable dyes red. Bases
are chemical compounds that, in solution, are soapy to
the touch and turn red vegetable dyes blue. When mixed,
acids and bases neutralize one another and produce salts,
substances with a salty taste and none of the characteristic
properties of either acids or bases.
Hints
An acid–base
reaction is a chemical reaction that
occurs between an
acid and a base.
It can be used to
determine pH.
Several theoretical
frameworks provide
alternative conceptions of the reaction
mechanisms and
their application
in solving related
problems; these are
called the acid–base
theories, for example, Brønsted–Lowry acid–base theory.
The idea that some substances are acids whereas
others are bases is almost as old as chemistry, and the
terms acid, base, and salt occur very early in the writings
of the medieval alchemists. Acids were probably the first
of these to be recognized, apparently because of their sour taste. The English word
acid, the French acide, the German Säure, and the Russian kislota are all derived
from words meaning sour. Other properties associated at an early date with acids
were their solvent, or corrosive, action; their effect on vegetable dyes; and the
effervescence resulting when they were applied to chalk (production of bubbles of
carbon dioxide gas). Bases (or alkalies) were characterized mainly by their ability to
neutralize acids and form salts, the latter being typified rather loosely as crystalline
substances soluble in water and having a saline taste.
In spite of their imprecise nature, these ideas served to correlate a considerable
range of qualitative observations, and many of the commonest chemical materials
that early chemists encountered could be classified as acids (hydrochloric, sulfuric,
nitric, and carbonic acids), bases (soda, potash, lime, ammonia), or salts (common
salt, sal ammoniac, saltpetre, alum, borax). The absence of any apparent physical
basis for the phenomena concerned made it difficult to make quantitative progress
in understanding acid–base behaviour, but the ability of a fixed quantity of acid
to neutralize a fixed quantity of base was one of the earliest examples of chemical
equivalence: the idea that a certain measure of one substance is in some chemical
55
Chemical Reactions
sense equal to a different amount of a second substance. In addition, it was found
quite early that one acid could be displaced from a salt with another acid, and this
made it possible to arrange acids in an approximate order of strength. It also soon
became clear that many of these displacements could take place in either direction
according to experimental conditions. This phenomenon suggested that acid–base
reactions are reversible—that is, that the products of the reaction can interact to
regenerate the starting material. It also introduced the concept of equilibrium to
acid–base chemistry: this concept states that reversible chemical reactions reach a
point of balance, or equilibrium, at which the starting materials and the products
are each regenerated by one of the two reactions as rapidly as they are consumed
by the other.
Apart from their theoretical interest, acids and bases play a large part in
industrial chemistry and in everyday life. Sulfuric acid and sodium hydroxide are
among the products manufactured in largest amounts by the chemical industry,
and a large percentage of chemical processes involve acids or bases as reactants
or as catalysts. Almost every biological chemical process is closely bound up with
acid–base equilibria in the cell, or in the organism as a whole, and the acidity or
alkalinity of the soil and water are of great importance for the plants or animals
living in them. Both the ideas and the terminology of acid–base chemistry have
permeated daily life, and the term salt is especially common.
3.3.1 Arrhenius Acids and Bases
The first person to define acids and bases in detail was the Swedish chemist
Svante Arrhenius. According to the Arrhenius definition, an acid is a substance like
hydrochloric acid that dissolves in water to produce H+ ions (protons; Equation 4 ),
and a base is a substance like sodium hydroxide that dissolves in water to produce
hydroxide (OH−) ions (Equation 5):
(4)
(5)
According to Arrhenius, the characteristic properties of acids and bases are due
exclusively to the presence of H+ and OH− ions, respectively, in solution. Although
Arrhenius’s ideas were widely accepted, his definition of acids and bases had two
major limitations:
•
•
First, because acids and bases were defined in terms of ions obtained
from water, the Arrhenius concept applied only to substances in aqueous
solution.
Second, and more important, the Arrhenius definition predicted that only
substances that dissolve in water to produce H+ and OH− ions should
exhibit the properties of acids and bases, respectively. For example,
56
University Chemistry
according to the Arrhenius definition, the reaction of ammonia (a base)
with gaseous HCl (an acid) to give ammonium chloride (Equation 6) is
not an acid–base reaction because it does not involve H+ and OH−:
(6)
3.3.2 Neutralization
For a strong acid and a strong base in water, the neutralization reaction is
between hydrogen and hydroxide ions—i.e., H3O+ + OH− ⇄ 2H2O. For a weak
acid and a weak base, neutralization is more appropriately considered to involve
direct proton transfer from the acid to the base. For example, the neutralization of
acetic acid by ammonia may be written as CH3CO2H + NH3 → CH3CO2− + NH4+.
This equation does not involve the solvent; it therefore also represents the process
of neutralization in an inert solvent, such as benzene, or in the complete absence
of a solvent. (If one of the reactants is present in large excess, the reaction is more
appropriately described as the dissociation of acetic acid in liquid ammonia or of
ammonia in glacial acetic acid.)
3.3.3 The Brönsted Definition
Because of the limitations of the Arrhenius definition,
a more general definition of acids and bases was needed.
One was proposed independently in 1923 by the Danish
chemist J. N. Brønsted (1879–1947) and the British
chemist T. M. Lowry (1874–1936), who defined acid–base
reactions in terms of the transfer of a proton (H+ ion)
from one substance to another.
Important
Oxidation
may be defined as
loss of electrons
from a substance,
the other definition of oxidation
reactions states
that the addition of
oxygen or the more
electronegative
element or removal
of hydrogen or the
more electropositive
element from a substance is called an
oxidation reaction.
According to Brønsted and Lowry, an acid (A
substance with at least one hydrogen atom that can
dissociate to form an anion and an H+ ion (a proton) in
aqueous solution, thereby forming an acidic solution) is
any substance that can donate a proton, and a base (a
substance that produces one or more hydroxide ions
(OH− and a cation when dissolved in aqueous solution,
thereby forming a basic solution) is any substance that
can accept a proton. The Brønsted–Lowry definition of an
acid is essentially the same as the Arrhenius definition,
except that it is not restricted to aqueous solutions. The
Brønsted–Lowry definition of a base, however, is far more
general because the hydroxide ion is just one of many substances that can accept a
proton. Ammonia, for example, reacts with a proton to form NH+4, so in Equation
3, NH3 is a Brønsted–Lowry base and HCl is a Brønsted–Lowry acid. Because of its
more general nature, the Brønsted–Lowry definition is used throughout this text
unless otherwise specified.
Chemical Reactions
57
3.4 REDOX REACTIONS
Redox reactions are oxidation-reduction chemical reactions in which the
reactants undergo a change in their oxidation states. The term ‘redox’ is a short
form of reduction-oxidation. All the redox reactions can be broken down into two
different processes – a reduction process and an oxidation process.
The oxidation and reduction reactions always occur simultaneously in the
redox reaction or Oxidation-Reduction reaction. The substance getting reduced in a
chemical reaction is known as the oxidizing agent, while a substance that is getting
oxidized is known as the reducing agent.
A redox reaction can be defined as a chemical reaction in which electrons are
transferred between two reactants participating in it. This transfer of electrons can
be identified by observing the changes in the oxidation states of the reacting species.
An illustration detailing the electron transfer between two reactants in a redox
reaction is provided below.
In the illustration provided below, it can be observed that the reactant, an
electron, was removed from reactant A and this reactant is oxidized. Similarly,
reactant B was handed an electron and was therefore reduced.
The loss of electrons and the corresponding increase in the oxidation state of
a given reactant is called oxidation. The gain of electrons and the corresponding
decrease in the oxidation state of a reactant is called reduction.
Electron-accepting species which tend to undergo a reduction in redox reactions
are called oxidizing agents. An electron-donating species which tends to hand over
electrons can be referred to as a reducing agent. These species tend to undergo
oxidation. It can be noted that any redox reaction can be broken down into two
half-reactions, namely the oxidation half-reaction and the reduction half-reaction.
When writing these half-reactions separately, each of them must be balanced in
a way that all the electrons are accounted for.
3.4.1 Electron Transfer
Electron transfer (ET) occurs when an electron relocates from an atom or
molecule to another such chemical entity. ET is a mechanistic description of certain
kinds of redox reactions involving transfer of electrons.
58
University Chemistry
Electrochemical processes are ET reaction. ET reactions are relevant to
photosynthesis and respiration. ET reactions commonly involve transition metal
complexes, In organic chemistry ET is a step in some commercial polymerization
reactions. It is foundational to photoredox catalysis.
This demonstration shows that the deep blue color of the CuSO4 solution (left
picture), which is caused by the Cu2+ ions, becomes light green due to Fe2+ ions
(right picture). As a result, a brown solid made of metallic copper forms and the
steel wool disintegrates as the Fe atoms disappear. The reaction has thus led Fe
to be converted to Fe2+ while the Cu2+ is converted to Cu. This indicates that the
oxidation state of copper has changed from +2 in the Cu2+ ions in solution to 0 in
the atoms comprising metallic copper. This is accomplished as each Fe atom gives
up two electrons, while each Cu2+ gains two electrons, as a result of which two
electrons are transferred from Fe atoms to Cu2+ ions in solution. This is an example
of an electron transfer reaction. The reaction is given by: Cu2+ + Fe → Cu + Fe2+.
Figure 1. An electron transfer reaction.
Mechanism of Electron Transfer Reactions
The process of electron transfer from one species to
another species leads to the oxidation of the donor and
reduction of the acceptor. The mechanism by which the
electron transfer occurs between inorganic complexes
can be classified in to two types: inner sphere electron
transfer mechanism and outer sphere electron transfer
mechanism.
Inner Sphere
Inner sphere electron transfer occurs between
complexes via a bridging ligand. At least one of the
complexes needs to be labile to allow the bridge to form.
Bonds are broken and formed.
Outer Sphere
Outer sphere electron transfer occurs between
two species that do not undergo substitution and do
Hints
In outersphere ET reactions,
the participating redox
centers are not linked
via any bridge during
the ET event. Instead,
the electron “hops”
through space from
the reducing center
to the acceptor. Outer
sphere electron transfer can occur between
different chemical
species or between
identical chemical species that differ only in
their oxidation state.
59
Chemical Reactions
not involve the incursion of significant covalent bond formation. It occurs when
none of the ligands can function as a bridge. It is faster than inner sphere because
the energetic demands are less. No new bonds are broken or formed. Interaction
between the two coordination spheres exist but is not as pronounced as for the
bridge complex in the inner sphere. Outer Adduct is held together by one of the
following: Electrostatic interactions, Vander Waals forces, or Hydrogen bonding.
3.4.2 The Activity Series
The activity series is a type of ordering system for elements, which ranks
how reactive a certain element is in relation to other elements. The activity series
determines the level of reactivity based on how well a certain element can displace
hydrogen gas from acidic solutions and water.
Single-replacement reactions only occur when the element that is doing the
replacing is more reactive than the element that is being replaced. Therefore, it
is useful to have a list of elements in order of their relative reactivity. The activity
series is a list of elements in decreasing order of their reactivity. Since metals replace
other metals, while nonmetals replace other nonmetals, they each have a separate
activity series. The table 1 below is an activity series of most common metals.
Table 1. Activity Series of Metal Elements
Elements, from most to
least reactive
Reaction Occurring
LiLi
KK
BaBa
SrSr
CaCa
NaNa
React with cold water, replacing hydrogen.
MgMg
AlAl
ZnZn
CrCr
FeFe
CdCd
React with steam, but not cold water, replacing
hydrogen.
CoCo
NiNi
SnSn
PbPb
Do not react with water. React with acids, replacing
hydrogen.
H2H2
CuCu
HgHg
AgAg
PtPt
AuAu
Unreactive with water or acids.
60
University Chemistry
For a single-replacement reaction, a given element is capable of replacing
an element that is below it in the activity series. This can be used to predict if a
reaction will occur. Suppose that small pieces of the metal nickel were placed into
two separate aqueous solutions: one of iron (III) nitrate and one of lead (II) nitrate.
Looking at the activity series, we see that nickel is below iron, but above lead.
Therefore, the nickel metal will be capable of replacing the lead in a reaction, but
will not be capable of replacing iron.
In the descriptions that accompany the activity series of metals, a given metal is
also capable of undergoing the reactions. For example, lithium will react with cold
water, replacing hydrogen. It will also react with steam and with acids, since that
requires a lower degree of reactivity.
3.4.3 Balancing Reactions by using Half Reactions
Another method for balancing redox reactions uses half-reactions. Recall that a
half-reaction is either the oxidation or reduction that occurs, treated separately. The
half-reaction method works better than the oxidation-number method when the
substances in the reaction are in aqueous solution. The aqueous solution is typically
either acidic or basic, so hydrogen ions or hydroxide ions are present.
In general, the half-reactions are first balanced by atoms separately. Electrons
are included in the half-reactions. These are then balanced so that the number
of electrons lost is equal to the number of electrons gained. Finally, the two halfreactions are added back together. The example is the oxidation of Fe2+ ions to Fe3+
ions by dichromate (Cr2O2−7) in acidic solution. The dichromate ions are reduced to
Cr3− ions.
Chemical Reactions
61
Step 1: Write the unbalanced ionic equation.
Notice that the equation is far from balanced, as there are no oxygen atoms on
the right side. This will be resolved by the balancing method.
Step 2: Write separate half-reactions for the oxidation and the reduction
processes. Determine the oxidation numbers first, if necessary.
Step 3: Balance the atoms in the half-reactions other than hydrogen and oxygen.
In the oxidation half-reaction above, the iron atoms are already balanced. The
reduction half-reaction needs to be balanced with the chromium atoms.
Step 4: Balance oxygen atoms by adding water molecules to the appropriate
side of the equation. For the reduction half-reaction above, seven H2O molecules
will be added to the product side.
Now the hydrogen atoms need to be balanced. In an acidic medium, add
hydrogen ions to balance. In this example, fourteen H+ ions will be added to the
reactant side.
Step 5: Balance the charges by adding electrons to each half-reaction. For the
oxidation half-reaction, the electrons will need to be added to the product side.
For the reduction half-reaction, the electrons will be added to the reactant side. By
adding one electron to the product side of the oxidation half-reaction, there is a 2+
total charge on both sides.
There is a total charge of 12+ on the reactant side of the reduction half-reaction
(14−2). The product side has a total charge of 6+ due to the two chromium ions
(2×3). To balance the charge, six electrons need to be added to the reactant side.
Now equalize the electrons by multiplying everything in one or both equations
by a coefficient. In this example, the oxidation half-reaction will be multiplied by
six.
62
University Chemistry
Step 6: Add the two half-reactions together. The electrons must cancel. Balance
any remaining substances by inspection. If necessary, cancel out H2O or H+ that
appear on both sides.
EXERCISE
Answer the following questions:
1. Explain the chemical reactions.
2. Write the balancing equations.
3. What is the net ionic equations?
4. Describe the using precipitation reactions in chemistry.
5. Identify the arrhenius acids and bases.
6. Define the neutralization.
7. Evaluate the electron transfer.
8. Define the activity series.
9. Focus on balancing reactions by using half-reactions.
MULTIPLE CHOICE QUESTIONS
Tick the correct answer.
1.
Magnesium ribbon is rubbed before burning because it has a coating of
a. basic magnesium carbonate
b. basic magnesium oxide
c. basic magnesium sulphide
d. basic magnesium chloride
2.
Which information is not conveyed by a balanced chemical equation?
a. Physical states of reactants and products
b. Symbols and formulae of all the substances involved in a particular
reaction
c. Number of atoms/molecules of the reactants and products formed
d. Whether a particular reaction is actually feasible or not
3.
Chemically rust is
a. hydrated ferrous oxide
b. only ferric oxide
Chemical Reactions
4.
5.
6.
7.
8.
9.
10.
c. hydrated ferric oxide
d. none of these
Reaction of ‘magnesium’ with air is
a. Exothermic reaction
b. Endothermic reaction
c. Reversible reaction
d. Substitution reaction
What chemicals are used in fireworks?
a. Copper chloride
b. Calcium chloride
c. Barium chloride
d. All of above
When a magnesium ribbon is burnt in air, the ash formed is
a. Black
b. White
c. Yellow
d. Pink
Color of magnesium oxide is
a. White
b. Blue
c. Grey
d. Pink
If magnesium is gently heated, it forms
a. Magnesium oxide
b. Magnesium sulfide
c. Magnesium nitrite
d. Magnesium carbonate
When carbon dioxide is passed through lime water,
a. Calcium hydroxide is formed
b . White precipitate of CaO is formed
c. Lime water turns milky
d. Color of lime water disappears.
When crystals of lead nitrate are heated strongly in a dry test tube
a. Crystals immediately melt
b. A brown residue is left
c. White fumes appear in the tube
d. A yellow residue is left
63
64
University Chemistry
ANSWER
1. (a)
2. (d)
3. (c)
4. (a)
5. (d)
6. (b)
7. (a)
8. (a)
9. (c)
10. (b)
REFERENCES
1.
2.
Elric, H., 2016. CHEMICAL REACTIONS. [online] Ric.edu.
Elschenbroich, Christoph (2008). Organometallchemie (6th ed.). Wiesbaden:
Vieweg+Teubner Verlag. p. 263.
3. Emig, Gerhard; Klemm, Elias (2005). Technical Chemistry (in German) (5th
ed.).
4. Fox, Marye Anne; Whitesell, James K. (2004). Organic chemistry (Third ed.).
Jones & Bartlett. p. 699.
5. Friedman, Leonard J.; Friedman, Samantha J. (2008). The History of the
Contact Sulfuric Acid Process (PDF). Boca Raton, Florida: Acid Engineering &
Consulting, Inc.
6. Latscha, Hans Peter; Kazmaier, Uli; Klein, Helmut Alfons (2008). Organische
Chemie: Chemie-basiswissen II (in German). Vol. 2 (6th ed.). Springer. p. 273.
7. Lechner, Manfred; Gehrke, Klaus; Nordmeier, Eckhard (2003). Macromolecular
Chemistry (3rd ed.). Basel: Birkhäuser. pp. 53–65.
8. Saunders, David Stanley (2002). Insect clocks (Third ed.). Amsterdam: Elsevier.
p. 179.
9. Stranges, Anthony N. (2000). “Germany’s synthetic fuel industry, 1935–1940”.
In Lesch, John E. (ed.). The German Chemical Industry in the Twentieth
Century. Kluwer Academic Publishers. p. 170.
10. Wingender, Jörg; Ortanderl, Stefanie (July 2009). “Ausfällung”. Römpp
Chemie-Lexikon.
65
Reactions Stoichiometry
CHAPTER 4
REACTIONS STOICHIOMETRY
OBJECTIVES
After studying this chapter, you will be able to:
1. Interpret stoichiometric coefficients
2. Understand stoichiometry of reactions in solution
INTRODUCTION
A balanced chemical equation provides a great deal of information in a very
succinct format. Chemical formulas provide the identities of the reactants and
products involved in the chemical change, allowing classification of the reaction.
Coefficients provide the relative numbers of these chemical species, allowing a
quantitative assessment of the relationships between the amounts of substances
consumed and produced by the reaction. These quantitative relationships are
known as the reaction’s stoichiometry, a term derived from the Greek words
stoicheion (meaning “element”) and metron (meaning “measure”).
The general approach to using stoichiometric relationships is similar in
concept to the way people go about many common activities. Food preparation,
66
University Chemistry
for example, offers an appropriate comparison. A recipe for making eight pancakes
calls for 1 cup pancake mix, cup milk, and one egg. The “equation” representing
the preparation of pancakes per this recipe is
If two dozen pancakes are needed for a big family breakfast, the ingredient
amounts must be increased proportionally according to the amounts given in the
recipe. For example, the number of eggs required to make 24 pancakes is
Balanced chemical equations are used in much the same fashion to determine
the amount of one reactant required to react with a given amount of another reactant,
or to yield a given amount of product, and so forth. The coefficients in the balanced
equation are used to derive stoichiometric factors that permit computation of the
desired quantity. To illustrate this idea, consider the production of ammonia by
reaction of hydrogen and nitrogen:
This equation shows ammonia molecules are produced from hydrogen
molecules in a 2:3 ratio, and stoichiometric factors may be derived using any
amount (number) unit:
These stoichiometric factors can be used to compute the number of ammonia
molecules produced from a given number of hydrogen molecules, or the number
of hydrogen molecules required to produce a given number of ammonia molecules.
Similar factors may be derived for any pair of substances in any chemical equation.
A stoichiometric chemical reaction is one where the quantities of the reactants
and products are such that all of the reactants are consumed and none remain after
completion of the chemical reaction. Stoichiometry is useful for measuring chemical
reactions such as those that occur in corrosion processes.
Stoichiometry is the technique used to calculate the required quantities of
chemical reactants and products using a balanced chemical equation. The reactants
and products must remain in the same proportion when increasing or decreasing the
quantity of the reactants or products for the chemical equation to remain balanced.
Stoichiometry is based on the law of conservation of mass, which states that
the total mass of the reactants are equal to the total mass of the products for the
chemical reaction equation to be balanced.
Reactions Stoichiometry
67
Stoichiometry can be of two types:
•
•
Composition stoichiometry
Gas stoichiometry
4.1 INTERPRETING STOICHIOMETRIC COEFFICIENTS
Stoichiometric coefficient or stoichiometric number is the number of molecules
that participate in the reaction. If you look at any balanced reaction you will
notice that there are an equal number of elements on both sides of the equation.
The stoichiometric coefficient is basically the number present in front of atoms,
molecules or ions.
Stoichiometric coefficients can be fractions as well as whole numbers. In essence,
the coefficients help us to establish the mole ratio between reactants and products.
In lay terms, the stoichiometric coefficient of any given component is the number
of molecules and/or formula units that participate in the reaction as written. A
related concept is the stoichiometric number (using IUPAC nomenclature), wherein
the stoichiometric coefficient is multiplied by +1 for all products and by -1 for all
reactants.
For example, in the reaction CH4 + 2 O2 → CO2 + 2 H2O, the stoichiometric
number of CH4 is −1, the stoichiometric number of O2 is −2, for CO2 it would be +1
and for H2O it is +2.
In more technically precise terms, the stoichiometric number in a chemical
reaction system of the ith component is defined as
∆N
νi = i
∆ξ
or
∆Ni = ν i ∆ξ
where Ni is the number of molecules of i, and ξ is the progress variable or
extent of reaction.
The stoichiometric number νi represents the degree to which a chemical
species participates in a reaction. The convention is to assign negative numbers
to reactants (which are consumed) and positive ones to products, consistent with
the convention that increasing the extent of reaction will correspond to shifting
the composition from reactants towards products. However, any reaction may be
viewed as going in the reverse direction, and in that point of view, would change in
the negative direction in order to lower the system’s Gibbs free energy. Whether a
reaction actually will go in the arbitrarily selected forward direction or not depends
on the amounts of the substances present at any given time, which determines the
68
University Chemistry
kinetics and thermodynamics, i.e., whether equilibrium lies to the right or the left
of the initial state,
In reaction mechanisms, stoichiometric coefficients for each step are always
integers, since elementary reactions always involve whole molecules. If one uses
a composite representation of an overall reaction, some may be rational fractions.
There are often chemical species present that do not participate in a reaction;
their stoichiometric coefficients are therefore zero. Any chemical species that is
regenerated, such as a catalyst, also has a stoichiometric coefficient of zero.
The simplest possible case is an isomerization
A → B
in which νB = 1 since one molecule of B is produced each time the reaction occurs,
while νA = −1 since one molecule of A is necessarily consumed. In any chemical
reaction, not only is the total mass conserved but also the numbers of atoms of
each kind are conserved, and this imposes corresponding constraints on possible
values for the stoichiometric coefficients.
There are usually multiple reactions proceeding simultaneously in any natural
reaction system, including those in biology. Since any chemical component can
participate in several reactions simultaneously, the stoichiometric number of the ith
component in the kth reaction is defined as
∂N
ν ik = i
∂ξk
so that the total (differential) change in the amount of the ith component is
dNi =
∑ν
k
ik dξ k .
Extents of reaction provide the clearest and most explicit way of representing
compositional change, although they are not yet widely used.
With complex reaction systems, it is often useful to consider both the
representation of a reaction system in terms of the amounts of the chemicals
present { Ni } (state variables), and the representation in terms of the actual
compositional degrees of freedom, as expressed by the extents of reaction { ξk }.
The transformation from a vector expressing the extents to a vector expressing
the amounts uses a rectangular matrix whose elements are the stoichiometric
numbers [ νi k ].
The maximum and minimum for any ξk occur whenever the first of the
reactants is depleted for the forward reaction; or the first of the «products» is
depleted if the reaction as viewed as being pushed in the reverse direction.
This is a purely kinematic restriction on the reaction simplex, a hyperplane in
composition space, or Nspace, whose dimensionality equals the number of linearlyindependent chemical reactions. This is necessarily less than the number of chemical
69
Reactions Stoichiometry
components, since each reaction manifests a relation
between at least two chemicals. The accessible region of
the hyperplane depends on the amounts of each chemical
species actually present, a contingent fact. Different such
amounts can even generate different hyperplanes, all
sharing the same algebraic stoichiometry.
Important
The stoichiometric coefficient
of any species that
does not participate
in a given chemical
reaction is zero.
In accord with the principles of chemical
kinetics and thermodynamic equilibrium, every chemical
reaction is reversible, at least to some degree, so that
each equilibrium point must be an interior point of the
simplex. As a consequence, extrema for the ξs will not occur unless an experimental
system is prepared with zero initial amounts of some products.
The number of physically-independent reactions can be even greater than the
number of chemical components, and depends on the various reaction mechanisms.
For example, there may be two (or more) reaction paths for the isomerism above.
The reaction may occur by itself, but faster and with different intermediates, in the
presence of a catalyst.
The (dimensionless) “units” may be taken to be molecules or moles. Moles are
most commonly used, but it is more suggestive to picture incremental chemical
reactions in terms of molecules. The Ns and ξs are reduced to molar units by
dividing by Avogadro’s number. While dimensional mass units may be used, the
comments about integers are then no longer applicable.
4.1.1 Mole Calculations
Atoms and molecules are extremely small in size and their numbers in a very
small amount of a substance are very large. Therefore, to represent atoms and
molecules in bulk, a mole concept was introduced. One mole of any substance
contains 6.022 x 1023 numbers of that substance. This number is also known as
Avogadro’s number.
The mass of one mole of a substance in grams is called molar mass. The molar
mass of one mole of a substance is numerically equal to the atomic/molecular
formula mass.
Let us take one example of a balanced chemical equation.
3Fe(s) + 4H2O(l) ⇾ Fe3O4 (s)+ 4H2 (g)
The quantitative information drawn from this balanced chemical equation is
•
•
3 mole of Fe reacts with 4 moles of H2O to yield one mole of Fe3O4 and 4
moles of H2.
168g ( 56×3) of Fe reacts with 72g( 18×4) of H20 to yield 231g of Fe3O4 and
8g of H2 gas.
70
University Chemistry
If the reactants and products are in gaseous form, then the molar volume is
taken into consideration. One mole of any gas occupies 22.4 Liters.
CH4(g) + 2O2(g)⇾ CO2(g)+ 2H20 (g)
In the above reaction, 22.4 Liters of CH4 reacts with 44.8 (2 x 22.4) liters of 02 to
yield 22.4 Liters of CO2 and 44.8 liters of H2O
In a chemical reaction, it is possible that one of the reactants is present in excess
amount. Some of these excess reactants will, therefore, be left over when the reaction
is complete; the reaction stops immediately as soon as one of the reactants is totally
consumed.
The substance that is totally consumed in a reaction is called the limiting
reagent.
Let us take one example of a chemical reaction to understand limiting reagent
concept.
N2 + 3H2 ➝ 2NH3
Suppose we have one mole of N2 reacting with one mole of H2. But from the
balanced chemical equation, one mole of N2 requires three moles of H2. So, the
limiting reagent in this reaction is H2.
4.1.2 Limiting Reactants
The limiting reagent (or limiting reactant or limiting agent) in a chemical
reaction is a reactant that is totally consumed when the chemical reaction is
completed. The amount of product formed is limited by this reagent, since the
reaction cannot continue without it. If one or more other reagents are present in
excess of the quantities required to react with the limiting reagent, they are described
as excess reagents or excess reactants (sometimes abbreviated as «xs”).
The limiting reagent must be identified in order to calculate the percentage
yield of a reaction since the theoretical yield is defined as the amount of product
obtained when the limiting reagent reacts completely. Given the balanced chemical
equation, which describes the reaction, there are several equivalent ways to identify
the limiting reagent and evaluate the excess quantities of other reagents.
Method 1: Comparison of reactant amounts
This method is most useful when there are only two reactants. One reactant (A)
is chosen, and the balanced chemical equation is used to determine the amount of
the other reactant (B) necessary to react with A. If the amount of B actually present
exceeds the amount required, then B is in excess and A is the limiting reagent. If the
amount of B present is less than required, then B is the limiting reagent.
71
Reactions Stoichiometry
Example for two reactants
Consider the combustion of benzene, represented by the following chemical
equation:
2C6 H6 (l) + 15O 2 (g) → 12CO 2 (g) + 6H 2 O(l)
This means that 15 moles of molecular oxygen (O2) is required to react with 2
moles of benzene (C6H6)
The amount of oxygen required for other quantities of benzene can be calculated
using cross-multiplication (the rule of three). For example, if 1.5 mol C6H6 is present,
11.25 mol O2 is required:
1.5 molC6 H6 ×
15 molO 2
2 molC6 H6
=
11.25 molO 2
If in fact 18 mol O2 are present, there will be an excess of (18 - 11.25) = 6.75 mol of
unreacted oxygen when all the benzene is consumed. Benzene is then the limiting
reagent.
This conclusion can be verified by comparing the mole ratio of O2 and
C6H6 required by the balanced equation with the mole ratio actually present:
molO
15 molO
molC6 H6
2 molC6 H6
•
2
2
Required:=
7.5 molO 2
=
•
2
2
12 molO 2
=
=
actual:
molC H
1.5 molC H
molO
6
18 molO
6
6
6
Since the actual ratio is larger than required, O2 is the reagent in excess, which
confirms that benzene is the limiting reagent.
Method 2: Comparison of product amounts which can be formed from
each reactant
In this method the chemical equation is used to calculate the amount of one
product which can be formed from each reactant in the amount present. The limiting
reactant is the one which can form the smallest amount of the product considered.
This method can be extended to any number of reactants more easily than the first
method.
Example
20.0 g of iron (III) oxide (Fe2O3) are reacted with 8.00 g aluminium (Al) in the
following thermite reaction:
Fe 2 O 3 (s) + 2Al(s)− > 2Fe(l) + Al 2 O 3 (s)
72
University Chemistry
Since the reactant amounts are given in grams, they must be first converted into
moles for comparison with the chemical equation, in order to determine how many
moles of Fe can be produced from either reactant.
Moles of Fe which can be produced from reactant Fe2O3
mol Fe 2 O 3 =
grams Fe 2 O 3
g / mol Fe 2 O 3
20.0 g
= 0.125 mol
159.7 g / mol
2 mol Fe
= 0.125 mol Fe 2 O 3 ×
= 0.250
mol Fe
1 mol Fe 2 O 3
=
•
Moles of Fe which can be produced from reactant A
grams Al
g / mol Al
8.00 g
= = 0.297 mol
26.98 g / mol
mol Al =
mol Fe
= 0.297 mol Al ×
2 mol Fe
= 0.297 mol Fe
2 mol Al
There is enough Al to produce 0.297 mol Fe, but only enough Fe2O3 to produce
0.250 mol Fe. This means that the amount of Fe actually produced is limited by the
Fe2O3 present, which is therefore the limiting reagent.
It can be seen from the example above that the amount of product (Fe) formed
from each reagent X (Fe2O3 or Al) is proportional to the quantity
Moles of Reagent X
Stoichiometric Coefficient of Reagent X
This suggests a shortcut which works for any number of reagents. Just calculate
this formula for each reagent, and the reagent that has the lowest value of this
formula is the limiting reagent.
4.1.3 Chemical Compositions from Measurements of Mass
Chemical composition refers to the arrangement, type, and ratio of atoms in
molecules of chemical substances. Chemical composition varies when chemicals
are added or subtracted from a substance, when the ratio of substances changes, or
when other chemical changes occur in chemicals.
The chemical composition of a pure substance corresponds to the relative
amounts of the elements that constitute the substance itself. It can be expressed
with a chemical formula, such as an empirical or molecular formula.
73
Reactions Stoichiometry
The molecular mass of a substance is the sum of
the average masses of the atoms in one molecule of a
substance. Calculations for formula mass and molecular
mass are described. Calculations involving conversions
between moles of a material and the mass of that material
are described. Calculations are illustrated for conversions
between mass and number of particles.
Chemists often need to know what elements are
present in a compound and in what percentage. The
percent composition is the percent by mass of each
element in a compound.
The percent composition of a compound can also
be determined from the formula of the compound. The
subscripts in the formula are first used to calculate the
mass of each element in one mole of the compound.
That is divided by the molar mass of the compound and
multiplied by 100%.
Note
The principles of
stoichiometry are
based upon the law
of conservation of
mass. Matter can
neither be created
nor destroyed, so
the mass of every
element present in
the product(s) of a
chemical reaction
must be equal to
the mass of each
and every element present in the
reactant(s).
4.2 THE STOICHIOMETRY OF
REACTIONS IN SOLUTION
Stoichiometry deals with the relative quantities of reactants and products in
chemical reactions. It can be used to find the quantities of the products from given
reactants in a balanced chemical reaction, as well as percent yield.
To calculate the quantity of a product, calculate the number of moles for each
reactant. Moles of a product are equal to the moles of a limiting reactant in one-toone reaction stoichiometry. To find product mass, moles must be multiplied by the
product’s molecular weight.
In stoichiometric calculations involving solutions, a given solution’s
concentration is often used as a conversion factor.
Concentration of Solutions
Recall that a solution consists of two components: solute (the dissolved material)
and solvent (the liquid in which the solute is dissolved). The amount of solute in a
given amount of solution or solvent is known as the concentration. The two most
common ways of expressing concentration are molarity and molality.
Molarity
The molar concentration (M) of a solution is defined as the number of moles
of solute (n) per liter of solution (i.e, the volume, Vsolution):
M=
N
Vsolution
74
University Chemistry
The units of molarity are mol/L, often abbreviated
as M.
For example, the number of moles of NaCl in 0.123L
of a 1.00M solution of NaCl can be calculated as follows:
0.123 L of solution ×
1.00 mole
=
0.123 moles NaCl
1.00 L of solution
Molality
The molal concentration (m) of a solution is defined
as the number of moles of solute (n) per kilogram
of solvent (i.e., the mass of the solvent, msolvent):
m=
Important
Stoichiometry is the study
and calculation of
quantitative (measurable) relationships
of the reactants and
products in chemical reactions (chemical equations)
n
m solvent
The units of molality are mol/kg, or m.
For example, the number of moles of NaCl dissolved in 0.123kg of H2O (the
solvent), in order to make a 1.00m solution of NaCl, can be calculated as follows:
0.123 kg of solvent ×
1.00 mole
t=
0.123 moles NaCl
1.00 kg of solven
Reaction Stoichiometry in Solutions
We can perform stoichiometric calculations for aqueous phase reactions just as
we can for reactions in solid, liquid, or gas phases. Almost always, we will use the
concentrations of the solutions as conversion factors in our calculations.
Example
•
123 mL of a 1.00 M solution of NaCl is mixed with 72.5 mL of a 2.71 M
solution of AgNO3. What is the mass of AgCl(s) formed in the precipitation
reaction?
First, we need to write out our balanced reaction equation:
AgNO 3 (aq ) + NaCl(aq ) → AgCl(s) + NaNO 3 (aq )
The next step, as in any calculation involving stoichiometry, is to determine our
limiting reactant. We can do this by converting both of our reactants into moles:
123 mL NaCl ×
1L
1.00 mol NaCl
×
=
0.123 mol NaCl
1000 mL
1L
72.5 mL AgNO 3 ×
2.71 mol AgNO 3
1L
×
=
0.196 mol AgNO 3
1000 mL
1L
75
Reactions Stoichiometry
We can see from our reaction equation that AgNO3 and NaCl react in a 1:1 ratio.
Because there are fewer moles of NaCl present in solution, NaCl is our limiting
reactant. We can now solve for the mass of AgCl formed:
123 mL NaCl ×
143 g
1L
1.00 mol NaCl 1 mol AgCl
×
×
×
=
17.6 g AgCl
1000 mL
1L
1 mol NaCl 1 mol AgCl
Therefore, 17.6 g AgCl(s) is formed in the reaction.
To sum up: we converted to each reactant’s moles by using the given
concentrations as conversion factors, expressing molarity as mol/L; once we found
our limiting reactant, we converted through to grams of AgCl formed.
4.2.1 Molar Concentration
Molar concentration is a measure of the concentration of a chemical species,
in particular of a solute in a solution, in terms of amount of substance per unit
volume of solution. In chemistry, the most commonly
used unit for molarity is the number of moles per liter,
having the unit symbol mol/L or mol⋅dm−3 in SI unit. A
Note
solution with a concentration of 1 mol/L is said to be 1
molar, commonly designated as 1 M.
Reaction stoichiom-
Definition
Molar concentration or molarity is most commonly
expressed in units of moles of solute per liter of solution.
For use in broader applications, it is defined as amount
of substance of solute per unit volume of solution, or
per unit volume available to the species, represented by
lowercase c:
=
c
etry describes the
quantitative relationship between
reactants and products within a given
chemical reaction.
n
N
C
=
=
.
V NA V NA
Here, n is the amount of the solute in moles, N is the number of constituent
particles present in volume V (in liters) of the solution, and NA is the Avogadro
N
constant, since 2019 defined as exactly 6.02214076×1023 mol−1. The ratio
is the
V
number density C.
In thermodynamics the use of molar concentration is often not convenient
because the volume of most solutions slightly depends on temperature due to
thermal expansion. This problem is usually resolved by introducing temperature
correction factors, or by using a temperature-independent measure of concentration
such as molality.
76
University Chemistry
The reciprocal quantity represents the dilution (volume) which can appear in
Ostwald’s law of dilution.
Formality or analytical concentration
If a molecular entity dissociates in solution, the concentration refers to the
original chemical formula in solution, the molar concentration is sometimes
called formal concentration or formality (FA) or analytical concentration (cA). For
example, if a sodium carbonate solution (Na2CO3) has a formal concentration
of c(Na2CO3) = 1 mol/L, the molar concentrations are c(Na+) = 2 mol/L and c(CO2−
) = 1 mol/L because the salt dissociates into these ions.
3
Units
In the International System of Units (SI) the coherent unit for molar concentration
is mol/m3. However, this is inconvenient for most laboratory purposes and most
chemical literature traditionally uses mol/dm3, which is the same as mol/L. This
traditional unit is often denoted by the letter M, optionally preceded by an SI
prefix as needed to denote sub-multiples, for example:
mol/m3 = 10−3 mol/dm3 = 10−3 mol/L = 10−3 M = 1 mmol/L = 1 mM.
The units millimolar and micromolar refer to mM and μM (10−3 mol/L and
10 mol/L), respectively.
−6
Name
Abbreviation
Concentration
(mol/L)
(mol/m3)
millimolar
mM
10−3
100
micromolar
μM
10−6
10−3
nanomolar
nM
10−9
10−6
picomolar
pM
10−12
10−9
femtomolar
fM
10
−15
10−12
attomolar
aM
10−18
10−15
zeptomolar
zM
10−21
10−18
yoctomolar
yM
10−24
(6 particles per 10 L)
10−21
Properties
Sum of molar concentrations – normalizing relations
The sum of molar concentrations gives the total molar concentration, namely
the density of the mixture divided by the molar mass of the mixture or by another
name the reciprocal of the molar volume of the mixture. In an ionic solution, ionic
strength is proportional to the sum of the molar concentration of salts.
Sum of products of molar concentrations and partial molar volumes
The sum of products between these quantities equals one:
77
Reactions Stoichiometry
∑ c V = 1.
i
i
i
Note
Dependence on volume
The molar concentration depends on the variation
of the volume of the solution due mainly to thermal
expansion. On small intervals of temperature, the
dependence is
ci =
c i,T
0
1 + α∆T
Before performing
any stoichiometric
calculation, we must
first have a balanced
chemical equation.
,
where ci,T0 is the molar concentration at a reference temperature, α is the
thermal expansion coefficient of the mixture.
4.2.2 The Volume of Solution required for Reaction
The mole is the unit for amount of substance. The number of particles in a
substance can be found using the Avogadro constant. The mass of product depends
upon the mass of limiting reactant.
Calculating amounts from concentration and volume
The amount in moles of a solute in a given volume of solution can be calculated
if the concentration of the solution is known. The mass of solute can then be
calculated.
Rearranging the equation for concentration below:
Concentration in mol / dm 3 =
amount of the solute in mol
volume in dm 3
Amount of solute in mol = concentration in mol/dm3 × volume in dm3
Example
Calculate the amount of sodium hydroxide, NaOH, in 25.0 cm3 of solution of
concentration 0.1 mol/dm3.
Converting the volume from cm3 to dm3, 25.0 cm3 = 25.0 ÷ 1000 = 0.025 dm3.
Amount of NaOH in mol = concentration in mol/dm3 × volume in dm3
= 0.1 mol/dm3 × 0.025 dm3
= 0.0025 mol
Calculating masses from concentration and volume
If the amount in mol of a solute in a given volume of solution is known, its mass
can also be calculated.
78
University Chemistry
4.2.3 Titrations
Titration is a common laboratory method of quantitative chemical analysis to
determine the concentration of an identified analyte (a substance to be analyzed). A
reagent, termed the titrant or titrator, is prepared as a standard solution of known
concentration and volume. The titrant reacts with a solution of analyte to determine
the analyte’s concentration. The volume of titrant that reacted with the analyte is
termed the titration volume.
Procedure
A typical titration begins with a beaker or Erlenmeyer flask containing a
very precise amount of the analyte and a small amount of indicator (such as
phenolphthalein) placed underneath a calibrated burette or chemistry pipetting
syringe containing the titrant. Small volumes of the titrant are then added to the
analyte and indicator until the indicator changes color in reaction to the titrant
saturation threshold, representing arrival at the endpoint of the titration, meaning
the amount of titrant balances the amount of analyte present, according to the
reaction between the two. Depending on the endpoint desired, single drops or less
than a single drop of the titrant can make the difference between a permanent and
temporary change in the indicator.
Preparation techniques
Typical titrations require titrant and analyte to be in a liquid (solution) form.
Though solids are usually dissolved into an aqueous solution, other solvents such
as glacial acetic acid or ethanol are used for special purposes (as in petrochemistry)
which specializes in petroleum. Concentrated analytes are often diluted to improve
accuracy.
Many non-acid–base titrations require a constant pH during the reaction.
Therefore, a buffer solution may be added to the titration chamber to maintain the
pH.
In instances where two reactants in a sample may react with the titrant and
only one is the desired analyte, a separate masking solution may be added to the
reaction chamber which eliminates the effect of the unwanted ion.
Some reduction-oxidation (redox) reactions may require heating the sample
solution and titrating while the solution is still hot to increase the reaction rate. For
instance, the oxidation of some oxalate solutions requires heating to 60 °C (140 °F)
to maintain a reasonable rate of reaction.
Titration Curves
A titration curve is a curve in graph the x-coordinate of which represents the
volume of titrant added since the beginning of the titration, and the y-coordinate
of which represents the concentration of the analyte at the corresponding stage of
Reactions Stoichiometry
79
the titration (in an acid–base titration, the y-coordinate usually represents the pH
of the solution).
In an acid–base titration, the titration curve represents the strength of
the corresponding acid and base. For a strong acid and a strong base, the
curve will be relatively smooth and very steep near the equivalence point.
Because of this, a small change in titrant volume near the equivalence point
results in a large pH change and many indicators would be appropriate (for
instance litmus, phenolphthalein or bromothymol blue).
If one reagent is a weak acid or base and the other is a strong acid or base, the
titration curve is irregular and the pH shifts less with small additions of titrant near
the equivalence point. For example, the titration curve for the titration between oxalic
acid (a weak acid) and sodium hydroxide (a strong base) is pictured. The equivalence
point occurs between pH 8-10, indicating the solution is basic at the equivalence
point and an indicator such as phenolphthalein would be appropriate. Titration
curves corresponding to weak bases and strong acids are similarly behaved, with
the solution being acidic at the equivalence point and indicators such as methyl
orange and bromothymol blue being most appropriate.
Titrations between a weak acid and a weak base have titration curves which are
very irregular. Because of this, no definite indicator may be appropriate and a pH
meter is often used to monitor the reaction.
The type of function that can be used to describe the curve is termed a sigmoid
function.
Types of Titrations
There are many types of titrations with different procedures and goals. The most
common types of qualitative titration are acid–base titrations and redox titrations.
Acid–base titration
Acid–base titrations depend on the neutralization between an acid and a base
when mixed in solution. In addition to the sample, an appropriate pH indicator
is added to the titration chamber, representing the pH range of the equivalence
point. The acid–base indicator indicates the endpoint of the titration by changing
color. The endpoint and the equivalence point are not exactly the same because
the equivalence point is determined by the stoichiometry of the reaction while the
endpoint is just the color change from the indicator. Thus, a careful selection of the
indicator will reduce the indicator error. For example, if the equivalence point is at
a pH of 8.4, then the phenolphthalein indicator would be used instead of Alizarin
Yellow because phenolphthalein would reduce the indicator error. When more
precise results are required, or when the reagents are a weak acid and a weak base,
a pH meter or a conductance meter are used.
80
University Chemistry
For very strong bases, such as organolithium reagent, metal amides, and
hydrides, water is generally not a suitable solvent and indicators whose pKa are in
the range of aqueous pH changes are of little use. Instead, the titrant and indicator
used are much weaker acids, and anhydrous solvents such as THF are used.
The approximate pH during titration can be approximated by three kinds of
calculations. Before beginning of titration, the concentration of [H+ ] is calculated in
aqueous solution of weak acid before adding any base. When the number of moles
of bases added equals the number of moles of initial acid or so called equivalence
point, one of hydrolysis and the pH is calculated in the same way that the conjugate
bases of the acid titrated was calculated. Between starting and end points, [H+ ]
is obtained from the Henderson-Hasselbalch equation and titration mixture is
considered as buffer. In Henderson-Hasselbalch equation the [acid] and [base] are
said to be the molarities that would have been present even with dissociation or
hydrolysis. In a buffer, [H+ ] can be calculated exactly but the dissociation of HA,
the hydrolysis of A − and self-ionization of water must be taken into account. Four
independent equations must be used:
[H + ][OH − ] = 10 −14
[H + ] = K a
[HA]
[A − ]
(n A + n B )
[HA] + [A − ] =
V
[H + ] +
nB
V
= [A − ] + [OH − ]
In the equations, n A and n B are the moles of acid (HA) and salt (XA where X is
the cation), respectively, used in the buffer, and the volume of solution is V. The
law of mass action is applied to the ionization of water and the dissociation of acid
to derived the first and second equations. The mass balance is used in the third
equation, where the sum of V[HA] and V[A − ] must equal to the number of moles of
dissolved acid and base, respectively. Charge balance is used in the fourth equation,
where the left hand side represents the total charge of the cations and the right
hand side represents the total charge of the anions: is the molarity of the cation
(e.g. sodium, if sodium salt of the acid or sodium hydroxide is used in making the
buffer).
Redox titration
Redox titrations are based on a reduction-oxidation reaction between an
oxidizing agent and a reducing agent. A potentiometer or a redox indicator is usually
used to determine the endpoint of the titration, as when one of the constituents is
Reactions Stoichiometry
81
the oxidizing agent potassium dichromate. The color change of the solution from
orange to green is not definite, therefore an indicator such as sodium diphenylamine
is used. Analysis of wines for sulfur dioxide requires iodine as an oxidizing agent.
In this case, starch is used as an indicator; a blue starch-iodine complex is formed in
the presence of excess iodine, signalling the endpoint.
Some redox titrations do not require an indicator, due to the intense color of
the constituents. For instance, in permanganometry a slight persisting pink color
signals the endpoint of the titration because of the color of the excess oxidizing
agent potassium permanganate. In iodometry, at sufficiently large concentrations,
the disappearance of the deep red-brown triiodide ion can itself be used as an
endpoint, though at lower concentrations sensitivity is improved by adding starch
indicator, which forms an intensely blue complex with triiodide.
Gas phase titration
Gas phase titrations are titrations done in the gas phase, specifically as methods
for determining reactive species by reaction with an excess of some other gas, acting
as the titrant. In one common gas phase titration, gaseous ozone is titrated with
nitrogen oxide according to the reaction
O3 + NO → O2 + NO2
After the reaction is complete, the remaining titrant and product are quantified
(e.g., by Fourier transform spectroscopy) (FT-IR); this is used to determine the
amount of analyte in the original sample.
Gas phase titration has several advantages over simple spectrophotometry.
First, the measurement does not depend on path length, because the same path
length is used for the measurement of both the excess titrant and the product.
Second, the measurement does not depend on a linear change in absorbance as a
function of analyte concentration as defined by the Beer–Lambert law. Third, it is
useful for samples containing species which interfere at wavelengths typically used
for the analyte.
Complexometric titration
Complexometric titrations rely on the formation of a complex between the
analyte and the titrant. In general, they require specialized complexometric
indicators that form weak complexes with the analyte. The most common example
is the use of starch indicator to increase the sensitivity of iodometric titration, the
dark blue complex of starch with iodine and iodide being more visible than iodine
alone. Other complexometric indicators are Eriochrome Black T for the titration of
calcium and magnesium ions, and the chelating agent EDTA used to titrate metal
ions in solution.
82
University Chemistry
Zeta potential titration
Zeta potential titrations are titrations in which the completion is monitored by
the zeta potential, rather than by an indicator, in order to characterize heterogeneous
systems, such as colloids. One of the uses is to determine the iso-electric point when
surface charge becomes zero, achieved by changing the pH or adding surfactant.
Another use is to determine the optimum dose for flocculation or stabilization.
Assay
An assay is a type of biological titration used to determine the concentration
of a virus or bacterium. Serial dilutions are performed on a sample in a fixed ratio
(such as 1:1, 1:2, 1:4, 1:8, etc.) until the last dilution does not give a positive test for
the presence of the virus. The positive or negative value may be determined by
inspecting the infected cells visually under a microscope or by an immunoenzymetric
method such as enzyme-linked immunosorbent assay (ELISA). This value is known
as the titer.
Measuring the endpoint of a titration
Different methods to determine the endpoint include:
•
•
•
•
•
•
Indicator: A substance that changes color in response to a chemical change.
An acid–base indicator (e.g., phenolphthalein) changes color depending
on the pH. Redox indicators are also used. A drop of indicator solution
is added to the titration at the beginning; the endpoint has been reached
when the color changes.
Potentiometer: An instrument that measures the electrode potential of the
solution. These are used for redox titrations; the potential of the working
electrode will suddenly change as the endpoint is reached.
pH meter: A potentiometer with an electrode whose potential depends
on the amount of H+ ion present in the solution. (This is an example of an
ion-selective electrode.) The pH of the solution is measured throughout
the titration, more accurately than with an indicator; at the endpoint there
will be a sudden change in the measured pH.
Conductivity: A measurement of ions in a solution. Ion concentration can
change significantly in a titration, which changes the conductivity. (For
instance, during an acid–base titration, the H+ and OH− ions react to form
neutral H2O.) As total conductance depends on all ions present in the
solution and not all ions contribute equally (due to mobility and ionic
strength), predicting the change in conductivity is more difficult than
measuring it.
Color change: In some reactions, the solution changes color without any
added indicator. This is often seen in redox titrations when the different
oxidation states of the product and reactant produce different colors.
Precipitation: If a reaction produces a solid, a precipitate will form during
the titration. A classic example is the reaction between Ag+ and Cl− to form
83
Reactions Stoichiometry
•
•
•
the insoluble salt AgCl. Cloudy precipitates usually make it difficult to
determine the endpoint precisely. To compensate, precipitation titrations
often have to be done as “back” titrations.
Isothermal titration calorimeter: An instrument that measures the heat
produced or consumed by the reaction to determine the endpoint. Used
in biochemical titrations, such as the determination of how substrates
bind to enzymes.
Thermometric titrimetry: Differentiated from calorimetric titrimetry
because the heat of the reaction (as indicated by temperature rise or fall)
is not used to determine the amount of analyte in the sample solution.
Instead, the endpoint is determined by the rate of temperature change.
Spectroscopy: Used to measure the absorption of light by the solution
during titration if the spectrum of the reactant, titrant or product is known.
The concentration of the material can be determined by Beer’s Law.
Amperometry: Measures the current produced by the titration reaction
as a result of the oxidation or reduction of the analyte. The endpoint is
detected as a change in the current. This method is most useful when the
excess titrant can be reduced, as in the titration of halides with Ag+.
Endpoint and equivalence point
Though the terms equivalence point and endpoint are often used interchangeably,
they are different terms. Equivalence point is the theoretical completion of the
reaction: the volume of added titrant at which the number of moles of titrant is
equal to the number of moles of analyte, or some multiple thereof (as in polyprotic
acids).
There is a slight difference between the endpoint and the equivalence point of
the titration. This error is referred to as an indicator error, and it is indeterminate.
Back titration
Back titration is a titration done in reverse; instead of titrating the original
sample, a known excess of standard reagent is added to the solution, and the excess
is titrated. A back titration is useful if the endpoint of the reverse titration is easier
to identify than the endpoint of the normal titration, as with precipitation reactions.
Back titrations are also useful if the reaction between the analyte and the titrant is
very slow, or when the analyte is in a non-soluble solid.
EXERCISE
Answer the following questions:
1. What do you understand by stoichiometric coefficients? Explain.
2. How to find limiting reagents? Describe.
3. Focus on chemical compositions from measurements of mass.
4. Define the term ‘molar concentration’.
84
University Chemistry
5.
6.
Discuss the volume of solution required for reaction.
Describe titrations.
MULTIPLE CHOICE QUESTIONS
Tick the correct answer
1.
In a particular reaction, one of the reactants limits the number of products
formed. That is called as ………
a Limiting reagent
b. Limiting product
c. Excessive reagent
d. Excessive reactant
2.
Which of the following is not true regarding balanced chemical equations?
a. They contain the same number of atoms on each side
b. Electrons are also balanced
c. An equal number of molecules on both the side
d. Follows the law of conservation of mass
3.
Which of the given reactions are counted as balanced reactions?
a. H2 + O2 → 2H2O
b. 4Al + 3O2 → 2Al2O3
c. Mg(OH)2 + 2HNO3 → 2Mg(NO3)2 + 2H2O
d. N2 + 3H2 → NH3
4.
What is the amount of water produced when 8g of hydrogen is reacted with
32g of oxygen?
a. 2moles
b. 1mole
c. 3 moles
d. 0.5mole
5.
Calculate the mass percent of magnesium in the formation of magnesium
oxide.
a. 0.3
b. 1.5
c. 0.67
d. 0.6
6.
A solution contains 8 moles of solute and the mass of the solution is 4 kg.
What’s the molality of this solution?
a. 5 mol/kg
b. 8 mol/kg
85
Reactions Stoichiometry
7.
8.
9.
10.
c. 4 mol/kg
d. 0.5 mol/kg
In a container, there are 4 moles of nitrogen, 3 moles of oxygen and 7 moles
of hydrogen; find out the mole fraction of oxygen in this reaction.
a. 0.2143
b. 0.2142
c. 0.1234
d. 0.2434
Find the amount of carbon dioxide produced by the combustion of 20g of
methane.
a. 44g
b. 20g
c. 66g
d. 22g
What’s the balanced equation of CO2 + H2O → C6H12O6 + O2?
a. CO2 + H2O → C6H12O6 + O2
b. 6 CO2 + 6 H2O → C6H12O6 + 6 O2
c. 6 CO2 + 6 H2O → C6H12O6 + 2 O2
d. 3 CO2 + 2H2O → C6H12O6 + O2
The molecules of a liquid which is in equilibrium with its vapor at its
boiling point on an average have equal in the two phases.
a. Potential energy
b. Intermolecular forces
c. Kinetic energy
d. Total energy
ANSWERS
1. (a)
2. (c)
3. (b)
4. (a)
5. (d)
6. (d)
7. (a)
8. (c)
9. (b)
10. (c)
REFERENCES
1.
2.
3.
Anbar, A. D. (2008) OCEANS: Elements and Evolution, Science, 322, 5907,
1481–1483.
Chopra, A. & Lineweaver, C. H. (2009) The major elemental abundance
differences between life, the oceans and the Sun, Reviewed Proc. of the 8th
Australian Space Sci. Conf., 49–55
J. C. Kotz P.M. Treichel, J. Townsend. Chemistry and Chemical Reactivity.
Brooks Cole, February 7, 2008.
86
University Chemistry
4.
Palsson, B. Ø. Systems Biology: Constraint–based Reconstruction and Analysis;
Cam- bridge Un. Press: Cambridge, UK, 2015.
Petrucci, harwood, Herring, Madura. General Chemistry Principles & Modern
Applications. Prentice Hall. New Jersey, 2007.
Pienta, N. J. Balancing Redox Equations. Journal of Chemical Education 2010,
87, 477.
Silberberg, M. S., & Amateis, P. (2018). Chemistry: the molecular nature of
matter and change. McGraw-Hill Education, New York
Smith, W. R.; Missen, R. W. Mass Conservation implications of a reaction
mechanism. J. Chem. Educ. 2003, 80, 833–838.
T. E. Brown, H.E LeMay, B. Bursten, C. Murphy. Chemistry: The Central
Science. Prentice Hall, January 8, 2008.
5.
6.
7.
8.
9.
Atomic Structure and the Periodic Table
87
CHAPTER 5
ATOMIC STRUCTURE AND THE PERIODIC TABLE
OBJECTIVES
After studying this chapter, you will be able to:
1. Understand light and spectroscopy
2. Explain the characteristics of light
3. Focus on quantization and photons
4. Describe the structure of the hydrogen atom
5. Learn the spectrum of atomic hydrogen
6. Discuss about particles and waves
7. Know the structure of many-electron atoms
8. Focus on orbital energies
9. Identify the building –up principle
10. Perform a survey of periodic table
11. Know blocks ,periods, and groups
12. Explain periodicity of physical properties
13. Describe trends in chemical properties
INTRODUCTION
An atom is composed of a nucleus and electrons that go around the former. The
nucleus is composed of protons with a positive charge and neutrons without charge,
and the number of protons (atomic number) determines the chemical properties of
the atom (element type).
88
University Chemistry
For example, carbon has six protons, but there are also types of carbon with
five, six, seven or eight neutrons. All of them have the same chemical properties.
When calling them distinctively, they are called Carbon 11, Carbon 12, Carbon
13 and Carbon 14, adding the nuclear number (total of protons and neutrons) after
the element name, which is a nominal designation that covers the same types of
atoms. Carbon 12 is the one that most commonly exists in nature.
Carbon 14 is a radionuclide which exists in nature and is made through a
process where a proton of Nitrogen 14 is hit and removed by a neutron originating
from cosmic rays. Carbon 14 has six protons and eight neutrons, and the state is
energetically unstable because of the unbalance of both numbers.
If one neutron of Carbon 14 changes to a proton, the element becomes stable
because the numbers of protons and neutrons are both seven. At this time, an
electron is emitted as extra energy. This is the identity of β (beta)-particles. In other
words, Carbon 14 returns to nitrogen having seven protons by emitting β-particles,
and becomes energetically stable.
5.1 LIGHT AND SPECTROSCOPY
To understand the processes in astronomy that generate light, we must realize
first that light acts like a wave. Light has particle-like properties too, so it’s actually
quite a twisted beast (which is why it took so many years to figure out). But right
now, let’s just explore light as a wave.
Picture yourself wading around on an ocean beach for a moment, and watch
the many water waves sweeping past you. Waves are disturbances, ripples on the
water, and they possess a certain height (amplitude), with a certain number of waves
rushing past you every minute (the frequency) and all moving at a characteristic
speed across the water (the wave speed). Notice the distance between successive
waves? That’s called the wavelength.
Source: http://loke.as.arizona.edu/~ckulesa/camp/spectroscopy_intro.html
Atomic Structure and the Periodic Table
89
Keeping this analogy in mind, let’s leave the ocean beach for a while and think
about light like a wave. The wave speed of a light wave is simply the speed of light,
and different wavelengths of light manifest themselves as different colors! The energy of
a light wave is inversely-proportional to its wavelength; in other words, low-energy
waves have long wavelengths, and high-energy light waves have short wavelengths.
Chemists study how different forms of electromagnetic radiation interact with
atoms and molecules. This interaction is known as spectroscopy. Just as there are
various types of electromagnetic radiation, there are various types of spectroscopy
depending on the frequency of light we are using. We will begin our discussion by
considering UV-Vis spectroscopy – that is, what occurs within atoms and molecules
when photons in the UV and visible ranges of the spectrum (wavelengths of about
10−700 nm are absorbed or emitted.
5.1.1 The Characteristics of Light
The Characteristics of light are:
•
•
•
•
•
•
•
•
•
•
Light is an electromagnetic wave.
Light travels in a straight line.
Light is a transverse wave, and does not need any medium to travel. Light
can travel through vaccum. Its speed through vaccum is 3 × 108 m/s.
The velocity of light changes when it travels from one medium to another.
The wavelength (λ) of light changes when it goes from one medium to
another.
The frequency (f) of the light wave remains the same in all media.
Light gets reflected back from polished surfaces, such as mirrors, polished
metal surfaces, etc.
Light undergoes refraction (bending) when it travels from one transparent
medium to another.
Light does not need a material medium to travel, that is, it can travel
through a vacuum too. Scientists have assigned a value of 299, 792, 458
m/s to the speed of light in vacuum.
According to current scientific theories, no material particle can travel at
a speed greater than that of light in vacuum.
5.1.2 Quantization and Photons
By the late 19th century, many physicists thought their discipline was well on
the way to explaining most natural phenomena. They could calculate the motions
of material objects using Newton’s laws of classical mechanics, and they could
describe the properties of radiant energy using mathematical relationships known
as Maxwell’s equations, developed in 1873 by James Clerk Maxwell, a Scottish
physicist. The universe appeared to be a simple and orderly place, containing
90
University Chemistry
matter, which consisted of particles that had mass and whose location and motion
could be accurately described, and electromagnetic radiation, which was viewed as
having no mass and whose exact position in space could not be fixed. Thus matter
and energy were considered distinct and unrelated phenomena. Soon, however,
scientists began to look more closely at a few inconvenient phenomena that could
not be explained by the theories available at the time.
Blackbody Radiation
One phenomenon that seemed to contradict the theories of classical physics was
blackbody radiation, which is electromagnetic radiation given off by a hot object.
The wavelength (i.e. color) of radiant energy emitted by a blackbody depends on
only its temperature, not its surface or composition. Hence an electric stove burner
or the filament of a space heater glows dull red or orange when heated, whereas the
much hotter tungsten wire in an incandescent light bulb gives off a yellowish light.
Figure 5.1: Blackbody Radiation. When heated, all objects emit electromagnetic radiation whose wavelength (and color) depends on the temperature of the object. A relatively low-temperature object, such as a horseshoe forged by a blacksmith, appears red,
whereas a higher-temperature object, such as the surface of the sun, appears yellow or
white.
Source: https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Chemistry_-_The_Central_Science_(Brown_et_al.)/06._Electronic_Structure_of_Atoms/6.2%3A_
Quantized_Energy_and_Photons
The intensity of radiation is a measure of the energy emitted per unit area. A
plot of the intensity of blackbody radiation as a function of wavelength for an object
at various temperatures is shown in Figure 5.2. One of the major assumptions of
classical physics was that energy increased or decreased in a smooth, continuous
manner. For example, classical physics predicted that as wavelength decreased, the
intensity of the radiation an object emits should increase in a smooth curve without
limit at all temperatures, as shown by the broken line for 6000 K in Figure 5.2. Thus
classical physics could not explain the sharp decrease in the intensity of radiation
emitted at shorter wavelengths (primarily in the ultraviolet region of the spectrum),
Atomic Structure and the Periodic Table
91
which was referred to as the “ultraviolet catastrophe.” In 1900, however, the
German physicist Max Planck (1858–1947) explained the ultraviolet catastrophe by
proposing (in what he called “an act of despair”) that the energy of electromagnetic
waves is quantized rather than continuous. This means that for each temperature,
there is a maximum intensity of radiation that is emitted in a blackbody object,
corresponding to the peaks in Figure 5.2, so the intensity does not follow a smooth
curve as the temperature increases, as predicted by classical physics. Thus energy
could be gained or lost only in integral multiples of some smallest unit of energy, a
quantum.
Figure 5.2: Relationship between the Temperature of an Object and the Spectrum of
Blackbody Radiation it Emits. At relatively low temperatures, most radiation is emitted at wavelengths longer than 700 nm, which is in the infrared portion of the spectrum. The dull red glow of the electric stove element in Figure 5.1 is due to the small
amount of radiation emitted at wavelengths less than 700 nm, which the eye can detect.
As the temperature of the object increases, the maximum intensity shifts to shorter
wavelengths, successively resulting in orange, yellow, and finally white light. At high
temperatures, all wavelengths of visible light are emitted with approximately equal
intensities. The white light spectrum shown for an object at 6000 K closely approximates the spectrum of light emitted by the sun (Figure 5.1). Note the sharp decrease
in the intensity of radiation emitted at wavelengths below 400 nm, which constituted
the ultraviolet catastrophe. The classical prediction fails to fit the experimental curves
entirely and does not have a maximum intensity.
Source: https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Chemistry_-_The_Central_Science_(Brown_et_al.)/06._Electronic_Structure_of_Atoms/6.2%3A_
Quantized_Energy_and_Photons
Although quantization may seem to be an unfamiliar concept, we encounter
it frequently. For example, US money is integral multiples of pennies. Similarly,
musical instruments like a piano or a trumpet can produce only certain musical
notes, such as C or F sharp. Because these instruments cannot produce a continuous
range of frequencies, their frequencies are quantized. Even electrical charge is
quantized: an ion may have a charge of −1 or −2 but not −1.33 electron charges.
92
University Chemistry
Planck postulated that the energy of a particular quantum of radiant energy
could be described by the equation
E=hu
(1)
where the proportionality constant h is called Planck’s constant, one of the most
accurately known fundamental constants in science. For our purposes, its value to
four significant figures is generally sufficient:
h=6.626×10−34J∙s (joule-segundos)
As the frequency of electromagnetic radiation increases, the magnitude of
the associated quantum of radiant energy increases. By assuming that energy
can be emitted by an object only in integral multiples of hν, Planck devised an
equation that fit the experimental data shown in Figure 5.2. We can understand
Planck’s explanation of the ultraviolet catastrophe qualitatively as follows: At
low temperatures, radiation with only relatively low frequencies is emitted,
corresponding to low-energy quanta. As the temperature of an object increases,
there is an increased probability of emitting radiation with higher frequencies,
corresponding to higher-energy quanta. At any temperature, however, it is simply
more probable for an object to lose energy by emitting a large number of lowerenergy quanta than a single very high-energy quantum that corresponds to
ultraviolet radiation. The result is a maximum in the plot of intensity of emitted
radiation versus wavelength, as shown in Figure 5.2, and a shift in the position of
the maximum to lower wavelength (higher frequency) with increasing temperature.
At the time he proposed his radical hypothesis, Planck could not explain why
energies should be quantized. Initially, his hypothesis explained only one set of
experimental data—blackbody radiation. If quantization were observed for a large
number of different phenomena, then quantization would become a law. In time,
a theory might be developed to explain that law. As things turned out, Planck’s
hypothesis was the seed from which modern physics grew.
The Photoelectric Effect
Only five years after he proposed it, Planck’s quantization hypothesis was used
to explain a second phenomenon that conflicted with the accepted laws of classical
physics. When certain metals are exposed to light, electrons are ejected from their
surface (Figure 5.3). Classical physics predicted that the number of electrons emitted
and their kinetic energy should depend on only the intensity of the light, not its
frequency. In fact, however, each metal was found to have a characteristic threshold
frequency of light; below that frequency, no electrons are emitted regardless of the
light’s intensity. Above the threshold frequency, the number of electrons emitted
was found to be proportional to the intensity of the light, and their kinetic energy
was proportional to the frequency. This phenomenon was called the photoelectric
effect (A phenomenon in which electrons are ejected from the surface of a metal that
has been exposed to light).
Atomic Structure and the Periodic Table
93
Figure 5.3: The Photoelectric Effect (a) Irradiating a metal surface with photons of sufficiently high energy causes electrons to be ejected from the metal. (b) A photocell that
uses the photoelectric effect, similar to those found in automatic door openers. When
light strikes the metal cathode, electrons are emitted and attracted to the anode, resulting in a flow of electrical current. If the incoming light is interrupted by, for example,
a passing person, the current drops to zero. (c) In contrast to predictions using classical
physics, no electrons are emitted when photons of light with energy less than Eo , such
as red light, strike the cathode. The energy of violet light is above the threshold frequency, so the number of emitted photons is proportional to the light’s intensity.
Source: https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Chemistry_-_The_Central_Science_(Brown_et_al.)/06._Electronic_Structure_of_Atoms/6.2%3A_
Quantized_Energy_and_Photons
Albert Einstein (1879–1955; Nobel Prize in Physics,
1921) quickly realized that Planck’s hypothesis about
Important
the quantization of radiant energy could also explain
the photoelectric effect. The key feature of Einstein’s
The fundahypothesis was the assumption that radiant energy arrives
mental building
at the metal surface in particles that we now call photons
blocks of energy are
(a quantum of radiant energy, each of which possesses a
quanta and of matparticular energy energy E given by Equation 1 Einstein
ter are atoms. The
properties of blackpostulated that each metal has a particular electrostatic
body radiation, the
attraction for its electrons that must be overcome before
radiation emitted by
an electron can be emitted from its surface (Eo=uo). If
hot objects, could
photons of light with energy less than Eo strike a metal
not be explained
surface, no single photon has enough energy to eject an
with classical physelectron, so no electrons are emitted regardless of the
ics.
intensity of the light. If a photon with energy greater
than Eo strikes the metal, then part of its energy is used
to overcome the forces that hold the electron to the metal
surface, and the excess energy appears as the kinetic energy of the ejected electron:
kinetic energy of ejected electron=E−Eo =hu−huo =h(u−uo)
When a metal is struck by light with energy above the threshold energy Eo, the
number of emitted electrons is proportional to the intensity of the light beam, which
corresponds to the number of photons per square centimeter, but the kinetic energy
of the emitted electrons is proportional to the frequency of the light. Thus Einstein
94
University Chemistry
showed that the energy of the emitted electrons depended on the frequency of the
light, contrary to the prediction of classical physics. Moreover, the idea that light
could behave not only as a wave but as a particle in the form of photons suggested
that matter and energy might not be such unrelated phenomena after all.
Figure 5.4: A Beam of Red Light Emitted by a Helium Neon laser reads a bar code.
Originally Helium neon lasers, which emit red light at a wavelength of 632.8 nm, were
used to read bar codes. Today, smaller, inexpensive diode lasers are used.
Source: https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Chemistry_-_The_Central_Science_(Brown_et_al.)/06._Electronic_Structure_of_Atoms/6.2%3A_
Quantized_Energy_and_Photons
Planck’s and Einstein’s postulate that energy is quantized is in many ways
similar to Dalton’s description of atoms. Both theories are based on the existence
of simple building blocks, atoms in one case and quanta of energy in the other. The
work of Planck and Einstein thus suggested a connection between the quantized
nature of energy and the properties of individual atoms.
5.2 THE STRUCTURE OF THE HYDROGEN ATOM
A hydrogen atom is, at its most general definition, a proton.
The subatomic proton is a hydrogen atom, but a positively charged one
That structure is, for diagram structure, a circle.
Hydrogen atoms can have neutrons, such as deuterium, which has one proton
and one neutron (two circles joined) that makes water heavy.
For electrons, hydrogen atoms at a neutral state, will have one electron in a 1s
orbital
A negatively charged hydrogen atom could have more electrons, in more
clouds, but generally, it will be a single proton (circle or sphere, depending if 2d or
3d nucleus) surrounded by a single orbital in a cloud, (circle on a circle or a sphere,
depending if 2d or 3d orbitals)
Atomic Structure and the Periodic Table
95
Source: https://www.quora.com/What-is-the-atomic-structure-of-a-hydrogen-atom
5.2.1 The Spectrum of Atomic Hydrogen
We all know that electrons in an atom or a molecule absorb energy and get
excited, they jump from a lower energy level to a higher energy level, and they emit
radiation when they come back to their original states. This phenomenon accounts
for the emission spectrum through hydrogen too, better known as the hydrogen
emission spectrum.
Source: https://byjus.com/chemistry/hydrogen-spectrum/
In the late 1800s, it was known that when a gas is excited using an electric
discharge and the light emitted is viewed through a diffraction grating; the spectrum
observed consists not of a continuous band of light, but of individual lines with
well-defined wavelengths. Experiments have shown that the wavelengths of the
lines were characteristic of the chemical element emitting the light. They were an
atomic fingerprint which resulted from the internal structure of the atom.
The hydrogen spectrum is an important piece of evidence to show the quantized
electronic structure of an atom. The hydrogen atoms of the molecule dissociate as
soon as an electric discharge is passed through a gaseous hydrogen molecule. It
96
University Chemistry
results in the emission of electromagnetic radiation initiated by the energetically
excited hydrogen atoms. The hydrogen emission spectrum comprises radiation of
discrete frequencies. These series of radiation are named after the scientists who
discovered them.
Hydrogen spectrum wavelength
When a hydrogen atom absorbs a photon, it causes the electron to experience
a transition to a higher energy level, for example, n = 1, n = 2. When a photon
is emitted through a hydrogen atom, the electron undergoes a transition from a
higher energy level to a lower, for example, n = 3, n = 2. During this transition from a
higher level to a lower level, there is the transmission of light occurs. The quantized
energy levels of the atoms, cause the spectrum to comprise wavelengths that reflect
the differences in these energy levels. For example, the line at 656 nm corresponds
to the transition n = 3 n = 2.
Source: https://byjus.com/chemistry/hydrogen-spectrum/
Hydrogen emission spectrum:
In the year 1885, on the basis of experimental observations, Balmer proposed
the formula for correlating the wave number of the spectral lines emitted and the
energy shells involved. This formula is given as:
 1 1 
=
v 109677  2 − 2 
2 n 
This series of the hydrogen emission spectrum is known as the Balmer series.
This is the only series of lines in the electromagnetic spectrum that lies in the visible
region. The value, 109,677 cm-1, is called the Rydberg constant for hydrogen. The
Balmer series is basically the part of the hydrogen emission spectrum responsible
for the excitation of an electron from the second shell to any other shell. Similarly,
other transitions also have their own series names. Some of them are listed below,
•
•
•
Transition from the first shell to any other shell – Lyman series
Transition from the second shell to any other shell – Balmer series
Transition from the third shell to any other shell – Paschen series
97
Atomic Structure and the Periodic Table
•
•
Transition from the fourth shell to any other shell – Bracket series
Transition from the fifth shell to any other shell – Pfund series
Source: https://byjus.com/chemistry/hydrogen-spectrum/
Johannes Rydberg, a Swedish spectroscopist, derived a general formula for
the calculation of wave number of hydrogen spectral line emissions due to the
transition of an electron from one orbit to another. The general formula for the
hydrogen emission spectrum is given by:
 1 1 
=
v 109677  2 − 2 
 n1 n2 
Where,
n1 = 1,2,3,4 …
n2 = n1 +1
Hints
Most light is
polychromatic and
contains light of
many wavelengths.
ν= wave number of electromagnetic radiation. The
value 109,677 cm-1 is known as Rydberg constant for
hydrogen.
5.2.2 Particles and Waves
Wave-particle duality, possession by physical entities (such as light and
electrons) of both wavelike and particle-like characteristics. On the basis of
experimental evidence, German physicist Albert Einstein first showed (1905) that
light, which had been considered a form of electromagnetic waves, must also be
thought of as particle-like, localized in packets of discrete energy. The observations
of the Compton effect (1922) by American physicist Arthur Holly Compton could
be explained only if light had a wave-particle duality. French physicist Louis de
Broglie proposed (1924) that electrons and other discrete bits of matter, which
until then had been conceived only as material particles, also have wave properties
such as wavelength and frequency. Later (1927) the wave nature of electrons was
experimentally established by American physicists Clinton Davisson and Lester
Germer and independently by English physicist George Paget Thomson. An
understanding of the complementary relation between the wave aspects and the
particle aspects of the same phenomenon was announced by Danish physicist Niels
Bohr in 1928.
98
University Chemistry
5.3 THE STRUCTURE OF MANY-ELECTRON ATOMS
The quantum mechanical model allowed us to determine the energies of the
hydrogen atomic orbitals; now we would like to extend this to describe the electronic
structure of every element in the Periodic Table. The process of describing each
atom’s electronic structure consists, essentially, of beginning with hydrogen and
adding one proton and one electron at a time to create the next heavier element
in the table; however, interactions between electrons make this process a bit more
complicated than it sounds. All stable nuclei other than hydrogen also contain one
or more neutrons. Because neutrons have no electrical charge, however, they can be
ignored in the following discussion. Before demonstrating how to do this, however,
we must introduce the concept of electron spin and the Pauli principle.
5.3.1 Orbital Energies
Unlike in hydrogen-like atoms with only one electron, in multielectron atoms
the values of quantum numbers n and l determine the energies of an orbital. The
energies of the different orbitals for a typical multielectron atom are shown in Figure
5.5. Within a given principal shell of a multielectron atom, the orbital energies
increase with increasing l. An ns orbital always lies below the corresponding np
orbital, which in turn lies below the nd orbital.
Figure 5.5: Orbital Energy Level Diagram for a Typical Multielectron Atom.
Source: https://chem.libretexts.org/Bookshelves/General_Chemistry/Map%3A_Chemistry_-_The_Central_Science_(Brown_et_al.)/06._Electronic_Structure_of_Atoms/6.7%3A_
Many-Electron_Atoms
Atomic Structure and the Periodic Table
99
These energy differences are caused by the effects
of shielding and penetration, the extent to which a given
Hints
orbital lies inside other filled orbitals. For example,
an electron in the 2s orbital penetrates inside a filled
Due to
1s orbital more than an electron in a 2p orbital does.
electron shielding,
Zeff increases more
Since electrons, all being negatively charged, repel each
rapidly going across
other, an electron closer to the nucleus partially shields
a row of the periodan electron farther from the nucleus from the attractive
ic table than going
effect of the positively charged nucleus. Hence in an
down a column.
atom with a filled 1s orbital, the effective nuclear charge
(Zeff) experienced by a 2s electron is greater than the
Zeff experienced by a 2p electron. Consequently, the 2s
electron is more tightly bound to the nucleus and has a lower energy, consistent
with the order of energies shown in Figure 5.5.
Notice in Figure 5.5 that the difference in energies between subshells can be
so large that the energies of orbitals from different principal shells can become
approximately equal. For example, the energy of the 3d orbitals in most atoms is
actually between the energies of the 4s and the 4p orbitals.
5.3.2 The Building –up Principle
It is not possible to proceed in an analogous manner for many-electron atoms
because the Schrödinger equation cannot be solved analytically for such systems.
However, one can think of the atom as a composition of the nucleus and all its
electrons, constructed in the following way:
•
•
•
The spatial arrangement of the atom is characterised by the same set of
orbitals as for the one-electron hydrogen atom.
The orbitals are filled successively with the atom’s electrons. In the ground
state of the atom each electron prefers to fill the orbital which is lowest in
energy. There are well-working rules (Madelung’s rule, Hund’s rules) of
thumb for the ordering of the orbitals and will be presented more detailed
below.
The number of electrons within a specific orbital is limited to a maximum
of two by the Pauli exclusion principle: A quantum state specified by the
four quantum numbers n, l, m and s can be possessed by at most one
particle at the same time. Due to the two spin quantum numbers s there
are two quantum states associated with each orbital. Therefore it can be
filled with up to two electrons.
In order to create ground state electron configurations for any element, it is
necessary to know the way in which the atomic sublevels are organized in order of
increasing energy. Figure 5.6 shows the order of increasing energy of the sublevels.
100
University Chemistry
Figure 5.6: Electrons are added to atomic orbitals in order from low energy (bottom
of the graph) to high (top of the graph), according to the Aufbau principle. Principle
energy levels are color coded, while sublevels are grouped together, and each circle
represents an orbital capable of holding two electrons.
Source: https://chem.libretexts.org/Bookshelves/Introductory_Chemistry/Book%3A_Introductory_Chemistry_(CK-12)/05%3A_Electrons_in_Atoms/5.14%3A_Aufbau_Principle
The lowest energy sublevel is always the 1s sublevel,
which consists of one orbital. The single electron of the
Important
hydrogen atom will occupy the 1s orbital when the
atom is in its ground state. As we proceed to atoms
The Aufbau
with multiple electrons, those electrons are added to the
principle gives the
next lowest sublevel: 2s, 2p, 3s, and so on. The Aufbau
order of electron filling in an atom.
principle states that an electron occupies orbitals in order
from lowest energy to highest. The Aufbau (German for
building up, construction) principle is sometimes referred
to as the “building up” principle. It is worth noting that in reality, atoms are not
built by adding protons and electrons one at a time, and that this method is merely
an aid to understand the end result.
As seen in the figure above, the energies of the sublevels in different principal
energy levels eventually begin to overlap. After the 3p sublevel, it would seem logical
Atomic Structure and the Periodic Table
101
that the 3d sublevel should be the next lowest in energy. However, the 4s sublevel
is slightly lower in energy than the 3d sublevel and thus fills first. Following the
filling of the 3d sublevel is the 4p, then the 5s and the 4d. Note that the 4f sublevel
does not fill until just after the 6s sublevel. Figure 5.7 is a useful and simple aid for
keeping track of the order of fill of the atomic sublevels.
Figure 5.7: The Aufbau principle is illustrated in the diagram by following each red arrow in order from top to bottom: 1s, 2s, 2p, 3s, etc.
Source: https://chem.libretexts.org/Bookshelves/Introductory_Chemistry/Book%3A_Introductory_Chemistry_(CK-12)/05%3A_Electrons_in_Atoms/5.14%3A_Aufbau_Principle
5.4 A SURVEY OF PERIODIC TABLE
Periodic table, in full periodic table of the elements, in chemistry, the organized
array of all the chemical elements in order of increasing atomic number—i.e., the
total number of protons in the atomic nucleus. When the chemical elements are thus
arranged, there is a recurring pattern called the “periodic law” in their properties,
in which elements in the same column (group) have similar properties. The initial
discovery, which was made by Dmitry I. Mendeleyev in the mid-19th century, has
been of inestimable value in the development of chemistry.
It was not actually recognized until the second decade of the 20th century
that the order of elements in the periodic system is that of their atomic numbers,
the integers of which are equal to the positive electrical charges of the atomic
nuclei expressed in electronic units. In subsequent years great progress was made
in explaining the periodic law in terms of the electronic structure of atoms and
molecules. This clarification has increased the value of the law, which is used as
much today as it was at the beginning of the 20th century, when it expressed the
only known relationship among the elements.
102
University Chemistry
Figure 5.8: Periodic table with each element’s atomic number, symbol, and atomic
weight.
Source: https://www.britannica.com/science/periodic-table
5.4.1 Blocks, Periods, and Groups
The period of an element corresponds to the principal quantum number of the
valence shell.
The block of an element corresponds to the type of orbital which receive the
last electron.
The group of an element is predicted from the number of electrons in the
valence shell or/and penultimate shell as follows:
•
•
•
For s block elements ,group number is equal to the number of valence
electrons.
For p block elements ,group number is equal to 10+number of electrons
in the valence shell.
For d block elements ,group number is equal to the number of electrons in
a (n-1) sub shell + the number of electrons in valence shell.
5.4.2 Periodicity of Physical Properties
The recurrence of similar properties of the elements when they are arranged
in the order of increasing atomic number, after certain regular intervals, is called
periodicity.
103
Atomic Structure and the Periodic Table
Cause of periodicity of elements
The modern periodic table is based on the electronic
configuration of the elements. The properties of an
element are determined largely by the electrons in its
outermost or valence shell. Valence electrons interact
with other atoms and take part in all chemical reactions,
while inner shell electrons have little influence on the
properties of elements.
Important
The elements
in the periodic table
are arranged in
order of increasing
atomic number.
When elements are placed in the order of their
increasing atomic number, the elements having the same number of valence shell
electron is repeated in such a way, so as to fall under the same group. Since, the
electronic configuration of the valence shell electrons is same they show similar
properties.
5.4.3 Trends in Chemical Properties
The systematic arrangement of elements in a periodic table discloses certain
periodic trends in the properties of elements. For example, atomic radii and ionic
radii decrease from left to right, moving in a period. Understanding the trends in
fundamental properties of elements (atomic and ionic radii, ionization enthalpy
and electron gain enthalpy) will let you conclude that the periodicity in properties
basically depends on the electronic configuration of an element. After studying the
concepts in detail, we can discover a relationship between chemical properties and
fundamental properties of elements.
Let us first understand the trends of fundamental properties of an element.
Atomic Radii and Ionic radii
The atomic radii and ionic radii of elements decrease while moving from left
to right in a period. They increase on moving from top to bottom in a group as the
number of shells increases with the increase in atomic number.
Ionization Enthalpy
While moving from left to right in a period, atomic radius decreases. So if the
size of an atom decreases, the attractive force between the nucleus and the outermost
electrons increases. Due to this, across a period in the periodic table, ionization
energy generally increases. However, when we see the trend of ionization enthalpy
in the groups, it decreases from top to bottom in a group. This is because the number
of shells increases down the group, due to which the outermost electrons will be
far away from the nucleus and consequently, the effective nuclear charge is less.
Secondly, the shielding effect also increases down the group with the increasing
number of shells, which in turn results in the decreasing ionization energy.
104
University Chemistry
Electron gain enthalpy
Electron gain enthalpy becomes more negative as we
move from left to right in a period.
•
•
Negative: When energy is released while
accepting an electron.
Positive: When energy is supplied to an atom
while adding an electron.
Hints
The modern
periodic table is
based on the law
that an element’s
properties are a
periodic function of
its atomic number.
These properties
are related to the
elements’ electronic
configuration.
Concluding the trends of properties in the periodic
table from above, we can say that the elements at the two
extremities of the periodic table are highly reactive (note:
noble gases have completely filled shells; hence they are
least reactive) and the elements at the center are the lowest
reactive element. The extreme left elements are the alkalis
which easily lose electrons to form cations. On the other
side are halogens – the elements on the extreme right which easily gain electrons
to form an anion. We can relate it to the metallic and non-metallic characteristics of
elements. Metallic property decreases while non-metallic characteristics increase
while moving from left to right in a period. In a group, metallic properties increases
and non-metallic property decreases down the group.
Chemical reactivity
Chemical reactivity of an element can be clearly understood by studying the
reaction of the element with oxygen and halogens. Elements combine with oxygen
to form oxides. Elements on the extreme left of the periodic table react with oxygen
to form basic oxides (Na2O), while elements on the extreme right form acidic oxides
on reaction with oxygen (Cl2O7). Oxides of the elements in the centre of the periodic
table are amphoteric (Al2O3). Amphoteric oxides are those which behave as acids
as well as bases.
EXERCISE
Answer the following questions.
1. What type of light is used in spectroscopy?
2. What is a spectroscope and what does it do to light?
3. What do you mean by quantization of energy and photon?
4. What is the fine structure of hydrogen atom?
5. What happens to particles in a wave?
6. How do you structure an electron?
7. How do you describe the periodic table?
8. What are the physical properties affected by periodicity?
9. What does chemical trend mean?
Atomic Structure and the Periodic Table
105
MULTIPLE CHOICE QUESTIONS
Tick the correct answer.
1.
Which of the following has a positive charge?
a. proton
b. neutron
c. anion
d. electron
e. atom
2.
Rutherford carried out experiments in which a beam of alpha particles was
directed at a thin piece of metal foil. From these experiments he concluded
that:
a. electrons are massive particles.
b. the positively charged parts of atoms are moving about with a velocity
approaching the speed of light.
c. the positively charged parts of atoms are extremely small and extremely
heavy particles.
d. the diameter of an electron is approximately equal to that of the nucleus.
e. electrons travel in circular orbits around the nucleus.
3.
Consider the species 72Zn, 75As and 74Ge. These species have:
a. the same number of electrons.
b. the same number of protons.
c. the same number of neutrons.
d. the same number of protons and neutrons.
e. the same mass number.
4.
The neutral atoms of all of the isotopes of the same element have
a. different numbers of protons.
b. equal numbers of neutrons.
c. the same number of electrons.
d. the same mass numbers.
e. the same masses.
5.
What is the atomic weight of a hypothetical element consisting of two
isotopes, one with mass = 64.23 amu (26.0%), and one with mass = 65.32
amu?
a. 65.3 amu
b. 64.4 amu
c. 64.9 amu
d. 65.0 amu
e. 64.8 amu
106
6.
University Chemistry
Naturally occurring rubidium consists of just two isotopes. One of the
isotopes consists of atoms having a mass of 84.912 amu; the other of 86.901
amu. What is the percent natural abundance of the heavier isotope?
a. 15%
b. 28%
c. 37%
d. 72%
e. 85%
ANSWERS
1. (a)
2. (c)
3. (c)
4. (c)
5. (d)
6. (b)
REFERENCES
1.
2.
3.
4.
5.
6.
7.
Alberty, Robert A. et. al, Physical Chemistry, 3rd Edition, 2001, John Wiley &
Sons, Inc, pg. 380.
Atkins, Peter et. al, Physical Chemistry, 7th Edition, 2002, W.H Freeman and
Company, New York, pg. 390.
J. Barrett, A. G. Davies, D. Phillips, E. W. Abel, J. Woollins Atomic Structure
and Periodicity The Royal Society of Chemistry 2007 (p. 39)
Kots, John C. et. al, Chemistry & Chemical Reactivity, 5th Edition, 2003,
Thomson Learning Inc, pg. 305-309.
Petrucci, Ralph H, et al. General Chemistry: Principles and Modern Applications.
9th Ed. New Jersey: Pearson, 2007.
Russo, Steve, and Mike Silver. Introductory Chemistry. San Francisco: Pearson,
2007.
S. H. Simon Oxford Solid State Basics Oxford 2013 (ch. 5.2)
107
The Chemical Bond
CHAPTER 6
THE CHEMICAL BOND
OBJECTIVES
After studying this chapter, you will be able to:
1. Discuss on ionic bonds
2. Learn about the covalent bonds
3. Evaluate the Lewis structures of polyatomic molecules
Chemical bonding, any of the interactions that account for the association of
atoms into molecules, ions, crystals, and other stable species that make up the
familiar substances of the everyday world. When atoms approach one another,
their nuclei and electrons interact and tend to distribute themselves in space in such
a way that the total energy is lower than it would be in any alternative arrangement.
If the total energy of a group of atoms is lower than the sum of the energies of the
component atoms, they then bond together and the energy lowering is the bonding
energy.
A chemical bond is a bond that holds atoms together. It is the force that
binds ions or molecules together. It helps form a chemical compound. Examples of
the chemical compounds that are of special interest to biologists are water, sodium
chloride, and carbon dioxide. These chemical compounds are formed by the
chemical bond that binds the constituent atoms. For instance, water is comprised of
two hydrogen atoms and an oxygen connected by chemical bonds.
108
University Chemistry
There are three types of chemical bonds that are biologically important: (1)
ionic bonds, (2) covalent bonds, and (3) hydrogen bonds. The ionic bond is a type of
chemical bond in which there is a complete transfer of an electron from one atom to
another. An example is the sodium chloride, which is made up of a cation (Na+) and
an anion (Cl–). The two ions of sodium chloride are held together by an ionic bond.
A covalent bond is a type of a chemical bond wherein electrons are shared between
atoms. The bond between hydrogen and oxygen atoms to form water is an example
of a covalent bond. The hydrogen bond is a low-energy electrostatic bond wherein
hydrogen serves as a bridge between two atoms.
6.1 IONIC BONDS
Ionic bond, also called electrovalent bond, type of linkage formed from the
electrostatic attraction between oppositely charged ions in a chemical compound.
Such a bond forms when the valence (outermost) electrons of one atom are
transferred permanently to another atom. The atom that loses the electrons becomes
a positively charged ion (cation), while the one that gains them becomes a negatively
charged ion (anion). A brief treatment of ionic bonds follows.
Ionic bonding results in compounds known as ionic, or electrovalent,
compounds, which are best exemplified by the compounds formed between
nonmetals and the alkali and alkaline-earth metals. In ionic crystalline solids of
this kind, the electrostatic forces of attraction between opposite charges and
repulsion between similar charges orient the ions in such a manner that every
positive ion becomes surrounded by negative ions and vice versa. In short, the ions
are so arranged that the positive and negative charges alternate and balance one
another, the overall charge of the entire substance being zero. The magnitude of the
electrostatic forces in ionic crystals is considerable. Accordingly, these substances
tend to be hard and nonvolatile.
An ionic bond is actually the extreme case of a polar covalent bond, the latter
resulting from unequal sharing of electrons rather than complete electron transfer.
Ionic bonds typically form when the difference in the electronegativity of the two
atoms is great, while covalent bonds form when the electronegativities are similar.
Compare covalent bond.
6.1.1 The Energetics of Ionic Bond Formation
Ionic compounds are formed when electrons are transferred between atoms or
groups of atoms to form charged ions, which then arrange in a crystalline lattice
structure due to electrostatic attraction. The formation of ionic compounds are
usually extremely exothermic. The strength of the electrostatic attraction between
ions with opposite charges is directly proportional to the magnitude of the charges
on the ions and inversely proportional to the internuclear distance. The total energy
of the system is a balance between the repulsive interactions between electrons on
adjacent ions and the attractive interactions between ions with opposite charges.
109
The Chemical Bond
Ionic bonds are formed when positively and negatively charged ions are
attracted by electrostatic forces. Consider a single pair of ions, one cation and one
anion. How strong will the force of their attraction be? We can rewrite Coulomb’s
Law (Equation 9.2.1) quantitatively for any two charged particles:
9.2.1
9.2.2
where each ion’s charge is represented by the symbol \(Q\) and the internuclear
distance between the particles is represented by (r). The proportionality constant k is
equal to 2.31 × 10−28 J·m. This value of \(k\) includes the charge of a single electron
(1.6022 × 10−19 C) for each ion. The equation can also be written using the charge of
each ion, expressed in coulombs (C), incorporated in the constant. In this case, the
proportionality constant, k, equals 8.999 × 109 J·m/C2. In the example given, Q1 =
+1(1.6022 × 10−19 C) and Q2 = −1(1.6022 × 10−19 C). If Q1 and Q2 have opposite signs
(as in NaCl, for example, where Q1 is +1 for Na+ and Q2 is −1 for Cl−), then E is
negative, which means that energy is released when oppositely charged ions are
brought together from an infinite distance to form an isolated ion pair.
As shown by the green curve in the lower half of Figure 1, the maximum energy
would be released when the ions are infinitely close to each other, at r = 0. Because
ions occupy space and have a structure with the positive nucleus being surrounded
by electrons, however, they cannot be infinitely close together. At very short distances,
repulsive electron–electron interactions between electrons on adjacent ions become
stronger than the attractive interactions between ions with opposite charges, as
shown by the red curve in the upper half of Figure 1. The total energy of the system
is a balance between the attractive and repulsive interactions. The purple curve in
Figure 1 shows that the total energy of the system reaches a minimum at r0, the point
where the electrostatic repulsions and attractions are exactly balanced. This distance
is the same as the experimentally measured bond distance.
Figure 1. A Plot of Potential Energy versus Internuclear Distance for the Interaction between a Gaseous Na+ Ion and a Gaseous Cl− Ion. The energy of the system reaches a minimum at
a particular distance (r0) when the attractive and repulsive interactions are balanced.
110
University Chemistry
Consider the energy released when a gaseous Na+ ion and a gaseous Cl− ion are
brought together from r = ∞ to r = r0. Given that the observed gas-phase internuclear
distance is 236 pm, the energy change associated with the formation of an ion pair
from an Na(g)+ ion and a Cl−(g) ion is as follows:
The negative value indicates that energy is released. The convention is that
if a chemical process provides energy to the outside world, the energy change is
negative. If it requires energy, the energy change is positive. To calculate the energy
change in the formation of a mole of NaCl pairs, we need to multiply the energy per
ion pair by Avogadro’s number:
9.2.3
This is the energy released when 1 mol of gaseous ion pairs is formed, not when
1 mol of positive and negative ions condenses to form a crystalline lattice. Because
of long-range interactions in the lattice structure, this energy does not correspond
directly to the lattice energy of the crystalline solid. However, the large negative
value indicates that bringing positive and negative ions together is energetically
very favorable, whether an ion pair or a crystalline lattice is formed.
Table 1. Lattice energies range from around 700 kJ/mol to 4000 kJ/mol
Compound
Lattice Energy (kJ/mol)
LiF
1024
LiI
744
NaF
911
NaCl
788
NaI
693
KF
815
KBr
682
KI
641
MgF2
2910
SrCl2
2130
MgO
3938
111
The Chemical Bond
We summarize the important points about ionic
bonding:
•
•
•
At r0, the ions are more stable (have a lower
potential energy) than they are at an infinite
internuclear distance. When oppositely
charged ions are brought together from r = ∞
to r = r0, the energy of the system is lowered
(energy is released).
Because of the low potential energy at r0, energy
must be added to the system to separate the
ions. The amount of energy needed is the bond
energy.
The energy of the system reaches a minimum
at a particular internuclear distance (the bond
distance).
Hints
Fluorine is
a halogen in group
17. Like all group 17
elements, fluorine
has seven valence
electrons. If fluorine
gains one electron,
it will also have a
full outer energy
level and the most
stable arrangement
of electrons.
6.1.2 Ionic Bond and the Periodic Table
Ionic bonds form only between metals and nonmetals. That’s because metals
“want” to give up electrons, and nonmetals “want” to gain electrons. Find sodium
(Na) in the Figure below. Sodium is an alkali metal in group 1. Like all group 1
elements, it has just one valence electron. If sodium loses that one electron, it will
have a full outer energy level, which is the most stable arrangement of electrons.
Now find fluorine (F) in the periodic table Figure below.
6.2 COVALENT BONDS
Covalent bond, in chemistry, the interatomic linkage that results from the sharing
of an electron pair between two atoms. The binding arises from the electrostatic
attraction of their nuclei for the same electrons. A covalent bond forms when the
bonded atoms have a lower total energy than that of widely separated atoms.
112
University Chemistry
Molecules that have covalent linkages include the inorganic substances
hydrogen, nitrogen, chlorine, water, and ammonia (H2, N2, Cl2, H2O, NH3) together
with all organic compounds. In structural representations of molecules, covalent
bonds are indicated by solid lines connecting pairs of atoms; e.g.,
A single line indicates a bond between two atoms (i.e., involving one electron
pair), double lines (=) indicate a double bond between two atoms (i.e., involving two
electron pairs), and triple lines (≡) represent a triple bond, as found, for example,
in carbon monoxide (C≡O). Single bonds consist of one sigma (σ) bond, double
bonds have one σ and one pi (π) bond, and triple bonds have one σ and two π
bonds.
Covalent bonds are directional, meaning that atoms so bonded prefer specific
orientations relative to one another; this in turn gives molecules definite shapes,
as in the angular (bent) structure of the H2O molecule. Covalent bonds between
identical atoms (as in H2) are nonpolar—i.e., electrically uniform—while those
between unlike atoms are polar—i.e., one atom is slightly negatively charged and
the other is slightly positively charged. This partial ionic character of covalent bonds
increases with the difference in the electronegativity of the two atoms.
When none of the elements in a compound is a metal, no atoms in the compound
have an ionization energy low enough for electron loss to be likely. In such a case,
covalence prevails. As a general rule, covalent bonds are formed between elements
lying toward the right in the periodic table (i.e., the nonmetals). Molecules of
identical atoms, such as H2 and buckminsterfullerene (C60), are also held together
by covalent bonds.
Elements having very high ionization energies are incapable of transferring
electrons and elements having very low electron affinity cannot take up electrons.
The atoms of such elements tend to share their electrons with the atoms of other
elements or with other atoms of the same element in a way that both the atoms
obtain octet configuration in their respective valence shell and thus achieve stability.
Such association through sharing of electron pairs among different or same kinds
is known as Covalent Bond.
The Chemical Bond
113
Covalent Bonding can be Achieved in two Ways:
•
•
Sharing of electrons between atoms of the same kind E.g. Formation of
H2, Cl2, O2, etc.
Sharing of electrons between atoms of different kind E.g. Formation of
CH4, H2O, NH3, etc.
Covalent Bonding in Carbon Atom
As per the electronic configuration of Carbon, it needs to gain or lose 4 electrons
to become stable, which seems impossible as:
•
•
Carbon cannot gain 4 electrons to become C4-, because it will be tough
for 6 protons to hold 10 electrons and so the atom will become unstable.
Carbon cannot lose 4 electrons to become C4+ because it would require a
large amount of energy to remove out 4 electrons and also the C4+ would
have only 2 electrons held by proton, which will again become unstable
Carbon cannot gain or donate electrons, so to complete its nearest noble gas
configuration, it shares electron to form a covalent bond.
Properties of Covalent Bond
If the normal valence of an atom is not satisfied by sharing a single electron
pair between atoms, the atoms may share more than one electron pair between
them. Some of the properties of covalent bonds are:
•
•
•
•
Covalent bonding does not result in the formation of new electrons. The
bond only pairs them.
They are very powerful chemical bonds that exist between atoms.
A covalent bond normally contains the energy of about ~80 kilocalories
per mole (kcal/mol).
Covalent bonds rarely break spontaneously after it is formed.
114
University Chemistry
•
•
•
•
•
Covalent bonds are directional where the atoms that are bonded showcase
specific orientations relative to one another.
Most compounds having covalent bonds exhibit relatively low melting
points and boiling points.
Compounds with covalent bonds usually have lower enthalpies of
vaporization and fusion.
Compounds formed by covalent bonding don’t conduct electricity due to
the lack of free electrons.
Covalent compounds are not soluble in water.
6.2.1 The Electron Pair Bond
An electron pair or Lewis pair consists of two electrons that occupy the same
molecular orbital but have opposite spins. Gilbert N. Lewis introduced the concepts
of both the electron pair and the covalent bond in a landmark paper he published
in 1916.
Because electrons are fermions, the Pauli Exclusion Principle forbids these
particles from having the same quantum numbers. Therefore, for two electrons to
occupy the same orbital, and thereby have the same orbital quantum number, they
must have different spin quantum number. This also limits the number of electrons
in the same orbital to two.
The pairing of spins is often energetically favorable, and electron pairs therefore
play a large role in chemistry. They can form a chemical bond between two atoms,
or they can occur as a lone pair of valence electrons. They also fill the core levels of
an atom.
Because the spins are paired, the magnetic moment of the electrons cancel one
another, and the pair’s contribution to magnetic properties is generally diamagnetic.
Although a strong tendency to pair off electrons can be observed in chemistry,
it is also possible that electrons occur as unpaired electrons.
In the case of metallic bonding the magnetic moments also compensate to a
large extent, but the bonding is more communal so that individual pairs of electrons
cannot be distinguished and it is better to consider the electrons as a collective
‘ocean’.
A very special case of electron pair formation occurs in superconductivity: the
formation of Cooper pairs.
6.3 LEWIS STRUCTURES OF POLYATOMIC
MOLECULES
The idea that two electrons can be shared between two atoms and serve as
The Chemical Bond
115
the link between them was first introduced in 1916 by the American chemist G.N.
Lewis, who described the formation of such bonds as resulting from the tendencies
of certain atoms to combine with one another in order for both to have the electronic
structure of a corresponding noble-gas atom.
In Lewis terms a covalent bond is a shared electron pair. The bond between a
hydrogen atom and a chlorine atom in hydrogen chloride is formulated as follows:
The bond between a hydrogen atom and a chlorine atom in hydrogen chloride
is formulated as per this process. (Lewis formulation of a covalent bond.)
In a Lewis structure of a covalent compound, the shared electron pair between
the hydrogen and chlorine ions is represented by a line. The electron pair is called a
bonding pair; the three other pairs of electrons on the chlorine atom are called lone
pairs and play no direct role in holding the two atoms together.
Each atom in the hydrogen chloride molecule attains a closed-shell octet of
electrons by sharing and hence achieves a maximum lowering of energy. In general,
an incomplete shell means that some attracting power of a nucleus may be wasted,
and adding electrons beyond a closed shell would entail the energetic disadvantage
of beginning the next shell of the atom concerned. Lewis’s octet rule is again
applicable and is seen to represent the extreme means of achieving lower energy
rather than being a goal in itself.
Lewis structures of more complex molecules can be constructed quite simply
by extending the process that has been described for hydrogen chloride. First,
the valence electrons that are available for bonding are counted (2 × 1 + 6 = 8 in
H2O, for example, and 4 + 4 × 7 = 32 in carbon tetrachloride, CCl4), and the chemical
symbols for the elements are placed in the arrangement that reflects which are
neighbors:
Next, one bonding pair is added between each linked pair of atoms:
The remaining electrons are then added to the atoms in such a way that each
atom has a share in an octet of electrons (this is the octet-rule part of the procedure):
116
University Chemistry
Hints
Finally, each bonding pair is represented by a dash:
(Note that Lewis structures do not necessarily show
the actual shape of the molecule, only the topological
pattern of their bonds.)
A covalent
bond forms if the
bonded atoms
have a lower total
energy than the
widely separated
atoms. The simplest
interpretation of the
decrease in energy
that occurs when
electrons are shared
is that both electrons
lie between two attracting centers (the
nuclei of the two
atoms linked by the
bond) and hence lie
lower in energy than
when they experience the attraction
of a single center.
In some older formulations of Lewis structures, a
distinction was made between bonds formed by electrons
that have been supplied by both atoms (as in H―Cl,
where one shared electron can be regarded as supplied
by the hydrogen atom and the other by the chlorine atom)
and covalent bonds formed when both electrons can be
regarded as supplied by one atom, as in the formation of
OH− from O2− and H+. Such a bond was called a coordinate
covalent bond or a dative bond and symbolized O → H−.
However, the difficulties encountered in the attempt to
keep track of the origin of bonding electrons and the suggestion that a coordinate
covalent bond differs somehow from a covalent bond (it does not) have led to this
usage falling into disfavor.
6.3.1 Advanced Aspects of Lewis structures
The Lewis structures illustrated so far have been selected for their simplicity. A
number of elaborations are given below.
Resonance
There is sometimes an ambiguity in the location of double bonds. This ambiguity
is illustrated by the Lewis structure for ozone (O3). The following are two possible
structures:
In such cases, the actual Lewis structure is regarded as a blend of these
contributions and is written:
The Chemical Bond
117
The blending together of these structures is actually a quantum mechanical
phenomenon called resonance. At this stage, resonance can be regarded as a blending
process that spreads double-bond character evenly over the atoms that participate
in it. In ozone, for instance, each oxygen-oxygen bond is rendered equivalent by
resonance, and each one has a mixture of single-bond and double-bond character
(as indicated by its length and strength).
Hypervalence
Lewis structures and the octet rule jointly offer a succinct indication of the type
of bonding that occurs in molecules and show the pattern of single and multiple
bonds between the atoms. There are many compounds, however, that do not
conform to the octet rule. The most common exceptions to the octet rule are the
so-called hypervalent compounds. These are species in which there are more atoms
attached to a central atom than can be accommodated by an octet of electrons. An
example is sulfur hexafluoride, SF6, for which writing a Lewis structure with six
S―F bonds requires that at least 12 electrons be present around the sulfur atom:
(Only the bonding electrons are shown here.) In Lewis terms, hypervalence
requires the expansion of the octet to 10, 12, and even in some cases 16 electrons.
Hypervalent compounds are very common and in general are no less stable than
compounds that conform to the octet rule.
The existence of hypervalent compounds would appear to deal a severe blow to
the validity of the octet rule and Lewis’s approach to covalent bonding if the expansion
of the octet could not be rationalized or its occurrence predicted. Fortunately, it
can be rationalized, and the occurrence of hypervalence can be anticipated. In
simple terms, experience has shown that hypervalence is rare in periods 1 and 2 of
the periodic table (through neon) but is common in and after period 3. Thus, the
octet rule can be used with confidence for carbon, nitrogen, oxygen, and fluorine,
but hypervalence must be anticipated thereafter. The conventional explanation of
this distinction takes note of the fact that in period-3 elements the valence shell
has n = 3, and this is the first shell in which d orbitals are available. (These orbitals
are occupied after the 4s orbitals have been filled and account for the occurrence
of the transition metals in period 4.) It is therefore argued that atoms of this and
subsequent periods can use the empty d orbitals to accommodate electrons beyond
an octet and hence permit the formation of hypervalent species.
In chemistry, however, it is important not to allow mere correlations to
masquerade as explanations. Although it is true that d orbitals are energetically
accessible in elements that display hypervalence, it does not follow that they are
responsible for it. Indeed, quantum mechanical theories of the chemical bond do
not need to invoke d-orbital involvement. These theories suggest that hypervalence
118
University Chemistry
is probably no more than a consequence of the greater radii of the atoms of period-3
elements compared with those of period 2, with the result that a central atom can
pack more atoms around itself. Thus, hypervalence is more a steric (geometric)
problem than an outcome of d-orbital availability.
6.3.2 Lewis Acids and Bases
Lewis acids and bases are described by the Lewis theory of acid-base reactions
as electron-pair acceptors and electron pair donors respectively. Therefore, a Lewis
base can donate a pair of electrons to a Lewis acid to form a product containing a
coordinate covalent bond. This product is also referred to as a Lewis adduct. An
illustration detailing the reaction between a Lewis acid and base leading to the
formation of a coordinate covalent bond between them is given below.
Lewis acids and bases are named after the American chemist Gilbert Newton
Lewis, who also made invaluable contributions in the fields of thermodynamics and
photochemistry.
Lewis Acid
Lewis Acids are the chemical species which have empty orbitals and are able to
accept electron pairs from Lewis bases. This term was classically used to describe
chemical species with a trigonal planar structure and an empty p-orbital. An
example of such a Lewis acid would be BR3 (where R can be a halide or an organic
substituent).
Water and some other compounds are considered as both Lewis acids and
bases since they can accept and donate electron pairs based on the reaction.
Examples of Lewis Acids
Some common examples of Lewis acids which can accept electron pairs include:
•
•
•
•
H+ ions (or protons) can be considered as Lewis acids along with onium
ions like H3O+.
The cations of d block elements which display high oxidation states can
act as electron pair acceptors. An example of such a cation is Fe3+.
Cations of metals such as Mg2+ and Li+ can form coordination compounds
with water acting as the ligand. These aquo complexes can accept electron
pairs and behave as Lewis acids.
Carbocations given by H3C+ and other trigonal planar species tend to
accept electron pairs.
The Chemical Bond
•
119
The Pentahalides of the following group 15 elements can act as Lewis
acids – Antimony, Arsenic, and Phosphorus.
Apart from these chemical compounds listed above, any electron-deficient π
system can act as an acceptor of electron pairs – enones, for example.
Lewis Base
Atomic or molecular chemical species having a highly localized HOMO (The
Highest Occupied Molecular Orbital) act as Lewis bases. These chemical species
have the ability to donate an electron pair to a given Lewis acid in order to form an
addict.
The most common Lewis bases are ammonia, alkyl amines, and other
conventional amines. Commonly, Lewis bases are anionic in nature and their base
strength generally depends on the pKa of the corresponding parent acid. Since
Lewis bases are electron-rich species that have the ability to donate electron-pairs,
they can be classified as nucleophiles. Similarly, Lewis acids can be classified as
electrophiles (since they behave as electron-pair acceptors).
Examples of Lewis Bases
Examples of Lewis bases which have an ability to donate an electron pair are
listed below.
•
•
•
•
Pyridine and the derivatives of pyridine have the ability to act as electron
pair donors. Thus, these compounds can be classified as Lewis bases.
The compounds in which Oxygen, Sulphur, Selenium, and Tellurium
(which belong to group 16 of the Periodic Table) exhibit an oxidation state
of -2 are generally Lewis bases. Examples of such compounds include
water and ketones.
The simple anions which have an electron pair can also act as Lewis bases
by donating these electrons. Examples of such anions include H– and F–.
Even some complex anions, such as the sulfate anion (SO42-) can donate
pairs of electrons.
The π-systems which are rich in electrons (such as benzene, ethyne, and
ethene) exhibit great electron pair donating capabilities.
Weak Lewis acids have strong conjugate Lewis bases. Apart from this, many
chemical species having a lone pair of electrons such as CH3– and OH– are identified
as Lewis bases due to their electron pair donating capabilities.
Applications of Lewis Acids and Bases
Some important applications of Lewis acids and bases are provided below.
Lewis acids play a vital role as a catalyst in the Friedel-Crafts reaction –
AlCl3 accepts a lone pair of electrons belonging to the chloride ion leading to the
formation of AlCl4– in the Friedel-Crafts alkylation process.
120
University Chemistry
This also leads to the formation of the highly electrophilic carbonium ion which
acts as a strong Lewis Acid. The chemical reaction can be written as follows.
RCl + AlCl3 ⟶ R+ + AlCl4–
In the field of organic chemistry, Lewis acids are widely used to encourage
many cationic or pseudo-cationic chemical reactions.
Lewis bases have immense applications in the modification of the selectivity
and the activity of metallic catalysts. For the production of pharmaceuticals,
asymmetric catalysis is an important part of enantioselective synthesis. In order to
enable asymmetric catalysis, chiral Lewis bases are often used to confer chirality on
catalysts.
Several Lewis bases have the ability to form many bonds with Lewis acids.
These compounds are also called ‘multidentate Lewis bases’ or ‘chelating agents’
and have a wide range of industrial and agricultural applications.
6.3.3 Exceptions to the Lewis Octet Rule
The octet rule is a bonding theory used to predict the molecular structure
of covalently bonded molecules. According to the rule, atoms seek to have eight
electrons in their outer—or valence—electron shells. Each atom will share, gain, or
lose electrons to fill these outer electron shells with exactly eight electrons. For many
elements, this rule works and is a quick and simple way to predict the molecular
structure of a molecule.
But, as the saying goes, rules are made to be broken. And the octet rule has
more elements breaking the rule than following it.
While Lewis electron dot structures help determine bonding in most compounds,
there are three general exceptions: molecules in which atoms have fewer than eight
electrons (boron chloride and lighter s- and p- block elements); molecules in which
atoms have more than eight electrons (sulfur hexafluoride and elements beyond
period 3); and molecules with an odd number of electrons (NO.)
Too Few Electrons: Electron Deficient Molecules
Hydrogen, beryllium, and boron have too few electrons to form an octet.
Hydrogen has only one valence electron and only one place to form a bond with
another atom. Beryllium has only two valence atoms, and can form only electron
pair bonds in two locations. Boron has three valence electrons. The two
molecules depicted in this picture show the central beryllium and boron atoms
with fewer than eight valence electrons.
Molecules, where some atoms have fewer than eight electrons, are called
electron deficient.
The Chemical Bond
121
Too Many Electrons: Expanded Octets
Elements in periods greater than period 3 on the periodic table have a d orbital
available with the same energy quantum number. Atoms in these periods may
follow the octet rule, but there are conditions where they can expand their valence
shells to accommodate more than eight electrons.
Sulfur and phosphorus are common examples of this behavior. Sulfur can
follow the octet rule as in the molecule SF2. Each atom is surrounded by eight
electrons. It is possible to excite the sulfur atom sufficiently to push valence atoms
into the d orbital to allow molecules such as SF4 and SF6. The sulfur atom in SF4 has
10 valence electrons and 12 valence electrons in SF6.
Lonely Electrons: Free Radicals
Most stable molecules and complex ions contain pairs of electrons. There is a
class of compounds where the valence electrons contain an odd number of electrons
in the valence shell. These molecules are known as free radicals. Free radicals contain
at least one unpaired electron in their valence shell. In general, molecules with an
odd number of electrons tend to be free radicals.
Nitrogen(IV) oxide (NO2) is a well-known example. Note the lone electron on
the nitrogen atom in the Lewis structure. Oxygen is another interesting example.
Molecular oxygen molecules can have two single unpaired electrons. Compounds
like these are known as biradicals.
122
University Chemistry
6.3.4 The Shapes of Molecules
Drawing a Lewis structure is the first steps towards predicting the threedimensional shape of a molecule. A molecule’s shape strongly affects its physical
properties and the way it interacts with other molecules, and plays an important
role in the way that biological molecules (proteins, enzymes, DNA, etc.) interact
with each other.
The approximate shape of a molecule can be predicted using the Valence-Shell
Electron-Pair Repulsion (VSEPR) model, which depicts electrons in bonds and lone
pairs as “electron groups” that repel one another and stay as far apart as possible:
•
•
•
Draw the Lewis structure for the molecule of interest and count the
number of electron groups surrounding the central atom. Each of the
following constitutes an electron group:
–
a single, double or triple bond (multiple bonds count as one electron
group)
–
a lone pair
–
an unpaired electron
Predict the arrangement of electron groups around each atom by assuming
that the groups are oriented in space as far away from one another as
possible.
The shapes of larger molecules having more than one central are a
composite of the shapes of the atoms within the molecule, each of which
can be predicted using the VSEPR model.
6.3.5 Electron Pair Repulsions
Electron pair repulsion is a theory that informs a wide variety of scientific
disciplines. Physics, engineering, and chemistry use this principle especially often.
The principle that electron pairs around a central atom tend to orient themselves
as far apart as possible. Electron pair repulsion is used to predict the geometry of a
molecule or a polyatomic ion.
According to the electron-pair repulsion theory, the general shape of a molecule
AXn may be predicted from the total number of electron pairs in the valency shell
of A. The extension of this simple theory to account for some of the finer details of
molecular shape is considered. The results of recent structure determinations on
SF4, PF5, CH3PF4, (CH3)2PF3, XeF6, TeBr6-2, and other molecules are discussed in terms
123
The Chemical Bond
of electron-pair repulsions. The apparently anomalous
trigonal prism molecules, such as Re[S2C2(C6H5)2]3.
6.3.6 Molecules with Multiple Bonds
In chemistry, a multiple bond is a chemical bond
where two or more electron pairs are shared between
two atoms. Double and triple bonds are multiple bonds.
In a double bond, four bonding electrons participate
in the bond rather than two electrons in a single bond.
Double bonds are found in azo compounds (N=N),
sulfoxides (S=O), and imines (C=N). The equal sign is
typically used to denote a double bond.
A triple bond involves six bonding electrons. The
triple bond is drawn using three parallel lines (≡). The
most common triple bond occurs in alkynes. Molecular
nitrogen (N2) is an excellent example of a compound
with a triple bond (N≡N).Triple bonds are stronger than
double or single bonds.
Important
Covalent
bonding is the
sharing of one or
more electron pairs.
In many covalent
bonding situations,
multiple chemical
bonds exist — more
than one electron
pair is shared. (In
hydrogen and the
other diatomic
molecules, only
one electron pair is
shared.)
Nitrogen is a diatomic molecule in the VA family on the periodic table. Nitrogen
has five valence electrons, so it needs three more valence electrons to complete its
octet.
A nitrogen atom can fill its octet by sharing three electrons with another
nitrogen atom, forming three covalent bonds, a so-called triple bond. The triple
bond formation of nitrogen is shown in the following figure.
Triple bond formation of nitrogen
It’s also why many explosive compounds (such as TNT and ammonium nitrate)
contain nitrogen. When these compounds break apart in a chemical reaction,
nitrogen gas is formed, and a large amount of energy is released.
Carbon dioxide is another example of a compound containing a multiple bond.
Carbon can react with oxygen to form carbon dioxide. Carbon has four valence
electrons, and oxygen has six. Carbon can share two of its valence electrons with
each of the two oxygen atoms, forming two double bonds. These double bonds are
shown in the following figure.
124
University Chemistry
The formation of carbon dioxide
EXERCISE
Answer the following questions
1. How to evaluate the energetics of ionic bond formation?
2. Find the relationship between ionic bond and the periodic table.
3. How does electron pair bond work?
4. What are the advanced aspects of Lewis structures?
5. Discuss on Lewis acids and bases.
6. What are the exceptions to the Lewis octet rule?
7. Examine the molecules with multiple bonds.
MULTIPLE CHOICE QUESTIONS
Tick the correct answer:
1.
C-O bond length is minimum in
a. CO2
b. CO32c. HCOO–
d. CO
2.
Molecules are held together in a crystal by
a. hydrogen bond
b. electrostatic attraction
c. Van der Waal’s attraction
d. dipole-dipole attraction
3.
Sp3d2 hybridization is present in [Co(NH3)63+], find its geometry
a. octahedral geometry
b. square planar geometry
c. tetragonal geometry
d. tetrahedral geometry
4.
Find the molecule with the maximum dipole moment
a. CH4
b. NH3
c. CO2
d. NF3
125
The Chemical Bond
5.
6.
7.
8.
9.
10.
MX6 is a molecule with octahedral geometry. How many X – M – X bonds are
at 180°?
a. four
b. two
c. three
d. six
Find the pair with sp2 hybridisation of the central molecule
a. NH3 and NO2–
b. BF3 and NH2–
c. BF3 and NO2–
d. NH2– and H2O
The formal charge and P-O bond order in PO43- respectively are
a. 0.6, -0.75
b. -0.75, 1.25
c. 1.0, -0.75
d. 1.25, -3
Which of the molecules does not have a permanent dipole moment?
a. SO3
b. SO2
c. H2S
d. CS2
pp – dp bonding is present in which molecule
a. SO32b. CO32c. NO3–
d. BO33Which one has a pyramidal shape?
a. SO3
b. PCl3
c. CO32d. NO3–
ANSWERS
1. (d)
2. (c)
3. (a)
4. (b)
5. (c)
6. (c)
7. (b)
8. (d)
9. (a)
10. (b)
126
University Chemistry
REFERENCES
1.
2.
3.
4.
5.
Chemguide.co.uk. 2016. ionic (electrovalent) bonding. [online] Available at:
<http://chemguide.co.uk/atoms/bonding/ionic.html> [Accessed 7 February
2016].
Chemwiki.ucdavis.edu. 2013. Ionic and Covalent Bonds - Chemwiki.
[online] Available at: <http://chemwiki.ucdavis.edu/Organic_Chemistry/
Fundamentals/Ionic_and_Covalent_Bonds> [Accessed 7 February 2016].
Gillespie, R.J. (2004), “Teaching molecular geometry with the VSEPR model”,
Journal of Chemical Education, 81 (3): 298–304, Bibcode:2004JChEd..81..298G,
Housecroft, Catherine E.; Sharpe, Alan G. (2005). Inorganic Chemistry (2nd
ed.). Pearson Prentice-Hal. p. 100.
Rud, Alexander D.; Kornienko, Nikolay E.; Kirian, Inna M.; Kirichenko,
Alexey N; Kucherov, O. P. (2018). “Local heteroallotropic structures of carbon”.
Materials Today: Proceedings. 5 (12): 26089–26095.
127
The Properties of Solutions
CHAPTER 8
THE PROPERTIES OF SOLUTIONS
OBJECTIVES
After studying this chapter, you will be able to:
1. Understand the measures of concentration
2. Define solubility
3. Describe the colligative properties
INTRODUCTION
Solutions are likely to have properties similar to those of their major component—
usually the solvent. However, some solution properties differ significantly from
those of the solvent. A solution possesses following properties –
•
•
•
•
•
•
A solution is a homogeneous mixture.
The constituent particles of a solution are smaller than 10-9 meter in
diameter.
Constituent particles of a solution cannot be seen by naked eyes.
Solutions do not scatter a beam of light passing through it. So, path of the
light beam is not visible in solutions.
Solute particles cannot be separated by filtration.
Solute or solvent particles do not settle down when left undisturbed.
128
University Chemistry
•
Solutions are stable at given temperature.
8.1 MEASURES OF CONCENTRATION
The concentration of a solution is a measure of the
amount of solute that has been dissolved in a given
amount of solvent or solution. A concentrated solution is
one that has a relatively large amount of dissolved solute.
A dilute solution is one that has a relatively small amount
of dissolved solute.
8.1.1 Molarity
Molarity is defined as the moles of a solute per
volume of total solution.
Concentration of a solution is often measured in
molarity (M), which is the number of moles of solute per
liter of solution. This molar concentration (ci) is calculated
by dividing the moles of solute (ni ) by the total volume
(V) of the:
ci =
Important
Molarity was
very useful for identifying the number
of solute particles
in a solution as you
could only measure
the total solution
(either its mass or
volume), but could
not directly measure
its components.
ni
V
The SI unit for molar concentration is mol/m3. However, mol/L is a more
common unit for molarity. A solution that contains 1 mole of solute per 1 liter of
solution (1 mol/L) is called “one Molar” or 1 M. The unit mol/L can be converted to
mol/m3 using the following equation:
1 mol/L = 1 mol/dm3 = 1 mol dm−3 = 1 M = 1000 mol/m3
Calculating Molarity
To calculate the molarity of a solution, the number of moles of solute must be
divided by the total liters of solution produced. If the amount of solute is given in
grams, we must first calculate the number of moles of solute using the solute’s molar
mass, then calculate the molarity using the number of moles and total volume.
Calculating Molarity Given Moles and Volume
If there are 10.0 grams of NaCl (the solute) dissolved in water (the solvent) to
produce 2.0 L of solution, what is the molarity of this solution?
First, we must convert the mass of NaCl in grams into moles. We do this by
dividing by the molecular weight of NaCl (58.4 g/mole).
The Properties of Solutions
129
Then, we divide the number of moles by the total solution volume to get
concentration.
ci =
ni
V
ci =
0.17 moles NaCl
2 liters solution
ci=0.1 M
The NaCl solution is a 0.1 M solution.
Calculating Moles Given Molarity
To calculate the number of moles in a solution given the molarity, we multiply
the molarity by total volume of the solution in liters.
How many moles of potassium chloride (KCl) are in 4.0 L of a 0.65 M solution?
ci =
ni
V
0.65 M =
ni
4.0 L
ni=(0.65 M)(4.0 L)=2.6 moles KCl
There are 2.6 moles of KCl in a 0.65 M solution that occupies 4.0 L.
Calculating Volume Given Molarity and Moles
We can also calculate the volume required to meet a specific mass in grams
given the molarity of the solution. This is useful with particular solutes that cannot
be easily massed with a balance. For example, diborane (B2H6) is a useful reactant
in organic synthesis, but is also highly toxic and flammable. Diborane is safer to use
and transport if dissolved in tetrahydrofuran (THF).
How many milliliters of a 3.0 M solution of BH3-THF are required to receive 4.0
g of BH3?
First we must convert grams of BH3 to moles by dividing the mass by the
molecular weight.
4.0 g BH
13.84g / mole BH
0.29 moles BH
Once we know we need to achieve 0.29 moles of BH3, we can use this and the
given molarity (3.0 M) to calculate the volume needed to reach 4.0 g.
ci =
ni
V
130
University Chemistry
3.0 M =
0.29moles BH 3
V
V=0.1L
Now that we know that there are 4.0 g of BH3 present in 0.1 L, we know that we
need 100 mL of solution to obtain 4.0 g of BH3.
8.1.2 Molality
Molality is a property of a solution that indicates the moles of solute per
kilogram of solvent.
Measurements of Mass (Molality) vs. Volume (Molarity)
Molality is an intensive property of solutions, and it is calculated as the moles
of a solute divided by the kilograms of the solvent. Unlike molarity, which depends
on the volume of the solution, molality depends only on the mass of the solvent.
Since volume is subject to variation due to temperature and pressure, molarity also
varies by temperature and pressure. In some cases, using weight is an advantage
because mass does not vary with ambient conditions. For example, molality is used
when working with a range of temperatures.
Defining Molality
The molality, b (or m), of a solution is defined as the amount of substance of
solute in moles, nsolute, divided by the mass in kg of the solvent, msolvent:
bM solute =
n solute
m solvent
Molality is based on mass, so it can easily be converted into a mass ratio,
denoted by w:
Compared to molar concentration or mass concentration, the preparation of
a solution of a given molality is easy because it requires only a good scale; both
solvent and solute are massed, rather than measured by volume. In many weak
aqueous solutions, the molarity and molality are similar because one kilogram of
water (the solvent) occupies one liter of volume at room temperature, and the small
amount of solute has little effect on the volume of the solvent.
Units
The SI unit for molality is mol/kg, or moles solute per kg of solvent. A solution
with a molality of 1 mol/kg is often described as “1 molal” or “1 m.” However,
following the SI system of units, the National Institute of Standards and Technology,
The Properties of Solutions
131
which is the United States’ authority on measurement, considers the term “molal”
and the unit symbol “m” to be obsolete, and suggests using mol/kg or another
related SI unit instead.
Calculating Molality
It is easy to calculate molality if we know the mass of solute and solvent in
a solution. Molality is an intensive property, and is therefore independent of the
amount being measured. This is true for all homogeneous solution concentrations,
regardless of if we examine a 1.0 L or 10.0 L sample of the same solution. The
concentration, or molality, remains constant.
Calculating Molality Given Mass
If we mass 5.36 g of KCl and dissolve this solid in 56 mL of water, what is the
molality of the solution? Remember that molality is moles of solute/kg per solvent.
KCl is our solute, while water is our solvent. We will first need to calculate the
amount of moles present in 5.36 g of KCl:
moles KCl =
5.36g × (
1 moles
)=
0.0719 moles KCl
74.5g
We also need to convert the the 56.0 mL of water to its equivalent mass in grams
by using the known density of water (1.0 g/mL):
56.0 mL × (
1.0g
)=
56.0 g
mL
56.0 g of water is equivalent to 0.056 kg of water. With this information, we can
divide the moles of solute by the kg of solvent to find the molality of the solution:
132
University Chemistry
moles
0.0719 moles KC
=
molality (=
) (
=
) 1.3 m
kg solvent
l0.056 kg water
The molality of our KCl and water solution is 1.3 m. Since the solution is very
dilute, the molality is almost identical to the molarity of the solution, which is 1.3
M.
Calculating Mass Given Molality
We can also use molality to find the amount of a substance in a solution. For
example, how much acetic acid, in mL, is needed to make a 3.0 m solution containing
25.0 g of KCN?
First, we must convert the sample of KCN from grams to moles:
moles KCN =
25.0g × (
1 moles
)=
0.38 moles
65.1g
The moles of KCN can then be used to find the kg of acetic acid. We multiply
the moles by the reciprocal of the given molality (3.0 moles/kg) so that our units
appropriately cancel. The result is the desired mass of acetic acid that we need to
make our 3 m solution:
0.38 moles KCl × (
kg acetic acid
)=
0.12 kg acetic acid
3.0 moles KCl
Once we have the mass of acetic acid in kg, we convert from kg to grams: 0.12
kg is equal to 120 g. Next, we use the density of acetic acid (1.05 g/mL at 20 oC) to
convert to the requested volume in mL. We must multiply by the reciprocal of the
density to accomplish this:
120.0 g acetic acid × (
mL
)=
114.0 mL acetic acid
1.05g
Therefore, we require 114 mL of acetic acid to make a 3.0 m solution that
contains 25.0 g of KCN.
8.2 SOLUBILITY
A solution is a homogeneous mixture of one or more solutes in a solvent.
Sugar cubes added to a cup of tea or coffee is a common example of a solution. The
property which helps sugar molecules to dissolve is known as solubility. Hence, the
term solubility can be defined as a property of a substance (solute) to dissolve in a
given solvent. A solute is any substance which can be either solid or liquid or gas
dissolved in a solvent.
The Properties of Solutions
133
8.2.1 Solubility Product
The term solubility product is generally applicable for sparingly soluble salts. It
is the maximum product of the molar concentration of the ions (raised to their
appropriate powers) which are produced due to dissociation of the compound.
At a given temperature the solubility product is constant. Lesser the value of
solubility product indicates lower solubility and higher value of solubility product
indicates greater solubility.
On the basis of solubility, the factors affecting solubility vary on the state of the
solute:
•
•
•
Liquids In Liquids
Solids In Liquids
Gases In Liquids
1. Solubility of Liquids in Liquids
Water is known as a universal solvent as it dissolves almost every solute except
for a few. Certain factors can influence the solubility of a substance.
Solubility is the new bond formation between the solute molecules and solvent
molecules. In terms of quantity, solubility is the maximum concentration of solute
that dissolves in a known concentration of solvent at a given temperature. Based on
the concentration of solute dissolves in a solvent, solutes are categorized into highly
soluble, sparingly soluble or insoluble. If a concentration of 0.1 g or more of a solute
can be dissolved in a 100ml solvent, it is said to be soluble. While a concentration
below 0.1 g is dissolved in the solvent it is said to be sparingly soluble. Thus, it is
said that solubility is a quantitative expression and expressed by the unit gram/litre
(g/L).
Based on solubility, different types of solution can be obtained. A saturated
solution is a solution where a given amount of solute is completely soluble in a
solvent at a given temperature. On the other hand, a supersaturated solution is
those where solute starts salting out or precipitate after a particular concentration
is dissolved at the same temperature.
Factors Affecting Solubility:
The solubility of a substance depends on the physical and chemical properties of
that substance. In addition to this, there are a few conditions which can manipulate
it. Temperature, pressure and the type of bond and forces between the particles are
few among them.
•
Temperature: By changing the temperature we can increase the soluble
property of a solute. Generally, water dissolves solutes at 20° C or 100° C.
Sparingly soluble solid or liquid substances can be dissolved completely
by increasing the temperature. But in the case of gaseous substance,
134
University Chemistry
•
•
temperature inversely influences solubility i.e.
as the temperature increases gases expand and
escapes from their solvent.
Forces and Bonds: Like dissolves in like. The
type of intermolecular forces and bonds vary
among each molecule. The chances of solubility
between two unlike substances are more
challengeable than the like substances. For
example, water is a polar solvent where a polar
solute like ethanol is easily soluble.
Pressure: Gaseous substances are much
influenced than solids and liquids by pressure.
When the partial pressure of gas increases, the
chance of its solubility is also increased. A soda
bottle is an example of where CO2 is bottled
under high pressure.
2. Solubility of Solids In Liquids
Hints
A state of dynamic equilibrium is
established between
these two processes
and at this point, the
number of solute
molecules entering
the solution becomes equal to the
number of particles
leaving the solution.
The concentration
of the solute in the
solution will remain
constant at a given
temperature and
pressure.
It has been observed that solid solubility depends
on the nature of the solute as well as the solvent. We
often see that substances like sugar, common salt (NaCl), etc. readily dissolve in
water while substances like naphthalene do not dissolve in water. From the various
observations and experimental results, it has been seen that only polar solutes tend
to dissolve in the polar solvent and non-polar solvents dissolve only non-polar
solutes. Hence, the nature of the solvent can be seen as one of the prominent factors
affecting solubility. The above observation led to the statement that like dissolves
like, that is polar solvents will dissolve polar solutes and non-polar solvents dissolve
non-polar solutes.
Now let us understand the process by which a solid dissolves in a solvent. Once
a solid solute is added to a solvent, the solute particles dissolve in the solvent and
this process is known as dissolution. Solute particles in the solution collide with
each other and some of these particles get separated out of the solution, this process
is called crystallization.
A solution in which no more solute can dissolve in the solvent at a given
temperature and pressure is said to be a saturated solution as the solution contains
the maximum amount of solute. The concentration of solute in such a solution is
called its solubility at that temperature and pressure. If more solute can be added
to a solution then it is called an unsaturated solution.
The Properties of Solutions
135
Factors Affecting Solubility:
•
•
Effect of Temperature: Apart from the nature of solute and solvent,
temperature also affects solid solubility considerably. If the dissolution
process is endothermic then the solubility should increase with an
increase in temperature in accordance with Le Chateliers Principle. If the
dissolution process is exothermic the solid solubility should decrease.
Effect of Pressure: Solid solubility hardly gets affected by changes
in pressure. This is due to the fact that solids and liquids are highly
incompressible and practically do not get affected by changes in pressure.
3. Solubility of Gases in Liquids
Gas solubility in liquids deals with the concept of gas dissolving in a solvent. Let
us first define solubility. For any substance, solubility is the maximum amount of
solute that can be dissolved in a given solvent at a particular temperature. Now our
concern is gas solubility in liquids. The gas solubility in liquids is greatly affected
by temperature and pressure as well as the nature of the solute and the solvent.
There are many gases that readily dissolve in water, while there are gases that
do not dissolve in water under normal conditions. Oxygen is only sparingly soluble
in water while HCl or ammonia readily dissolves in water.
Factors Affecting Solubility:
•
Effect of Pressure: It has been found that the gas solubility in liquids
increases with increase in pressure. To have a better understanding of the
effect of pressure on gas solubility let us consider a system of a gas solution
in a solvent in a closed container in a state of dynamic equilibrium. Now
the solution is in equilibrium and hence the rate of gaseous molecules
entering the solution is equal to the rate of gaseous molecules leaving the
solution.
Now suppose we increase the pressure of the system by compressing the
gas molecules present in the solution. As a result of an increase in pressure, the
gases molecules will now be concentrated in a smaller volume. This will result
in an increase in the number of gas molecules per unit volume available above
136
University Chemistry
the solution. Since the number of gas molecules presents above the solution has
increased, the rate with which the gas molecules will be entering the solution will
also increase. The end result is an increase in the number of gas molecules in the
solution until a new equilibrium point is attained. Thus the solubility of gases
increases with an increase in the pressure of a gas above the solution.
•
The solubility of gases in liquids: Henry’s Law gives a quantitative relation
between pressure and gas solubility in a liquid. It states that:
The solubility of a gas in a liquid is directly proportional to the partial pressure
of the gas present above the surface of liquid or solution.
The most general way of using Henry’s Law is that the partial pressure of a gas
above a solution is proportional to the mole fraction of the gas in the solution.
P = K Hx
Where, p = partial pressure of the gas
x = mole fraction of the gas in solution
KH = Henry’s law constant
•
Effect of Temperature: Gas solubility in liquids is found to decrease with
increase in temperature. The gas molecules in a liquid are dissolved by
the process of dissolution. During this process, heat is evolved. According
to Le Chatelier’s Principle which states that when the equilibrium of a
system is disturbed, the system readjusts itself in such a way that the effect
that has caused the change in equilibrium is countered. So, as we know
that dissolution is an exothermic process, the solubility should decrease
with an increase in temperature to validate Le Chatelier’s Principle.
8.2.2 Effects of Temperature and Pressure on Solubility
Experimentally it is found that the solubility of most compounds depends
strongly on temperature and, if a gas, on pressure as well. The ability to manipulate
the solubility by changing the temperature and pressure has several important
consequences.
Effect of Temperature on the Solubility of Solids
Figure 1 shows plots of the solubilities of several organic and inorganic
compounds in water as a function of temperature. Although the solubility of a solid
generally increases with increasing temperature, there is no simple relationship
between the structure of a substance and the temperature dependence of its
solubility. Many compounds (such as glucose and CH3CO2Na exhibit a dramatic
increase in solubility with increasing temperature. Others (such as NaCl and K2SO4)
exhibit little variation, and still others (such as Li2SO4) become less soluble with
increasing temperature.
The Properties of Solutions
137
Figure 1. Solubilities of Several Inorganic and Organic Solids in Water as a Function of
Temperature. Solubility may increase or decrease with temperature; the magnitude of
this temperature dependence varies widely among compounds.
Notice in particular the curves for NH4NO3 and CaCl2. The dissolution
of ammonium nitrate in water is endothermic (ΔHsoln=+25.7kJ/mol), whereas
the dissolution of calcium chloride is exothermic (ΔHsoln=−68.2kJ/mol), yet
Figure 1 shows that the solubility of both compounds increases sharply with
increasing temperature. In fact, the magnitudes of the changes in both enthalpy
and entropy for dissolution are temperature dependent. Because the solubility of a
compound is ultimately determined by relatively small differences between large
numbers, there is generally no good way to predict how the solubility will vary
with temperature.
Chemists are often able to use this information to separate the components of
a mixture by fractional crystallization, the separation of compounds on the basis of
their solubilities in a given solvent. For example, if we have a mixture of 150 g of
sodium acetate (CH3CO2Na) and 50 g of KBr, we can separate the two compounds
by dissolving the mixture in 100 g of water at 80°C and then cooling the solution
slowly to 0°C. According to the temperature curves in Figure 1, both compounds
dissolve in water at 80°C, and all 50 g of KBr remains in solution at 0°C. Only about
36 g of CH3CO2Na are soluble in 100 g of water at 0°C, however, so approximately
114 g (150 g − 36 g) of CH3CO2Na crystallizes out on cooling. The crystals can then
be separated by filtration. Thus fractional crystallization allows us to recover about
75% of the original CH3CO2Na in essentially pure form in only one step.
Fractional crystallization is a common technique for purifying compounds
as diverse as those shown in Figure 1 and from antibiotics to enzymes. For the
technique to work properly, the compound of interest must be more soluble at high
temperature than at low temperature, so that lowering the temperature causes it
to crystallize out of solution. In addition, the impurities must be more soluble than
138
University Chemistry
the compound of interest (as was KBr in this example) and preferably present in
relatively small amounts.
8.2.3 Effect of Temperature on the Solubility of Gases
The solubility of gases in liquids decreases with increasing temperature,
as shown in Figure 2. Attractive intermolecular interactions in the gas phase are
essentially zero for most substances. When a gas dissolves, it does so because its
molecules interact with solvent molecules. Because heat is released when these
new attractive interactions form, dissolving most gases in liquids is an exothermic
process (ΔHsoln<0). Conversely, adding heat to the solution provides thermal energy
that overcomes the attractive forces between the gas and the solvent molecules,
thereby decreasing the solubility of the gas. The phenomenon is similar to that
involved in the increase in the vapor pressure of a pure liquid with increasing
temperature. In the case of vapor pressure, however, it is attractive forces between
solvent molecules that are being overcome by the added thermal energy when the
temperature is increased.
Figure 2. Solubilities of Several Common Gases in Water as a Function of Temperature
at Partial Pressure of 1 atm. The solubilities of all gases decrease with increasing temperature.
The decrease in the solubilities of gases at higher temperatures has both practical
and environmental implications. Anyone who routinely boils water in a teapot or
electric kettle knows that a white or gray deposit builds up on the inside and must
eventually be removed. The same phenomenon occurs on a much larger scale in the
giant boilers used to supply hot water or steam for industrial applications, where
it is called “boiler scale,” a deposit that can seriously decrease the capacity of hot
water pipes (Figure 3). The problem is not a uniquely modern one: aqueducts that
The Properties of Solutions
139
were built by the Romans 2000 years ago to carry cold water from alpine regions
to warmer, drier regions in southern France were clogged by similar deposits. The
chemistry behind the formation of these deposits is moderately complex and will
be described elsewhere, but the driving force is the loss of dissolved CO2 from
solution. Hard water contains dissolved Ca2+ and HCO −3 (bicarbonate) ions. Calcium
bicarbonate (Ca(HCO3)2 is rather soluble in water, but calcium carbonate (CaCO3)
is quite insoluble. A solution of bicarbonate ions can react to form carbon dioxide,
carbonate ion, and water:
2HCO −3 (aq ) → CO 23 − (aq ) + H 2 O(l) + CO 2 (aq )
Heating the solution decreases the solubility of CO2, which escapes into the
gas phase above the solution. In the presence of calcium ions, the carbonate ions
precipitate as insoluble calcium carbonate, the major component of boiler scale.
Figure 3. Calcium carbonate deposits (left) Calcium carbonate (CaCO3) deposits in hot
water pipes can significantly reduce pipe capacity. These deposits, called boiler scale,
form when dissolved CO2 is driven into the gas phase at high temperatures. (right)
Highly calcified remains of Eiffel aqueduct near Euskirchen-Kreuzweingarten, Germany.
Effect of Pressure on the Solubility of Gases: Henry’s Law
External pressure has very little effect on the solubility of liquids and solids.
In contrast, the solubility of gases increases as the partial pressure of the gas above
a solution increases. This point is illustrated in Figure 4, which shows the effect
of increased pressure on the dynamic equilibrium that is established between the
dissolved gas molecules in solution and the molecules in the gas phase above the
solution. Because the concentration of molecules in the gas phase increases with
increasing pressure, the concentration of dissolved gas molecules in the solution at
equilibrium is also higher at higher pressures.
140
University Chemistry
Figure 4. A Model Depicting Why the Solubility of a Gas Increases as the Partial Pressure Increases at Constant Temperature. (a) When a gas comes in contact with a pure
liquid, some of the gas molecules (purple spheres) collide with the surface of the liquid
and dissolve. When the concentration of dissolved gas molecules has increased so
that the rate at which gas molecules escape into the gas phase is the same as the rate
at which they dissolve, a dynamic equilibrium has been established, as depicted here.
This equilibrium is entirely analogous to the one that maintains the vapor pressure of
a liquid. (b) Increasing the pressure of the gas increases the number of molecules of gas
per unit volume, which increases the rate at which gas molecules collide with the surface of the liquid and dissolve. (c) As additional gas molecules dissolve at the higher
pressure, the concentration of dissolved gas increases until a new dynamic equilibrium
is established.
The relationship between pressure and the solubility of a gas is described
quantitatively by Henry’s law, which is named for its discoverer, the English
physician and chemist, William Henry (1775–1836):
C=kP
where
C is the concentration of dissolved gas at equilibrium,
P is the partial pressure of the gas, and
k is the Henry’s law constant, which must be determined experimentally for
each combination of gas, solvent, and temperature.
Although the gas concentration may be expressed in any convenient units, we
will use molarity exclusively. The units of the Henry’s law constant are therefore
mol/(L·atm) = M/atm. Values of the Henry’s law constants for solutions of several
gases in water at 20°C are listed in Table 1.
As the data in Table 1 demonstrate, the concentration of a dissolved gas in
water at a given pressure depends strongly on its physical properties. For a series of
related substances, London dispersion forces increase as molecular mass increases.
Thus among the Group 18 elements, the Henry’s law constants increase smoothly
from He to Ne to Ar.
141
The Properties of Solutions
Table 1. Henry’s Law Constants for Selected Gases in Water at 20°C
Gas
Henry’s Law Constant [mol/(L·atm)] × 10−4
He
3.9
Ne
4.7
Ar
15
H2
8.1
N2
7.1
O2
14
CO2
392
Gases that react chemically with water, such as HCl and the other hydrogen
halides, H2S, and NH3, do not obey Henry’s law; all of these gases are much more
soluble than predicted by Henry’s law. For example, HCl reacts with water to
give H+(aq) and Cl−(aq), not dissolved HCl molecules, and its dissociation into ions
results in a much higher solubility than expected for a neutral molecule.
Henry’s law has important applications. For example, bubbles of CO2 form as
soon as a carbonated beverage is opened because the drink was bottled under CO2 at
a pressure greater than 1 atm. When the bottle is opened, the pressure of CO2 above
the solution drops rapidly, and some of the dissolved gas escapes from the solution
as bubbles. Henry’s law also explains why scuba divers have to be careful to ascend
to the surface slowly after a dive if they are breathing compressed air. At the higher
pressures under water, more N2 from the air dissolves in the diver’s internal fluids.
If the diver ascends too quickly, the rapid pressure change causes small bubbles
of N2 to form throughout the body, a condition known as “the bends.” These bubbles
can block the flow of blood through the small blood vessels, causing great pain and
even proving fatal in some cases.
Due to the low Henry’s law constant for O2 in water, the levels of dissolved
oxygen in water are too low to support the energy needs of multicellular organisms,
including humans. To increase the O2 concentration in internal fluids, organisms
synthesize highly soluble carrier molecules that bind O2 reversibly. For example,
human red blood cells contain a protein called hemoglobin that specifically
binds O2 and facilitates its transport from the lungs to the tissues, where it is used to
oxidize food molecules to provide energy. The concentration of hemoglobin in normal
blood is about 2.2 mM, and each hemoglobin molecule can bind four O2 molecules.
Although the concentration of dissolved O2 in blood serum at 37°C (normal body
temperature) is only 0.010 mM, the total dissolved O2 concentration is 8.8 mM,
almost a thousand times greater than would be possible without hemoglobin.
Synthetic oxygen carriers based on fluorinated alkanes have been developed for
use as an emergency replacement for whole blood. Unlike donated blood, these
142
University Chemistry
“blood substitutes” do not require refrigeration and have
a long shelf life. Their very high Henry’s law constants
for O2 result in dissolved oxygen concentrations
comparable to those in normal blood.
8.3 COLLIGATIVE PROPERTIES
Important
The word
“colligative” has
been adapted or
taken from the Latin
word “colligatus”
which translates to
“bound together”.
A colligative property is a property of a solution
that is dependent on the ratio between the total number
of solute particles (in the solution) to the total number
of solvent particles. Colligative properties are not
dependent on the chemical nature of the solution’s
components. Thus, colligative properties can be linked to several quantities that
express the concentration of a solution, such as molarity, normality, and molality.
The four colligative properties that can be exhibited by a solution are:
•
•
•
•
Boiling point elevation
Freezing point depression
Relative lowering of vapor pressure
Osmotic pressure
8.3.1 Vapor Pressure Depression
Physical properties can be divided into two categories. Extensive properties (such
as mass and volume) depend on the size of the sample. Intensive properties (such as
density and concentration) are characteristic properties of the substance; they do
not depend on the size of the sample being studied. This third category, known
as colligative properties, can only be applied to solutions. By definition, one of the
properties of a solution is a colligative property if it depends only on the ratio of the
number of particles of solute and solvent in the solution, not the identity of the
solute.
Very few of the physical properties of a solution are colligative properties. As
an example of this limited set of physical properties, let’s consider what happens
to the vapor pressure of the solvent when we add a solute to form a solution. We’ll
define Po as the vapor pressure of the pure liquid
the solvent
and P as the
vapor pressure of the solvent after a solute has been added.
Po = vapor pressure of the pure liquid, or solvent
P = vapor pressure of the solvent in a solution
When the temperature of a liquid is below its boiling point, we can assume that
the only molecules that can escape from the liquid to form a gas are those that lie
near the surface of the liquid.
The Properties of Solutions
143
When a solute is added to the solvent, some of the solute molecules occupy the
space near the surface of the liquid, as shown in the figure below. When a solute is
dissolved in a solvent, the number of solvent molecules near the surface decreases,
and the vapor pressure of the solvent decreases.
This has no effect on the rate at which solvent molecules in the gas phase
condense to form a liquid. But it decreases the rate at which the solvent molecules
in the liquid can escape into the gas phase. As a result, the vapor pressure of the
solvent escaping from a solution should be smaller than the vapor pressure of the
pure solvent.
Between 1887 and 1888, Francois-Marie Raoult showed that the vapor pressure
of a solution is equal to the mole fraction of the solvent times the vapor pressure of
the pure liquid.
This equation, which is known as Raoult’s law, is easy to understand. When the
solvent is pure, and the mole fraction of the solvent is equal to 1, P is equal to Po. As
the mole fraction of the solvent becomes smaller, the vapor pressure of the solvent
escaping from the solution also becomes smaller.
Let’s assume, for the moment, that the solvent is the only component of the
solution that is volatile enough to have a measurable vapor pressure. If this is true,
the vapor pressure of the solution will be equal to the vapor pressure of the solvent
144
University Chemistry
escaping from the solution. Raoult’s law suggests that the difference between the
vapor pressure of the pure solvent and the solution increases as the mole fraction
of the solvent decreases.
The change in the vapor pressure that occurs when a solute is added to a solvent
is therefore a colligative property. If it depends on the mole fraction of the solute,
then it must depend on the ratio of the number of particles of solute to solvent in
the solution but not the identity of the solute.
8.3.2 Boiling Point Elevation and Freezing Point Depression
The figure below shows the consequences of the fact that solutes lower the
vapor pressure of a solvent. The solid line connecting points B and C in this phase
diagram contains the combinations of temperature and pressure at which the pure
solvent and its vapor are in equilibrium. Each point on this line therefore describes
the vapor pressure of the pure solvent at that temperature. The dotted line in this
figure describes the properties of a solution obtained by dissolving a solute in the
solvent. At any given temperature, the vapor pressure of the solvent escaping from
the solution is smaller than the vapor pressure of the pure solvent. The dotted line
therefore lies below the solid line.
The decrease in the vapor pressure of the solvent that occurs when a solute is added to
the solvent causes an increase in the boiling point and decrease in the melting point of
the solution.
According to this figure, the solution can’t boil at the same temperature as
the pure solvent. If the vapor pressure of the solvent escaping from the solution
is smaller than the vapor pressure of the pure solvent at any given temperature,
the solution must be heated to a higher temperature before it boils. The lowering
of the vapor pressure of the solvent that occurs when it is used to form a solution
therefore increases the boiling point of the liquid.
The Properties of Solutions
145
When phase diagrams were introduced, the triple point was defined as the only
combination of temperature and pressure at which the gas, liquid, and solid can
exist at the same time. The figure above shows that the triple point of the solution
occurs at a lower temperature than the triple point of the pure solvent. By itself,
the change in the triple point is not important. But it results in a change in the
temperature at which the solution freezes or melts. To understand why, we have
to look carefully at the line that separates the solid and liquid regions in the phase
diagram. This line is almost vertical because the melting point of a substance is not
very sensitive to pressure.
Adding a solute to a solvent doesn’t change the way the melting point depends
on pressure. The line that separates the solid and liquid regions of the solution is
therefore parallel to the line that serves the same function for the pure solvent. This
line must pass through the triple point for the solution, however. The decrease in the
triple point that occurs when a solute is dissolved in a solvent therefore decreases
the melting point of the solution.
The figure above shows how the change in vapor pressure that occurs when a
solute dissolves in a solvent leads to changes in the melting point and the boiling
point of the solvent as well. Because the change in vapor pressure is a colligative
property, which depends only on the relative number of solute and solvent
particles, the changes in the boiling point and the melting point of the solvent are
also colligative properties.
8.3.3 Colligative Properties Calculations
The best way to demonstrate the importance of colligative properties is to
examine the consequences of Raoult’s law. Raoult found that the vapor pressure
of the solvent escaping from a solution is proportional to the mole fraction of the
solvent.
P = Xsolvent Po
But the vapor pressure of a solvent is not a colligative property. Only the change
in the vapor pressure that occurs when a solute is added to the solvent can be included
among the colligative properties of a solution.
Because pressure is a state function, the change in the vapor pressure of the
solvent that occurs when a solute is added to the solvent can be defined as the
difference between the vapor pressure of the pure solvent and the vapor pressure
of the solvent escaping from the solution.
P = Po - P
Substituting Raoult’s law into this equation gives the following result.
P = Po - Xsolvent Po = (1 - Xsolvent) Po
146
University Chemistry
This equation can be simplified by remembering the relationship between the
mole fraction of the solute and the mole fraction of the solvent.
Xsolute + Xsolvent = 1
Substituting this relationship into the equation that defines P gives another
form of Raoult’s law.
P = Xsolute Po
This equation reminds us that the change in the vapor pressure of the solvent
that occurs when a solute is added to the solvent is proportional to the mole fraction
of the solute. As more solute is dissolved in the solvent, the vapor pressure of the
solvent decreases, and the change in the vapor pressure of the solvent increases.
Because changes in the boiling point of the solvent (∆TBP) that occur when a
solute is added to a solvent result from changes in the vapor pressure of the solvent,
the magnitude of the change in the boiling point is also proportional to the mole
fraction of the solute.
∆TBP = kb solute
In dilute solutions, the mole fraction of the solute is proportional to the molality
of the solution, as shown in the figure below.
The equation that describes the magnitude of the boiling point elevation that
occurs when a solute is added to a solvent is therefore often written as follows.
∆TBP = kbm
Here, ∆ TBP is the boiling point elevation -- the change in the boiling point that
occurs when a solute dissolves in the solvent
and kb is a proportionality constant
known as the molal boiling point elevation constant for the solvent.
A similar equation can be written to describe what happens to the freezing
point (or melting point) of a solvent when a solute is added to the solvent.
∆TFP = -kf m
147
The Properties of Solutions
In this equation, ∆TFP is the freezing point depression
the change in the
freezing point that occurs when the solute dissolves in the solvent -- and kf is
the molal freezing point depression constant for the solvent. A negative sign is used
in this equation to indicate that the freezing point of the solvent decreases when a
solute is added.
Values of kf and kb as well as the freezing points and boiling points for a number
of pure solvents are given in the tables below.
Freezing Point Depression Constants:
Compound
Freezing Point (oC)
water
0
acetic acid
16.66
3.90
benzene
5.53
5.12
p-xylene
13.26
4.3
naphthalene
80.29
6.94
cyclohexane
6.54
20.0
carbon tetrachloride
-22.95
29.8
camphor
178.75
37.7
kf (oC/m)
1.853
Boiling Point Elevation Constants:
Compound
Boiling Point (oC)
water
100
ethyl ether
34.55
1.824
carbon disulfide
46.23
2.35
benzene
80.10
2.53
carbon tetrachloride
76.75
4.48
camphor
207.42
5.611
kb (oC/m)
0.515
8.3.4 Osmotic Pressure
In 1784, the French physicist and clergyman Jean Antoine Nollet discovered
that a pig’s bladder filled with a concentrated solution of alcohol in water expanded
when it was immersed in water. The bladder acted as a semipermeable membrane,
which allowed water molecules to enter the solution, but kept alcohol molecules
from moving in the other direction. Movement of one component of a solution
through a membrane to dilute the solution is called osmosis, and the pressure this
produces is called the osmotic pressure (π).
Osmotic pressure can be demonstrated with the apparatus shown in the figure
below. A semipermeable membrane is tied across the open end of a thistle tube.
The tube is then partially filled with a solution of sugar or alcohol in water and
immersed in a beaker of water. Water will flow into the tube until the pressure
148
University Chemistry
on the column of water due to the force of gravity balances the osmotic pressure
driving water through the membrane.
Water flows through the semipermeable membrane to dilute the alcohol solution until
the force of gravity pulling down on the column of this solution balances the osmotic
pressure pushing the water through the membrane.
The same year that Raoult discovered the relationship
between the vapor pressure of a solution and the vapor
pressure of a pure solvent, Jacobus Henricus van’t Hoff
found that the osmotic pressure of a dilute solution ( )
obeyed an equation analogous to the ideal gas equation.
p=
nRT
V
This equation suggests that osmotic pressure is
another example of a colligative property, because this
pressure depends on the ratio of the number of solute
particles to the volume of the solution
n/V
not
the identity of the solute particles. It also reminds us of
the magnitude of osmotic pressure. According to this
equation, a 1.00 M solution has an osmotic pressure of
22.4 atm at 0oC.
p
(1.00 mol)(0.08206 L atm / mol K) (273 K)
= 2.24 atm
(1.00 L)
This means that a 1.00 M solution should be able to
support a column of water 670 inches, or almost 56 feet,
tall!
Hints
The concentration of an
isotonic sodium
chloride (NaCl)
solution is only half
that of an isotonic
glucose (C6H12O6)
solution because
NaCl produces two
ions when a formula unit dissolves,
while molecular
C6H12O6 produces
only one particle
when a formula
unit dissolves. The
osmolarities are
therefore the same
even though the
concentrations of
the two solutions
are different.
Biologists and biochemists often take advantage of
osmotic pressure when they isolate the components of
a cell. When a cell is added to an aqueous solution that contains a much higher
concentration of ions than the liquid within the cell, water leaves the cell by flowing
through the cell membrane until the cell shrinks so much that the membrane breaks.
Alternatively, when a cell is placed in a solution that has a much smaller ionic
149
The Properties of Solutions
strength, water pours into the cell, and the cell expands until the cell membrane
bursts.
EXERCISE
Answer the following questions:
1. What is Solution?
2. What are the colligative properties of solutions?
3. Explain how the following properties of solutions differ from those of the
pure solvent: vapor pressure, boiling point, freezing point, and osmotic
pressure.
4. Why are Colligative properties important?
5. Why is Molality used in Colligative properties?
6. Which solution’s freezing point deviates more from that of pure water—a
1 M solution of NaCl or a 1 M solution of CaCl2?
7. A 0.50 M NaCl aqueous solution and a 0.30 M Ca(NO3)2 aqueous solution
are placed on opposite sides of a semipermeable membrane. Determine
the osmolarity of each solution and predict the direction of solvent flow.
MULTIPLE CHOICE QUESTIONS
Tick the correct answer.
1.
The total volume of the solution may not be equal to the sum of volumes of
solute and solvent. This is because.....
a. volume depends on temperature
b. solute particles may occupy empty space structure of liquids
c. solute particles are larger in size than solvent
d. all of these
2.
6g of urea was dissolved in 500 g of water. The percentage by mass of urea
in solution is...
a. 0.118 %
b. 1.18 %
c. 2.01 %
d. 1.45 %
3.
Which of the following units is useful in relating the concentration of a
solution with its vapor pressure?
a. mole fraction
b. parts per million
c. mass percentage
d. molality
150
4.
5.
University Chemistry
Low concentration of oxygen in the blood and tissues of people living at
high altitude is due to ……………...
a. low temperature
b. low atmospheric pressure
c. high atmospheric pressure
d. both low temperature and high atmospheric pressure
Colligative properties depend on ……………...
a. the nature of the solute particles dissolved in solution.
b. the number of solute particles in solution.
c. the physical properties of the solute particles dissolved in solution.
d. the nature of solvent particles.
ANSWER
1. (b)
2. (b)
3. (a)
4. (b)
5. (b)
REFERENCES
1.
2.
3.
4.
5.
6.
7.
8.
Herbst, Eric (12 May 2005). “Chemistry of Star-Forming Regions”. Journal of
Physical Chemistry A. 109 (18): 4017–4029.
Hill, J.W.; Petrucci, R.H.; McCreary, T.W.; Perry, S.S. (2005). General Chemistry
(4th ed.). Upper Saddle River, New Jersey: Pearson Prentice Hall. p. 37.
McQuarrie, Donald, et al. Colligative properties of Solutions” General
Chemistry Mill Valley: Library of Congress, 2011.
Simpson, David (29 June 2005). “Lucretius (c. 99–55 BCE)”. The Internet History
of Philosophy. Retrieved 10 November 2020.
Strodach, George K. (2012). The Art of Happiness. New York: Penguin Classics.
pp. 7–8.
The Lewis Acid–Base Concept. Apsidium. 19 May 2003. Archived from the
original on 27 May 2008. Retrieved 31 July 2010.
Tro, Nivaldo J. (2018). Chemistry; Structure and Properties (Textbook.) (2nd
ed.). Pearson Education. pp. 563–566.
Tullo, Alexander H. (28 July 2014). “C&EN’s Global Top 50 Chemical Firms For
2014”. Chemical & Engineering News. American Chemical Society. Retrieved
22 August 2014.
151
Chemical Equilibrium
CHAPTER 9
CHEMICAL EQUILIBRIUM
OBJECTIVES
After studying this chapter, you will be able to:
1. Description of chemical equilibrium
2. Derivate the reactions at equilibrium
3. Define the equilibrium constant
4. Explain heterogeneous equilibria
5. Discuss about the equilibrium calculations
6. Explain the response of equilibria to the reaction conditions
Chemical equilibrium definition refers to the state of a system where the
concentration of the reactant and the concentration of the products do not change
with respect to time and the system does not display any further change in
properties.
Chemical equilibrium is said to be achieved by the system when the rate of the
forward reaction is equal to the rate of the reverse reaction. When there is no further
change in the concentrations of the reactants and the products due to the equal
rates of the forward and reverse reactions, at the time point of time the system is
said to be in a dynamic state of equilibrium.
152
University Chemistry
In a chemical reaction, chemical equilibrium is the state in which the forward
reaction rate and the reverse reaction rate are equal. The result of this equilibrium is
that the concentrations of the reactants and the products do not change. However,
just because concentrations aren’t changing does not mean that all chemical reaction
has ceased. Just the opposite is true; chemical equilibrium is a dynamic state in which
reactants are being converted into products at all times, but at the exact rate that
products are being converted back into reactants.
9.1 DESCRIPTION OF CHEMICAL EQUILIBRIUM
Chemical equilibrium is the thermodynamic equilibrium in a system where
direct and reverse chemical reactions are possible. If chemical equilibrium takes
place in the system, the rates of all reactions proceeding in two opposite directions
are equal. Therefore, the macroscopic parameters of the system do not change and
the relationship between concentrations of reacting substances remains constant
at a given temperature. Equilibrium for any chemical reaction is expressed by an
equality ∑νiμi = 0, where μi is the chemical potential of each reagent (i = 1,2, . . .)
and νi is the stoichiometric coefficient of each substance in an equation of chemical
reaction (it is positive for initial substances and negative for products of a reaction).
The dependence of chemical equilibrium on external conditions is expressed by
the Le Chatelier-Braun principle (1885-1886). It consists of the following correlation:
Let equilibrium take place and then influence the system, changing some external
conditions (temperature, pressure, concentrations of reacting substances). The
equilibrium of a reaction tends to follow such direction that allows the reduction
of an external influence. A temperature increase will cause a displacement of the
equilibrium to the direction of such reaction that proceeds with heat absorption.
A pressure increase will cause equilibrium displacement to follow the direction of
such reaction that leads to a volume decrease. The introduction of any additional
reagent in the system will propel equilibrium displacement to a direction where
this reagent is consumed.
The total Gibbs energy change of a chemical reaction aA + bB = cC + dD (when
temperature and pressure are constant) is expressed by the equation
∆ r G p ,T =
∆ r G 0p ,T + RT ln(α Cc α Dd α Aaα Bb ),
where R is the gas constant, p is the pressure, T is the absolute temperature,
αi refers to the activities of the reacting substances and ∆ r G 0p ,T is the standard Gibbs
energy›s change of that reaction (αi = 1). The value of ∆ r G 0p ,T can be calculated on the basis
of standard values of the Gibbs energies of formation (ΔfG0) of the reagents at 298.15 K
and of known thermodynamic relationships that determine the temperature and
pressure dependencies of Gibbs energy change.
153
Chemical Equilibrium
If equilibrium is attained, then
0
∆ r G p ,T =
− RT ln
(α Cequil )c (α Dequil ) d
(α Aequil ) a (α Bequil )b
=
− RT ln Kα
Here, αequil are the activities corresponding to the equilibrium state and
Ka is the equilibrium constant expressed in terms of activities. Hence, it follows
=
∆G
RT [ln(α Cc α Dd / α Aaα Bb ) − ln Kα ]
that r p ,T
. The last relationship is the van’t Hoff
isotherm equation (or van’t Hoff equation). It permits the determination of a probable
direction of the reaction under given conditions. The process will take place when
K > (α Cc α Dd / α Aaα Bb ) . Analogous relationships can be obtained
ΔrGP,T < 0, i.e., when α
when the equilibrium constant (Kp) is expressed in terms of partial pressures (Pi) of
the reagents:
 Pc Pd

∆ r G p ,T =
− RT ln K p .
RT ln Ca Db − ln K p  ; ∆ r G 0p ,T =
P
P


A B
The “equilibrium constant of reaction” is the result of the mass action law,
which determines a correlation between the masses of reacting substances under
equilibrium. According to this law, the reaction’s rate depends on the concentrations
of reacting substances. The rate constant of a given reaction at fixed temperature
is a constant value; therefore, the relationship of the rate
constants of direct and reverse reactions is a constant
value too. This relationship is a function of temperature
Important
only.
The equations that express a relationship between the
In a chemical equilibrium, the
forward and reverse
reactions occur at
equal rates, and
the concentrations
of products and
reactants remain
constant.
∆ r G 0p ,T -value and equilibrium constant of reaction allow
the calculation of the equilibrium of chemical reactions,
avoiding expensive and prolonged experiments. For
such calculations, it is necessary to have reliable values
of thermodynamic functions for all reacting substances.
Various experimental methods are used to determine
equilibrium constants of chemical reactions. There are
static and dynamic methods as well as the circulation
method. The last is a specific combination of the static and dynamic methods. When
static methods are used, the reaction mixture stays at a given temperature until an
adjustment of the equilibrium takes place. Then “tempering” and chemical analysis
of the reaction mixture are carried out. “Equilibrium tempering” is the fast-cooling
of the reaction mixture to a low temperature where the rate of reaction is very small.
The more common dynamic method of defining equilibrium constants has
154
University Chemistry
often been called the transportation method. A steady stream of inert gas is passed
over the mixture of substances that is maintained at a constant temperature. This
“carrier” gas removes the volatile components of the reaction at a rate that depends
on the rate of gas flow. The vapors of the reagents are condensed or collected by
absorption or chemical combination at the colder portion of the apparatus. The
experiments are carried out at different rates of gas flow. The equilibrium pressures
of volatile reagents are determined by extrapolation of the results up to zero rate of
the carrier gas.
A modification of the dynamic method used for investigating heterogeneous
equilibria is the circulation method. The gas mixture is circulated in a closed
space; circulation is carried out by means of electromagnetic pump. Equilibrium
is attained when passing this mixture many times over the solid phase into the
furnace. Tempering of the gas mixture is done when it is taken out from the hot
zone and passed through a capillary. In view of the large linear rate of gas flow, this
mixture becomes cold rapidly and its composition is not changed.
The most direct way of measuring equilibrium constants of chemical reactions is
through the measurement of electromotive forces (the e.m.f. method). For example,
the reaction
Zn(cryst ) + CuSo4 ( solution) = ZnSo4 ( solution) + Cu (cryst.)
is a process of potential generation for the Daniel galvanic element:Zn0/Zn2+//
Cu /Cu2+ A zinc plate (one electrode) is immersed into a solution of zinc sulfate and
a copper plate (the other electrode) is immersed into a solution of copper sulfate.
A galvanic element (source of electromotive force) can be created if both electrodes
are connected by a tube that contains a solution-conductor. The dissolution of zinc
(process: Zn0 = Zn2+ + 2e) takes place at one electrode; the precipitation of copper
(process: Cu2+ + 2e = Cu0) takes place at the second electrode. Therefore, the common
potential forming reaction is: Zn0 + Cu2+ = Zn2+ + Cu0 The Gibbs energy change
0
for such reaction is given by the formula ∆ r GT0 =
−nFET , where n is the number
of gramme-equivalents of reagent; F is Faraday’s constant (nF is the number of
coulombs of electricity passed); and ET is the electromotive force of the galvanic
element at a given temperature. The value of the Gibbs energy of reaction can be
used for calculating its equilibrium constant (K): ln K = −∆ r GT0 / RT = nFET / RT .
The equilibrium state is a thermodynamic state of a system that is permanent
in time. This invariability is not connected with some external process taking
place. There are different kinds of equilibria. If the equilibrium is “steady,” then
any adjacent states of the system are less steady. It would be necessary to spend
external work for transition from the equilibrium state to these adjacent states.
It is also typical that steady equilibrium can be approached from two opposite
directions. However, this discussion is concerned with steady equilibria only or
“chemical equilibria.” From the physicist’s point of view, steady equilibrium is
Chemical Equilibrium
155
dynamic. It is attained when the rates of direct and reverse reactions are equal, but
not under conditions when the process is stopped in general. The equality dG = 0
is a general condition for “steady” and “unsteady” equilibria, but the value of the
second differential of Gibbs energy is positive under steady equilibrium (d2G > 0)
and negative under unsteady equilibrium (d2G < 0). The conditions of stability of
the equilibrium can be deduced using the second law of thermodynamics. These are:
1) the pressure increases at a constant temperature if volume decreases [(dP/dV)T <
0]; and 2) the value of heat capacity is positive (Cp > 0).
The degree of stability of the different states of chemical systems can vary. States
which possess some relative stability are called “metastable” states. Such states
have often arisen due to kinetic factors, which create difficulties for the transition of
a system from the metastable (unsteady) state to a steady equilibrium state.
9.2 REACTIONS AT EQUILIBRIUM
An equilibrium reaction is a chemical reaction between the reactants that
stays in a stable state before and after the completion of the reaction (i.e., in a
thermodynamic equilibrium state). A reaction is said to be in a thermodynamic
equilibrium state when it satisfies all three types of equilibrium:
•
•
•
Thermal equilibrium
Chemical equilibrium
Mechanical equilibrium
The product obtained from an equilibrium reaction also remains in equilibrium
with the reactants.
An equilibrium reaction abides by the so-called “minus first” or “zeroth” law of
thermodynamics. Per this law of thermodynamics, if the first object is in equilibrium
with the second object and the second object is in equilibrium with the third object,
then the first and third objects will also be in equilibrium. An equilibrium reaction
satisfies thermal, chemical and mechanical equilibrium states.
•
•
•
Thermal equilibrium is a state in which two substances or objects in
physical contact have no difference in their temperatures (i.e., either the
objects have the same temperature or both objects are connected by a
permeable barrier that does not allow the transfer of heat between the
two objects). Thermal equilibrium is mainly associated with the laws of
thermodynamics in physics and mechanics.
Chemical equilibrium is a state where there is no chemical reaction
occurring between the various objects, or any transfer of matter from one
part of the system to another part due to any kind of diffusion. Thus,
during a chemical equilibrium the chemical potential of the various
systems remains the same.
Mechanical equilibrium is a state where there are no unbalanced forces
acting within the system or between the system and its surroundings.
156
University Chemistry
Thus, the pressure throughout the system remains the same or constant.
Reactions don’t stop when they come to equilibrium. But the forward and
reverse reactions are in balance at equilibrium, so there is no net change in the
concentrations of the reactants or products, and the reaction appears to stop on
the macroscopic scale. Chemical equilibrium is an example of a dynamic balance
between opposing forces
the forward and reverse reactions
not a static
balance.
Let’s look at the logical consequences of the assumption that the reaction
between ClNO2 and NO eventually reaches equilibrium.
ClNO2(g) +
NO(g)
NO2(g)
+
ClNO(g)
The rates of the forward and reverse reactions are the same when this system
is at equilibrium.
At equilibrium:
rateforward
=
ratereverse
Substituting the rate laws for the forward and reverse reactions into this equality
gives the following result.
At equilibrium:
kf(ClNO2)(NO)
=
kr(NO2)(ClNO)
But this equation is only valid when the system is at equilibrium, so we should
replace the (ClNO2), (NO), (NO2), and (ClNO) terms with symbols that indicate
that the reaction is at equilibrium. By convention, we use square brackets for this
purpose. The equation describing the balance between the forward and reverse
reactions when the system is at equilibrium should therefore be written as follows.
At equilibrium:
kf[ClNO2][NO]
=
kr[NO2][ClNO]
Rearranging this equation gives the following result.
kf
kr
=
[ NO2 ][ClNO]
[ClNO2 ][ NO]
Since kf and kr are constants, the ratio of kf divided by kr must also be a
constant. This ratio is the equilibrium constant for the reaction, Kc. The ratio of the
concentrations of the reactants and products is known as the equilibrium constant
expression.
K=
c
k f [ NO2 ][ClNO]
=
kr [ClNO2 ][ NO]
No matter what combination of concentrations of reactants and products we
start with, the reaction will reach equilibrium when the ratio of the concentrations
157
Chemical Equilibrium
defined by the equilibrium constant expression is equal to the equilibrium constant
for the reaction. We can start with a lot of ClNO2 and very little NO, or a lot of NO
and very little ClNO2. It doesn’t matter. When the reaction reaches equilibrium, the
relationship between the concentrations of the reactants and products described by
the equilibrium constant expression will always be the same. At 25oC, this reaction
always reaches equilibrium when the ratio of these concentrations is 1.3 x 104.
K=
c
k f [ NO2 ][ClNO]
=
= 1.3 ×104
kr [ClNO2 ][ NO]
The procedure used in this section to derive the equilibrium constant expression
only works with reactions that occur in a single step, such as the transfer of a
chlorine atom from ClNO2 to NO. Many reactions take a number of steps to convert
reactants into products. But any reaction that reaches equilibrium, no matter how
simple or complex, has an equilibrium constant expression that satisfies the rules
in the following section.
9.2.1 Altering or Combining Equilibrium Reactions
What happens to the magnitude of the equilibrium constant for a reaction when
we turn the equation around? Consider the following reaction, for example.
ClNO2(g)
+
NO(g)
NO2(g)
+
ClNO(g)
The equilibrium constant expression for this equation is written as follows.
[ NO2 ][ClNO]
Kc =
=
1.3 ×104 (at 25°C )
[ClNO2 ][ NO]
Because this is a reversible reaction, it can also be represented by an equation
written in the opposite direction.
NO2(g)
+
ClNO(g)
ClNO2(g) +
NO(g)
The equilibrium constant expression is now written as follows.
K c' =
[ NO2 ][ClNO]
[ClNO2 ][ NO]
Each of these equilibrium constant expressions is the inverse of the other. We
can therefore calculate Kc by dividing Kc into 1.
'
K=
c
1
1
7.7 ×10−5
=
=
4
K c 1.3 ×10
158
University Chemistry
We can also calculate equilibrium constants by combining two or more
reactions for which the value of Kc is known. Assume, for example, that we know
the equilibrium constants for the following gas-phase reactions at 200oC.
N2(g)
2 NO(g)
+
+
O2(g)
2 NO(g)
O2(g)
2 NO2(g)
Kc1 = 2.3 x 10-19
Kc2 = 3 x 106
We can combine these reactions to obtain an overall equation for the reaction
between N2 and O2 to form NO2.
+
N2(g)
+
2 NO(g)
+
N2(g)
+
O2(g)
2 NO(g)
2 O2(g)
2 NO2(g)
O2(g)
2 NO2(g)
It is easy to show that the equilibrium constant expression for the overall
reaction is equal to the product of the equilibrium constant expressions for the two
steps in this reaction.
[ NO2 ]
[ NO2 ]
[ NO2 ]2
=
×
[ N 2 ][O2 ]2 [ N 2 ][O2 ] [ NO]2 [O2 ]2
The equilibrium constant for the overall reaction is therefore equal to the
product of the equilibrium constants for the individual steps.
Kc = Kc1 x Kc2 = (2.3 x 10-19)(3 x 106) = 7 x 10-13
9.3 THE EQUILIBRIUM CONSTANT
The expression of equilibrium constant depends upon
the manner in which the chemical equation representing it
is written. For the reaction
The equilibrium constant K is given by
When the same reaction is written as
the corresponding equilibrium constant K1 is given by
Important
The concentrations of pure
solids, pure liquids,
and solvents are
omitted from equilibrium constant
expressions because
they do not change
significantly during reactions when
enough is present to
reach equilibrium.
Chemical Equilibrium
159
It may be noted that equilibrium constants K and K1 are related as K1 = K
(b) When the reaction is written as reverse
Here it can be seen that
Similar relationship is also observed in the pressure equilibrium constant
Kp . Thus the expression of equilibrium constant depends on how the reaction is
expressed in the form of a chemical equation.
9.3.1 Units of Equilibrium Constant
Units of equilibrium constant Kc or Kp depend upon the fact whether during
the reactions there is any change in the moles of substance or not.
(a) The reactions in which there is no change in moles of substance i.e. ∆ n =
0. The equilibrium constant for such reaction has no units. For example in
the reaction between H2 and I2
∴ Hence Kp and Kc have no units in such cases.
(b) The reaction where there is change in the moles of substance i.e. ∆ n ≠
0. The equilibrium constant for such reactions has units which depend
upon the change in moles of substances.
For example :
160
University Chemistry
The units of Kc for this reaction would be (mol L–1) –2 or L2 mol–2 and those of Kp
would be bar–2 as shown below :
The equilibrium constant for such reactions are
9.4 HETEROGENEOUS EQUILIBRIA
When the state of equilibrium in a system has
components in more than one phase it is termed as a
heterogeneous equilibrium. For example, if we take a
container with ice and water at a temperature that is
allowing the existence of both the phases simultaneously,
both ice and water are present in a state of equilibrium.
This state is termed as heterogeneous equilibrium.
9.4.1 Equilibrium in Heterogeneous System
Important
The equilibrium constant for
a reaction written
in reverse is the inverse of the equilibrium constant
for the reaction as
written originally.
In a chemical reaction the reacting species combine
with each other to form products. If such a reaction
is carried in a closed container, in many cases, we
find that it is not completed and after some time , the
reaction mixture contains both the reactant and product species. Moreover, the
Chemical Equilibrium
161
concentration of these species also becomes constant after some time , provided
both the temperature and the pressure are constant . Under the conditions, the
reaction is said to be equilibrium.
In a heterogeneous system involving a reversible reaction, the reactants and
products are in the different phases.
Let us illustrate the following reversible reaction involving decomposition of
CaCO3 (s). Calcium carbonate on decomposition to produce CaO and CO2 . At a
time the rate at which CaCO3 (s) decomposes is the same at the product combine to
give the reactant. At this time equilibrium takes place.
CaCO3  CaO( s ) + CO2 ( g )
9.4.2 Examples of Heterogeneous Equilibrium
The examples of a heterogeneous equilibrium are
•
•
•
Reaction of solid Ferrous oxide with Gaseous carbon monoxide produces
solid iron and gaseous carbon dioxide.
FeO (s) + CO (g) ⇌ Fe (s) + CO2(g)
Reaction of steam with red hot carbon produces hydrogen gas and carbon
monoxide gas.
H2O (g) + C (s) ⇌ H2(g) + CO(g)
Reaction of carbon dioxide gas with solid carbon at equilibrium produces
carbon monoxide gas.
CO2 (g) + C (s) ⇌ 2 CO(g)
9.4.3 Equilibrium Constant in Heterogeneous System
Writing the equilibrium constant for heterogeneous reactions is different from
that of the homogeneous reactions. For example, consider the thermal dissociation
of calcium carbonate into calcium oxide and carbon dioxide. Here, we can see that
the equilibrium constant for the dissociation of the reactant into its products is only
dependent on the gaseous component, as the solid and the liquid reactants are
considered to be constant.
CaCO3 ( s )  CaO( s ) + CO2 ( g )
Here, the components CaCO3 and CaO are solids, so their molar concentration
remains constant throughout the reaction. Therefore, the equilibrium constant can
be written as,
162
University Chemistry
Kc =
[CaO][CO2 ]
[CaCO3 ]
K c = [CO2 ]
Also, in terms of Kp, we can write
K p = pCO2
Where p denotes the partial pressure. In other words, we can state that, at a
given temperature, there is a constant concentration or partial pressure of CO2 in
the equilibrium reaction with CaO and CaCO3.
Similarly: for equation CO2 (g) + C (s) ⇌ 2 CO(g)
Equilibrium constant Kc = [ CO]2 / [CO2]
9.4.4 Equilibrium in Physical Change
The phase changes are also heterogeneous equilibrium systems. We know that
solid , liquid and gas are the state of substance so that types of equilibrium are
possible between
Solid – Liquid Equilibrium (solid ⇋ liquid)
Liquid – Gas Equilibrium (liquid ⇋ Gas)
Solid – Gas Equilibrium (solid ⇋ Gas)
Solid – Liquid Equilibrium:
The equilibrium between ice and water at 0℃ and 1 atm pressure in a thermo
flask . In this case the rate of melting of ice and freezing water are taking place at
the same rate.
In terms of its equation, this can be written as:
H 2O( s )  H 2O(l )
Liquid – Gas Equilibrium:
The liquid -gas equilibrium can be illustrated with the help of evaporation of
water in a closed container at room temperature.
H2O (l) ⇋ H2O (g)
Solid – Gas Equilibrium:
This state of equilibrium takes place in case of substances like ammonium
chloride, iodine, camphor etc. which undergo sublimation upon heating or even
at room temperature. They directly change to the vapour or gaseous state without
Chemical Equilibrium
163
passing through the liquid state.
I2 (s) ⇋ I2 (g)
NH4Cl (s) ⇋ NH4Cl (g)
Camphor (s) ⇋ Camphor (g)
9.5 EQUILIBRIUM CALCULATIONS
Having covered the essential concepts of chemical equilibria in the preceding
sections of this chapter, this final section will demonstrate the more practical
aspect of using these concepts and appropriate mathematical strategies to perform
various equilibrium calculations. These types of computations are essential to
many areas of science and technology—for example, in the formulation and dosing
of pharmaceutical products. After a drug is ingested or injected, it is typically
involved in several chemical equilibria that affect its ultimate concentration in the
body system of interest. Knowledge of the quantitative aspects of these equilibria is
required to compute a dosage amount that will solicit the desired therapeutic effect.
Many of the useful equilibrium calculations that will be demonstrated here
require terms representing changes in reactant and product concentrations.
These terms are derived from the stoichiometry of the reaction, as illustrated by
decomposition of ammonia:
2NH3(g)⇌N2(g)+3H2(g)
As shown earlier in this chapter, this equilibrium may be established within
a sealed container that initially contains either NH3 only, or a mixture of any two
of the three chemical species involved in the equilibrium. Regardless of its initial
composition, a reaction mixture will show the same relationships between changes
in the concentrations of the three species involved, as dictated by the reaction
stoichiometry (see also the related content on expressing reaction rates in the chapter
on kinetics). For example, if the nitrogen concentration increases by an amount x:
Δ[N2]=+x
the corresponding changes in the other species concentrations are
 3molH 2 
∆[ H 2 ] =
∆[ N 2 ] 
+3 x
=
 1molN 2 
 2molNH 3 
∆[ NH 3 ] =
∆[ N 2 ] 
−2 x
=
 1molN 2 
where the negative sign indicates a decrease in concentration.
164
University Chemistry
9.5.1 Calculations of an Equilibrium Constant
The equilibrium constant for a reaction is calculated
from the equilibrium concentrations (or pressures) of
its reactants and products. If these concentrations are
known, the calculation simply involves their substitution
into the K expression, as was illustrated by Example 2.
A slightly more challenging example is provided next,
in which the reaction stoichiometry is used to derive
equilibrium concentrations from the information
provided. The basic strategy of this computation is
helpful for many types of equilibrium computations and
relies on the use of terms for the reactant and product
concentrations initially present, for how they change as
the reaction proceeds, and for what they are when
the system reaches equilibrium. The acronym ICE is
commonly used to refer to this mathematical approach,
and the concentrations terms are usually gathered in a
tabular format called an ICE table.
Important
Calculating
values for equilibrium constants
and/or equilibrium
concentrations is of
practical benefit to
many applications.
A mathematical
strategy that uses
initial concentrations, changes in
concentrations, and
equilibrium concentrations (and goes
by the acronym ICE)
is useful for several
types of equilibrium
calculations.
Example: Calculation Of An Equilibrium Constant
Iodine molecules react reversibly with iodide ions to
produce triiodide ions.
I 2 (aq ) + I − (aq )  I 3− (aq )
If a solution with the concentrations of I2 and I− both equal to 1.000 × 10−3M before
reaction gives an equilibrium concentration of I2 of 6.61 × 10−4M, what is the
equilibrium constant for the reaction?
9.5.2 Calculation of a Missing Equilibrium Concentration
When the equilibrium constant and all but one equilibrium concentration are
provided, the other equilibrium concentration(s) may be calculated.
9.5.3 Calculation of Changes in Concentration
Perhaps the most challenging type of equilibrium calculation can be one in
which equilibrium concentrations are derived from initial concentrations and an
equilibrium constant. For these calculations, a four-step approach is typically useful:
•
•
•
Identify the direction in which the reaction will proceed to reach
equilibrium.
Develop an ICE table.
Calculate the concentration changes and, subsequently, the equilibrium
concentrations.
165
Chemical Equilibrium
•
Confirm the calculated equilibrium concentrations.
9.5.4 Specific Initial Concentrations
Calculating the initial concentration of a solution – otherwise known as molarity
– is an important process commonly found in the chemical and biochemical world.
Molarity is the number of moles of solute per liter of solution. Therefore, you need
to determine how many moles of a solute are in the solution and the total volume
of the solution.
Step 1. Weigh the amount of solute (the compound being dissolved) in grams.
Then determine how many grams are in a mole of the
solute. There are 40 g per mole in sodium hydroxide
(NaOH). Therefore, 20 g of NaOH would equal 0.50 mol
Important
of NaOH. The equation looks like this:
mol NaOH = 20.0g NaOH x 1 mol NaOH/40.0 g
NaOH.
Step 2. Measure the amount of the solvent that
you have. If it is less than a liter, convert the number of
milliters into liters. There are 1000mL in 1L. For example,
if you have 500 mL:
500 mL x 1L/1000mL = 0.500 L solvent.
Step 3. Divide the moles of solute found in Step 1
by the liters of solvent found in Step 2 to find the initial
concentration of a solution. The equation looks like this:
Keep track
of your units so
that you can make
a clean conversion
to moles and the
liters of the solvent.
Not keeping track
of units can make
it difficult when
dealing with conversions from very
small amounts to
moles.
M = 0.50 mol NaOH/0.500 L solvent = 1 M NaOH.
In this example, the molarity (M) of the NaOH in the solvent is one mole. As
more of the solvent is removed, the concentration of the NaOH would continue to
rise. With acids and bases, the higher the concentration, the stronger it becomes.
9.6 THE RESPONSE OF EQUILIBRIA TO THE REACTION
CONDITIONS
A system is at equilibrium when the rates of the forward and reverse reactions
are equal. If additional reactant is added the rate of the forward reaction increases.
As the rate of the reverse reaction is initially unchanged, the equilibrium appears
to shift toward the product, or right, side of the equation. As the additional reactant
is consumed the forward rate slows. When the rates of the forward and reverse
reactions are again equal, the system has returned to equilibrium. It is convenient
to think of this as the system shifting to the right to remove the added reactant.
166
University Chemistry
While a disturbed system will return to equilibrium, it is not exactly the same
equilibrium that existed before the stress was applied. One difference will be in the
concentrations of the reactants and products. Consider the system below:
H2 + I2  2 HI
If we add some hydrogen to the system at equilibrium, the system will
shift to the right in an attempt to remove the extra hydrogen. It can be shown
mathematically that it will be unable to remove all of it. Therefore, the new
equilibrium concentration of hydrogen will be higher than it was in the original
equilibrium. The only way the system has to remove the hydrogen is by reaction
with iodine. Therefore, the new equilibrium concentration of iodine will be lower
than it was in the original equilibrium. When hydrogen and iodine react they form
hydrogen iodide, so the concentration of hydrogen iodide will be higher than it was
in the original equilibrium conditions. We could summarize the effects in a table
like the one below in which brackets are used to represent concentrations.
Applied
stress
Direc- effect on
tion
[H2]
of shift
effect on
[I2]
H2 added
right
decreased increased
increased
effect on
[HI]
If hydrogen had been removed from the system, the equilibrium would shift
to the left in an attempt to make more of it. We could prove mathematically that
it will be unable to replenish all of it and that the new equilibrium concentration
of hydrogen would be lower than it was in the original equilibrium. In order to
make more hydrogen, the system must use up hydrogen iodide causing its new
equilibrium concentration to be lower than it was in the original equilibrium.
When hydrogen iodide decomposes it forms both hydrogen and iodine. The
new equilibrium concentration of iodine will be higher than it was in the original
equilibrium.
Applied
stress
Direc- effect on
tion
[H2]
of shift
effect on
[I2]
effect on
[HI]
H2 removed
left
increased
decreased
decreased
Practice using LeChatelier’s Principle by predicting what should happen if (1)
iodine is added to the system at equilibrium, (2) iodine is removed from the system
at equilibrium, (3) hydrogen iodide is added to the system at equilibrium and (4)
hydrogen iodide is removed from the system at equilibrium.
9.6.1 Le Chatelier’s Principle
Le Chatelier’s principle states that changes to an equilibrium system will result
in a predictable shift that will counteract the change.
167
Chemical Equilibrium
Le Chatelier’s principle is an observation about
chemical equilibria of reactions. It states that changes in
the temperature, pressure, volume, or concentration of a
system will result in predictable and opposing changes
in the system in order to achieve a new equilibrium
state. Le Chatelier’s principle can be used in practice to
understand reaction conditions that will favor increased
product formation. This idea was discovered and
formulated independently by Henri Louis Le Chatelier
and Karl Ferdinand Braun.
9.6.2 Changes in Concentration
Important
Le Chatelier’s principle
implies that the
addition of heat to
a reaction will favor
the endothermic
direction of a reaction as this reduces
the amount of heat
produced in the
system.
According to Le Chatelier’s principle, adding
additional reactant to a system will shift the equilibrium
to the right, towards the side of the products. By the same logic, reducing the
concentration of any product will also shift equilibrium to the right.
The converse is also true. If we add additional product to a system, the
equilibrium will shift to the left, in order to produce more reactants. Or, if we
remove reactants from the system, equilibrium will also be shifted to the left.
Thus, according to Le Chatelier’s principle, reversible reactions are self-correcting;
when they are thrown out of balance by a change in concentration, temperature, or
pressure, the system will naturally shift in such a way as to “re-balance” itself after
the change.
This can be illustrated by the equilibrium of this reaction, where carbon
monoxide and hydrogen gas react to form methanol:
CO+2H2⇌CH3OH
Suppose we were to increase the concentration of CO in the system. By Le
Chatelier’s principle, we can predict that the amount of methanol will increase,
thereby decreasing the total change in CO. If we add a species to the overall reaction,
the reaction will favor the side opposing the addition of the species. Likewise, the
subtraction of a species would cause the reaction to fill the “gap” and favor the side
where the species was reduced.
This observation is supported by the collision theory. As the concentration of CO
is increased, the frequency of successful collisions of that reactant would increase
as well, allowing for an increase in the forward reaction, and thus the generation of
the product. Even if a desired product is not thermodynamically favored, the endproduct can be obtained if it is continuously removed from the solution.
168
University Chemistry
9.6.3 Changes in Pressure
A change in pressure or volume will result in an attempt to restore equilibrium
by creating more or less moles of gas. For example, if the pressure in a system
increases, or the volume decreases, the equilibrium will shift to favor the side of
the reaction that involves fewer moles of gas. Similarly, if the volume of a system
increases, or the pressure decreases, the production of additional moles of gas will
be favored.
Consider the reaction of nitrogen gas with hydrogen gas to form ammonia:
N2+3H2⇌2NH3ΔH=−92kJ mol−1
Note the number of moles of gas on the left-hand side and the number of
moles of gas on the right-hand side. When the volume of the system is changed, the
partial pressures of the gases change. If we were to decrease pressure by increasing
volume, the equilibrium of the above reaction would shift to the left, because the
reactant side has greater number of moles than the product side. The system tries to
counteract the decrease in partial pressure of gas molecules by shifting to the side
that exerts greater pressure.
Similarly, if we were to increase pressure by decreasing volume, the equilibrium
would shift to the right, counteracting the pressure increase by shifting to the side
with fewer moles of gas that exert less pressure.
Lastly, for a gas-phase reaction in which the number of moles of gas on both
sides of the equation are equal, the system will be unaffected by changes in pressure,
since Δn=0.
9.6.4 Addition of an Inert Gas
What would happen to the equilibrium position of the reaction if an inert gas,
such as krypton or argon, were added to the reaction vessel? Answer: nothing
at all. Remember that the system will always shift so that the ratio of products
and reactants remains equal to Kp or Kc. An inert gas will not react with either the
reactants or the products, so it will have no effect on the product/reactant ratio, and
therefore, it will have no effect on equilibrium.
9.6.5 Changes in Temperature
The effect of temperature on equilibrium has to do with the heat of reaction.
Recall that for an endothermic reaction, heat is absorbed in the reaction, and the
value of ΔH is positive. Thus, for an endothermic reaction, we can picture heat as
being a reactant:
heat+A⇌BΔH=+
For an exothermic reaction, the situation is just the opposite. Heat is released in
the reaction, so heat is a product, and the value of ΔHΔH is negative:
Chemical Equilibrium
169
A⇌B+heatΔH=−
If we picture heat as a reactant or a product, we can apply Le Chatelier’s
principle just like we did in our discussion on raising or lowering concentrations.
For instance, if we raise the temperature on an endothermic reaction, it is essentially
like adding more reactant to the system, and therefore, by Le Chatelier’s principle,
the equilibrium will shift the right. Conversely, lowering the temperature on an
endothermic reaction will shift the equilibrium to the left, since lowering the
temperature in this case is equivalent to removing a reactant.
For an exothermic reaction, heat is a product. Therefore, increasing the
temperature will shift the equilibrium to the left, while decreasing the temperature
will shift the equilibrium to the right.
9.6.6 The Effect of a Catalyst
Catalysts speed up the rate of a reaction, but do not have an affect on the
equilibrium position.
Reactions can be sped up by the addition of a catalyst, including reversible
reactions involving a final equilibrium state. Recall that for a reversible reaction, the
equilibrium state is one in which the forward and reverse reaction rates are equal. In
the presence of a catalyst, both the forward and reverse reaction rates will speed up
equally, thereby allowing the system to reach equilibrium faster. However, it is very
important to keep in mind that the addition of a catalyst has no effect whatsoever on
the final equilibrium position of the reaction. It simply gets it there faster.
Recall that catalysts are compounds that accelerate the progress of a reaction
without being consumed. Common examples of catalysts include acid catalysts
and enzymes. Catalysts allow reactions to proceed faster through a lower-energy
transition state. By lowering the energy of the transition state, which is the ratelimiting step, catalysts reduce the required energy of activation to allow a reaction
170
University Chemistry
to proceed and, in the case of a reversible reaction, reach equilibrium more rapidly.
To reiterate, catalysts do not affect the equilibrium state of a reaction. In the
presence of a catalyst, the same amounts of reactants and products will be present at
equilibrium as there would be in the uncatalyzed reaction. To state this in chemical
terms, catalysts affect the kinetics, but not the thermodynamics, of a reaction. If the
addition of catalysts could possibly alter the equilibrium state of the reaction, this
would violate the second rule of thermodynamics; we would be getting “something
for nothing,” which is physically impossible.
Interactive: Catalysis: The model contains reactants which will form the
reaction: A₂ + B₂ –> 2 AB. In this case the model has been set so the activation energy
is high. Try running the reaction with and without a catalyst to see the effect catalysts
have on chemical reactions. 1. Run the model to observe what happens without a
catalyst. 2. Pause the model. 3. Add a few (3 – 4) catalyst atoms to the container by
clicking the button. 4. Run the model again, and observe how the catalyst affects
the reaction.
EXERCISE
Answer the following questions:
1. What is chemical equilibrium?
2. What is Le Chatelier’ principle?
3. What will happen to solid-vapour equilibrium when the temperature and
pressure are decreased?
4. Expression of equilibrium constant depends upon how the chemical equation is written for the reaction.
5. What are Kp and Kc? Derive a relation between them.
MULTIPLE CHOICE QUESTIONS
Tick the correct answer:
1.
Highest pH will be recorded for which of the following solutions if they are
equimolar
a. AlCl3
b. BaCl2
c. BeCl2
d. LiCl
2. What will be the pH of a buffer solution having an equal concentration of
B– and HB (Kb = 10-10 for B–)
a. 7
b. 4
c. 10
d. 6
171
Chemical Equilibrium
3.
4.
5.
On increasing the concentration of reactants in a reversible reaction, then
equilibrium constant will
a. depend on the concentration
b. increase
c. unchanged
d. decrease
Find the conjugate acid of NH2–
a. NH3
b. NH4OH
c. NH4+
d. NH2–
What is the equilibrium constant for a reaction that has a value of ∆ Go =
-41.8 kJ at 100oC?
a. 1.01
b. 7.1 x 105
c. -5.87
d. 1.4 x 10-6
e. 13.5
ANSWERS
1. (b)
2. (b)
3. (c)
4. (a)
5. (b)
REFERENCES
1.
2.
3.
4.
5.
6.
7.
“Chemical Equilibrium with Applications”. NASA. Archived from the original
on September 1, 2000. Retrieved October 5, 2019.
Atkins, Peter W.; Jones, Loretta (2008). Chemical Principles: The Quest for
Insight (2nd ed.). ISBN 978-0-7167-9903-0.
Atkins, Peter; De Paula, Julio (2006). Atkins’ Physical Chemistry (8th ed.). W.
H. Freeman. pp. 200–202. ISBN 0-7167-8759-8.
Ernest, Z., 2014. How is Le Chatelier’s principle used to treat CO poisoning? |
Socratic. [online] Socratic.org. Available at: <https://socratic.org/questions/
how-is-le-chatelier-s-principle-used-to-treat-co-poisoning>.
IUPAC, Compendium of Chemical Terminology, 2nd ed. (the “Gold Book”)
(1997). Online corrected version: (2006–) “chemical equilibrium”.
Lower, S., 2017. Le Chatelier principle. [online] Chem1.com. Available at: <http://
www.chem1.com/acad/webtext/chemeq/Eq-02.html>.
n.d. EXPERIMENT 8: DETERMINATION OF EQUILIBRIUM CONSTANT.
[ebook] University of Missouri, p.10. Available at: <https://chemistry.missouri.
edu/sites/default/files/class-files/use_det_eq_const_1.pdf>.
172
8.
University Chemistry
Scienceclarified. n.d. Real-life applications - Chemical Equilibrium. [online]
Available at: <http://www.scienceclarified.com/everyday/Real-Life-ChemistryVol-2/Chemical-Equilibrium-Real-life-applications.html> [Accessed 12 March
2019].
Introduction to Functional Groups and their Typical Reactions
173
CHAPTER 10
INTRODUCTION TO FUNCTIONAL GROUPS
AND THEIR TYPICAL REACTIONS
OBJECTIVES
After studying this chapter, you will be able to:
1. Introduction to Functional Groups
2. Focus on alkanes, alkenes and alkynes
3. Explain the aromatic compounds
4. Focus on alcohols
5. Disuses about aldehydes and ketones
6. Explain the carboxylic acids and their derivatives
7. Focus on ethers
8. Discuss about amines
10.1 INTRODUCTION TO FUNCTIONAL GROUPS
Chemists observed early in the study of organic compounds that certain groups
of atoms and associated bonds, known as functional groups, confer specific reactivity
patterns on the molecules of which they are a part. Although the properties of each
of the several million organic molecules whose structure is known are unique in
some way, all molecules that contain the same functional group have a similar
pattern of reactivity at the functional group site. Thus, functional groups are a key
174
organizing feature of organic chemistry. By focusing
on the functional groups present in a molecule (most
molecules have more than one functional group), several
of the reactions that the molecule will undergo can be
predicted and understood. Because carbon-to-carbon
and carbon-to-hydrogen bonds are extremely strong
and the charge of the electrons in these covalent bonds
is spread more or less evenly over the bonded atoms,
hydrocarbons that contain only single bonds of these two
types are not very reactive. The reactivity of a molecule
increases if it contains one or more weak bonds or bonds
that have an unequal distribution of electrons between the
two atoms. If the two electrons of a covalent bond are, for
one reason or another, drawn more closely to one of the
bonded atoms, that atom will develop a partial negative
charge and the atom to which it is bonded will develop
a partial positive charge. A covalent bond in which the
electron pair linking the atoms is shared unequally is
known as a polar bond.
University Chemistry
Note
Polar bonds, and
any other bonds that
have unique electronic properties,
confer the potential
for chemical reaction on the molecule
in which they are
present. This is because, for every reaction, one or more
bonds of a molecule
must be broken and
new bonds formed.
The presence of a partial negative charge (a region of high electron density) will
draw to itself other atoms or groups of atoms that are deficient in electron density.
This initiates the process of bond breaking that is a prerequisite for a chemical
reaction. For these reasons, molecules with regions of increased or decreased
electron density are especially important for chemical change. There are two major
bonding features that generate the reactive sites of functional groups. The first,
already mentioned, is the presence of multiple bonds. Both double and triple bonds
have regions of high electron density lying outside the atom-to-atom bond axis.
Double and triple bonds are known as functional groups, a term that is used to
identify atoms or groups of atoms within a molecule that are sites of comparatively
high reactivity. A second type of reactive site results when an atom other than
carbon or hydrogen (termed a heteroatom) is bonded to carbon. All heteroatoms
have a greater or lesser attraction for electrons than doe’s carbon. Thus, each bond
between a carbon and a heteroatom is polar, and the degree of polarity depends on
the difference between the electron-attracting properties of the two atoms. The most
important atomic groupings that contain such reactive polar bonds are also able
to generate functional groups. To emphasize the generality of reactions between
molecules that contain the same functional group, chemists often represent the less
reactive portions of a molecule by the symbol R. Thus, all molecules that contain a
double bond, however complicated, can be represented by the general formula for
an alkene—i.e.,
Introduction to Functional Groups and their Typical Reactions
175
This type of formula suggests that the molecule will undergo those reactions
that are common to double bonds and that the reaction will occur at the double
bond. The rest of the molecule, represented by the four R groups, will remain
unchanged by the reaction occurring at the functional group site.
10.2 ALKANES, ALKENES AND ALKYNES
Molecules with more than one functional group, called polyfunctional,
may have more complicated properties that result from the identity—and
interconnectedness—of the multiple functional groups. Many natural products
contain several functional groups located at specific sites within a large, complicated,
three-dimensional structure.
A brief overview of the principal functional groups is presented here.
10.2.1 Alkanes
Alkanes are compounds that consist entirely of atoms of carbon and hydrogen (a
class of substances known as hydrocarbons) joined to one another by single bonds.
The shared electron pair in each of these single bonds occupies space directly
between the two atoms; the bond generated by this shared pair is known as a sigma
(σ) bond. Both carbon-carbon and carbon-hydrogen sigma bonds are single strong,
nonpolar covalent bonds that are normally the least reactive bonds in organic
molecules. Alkane sequences form the inert framework of most organic compounds.
For this reason, alkanes are not formally considered a functional group. When a
hydrocarbon chain is connected as a substituent to a more fundamental structural
unit, it is termed an alkyl group. The simplest examples of alkanes are methane (CH4;
the principal constituent of natural gas), ethane (C2H6), propane (C3H8; widely
used as a barbecue fuel), and butane (C4H10; the liquid fuel in pocket lighters).
Hydrocarbon chains commonly occur in cyclic forms, or rings; the most common
example is cyclohexane (C6H12). (For a more detailed examination of these
compounds, see hydrocarbon.)
10.2.2 Alkenes
Organic compounds are termed alkenes if they contain a carbon-carbon double
bond. The shared electron pair of one of the bonds is a σ bond. The second pair of
electrons occupies space on both sides of the σ bond; this shared pair constitutes a pi
(π) bond. A π bond forms a region of increased electron density because the electron
pair is more distant from the positively charged carbon nuclei than is the electron
pair of the σ bond (see chemical bonding: The quantum mechanics of bonding).
176
Even though a carbon-carbon double bond is very strong,
a π bond will draw to itself atoms or atomic groupings
that are electron-deficient, thereby initiating a process
of bond-breaking that can lead to rupture of the π bond
and formation of new σ bonds. A simple example of an
alkene reaction, which illustrates the way in which the
electronic properties of a functional group determine its
reactivity, is the addition of molecular hydrogen to form
alkanes, which contain only σ bonds.
University Chemistry
Important
Alkanes,
which cannot be
transformed by addition reactions into
molecules with a
greater number of σ
bonds, are said to be
saturated.
Such reactions, in which the π bond of an alkene reacts to form two new σ
bonds, are energetically favourable because the new bonds formed (two carbonhydrogen σ bonds) are stronger than the bonds broken (one carbon-carbon π bond
and one hydrogen-hydrogen σ bond). Because the addition of atoms to the π bond
of alkenes to form new σ bonds is a general and characteristic reaction of alkenes,
alkenes are said to be unsaturated.
The alkene functional group is an important one in chemistry and is widespread
in nature. Some common examples (shown here) include ethylene (used to
make polyethylene), 2-methyl-1,3-butadieneisoprene (used to make rubber),
and vitamin A (essential for vision).
For ethene, both the carbon atoms of an alkene and the four atoms connected to
the double bond lie in a single plane.
10.2.3 Alkynes
Molecules that contain a triple bond between two carbon atoms are known as
alkynes. The triple bond is made up of one σ bond and two π bonds. As in alkenes,
the π bonds constitute regions of increased electron density lying parallel to the
carbon-carbon bond axis. Carbon-carbon triple bonds are very strong bonds, but
reactions do occur that break the π bonds to form stronger σ bonds.
Introduction to Functional Groups and their Typical Reactions
177
The most common example of an alkyne is ethyne (also known as acetylene),
used as a fuel for oxyacetylene torches in welding applications. Alkynes are not
abundant in nature, but the fungicide capillan contains two alkyne functional
groups.
Aromatic hydrocarbons (arenes)
A distinctive set of physical and chemical properties is imparted to molecules
that contain a functional group composed of three pairs of doubly bonded atoms
(usually all carbon atoms) bonded together in the shape of a regular planar (flat)
hexagon. The hexagonal ring is usually drawn with an alternating sequence of
single and double bonds. The molecule benzene, C6H6, first discovered by English
physicist and chemist Michael Faraday in 1825, is the smallest molecule that
can contain this functional group, and arenes contain one or more benzene (or
structurally similar) rings. Because benzene and many larger arenes have a strong
odour, they have long been known as aromatic hydrocarbons. Benzene, and all the
larger arenes, have a characteristic planar structure forced on them by the electronic
requirements of the six (or more) pi electrons. When named as substituents on other
structural units, the aromatic units are called aryl substituents. Naphthalene, the
active component of mothballs, contains two fused benzene rings. Benzo[a]pyrene,
an aromatic hydrocarbon produced in small amounts by the combustion of organic
substances, contains five fused benzene rings. Like several other polycyclic aromatic
hydrocarbons, it is carcinogenic. Aromatic compounds are widely distributed in
nature. Benzaldehyde, anisole, and vanillin, for example, have pleasant aromas.
Figure 1: Chemical bonding in benzene.
Benzene is the smallest of the organic aromatic hydrocarbons. It contains sigma
bonds (represented by lines) and regions of high-pi electron density, formed by the
178
University Chemistry
overlapping of p orbitals (represented by the dark yellow shaded area) of adjacent
carbon atoms, which give benzene its characteristic planar structure.
10.3 AROMATIC COMPOUNDS
Aromatic compounds are chemical compounds that consist of conjugated
planar ring systems accompanied by delocalized pi-electron clouds in place of
individual alternating double and single bonds. They are also called aromatics or
arenes. The best examples are toluene and benzene. Aromatics require satisfying
Huckel’s rule. Plants and micro-organisms have an exclusive route to benzene-ring
compounds. The great majority of aromatic compounds in nature, therefore, are
produced by plants and micro-organisms, and animals are dependant upon plants
for many aromatic compounds either directly or indirectly.
Aromatic Compounds Examples
Aromatic hydrocarbon, are hydrocarbons containing sigma bonds and
delocalized pi electrons between carbon atoms in a ring. For example, benzene.
They are known as aromatic due to their pleasant smell.
Aromatic compounds are broadly divided into two categories: benzenoids (one
containing benzene ring) and non-benzenoids (those not containing a benzene ring)
for example, furan. Any hydrocarbon can be classified as an aromatic compound
provided they follow the Huckel rule. According to Huckel rule, for a ring to be
aromatic it should have the following properties:
•
•
•
Planarity
Complete delocalization of the π electrons in the ring
Presence of (4n + 2) π electrons in the ring where n is an integer (n = 0, 1,
2, . . .)
Introduction to Functional Groups and their Typical Reactions
179
Huckel’s Rule of Aromaticity
Huckel’s rule states that only planar, fully conjugated monocyclic polyenes
having 4n + 2 π electrons, where n is an integer, that is, n = 0, 1, 2, 3, 4, etc., should
possess aromatic stability. An aromatic compound must be planar and contain a
cyclic cloud of π electrons below and above the plane of the molecule. It contains
SP2 hybridized carbon atoms and must obey the Huckel rule.
According to this rule, the ring system must have (4n+2) π electrons, where n is
any whole number (0, 1, 2, 3, etc). On this basis the ring systems which have 2(n=0),
6(n=1), 10(n=2), 14(n=3) etc pi electrons are aromatic. Typical examples of aromatic
compounds are benzene, naphthalene, and anthracene.
10.3.1 Properties of Aromatic Compounds
Arenes are mostly nonpolar and non-miscible in water. These compounds are
usually unreactive and are used as solvents for various other nonpolar compounds.
Their carbon to hydrogen ratio is high therefore, they are characterized by sooty
yellow flame.
10.3.2 Classification of Aromatic Compounds
The classification of arenes is based on the position of the functional group.
They are classified into two and we have discussed below:
1. Nuclear Substituted Compounds
In any aromatic compound whenever any substituent or functional group, is
directly linked to the benzene ring, it is known as a nuclear-substituted compound.
2. Side chain Substituted Compounds
In any aromatic compound if the functional group is available in the side chain
of the ring then it is known as a side chain substituted compound. These compounds
are named as the phenyl derivatives of the relative aliphatic compounds.
IUPAC Nomenclature of Aromatic Compounds
Earlier, most of the compounds with the same structural formula were known
by different names depending on the regions where they were synthesized. This
180
University Chemistry
naming system was very trivial since it raised a lot of confusion. Finally, a common
naming system enlisting standard rules was set up by IUPAC (International Union
for Pure and Applied Chemistry) for the naming of compounds. This method of
naming is IUPAC naming or IUPAC nomenclature.
IUPAC nomenclature of aromatic hydrocarbons is explained below:
1.
According to IUPAC nomenclature of substituted aromatic compounds,
the substituent name is placed as a prefix to the name of aromatic
compounds. For example, a benzene ring attached to a one-nitro group is
named as nitrobenzene.
2.
When more than one similar substituent group is present in the ring,
they are labelled with the Greek numerical prefixes such as di, tri, tetra
to denote the number of similar substituent groups attached to the ring.
If two bromo- groups are attached to the adjacent carbon atoms of the
benzene ring, it is named 1,2-dibromobenzene.
3.
When different substituted groups are attached to the aromatic compounds,
the substituent of the base compound is assigned number one and then
the direction of numbering is chosen such that the next substituent gets
the lowest number. Substituents are named in alphabetical order. For
example: when chloro and nitro groups are attached to the benzene ring,
we first locate the chloro group then nitro groups.
4.
In the case of multiple substituted aromatic compounds, sometimes terms
like ortho (o), meta (m) and para (p) are also used as prefixes to indicate
the relative positions 1,2-; 1,3- and 1,4- respectively. For example, 1,2-Dibromo-benzene can be named as o-di-bromo-benzene.
181
Introduction to Functional Groups and their Typical Reactions
5.
When an alkane with a functional group is attached to an aromatic
compound, the aromatic compound is considered as a substituent, instead
of a parent. For example: when a benzene ring is attached to an alkane
with a functional group, it is considered as a substituent named phenyl,
denoted by Ph-.
10.4 ALCOHOLS
Alcohol, any of a class of organic compounds characterized by one or
more hydroxyl (―OH) groups attached to a carbon atom of an alkyl group
(hydrocarbon chain). Alcohols may be considered as organic derivatives
of water (H2O) in which one of the hydrogen atoms has been replaced by an alkyl
group, typically represented by R in organic structures. For example, in ethanol
(or ethyl alcohol) the alkyl group is the ethyl group, ―CH2CH3.
Alcohols are among the most common
organic compounds. They are used as sweeteners and
in making perfumes, are valuable intermediates in the
synthesis of other compounds, and are among the most
abundantly produced organic chemicals in industry.
Perhaps the two best-known alcohols are ethanol
and methanol (or methyl alcohol). Ethanol is used in
toiletries, pharmaceuticals, and fuels, and it is used to
sterilize hospital instruments. It is, moreover, the alcohol
in alcoholic beverages. The anesthetic ether is also
made from ethanol. Methanol is used as a solvent, as a
raw material for the manufacture of formaldehyde and
special resins, in special fuels, in antifreeze, and for
cleaning metals.
Note
Alcohols may be
classified as primary, secondary, or
tertiary, according
to which carbon of
the alkyl group is
bonded to the hydroxyl group. Most
alcohols are colorless liquids or solids at room temperature.
Alcohols of low molecular weight are highly
soluble in water; with increasing molecular weight, they
become less soluble in water, and their boiling points, vapour pressures, densities,
and viscosities increase.
This article covers the structure and classification, physical properties,
commercial importance, sources, and reactions of alcohols. For more information
about closely related compounds, see chemical compound, phenol, and ether.
182
University Chemistry
10.4.1 Structure and classification of alcohols
Similar to water, an alcohol can be pictured as having an sp3 hybridized
tetrahedral oxygen atom with nonbonding pairs of electrons occupying two
of the four sp3 hybrid orbitals. (See chemical bonding for a discussion of hybrid
orbitals.) Alkyl groups are generally bulkier than hydrogen atoms, however, so the
R―O―H bond angle in alcohols is generally larger than the 104.5° H―O―H bond
angle in water. For example, the 108.9° bond angle in methanol shows the effect of
the methyl group, which is larger than the hydrogen atom of water.
One way of classifying alcohols is based on which carbon atom is bonded to
the hydroxyl group. If this carbon is primary (1°, bonded to only one other carbon
atom), the compound is a primary alcohol. A secondary alcohol has the hydroxyl
group on a secondary (2°) carbon atom, which is bonded to two other carbon atoms.
Similarly, a tertiary alcohol has the hydroxyl group on a tertiary (3°) carbon atom,
which is bonded to three other carbons. Alcohols are referred to as allylic or benzylic
if the hydroxyl group is bonded to an allylic carbon atom (adjacent to a C=C double
bond) or a benzylic carbon atom (next to a benzene ring), respectively.
Nomenclature
As with other types of organic compounds, alcohols are named by both formal
and common systems. The most generally applicable system is that adopted at a
Introduction to Functional Groups and their Typical Reactions
183
meeting of the International Union of Pure and Applied Chemistry (IUPAC) in
Paris in 1957. Using the IUPAC system, the name for an alcohol uses the -ol suffix
with the name of the parent alkane, together with a number to give the location of
the hydroxyl group. The rules are summarized in a three-step procedure:
•
•
•
Name the longest carbon chain that contains the carbon atom bearing the
―OH group. Drop the final -e from the alkane name, and add the suffix
-ol.
Number the longest carbon chain starting at the end nearest the ―OH
group, and use the appropriate number, if necessary, to indicate the
position of the ―OH group.
Name the substituents, and give their numbers as for an alkane or alkene.
The first example below has a longest chain of six carbon atoms, so the root
name is hexanol. The ―OH group is on the third carbon atom, which is indicated
by the name 3-hexanol. There is a methyl group on carbon 3 and a chlorine atom
on carbon 2. The complete IUPAC name is 2-chloro-3-methyl-3-hexanol. The prefix
cyclo- is used for alcohols with cyclic alkyl groups. The hydroxyl group is assumed
to be on carbon 1, and the ring is numbered in the direction to give the lowest possible
numbers to the other substituents, as in, for example, 2,2-dimethylcyclopentanol.
Common names
The common name of an alcohol combines the name of the alkyl group with the
word alcohol. If the alkyl group is complex, the common name becomes awkward
and the IUPAC name should be used. Common names often incorporate obsolete
terms in the naming of the alkyl group; for example, amyl is frequently used instead
of pentyl for a five-carbon chain.
184
University Chemistry
10.4.2 Physical properties of alcohols
Most of the common alcohols are colourless liquids at room temperature.
Methyl alcohol, ethyl alcohol, and isopropyl alcohol are free-flowing liquids with
fruity odours. The higher alcohols—those containing 4 to 10 carbon atoms—are
somewhat viscous, or oily, and they have heavier fruity odours. Some of the highly
branched alcohols and many alcohols containing more than 12 carbon atoms are
solids at room temperature.
Physical properties of selected alcohols
IUPAC name
common name
formula
mp (°C)
*Ph represents the phenyl group, C6H5—.
methanol
methyl alcohol
CH3OH
−97
ethanol
ethyl alcohol
CH3CH2OH
−114
1-propanol
n-propyl alcohol
CH3CH2CH2OH
−126
2-propanol
isopropyl alcohol
(CH3)2CHOH
−89
1-butanol
n-butyl alcohol
CH3(CH2)3OH
−90
2-butanol
sec-butyl alcohol
(CH3)CH(OH)CH2CH3
−114
2-methyl-1-propanol
isobutyl alcohol
(CH3)2CHCH2OH
−108
2-methyl-2-propanol
t-butyl alcohol
(CH3)3COH
25
1-pentanol
n-pentyl alcohol
CH3(CH2)4OH
−79
3-methyl-1-butanol
isopentyl alcohol
(CH3)2CHCH2CH2OH
−117
2,2-dimethyl-1-propanol
neopentyl alcohol
(CH3)3CCH2OH
52
cyclopentanol
cyclopentyl alcohol
cyclo-C5H9OH
−19
1-hexanol
n-hexanol
CH3(CH2)5OH
−52
cyclohexanol
cyclohexyl alcohol
cyclo-C6H11OH
25
1-heptanol
n-heptyl alcohol
CH3(CH2)6OH
−34
1-octanol
n-octyl alcohol
CH3(CH2)7OH
−16
1-nonanol
n-nonyl alcohol
CH3(CH2)8OH
−6
1-decanol
n-decyl alcohol
CH3(CH2)9OH
6
2-propen-1-ol
allyl alcohol
H2C=CH−CH2OH
−129
phenylmethanol
benzyl alcohol
Ph−CH2OH*
−15
diphenylmethanol
diphenylcarbinol
Ph2CHOH*
69
triphenylmethanol
triphenylcarbinol
Ph3COH*
162
IUPAC name
bp (°C)
density (grams per millilitre)
solubility in
water
methanol
65
0.79
miscible
ethanol
78
0.79
miscible
1-propanol
97
0.80
miscible
2-propanol
82
0.79
miscible
1-butanol
118
0.81
9.1%
185
Introduction to Functional Groups and their Typical Reactions
2-butanol
100
0.81
7.7%
2-methyl-1-propanol
108
0.80
10.0%
2-methyl-2-propanol
83
0.79
miscible
1-pentanol
138
0.82
2.7%
3-methyl-1-butanol
132
0.81
2.0%
2,2-dimethyl-1-propanol
113
0.81
3.5%
cyclopentanol
141
0.95
1-hexanol
156
0.82
0.6%
cyclohexanol
162
0.96
3.6%
1-heptanol
176
0.82
0.1%
1-octanol
194
0.83
1-nonanol
214
0.83
1-decanol
233
0.83
2-propen-1-ol
97
0.86
phenylmethanol
205
1.05
diphenylmethanol
298
triphenylmethanol
380
1.20
10.5 ALDEHYDES AND KETONES
Aldehydes and ketones contain the carbonyl group. Aldehydes are considered
the most important functional group. They are often called the formyl or methanoyl
group. Aldehydes derive their name from the dehydration of alcohols. Aldehydes
contain the carbonyl group bonded to at least one hydrogen atom. Ketones contain
the carbonyl group bonded to two carbon atoms.
Aldehydes and ketones are organic compounds which incorporate a carbonyl
functional group, C=O. The carbon atom of this group has two remaining bonds
that may be occupied by hydrogen, alkyl or aryl substituents. If at least one of these
substituents is hydrogen, the compound is an aldehyde. If neither is hydrogen, the
compound is a ketone.
Naming Aldehydes
The IUPAC system of nomenclature assigns a characteristic suffix -al to
aldehydes. For example, H2C=O is methanal, more commonly called formaldehyde.
Also, there is a common method for naming aldehydes and ketones. For
aldehydes common parent chain names, similar to those used for carboxylic
acids, are used and the suffix –aldehyde is added to the end. In common names
186
University Chemistry
of aldehydes, carbon atoms near the carbonyl group are
often designated by Greek letters. The atom adjacent to
the carbonyl function is alpha, the next removed is beta
and so on.
If the aldehyde moiety
(-CHO) is attached to a
ring the suffix –carbaldehyde is added to the name of the
ring. The carbon attached to this moiety will get the #1
location number in naming the ring.
Summary of Aldehyde Nomenclature rules
•
•
•
•
Note
Since an aldehyde
carbonyl group
must always lie at
the end of a carbon
chain, it is always is
given the #1 location
position in numbering and it is not
necessary to include
it in the name. There
are several simple
carbonyl containing
compounds which
have common
names which are
retained by IUPAC.
Aldehydes take their name from their parent
alkane chains. The -e is removed from the end
and is replaced with -al.
The aldehyde funtional group is given the #1
numbering location and this number is not included in the name.
For the common name of aldehydes start with the common parent chain
name and add the suffix -aldehyde. Substituent positions are shown with
Greek letters.
When the -CHO functional group is attached to a ring the
suffix -carbaldehyde is added, and the carbon attached to that group is
C1.
Example 1
The IUPAC system names are given on top while the common name is given on
the bottom in parentheses.
Introduction to Functional Groups and their Typical Reactions
187
Aldehyde Common Names to Memorize
There are some common names that are still used and need to be memorized.
Recognizing the patterns can be helpful.
10.5.1 Naming Ketones
The IUPAC system of nomenclature assigns a characteristic suffix of -one to
ketones. A ketone carbonyl function may be located anywhere within a chain or
ring, and its position is usually given by a location number. Chain numbering
normally starts from the end nearest the carbonyl group. Very simple ketones,
such as propanone and phenylethanone do not require a locator number, since
there is only one possible site for a ketone carbonyl function. The common names
for ketones are formed by naming both alkyl groups attached to the carbonyl then
adding the suffix -ketone. The attached alkyl groups are arranged in the name
alphabetically.
Ketone Nomenclature rules
•
•
•
Ketones take their name from their parent alkane chains. The ending -e is
removed and replaced with -one.
The common name for ketones are simply the substituent groups listed
alphabetically + ketone.
Some common ketones are known by their generic names. Such as the
fact that propanone is commonly referred to as acetone.
Example 2
The IUPAC system names are given on top while the common name is given on
the bottom in parentheses.
188
University Chemistry
Ketone Common Names to Memorize
There are some common names that are still used and need to be memorized.
Recognizing the patterns can be helpful.
Naming Aldehydes and Ketones in the Same Molecule
As with many molecules with two or more functional groups, one is given
priority while the other is named as a substituent. Because aldehydes have a higher
priority than ketones, molecules which contain both functional groups are named
as aldehydes and the ketone is named as an «oxo” substituent. It is not necessary
to give the aldehyde functional group a location number, however, it is usually
necessary to give a location number to the ketone.
Example 3
Naming Dialdehydes and Diketones
For dialdehydes the location numbers for both carbonyls are omitted because
the aldehyde functional groups are expected to occupy the ends of the parent chain.
189
Introduction to Functional Groups and their Typical Reactions
The ending –dial is added to the end of the parent chain name.
Example 4
For diketones both carbonyls require a location
ending -dione or -dial is added to the end of the parent chain.
number.
The
Example 5
Naming Cyclic Ketones and Diketones
In cyclic ketones the carbonyl group is assigned location position #1, and
this number is not included in the name, unless more than one carbonyl group is
present. The rest of the ring is numbered to give substituents the lowest possible
location numbers. Remember the prefix cyclo is included before the parent chain
name to indicate that it is in a ring. As with other ketones the –e ending is replaced
with the –one to indicate the presence of a ketone.
With cycloalkanes which contain two ketones both carbonyls need to be given a
location numbers. Also, an –e is not removed from the end, but the suffix –dione is
added.
Example 6
190
University Chemistry
Naming Carbonyls and Hydroxyls in the Same Molecule
When and aldehyde or ketone is present in a molecule which also contains
an alcohol functional group the carbonyl is given nomenclature priority by the
IUPAC system. This means that the carbonyl is given the lowest possible location
number and the appropriate nomenclature suffix is included. In the case of alcohols
the OH is named as a hydroxyl substituent. However, the l in hydroxyl is generally
removed.
Example 7
Naming Carbonyls and Alkenes in the Same Molecule
When and aldehyde or ketone is present in a molecule which also contains
analkene functional group the carbonyl is given nomenclature priority by the
IUPAC system. This means that the carbonyl is given the lowest possible location
number and the appropriate nomenclature suffix is included. When carbonyls are
included with an alkene the following order is followed:
(Location number of the alkene)-(Prefix name for the longest carbon chain minus
the -ane ending)-(an -en ending to indicate the presence of an alkene)-(the location number
of the carbonyl if a ketone is present)-(either an –one or and -anal ending).
Remember that the carbonyl has priority so it should get the lowest possible
location number. Also, remember that cis/tran or E/Z nomenclature for the alkene
needs to be included if necessary.
Example 8
Introduction to Functional Groups and their Typical Reactions
191
Example 9
Additional Examples of Carbonyl Nomenclature
1) Please give the IUPAC name for each compound:
10.6 CARBOXYLIC ACIDS AND THEIR DERIVATIVES
The functional groups at the heart of this chapter are called carboxylic acid
derivatives: they include carboxylic acids themselves, carboxylates (deprotonated
carboxylic acids), amides, esters, thioesters, and acyl phosphates.
192
University Chemistry
Cyclic esters and amides are referred to as lactones and lactams, respectively.
Carboxylic acid anyhydrides and acid chlorides, which also fall under the
carboxylic acid derivative category, are not generally found in biomolecules but
are useful intermediates in laboratory synthesis. They are discussed in a section on
laboratory reactions at the end of this chapter.
Carboxylic acid derivatives can be distinguished from aldehydes and ketones
by the presence of a group containing an electronegative heteroatom - usually
oxygen, nitrogen, or sulfur – bonded directly to the carbonyl carbon. You can think
of a carboxylic acid derivative as having two sides. One side is the acyl group, which
is the carbonyl plus the attached alkyl (R) group. In the specific cases where R is a
hydrogen or methyl, chemists use the terms formyl and acetyl group, respectively.
One the other side is the heteroatom-linked group: in this text, we will sometimes
refer to this component as the ‘acyl X’ group (this, however, is not a standard term
in organic chemistry).
Introduction to Functional Groups and their Typical Reactions
193
Notice that the acyl X groups are simply deprotonated forms of other functional
groups linked to the acyl group: in an amide, for example, the acyl X group is an
amine, while in an ester the acyl X group is an alcohol.
10.7 ETHERS
Ether, any of a class of organic compounds characterized by an oxygen atom
bonded to two alkyl or aryl groups. Ethers are similar in structure to alcohols,
and both ethers and alcohols are similar in structure to water. In an alcohol one
hydrogen atom of a water molecule is replaced by an alkyl group, whereas in an
ether both hydrogen atoms are replaced by alkyl or aryl groups.
At room temperature, ethers are pleasant-smelling colourless liquids. Relative
to alcohols, ethers are generally less dense, are less soluble in water, and have
lower boiling points. They are relatively unreactive, and as a result they are useful
as solvents for fats, oils, waxes, perfumes, resins, dyes, gums, and hydrocarbons.
Vapors of certain ethers are used as insecticides, miticides, and fumigants for soil.
Ethers are also important in medicine and pharmacology, especially for use
as anesthetics. For example, ethyl ether (CH3CH2―O―CH2CH3), simply known as
ether, was first used as a surgical anesthetic in 1842. Codeine, a potent pain-relieving
drug, is the methyl ether of morphine. Because ether is highly flammable, it has
largely been replaced by less-flammable anesthetics, including nitrous oxide (N2O)
and halothane (CF3―CHClBr).
194
University Chemistry
Ethyl ether is an excellent solvent for extractions and for a wide variety
of chemical reactions. It is also used as a volatile starting fluid for diesel
engines and gasoline engines in cold weather. Dimethyl ether is used as a spray
propellant and refrigerant. Methyl t-butyl ether (MTBE) is a gasoline additive that
boosts the octane number and reduces the amount of nitrogen-oxide pollutants in
the exhaust. The ethers of ethylene glycol are used as solvents and plasticizers.
Nomenclature of ethers
Common names of ethers simply give the names of the two alkyl groups bonded
to oxygen and add the word ether. The current practice is to list the alkyl groups in
alphabetical order (t-butyl methyl ether), but older names often list the alkyl groups
in increasing order of size (methyl t-butyl ether). If just one alkyl group is described
in the name, it implies two identical groups, as in ethyl ether for diethyl ether.
Systematic (IUPAC) names for ethers use the more complex group as the
root name, with the oxygen atom and the smaller group named as an alkoxy
substituent. Examples given above are ethoxyethane (diethyl ether), methoxyethane
(methyl ethyl ether), 2-methoxy-2-methylpropane (MTBE), and phenoxybenzene
(diphenyl ether). The IUPAC nomenclature works well for compounds with
additional functional groups, because the other functional groups can be described
in the root name.
10.7.1 Physical Properties of Ethers
Ethers lack the hydroxyl groups of alcohols. Without the strongly polarized
O―H bond, ether molecules cannot engage in hydrogen bonding with each other.
Introduction to Functional Groups and their Typical Reactions
195
Ethers do have nonbonding electron pairs on their oxygen atoms, however, and they
can form hydrogen bonds with other molecules (alcohols, amines, etc.) that have
O―H or N―H bonds. The ability to form hydrogen bonds with other compounds
makes ethers particularly good solvents for a wide variety of organic compounds
and a surprisingly large number of inorganic compounds. (For more information
about hydrogen bonding, see chemical bonding: Intermolecular forces.)
Complexes of ethers with reagents
The unique properties of ethers (i.e., that they are strongly polar, with
nonbonding electron pairs but no hydroxyl group) enhance the formation and use
of many reagents. For example, Grignard reagents cannot form unless an ether is
present to share its lone pair of electrons with the magnesium atom. Complexation
of the magnesium atom stabilizes the Grignard reagent and helps to keep it in
solution.
Electron-deficient reagents are also stabilized by ethers. For example, borane (BH3)
is a useful reagent for making alcohols. Pure borane exists as its dimer, diborane
(B2H6), a toxic gas that is inconvenient and hazardous to use. Borane forms stable
complexes with ethers, however, and it is often supplied and used as its liquid
complex with tetrahydrofuran (THF). Similarly, gaseous boron trifluoride (BF3) is
more easily used as its liquid complex with diethyl ether, called BF3 etherate, rather
than as the toxic, corrosive gas.
Crown ethers are specialized cyclic polyethers that surround specific metal ions
to form crown-shaped cyclic complexes. They are named by using the parent
196
University Chemistry
name crown preceded by a number describing the size of the ring and followed
by the number of oxygen atoms in the ring. In the crown-ether complex, the metal
ion fits into the cavity of the crown ether and is solvated by the oxygen atoms.
The exterior of the complex is nonpolar, masked by the alkyl groups of the crown
ether. Many inorganic salts can be made soluble in nonpolar organic solvents by
complexing them with an appropriate crown ether.
Potassium ions (K+) are complexed by 18-crown-6 (an 18-membered ring with
6 oxygen atoms), sodium ions (Na+) by 15-crown-5 (15-membered ring, 5 oxygens),
and lithium ions (Li+) by 12-crown-4 (12-membered ring, 4 oxygens).
In each of these crown-ether complexes, only the cation is solvated by the crown
ether. In a nonpolar solvent, the anion is not solvated but is dragged into solution
by the cation. These “bare” anions in nonpolar solvents can be much more reactive
than they are in polar solvents that solvate and shield the anion. For example, the
18-crown-6 complex of potassium permanganate, KMnO4, dissolves in benzene to
give “purple benzene,” with a bare MnO4− ion acting as a powerful oxidizing agent.
Similarly, the bare −OH ion in sodium hydroxide (NaOH), made soluble in hexane
(C6H14) by 15-crown-5, is a more powerful base and nucleophile than it is when
solvated by polar solvents such as water or an alcohol.
Synthesis of ethers
Williamson ether synthesis
The most versatile method for making ethers is the Williamson ether synthesis,
named for English chemist Alexander Williamson, who devised the method in
the 19th century. It uses an alkoxide ion to attack an alkyl halide, substituting the
alkoxy (―O―R) group for the halide. The alkyl halide must be unhindered (usually
primary), or elimination will compete with the desired substitution.
Introduction to Functional Groups and their Typical Reactions
197
10.8 AMINES
Amines are one of the most important classes of organic compounds which can
be derived when we replace one or more hydrogen atoms of ammonia molecules
with an alkyl group. An amine is generally a functional group with a nitrogen atom
having a lone pair. Amines resemble ammonia structurally where nitrogen can bond
up to 3 hydrogen atoms. It is also characterized by various properties that are based
on carbon connectivity. Compounds of nitrogen connected to a carbonyl group are
called as amides, they have a structure R–CO–NR′R″ and vary in properties with
amines.
Amines are organic compounds that contain nitrogen atoms with a lone pair.
Basically, they are derived from ammonia (NH3) in which one or more hydrogen
atom is replaced by an alkyl or aryl group, and so they are known as alkylamines
and arylamines respectively.
198
University Chemistry
10.8.1 Amine Structure
Nitrogen has 5 valence electrons and so is trivalent with a lone pair. As
per VSEPR theory, nitrogen present in amines is sp3 hybridized and due to the
presence of lone pair, it is pyramidal instead of tetrahedral shape which is a general
structure for most sp3 hybridized molecules. Each of the three sp3 hybridized
orbitals of nitrogen overlap with orbitals of hydrogen or carbon depending upon
the configuration of amines. Due to the presence of lone pair, the C-N-H angle
in amines is less than 109 degrees which is a characteristic angle of tetrahedral
geometry. The angle of amines is near about 107 degrees.
Occurrence of Amines
Naturally, amines occur in proteins, vitamins, hormones, etc. and they are also
prepared synthetically to make polymers, drugs, and dyes.
10.8.2 Types of Amines
On the basis of how the hydrogen atoms are replaced by an ammonia
molecule, amines can be divided into 4 types.
1.
Primary Amines
When one of the hydrogen atoms of the ammonia molecule is replaced by an
alkyl or aryl group.
Eg: Methylamine CH3NH2, Aniline C6H5NH2
2.
Secondary Amines
Two organic substituents replace the hydrogen atoms of the ammonia molecule
forming an amine.
Eg: Dimethylamine (CH3)2NH, Diphenylamine (C6H5)2NH
Introduction to Functional Groups and their Typical Reactions
3.
199
Tertiary Amines
When all 3 of the hydrogen atoms are replaced by an organic substituent, it
could be an aryl or aromatic group.
Eg: Trimethylamine N(CH3)3, Ethylenediaminetetraacetic acid (EDTA)
4.
Cyclic Amines
These are secondary or tertiary amines in an aromatic ring structure. Eg:
Piperidine (CH2)5NH, Aziridines C2H5N
10.8.3 Preparation of Amines
Some processes for preparing primary amines are mentioned below.
Preparation of Primary Amines
1. Making of amines from halogenoalkanes
This process will be carried out in a sealed tube. Here haloalkanes will be
heated with the concentrated solution of ammonia in ethanol. The mixture cannot
be heated under the reflux as ammonia would move out in the form of gas from a
container.
Now coming to the preparation of primary amine from halogenoalkane the
reaction takes place in two stages. Salt will be formed at the first stage. Here ethyl
ammonium bromide is the salt. It is similar to ammonium bromide except for the
fact that one of the hydrogens in the ammonium atom is replaced by an ethyl group.
A reverse reaction can occur between ammonia and the salt. It is illustrated in
the above reaction.
2. Reduction of nitriles
We can get primary amines when nitriles are reduced with lithium aluminium
hydride. This method is mainly used for the preparation of amines which contain
one carbon atom more than the starting amine.
3. Gabriel phthalimide synthesis
We can get primary amines easily by Gabriel synthesis. In this process, on the
treatment of phthalimide with ethanolic potassium hydroxide, we get potassium
200
University Chemistry
salts of phthalimide. When this is further heated with alkyl halide followed by
alkaline hydrolysis then primary amine is produced. We cannot prepare aromatic
primary amines because aryl halides do not undergo nucleophilic substitution with
the anion which is formed by phthalimide.
10.8.4 Basicity of Amines
Similar to ammonia, primary & secondary amines have protic hydrogens and
thus they showcase a degree of acidity. Whereas tertiary amines have no protic
hydrogen and thus do not possess a degree of acidity.
pKa value for primary & secondary amines is about 38, which makes them a real
weak acid. Whereas if we take the pKb, it is about 4. This makes the amines much
more basic than acidic. Thus, an aqueous solution of an amine is strongly alkaline.
10.8.5 Uses of Amines
Amines have a widespread application in our daily lives. Some uses of
amines are listed below:
•
•
•
•
It is used in water purification, medicine manufacturing and development
of insecticides and pesticides.
It is involved in the production of amino acids which is the building block
of proteins in living beings. Many varieties of vitamins are also made by
amines.
Serotonin is an important amine that functions as one of the primary
neurotransmitters. It controls the feelings of hunger and is critical for the
speed with which the brain operates in general.
Pain-relieving medicines such as Morphine and Demerol which are also
known as analgesics are made from amines.
EXERCISE
Answer the following questions
1. Describe the functional group of alkane alkene and alkyne.
2. What is the reactivity order of alkane’s alkenes and alkynes?
3. Defines an aromatic compound.
4. Why is it called aromatic compound?
5. Explain the characteristics of an aromatic compound.
Introduction to Functional Groups and their Typical Reactions
6.
7.
8.
9.
201
What are examples of aldehydes and ketones?
What is the difference structurally between aldehydes and ketones?
How each of the acid derivatives is produced from carboxylic acid?
What are the examples of ethers?
MULTIPLE CHOICE QUESTIONS
Tick the correct answer:
1.
An alcohol has a higher boiling point than an alkane with the same length
carbon chain. True or false?
a. True
b. False
2.
3.
What is the name of this compound?
a. Propanol
b. Ethoxymethane
c. Methoxyethane
d. Propanone
Which of these compounds is propanoic acid?
a.
b.
c.
4.
d.
Which of these compounds is a ketone?
a.
b.
c.
202
University Chemistry
d.
5.
6.
What is the name of this compound?
a. Propane
b. Propanone
c. Propanal
d. Propanoic acid
Which of the following compounds is an ester?
a.
b.
c.
7.
8.
9.
d.
Butanal contains which one of the following functional groups?
a. Carboxyl group
b. Carbonyl group
c. Hydroxyl group
d. Phenyl group
Esters have a higher boiling point than their equivalent carboxylic acids.
True or false?
a. True
b. False
Which of the following compounds is a secondary amine?
a.
b.
203
Introduction to Functional Groups and their Typical Reactions
c.
d.
10.
e.
Which of the following families of organic compound is the least soluble in
water?
a. The ethers
b. The alcohols
c. The carboxylic acids
d. The primary amines
ANSWERS
1. (a)
2. (c)
3. (d)
4. (a)
5. (s)
6. (c)
7. (b)
8. (b)
9. (d)
10. (a)
REFERENCES
1.
2.
3.
4.
5.
6.
7.
8.
Allen, K. N., & Dunaway-Mariano, D. (2004). Phosphoryl group transfer:
evolution of a catalytic scaffold. Trends in Biochemical Sciences, 29(9), 495–503.
Brown, Theodore (2002). Chemistry: the central science. Upper Saddle River,
NJ: Prentice Hall. p. 1001. ISBN 0130669970.
Fushinobu, S., Nishimasu, H., Hattori, D., Song, H.-J., & Wakagi, T. (2011).
Structural basis for the bifunctionality of fructose-1,6-bisphosphate aldolase/
phosphatase. Nature, 478(7370), 538– 541.
Kirchberg, K., Kim, T.-Y., Haase, S. & Alexiev, U. (2010). Functional interaction
structures of the phochromic retinal protein rhodopsin. Photochemical &
Photobiological Sciences, 9, 226-233.
Kitzing, K., Auweter, S., Amrhein, N. & Macheroux, P. (2004). Mechanism of
Chorismate Synthase. The Journal of Biological Chemistry, 279(5), 9451-9461.
Lietzan, A. D., Lin, Y. & St. Maurice, M. (2014). The role of biotin and oxamate
in the carboxyltransferase reaction of pyruvate carboxylase. Archives of
Biochemistry and Biophysics, 562, 70-79.
McMurry, J. E. & Bagley, T. P. (2005). The Organic Chemistry of Biological
Pathways. Robeerts and Company Publishers.
Yuan, H. & Marmorstein, R. (2012). Structural basis for sirtuin activity and
inhibition. Journal of Biological Chemistry, 287(14), 42428-42435.
INDEX
Acid-base reaction 54
Activation energy 48
Activity series 59, 60, 62
alkene reaction 176
alkyne functional groups 177
argon 32
Aromatic compounds 177, 178
aromatic units 177
Arrhenius definition 55, 56
atom 87, 94, 95, 96, 98, 99, 100, 103, 104,
105
atomic hydrogen 87
atomic number 32, 35, 36
Atoms 31, 35, 36, 37, 38, 41, 46
Aufbau principle 100, 101
B
Back titration 83
Balanced chemical equation 65
Balance oxygen atom 61
Beer–Lambert law 81
Benzene 177
biology 35
Blackbody Radiation 90, 91
blocks 87, 93, 94
Brönsted Definition 56
C
Californium 33
carbon dioxide 32
Chemical bonding 107
chemical change 4, 6, 27
Chemical changes 72
Chemical composition 72
Chemical compound 107, 108
Chemical equation 66, 72
Chemical equations 48
Chemical equilibrium 151, 152, 155, 156
Chemical Mixtures 8, 9
chemical property 1, 5, 6
Chemical reaction 47, 49, 50, 51, 52, 54,
57, 68
Chemical species 65
chemical substance 8
Chemical substances 8
Chemical transformation 47
chemistry 8
chemists 174, 192
Chemists 173
Colligative property 142, 144, 145, 148
combustion of magnesium metal 4
Complex reaction systems 68
Coulomb’s Law 109
Covalent bond 111
Crystalline lattice 108, 110
Crystalline solids 53
Crystallization 134, 137
cutting 3
D
denting 3
E
electromagnetic radiation 89, 90, 92, 96,
97
electron-attracting properties 174
Electronegativity 108, 112
Electronic configuration 113
electronic properties 174, 176
Electron pair 111, 112, 113, 114, 115, 118,
206
University Chemistry
119, 120, 123, 124
electrons 31, 35, 36, 37, 42
Electron transfer (ET) 57
Electrostatic attraction 108, 111, 124
Electrostatic bond 108
Energy change 110
Equilibrium constant 151, 153, 154, 156,
157, 158, 159, 160, 161, 164, 170, 171
Equilibrium reaction 155, 162
Equilibrium state 153, 154, 155, 167, 169,
170
erbium 33
F
forms of combustion 6
Fourier transform spectroscopy 81
functional group 173, 175, 176, 177, 179,
181, 185, 186, 188, 190, 197, 200
G
Gases 31, 32
Gas molecules 135, 136, 139, 140
Gas phase titration 81
Gas solubility 135, 136
geology 35
gold 32, 33, 44
greater density 2
groups 87, 103
H
Half-reactions 57, 60, 61, 62
Heterogeneous Mixtures 9
Homogenous Mixtures 9
hydrocarbon 175, 177, 178, 181
Hydrochloric 54, 55
Hydrogen gas 48, 59
I
intensive properties 2, 16
Ionic bonds 108, 109, 111
Ionic crystal 108
IUPAC nomenclature. 180
K
kinetic energy 92, 93
L
light 87, 88, 89, 90, 91, 92, 93, 94, 95, 96,
97, 104, 105
Limiting reagent 70
Liquids 31, 32
M
many-electron atoms 87, 99
mass 1, 2, 4, 7, 14, 15, 19, 24, 25, 26, 27, 31,
35, 36, 37, 38, 39, 40, 41, 44
materials science 35
Matter 31
Mechanical equilibrium 155
metals 5, 7
minor discomfort 7
mixture 32, 37, 39, 40, 44, 45
Molality 130, 131, 132, 149
Molar concentration 75, 128, 130, 133
Molarity 128, 129, 130
Molecular entity 76
Molecular mass 73
molecule 174, 175, 177, 179, 190, 193, 198
molecules 31, 36, 38, 39, 40, 41, 42, 43,
173, 175, 176, 177, 188, 194, 197, 198
Molecules 175, 176
Money 10, 12
N
Net ionic equation 51, 52, 53
Neutralization 56
neutrons 31, 35, 36, 37
nitric acid 1, 6
nitrogen 32, 33, 34
Nitrogen atoms 49
Non-soluble solid 83
Nuclear chemistry 47
Nuclear Substituted Compounds 179
O
orbital energies 87, 98
Organic compounds 175
organometallic compounds 8
Osmotic pressure 142, 147
oxygen 4, 5, 6, 8, 25, 26, 32, 33, 34, 39, 40,
41, 42
207
Index
P
particles 31, 32, 36, 37, 44, 87, 88, 90, 93,
97, 104, 105
Pauli exclusion principle 99
periodicity 87, 102, 103, 104
periodic table 87, 99, 101, 103, 104
periods 87
Photoelectric Effect 92, 93
physical and chemical properties 1
physical change 3
Physical deformation 3
Physical properties 1, 2
Physical quantities 10
Potassium 33
Precipitation Reactions 51, 53
protons 31, 32, 35, 36, 37, 42
pure substance 32
Q
Spectator ion 52
spectrum 87, 89, 90, 91, 95, 96, 97
Stoichiometric calculations 74
Stoichiometric chemical reaction 66
Stoichiometric coefficient 67
Stoichiometric factors 66
Stoichiometric number 67
Stoichiometry 65, 66, 67, 73, 74
stretching 3
study of chemistry 7
Sulfuric 54
T
terbium 33
Tetrahydrofuran (THF) 129
Thermal equilibrium 155
Thermodynamic relationships 152
Titration 78, 79
Titration curve 78
quantization 87, 91, 92, 93, 104
R
Raoult’s law 143, 144, 145, 146
Reaction mechanism 48
reactivity 173, 176, 200
Redox reaction 57
S
silver 32, 33
Solids 31
Solubility 132, 133, 134, 135, 136, 137,
138, 139, 140
V
vanadium (V) 33
vitamin A 176
volume 1, 2, 7, 15, 18
W
water 32, 39, 41, 42, 45
waves 87, 88, 89, 91, 97
Wax undergoes 5
Y
ytterbium 33
yttrium 33, 34