GENERAL CHEMISTRY WEEK 1
Introduction to Chemistry
(Chemistry and Engineering)
Chemistry is the study of substances, their composition,
and the changes they undergo. A deep understanding of
chemistry is beneficial for an engineer in choosing the
design of an innovative process. a good example is the
selection of materials –its availability, effects on the
environment, external factors that affect the material such
as temperature, light, time, etc.
The Study of Chemistry
The importance of Chemistry to different scientific fields
and studies gave it the coined term “central science.”
Therefore, it is most likely that you have seen chemistry
even if you haven’t had any chemistry course. This course
will help you gain an in-depth understanding of chemistry
and its relation to the natural world. The appreciation of
the chemical viewpoint helped engineers devise strategies
in approaching the problems encountered in the
application of chemistry in innovation.
These viewpoints are classified into three levels of
understanding:
Macroscopic –immediately seen in substances and their
reactions.
Microscopic – focuses on the smallest unit of the system.
Symbolic Perspective –allows the communication of
concepts efficiently.
The Macroscopic Perspective
The chemical reactions seen in our environment are being
observed at the macroscopic level.
Matter –anything observable, occupies space, and has mass
– is constantly in contact with us that its existence can
easily be proved through our intuitive feel.
The observable changes in matter are classified into two.
Physical change does not affect a substance’s chemical
properties, thus maintaining its composition.
Chemical change, on the other hand, affects the chemical
structure of a substance and thus, produces another
substance. These changes are greatly affected by the
properties of the material.
Physical properties of matter are typically perceivable
through the five senses, e.g., taste, odor, texture, color, and
state (solid, liquid, or gas). Other properties might need
instrumentation to observe, like volume, mass, density,
melting point, boiling point, etc.
Chemical properties of matter refer to the ability to
produce another substance through different chemical
reactions. Corrosivity, combustibility, flammability, and
reactivity are examples of the chemical properties of
matter.
The Microscopic or Particulate Perspective
All matter consists of atoms and molecules. It is the most
fundamental concept in chemistry. Therefore, chemists
consider everything as “a chemical” in a way. Most of the
time, the matter we encounter is a diverse chemical
mixture, and each component is a chemical substance.
Elements are the building blocks of all substances of
matter. Each element is made up of
atoms –particles that cannot be made any smaller but still
behave like a chemical system.
Molecules are group of atoms held together by a force
called “chemical bonds”.
All matter exists as either a mixture or pure substance. Pure
substances will always have the same composition while
mixtures are composed of variable substances as
components and can be separated into pure substances.
Pure substances are either elements or compounds.
Elements are substances having only one type of atom.
Compounds are substances having two or more atoms of
different elements.
Mixtures occur as either homogeneous (uniform in
appearance) or heterogeneous (different distinguishable
parts). A homogeneous mixture is often referred to as
solution. Heterogeneous mixture is classified into either a
suspension or colloid.
Symbolic Representation
Symbolic representation uses symbols to represent
atoms, molecules, composition, and reaction. This
viewpoint gives way for the discussion of most of the
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abstract aspects of chemistry. This level of understanding is
fundamental in the interaction of ideas at the particulate
level.
Numbers and Measurements
Observation is significant in scientific processes. A
qualitative observation is made by looking at the
perceivable properties using the five senses -color, taste,
odor, etc. A quantitative observation is made by looking
at the measurable properties of a substance, thus the
term measurement.
Numbers and Significant Figures
Measurements could either be very small or very large. In
these cases, scientific notation is helpful. Scientific Notation
factors out all powers of ten and writes separately. Large
numbers use positive, while small numbers use negative
powers of ten. Significant figures give a reliable amount
of information derived from observations.
Prefixes Used in SI System
Measurement is an integral part of our daily lives. Most of
the processes we carry throughout the day are measured.
Utilities like water and electricity use measurements to
provide services. The construction sector also uses
measurement to carry out projects and plans for
innovation and development.
Measurements are made with the use of measuring
devices. Each measurement consists of a number and
unit.
Units
Comparative Scale of Macro, Micro, and Nanoscale
Unit defines the scale or standard used to represent the
results
of
a measurement. Units are used to
determine the standards followed relevant to other
measurements. Units are always determined and uniform.
This is to provide stability across the system. If each
scientist has their own way of determining units of
measurement, it will be complete chaos.
Standard systems are present and used around the world.
However, there are different systems adopted. Two of
the most widely used are the English system -used in the
United States and Metric System -preferred in the industrial
sector. In scientific fields, the metric system is the unit of
choice. A comprehensive system of units was established in
the 1960s. This is the International System (SI) -le Système
Internationale in French.
READING MATERIALS
GENERAL CHEMISTRY CM1: MATTER AND ITS PROPERTIES
Phases of Matter
As we recall our high school chemistry, matter is anything
that can occupy space and has mass. Matter can undergo
phase changes, which is an example of physical change.
Phases of matter, namely solid, liquid and gas, can be
differentiated based on these properties:
Solids: with fixed shape due to particles being closed
together. They are incompressible and slightly expand
when heated
Liquids: have particles that are moving slightly in random
motion. Liquids are slightly compressible, can flow, and
copies the shape of their container
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Gas: have particles moving randomly and far apart. Gases
have low density, can flow and copies the shape of their
container, just like liquids
Particle: an entity that comprises matter. These are
molecules, atoms and respective sub-atomic particles.
GENERAL CHEMISTRY CM2: CLASSIFICATION OF MATTER
Mixtures
Mixtures are defined as a physical combination of matter,
which the matter components can be distinguish
individually (Ilao et. al, 2016). Mixtures can be separated
easily. We will be discussed later on different separation
techniques in this Course Material.
Properties of Matter
Mixtures can be classified based on two categories:
1. Physical property: can be observed and measured
without changing the identity and composition of the
substances
Example: freezing point, color, state/phase of matter.
2. Chemical property: ability of a substances to undergo
changes, which produce another substance
Example: flammability.
Based on these following properties, which does not belong
to the group:
· Luster (ability to reflect light)
· Combustibility (ability to burst into flames due to a
chemical reaction).
· Rusting (ability to form rust).
3.Extensive property: property that depends on the
amount of matter present.
4. Intensive property: opposite of extensive property; does
not depend on the amount of matter, but on the nature of
the matter
Here are some examples of extensive and intensive
properties:
Extensive Property
Intensive Property
· Mass
· Boiling point
· Volume
· Melting point
· Amount of energy in a
· Conductivity (ability to
substance
conduct electricity)
· Heat capacity.
· Density (mass per volume)
Key points:
Density is the ratio of mass and volume, two of the
extensive property mentioned. But, density is unique to
every substance and stays constant even the amount of
substance is increased. Therefore, density is intensive
property.
Heat capacity describes as the amount of energy required
to increase the temperature of a certain substance by 1
Kelvin. It is extensive property because heat capacity
increases as the amount of matter increases.
Heterogenous mixture: the compositions of these mixtures
are not uniform, and can be observed with our eyes or with
the use of microscope. Such example of these mixtures are
the food dishes we eat, milk and blood.
Homogenous mixture: the composition of these mixture is
uniform in appearance. This uniformity is due to the
molecular/atomic level size of the individual components,
which is very small to be seen.
Usual examples of homogenous mixtures are solutions with
water as solvent, which are called “aqueous solutions.”
Another example of homogenous of mixture is air, which is
composed of mixture of different gases and particles.
In solution, we describe the following:
Solute: the substance that was dissolved by the solvent.
Solvent: substance that dissolves the solute.
Separation Techniques for Mixtures
One of the properties of mixtures is their components can
be easily separated to each other. Here are some of the
separatory techniques used in the laboratory to separate
mixtures:
1.Evaporation: used to separate solutions/homogenous
mixtures, given there is a soluble solid (solute) from a liquid.
Liquid is evaporated in this process, leaving the solid
(solute) in the set-up
2. Filtration: used to separate heterogenous mixtures,
which composes an insoluble solid from a liquid.
3.Chromatography: used for separating different
substances with different solubility to a given solvent
(examples are pigments and dyes). Based on figure 2, some
substances travel further up the paper (chromatogram)
because they are more soluble to the solvent.
4. Simple Distillation: used for separating mixture with
different boiling points or separating volatile liquid from
non-volatile liquid (Ilao et. al, 2016). Liquid components are
separated by evaporating and condensing the volatile liquid
(distillate) in the set-up.
5. Decantation: used for separating heterogenous mixtures
with different density. Separation of mixtures occurs with
the help of gravity (Berk, 2009). For liquids, decantation
occurs with the help of separatory funnel.
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6. Centrifugation: like decantation, this is used for
separating heterogenous mixtures with different density,
but with the help of centrifugal forces created by the
centrifuge machine (Berk, 2009). This is used to separate
blood components.
the measurement values and their calculations. Significant
figures in the measurement reading consist of all
known/certain digit plus one uncertain/estimated digit.
Pure Substances - Unlike mixture, pure substances can be
separated by chemical means. Pure substances have fixed
and characteristic elemental composition and properties.
Types of pure substance:
1. All nonzero digits are ALWAYS significant.
Examples:
- 274 -- 3 sig. fig.
- 25.632 -- 5 sig. fig.
Element: simplest type of matter, which consist of one type
of atom only. Elements can be classified into the following.
➢ Metals: usually shiny, malleable substance and
good conductor of electricity and heat. Metals
usually gives electrons to non-metals when
forming ionic compounds.
➢ Non-metals: usually gases or brittle solids. Nonmetals are poor conductor of electricity and heat.
Most non-metals gain electrons from metals when
forming ionic compounds, and tend to share
electrons when forming covalent compounds.
➢ Metalloids: have properties of metals and nonmetals. Metalloids in periodic table can be found at
the staircase line.
2. All zeros between significant digits (captive) are
ALWAYS significant.
Example:
- 1.008 has 4 significant figures.
Compounds: pure substances, which composed of two or
more elements combined in fixed parts by mass.
Compounds have different properties from its element
components.
Rules in Evaluating Significant Figures.
3. Leading zeroes before all the nonzero digits DO NOT
count as significant figures.
Example:
- 0.0025 mL has 2 s.f.
4. All FINAL/TRAILING zeros to the right of the decimal ARE
significant.
Example:
- 105.0020 has 7 s.f.
- 100 has 1 s.f.
5.All zeros that act as place holders are NOT significant.
Example:
- 320 (place holder: ones place)
• 0.025 (place holder: ones, tenths place)
➢ Ionic compound: composed of opposite charged
ions. Bonds formed within this compound is due to
the give and take of electrons
➢ Covalent compounds: composed of elements with
bond formed due to sharing of electrons. Usually
composed of non-metal elements
Ions: composed of one (monoatomic) or more (polyatomic)
elements, which exhibit charges due to the excess or lack of
electrons.
➢ Cations: gives electrons to the anions; positive
charge ions.
➢ Anions: gains electron from cations; negatively
charge
GENERAL CHEMISTRY CM3: SIGNIFICANT FIGURES AND
SCIENTIFIC NOTATION
Significant Figures
Different measurement devices have different levels of
uncertainty due to their limitation of their measurement
capabilities. In science, we compensate with the different
accuracies by noting significant figures (sig. fig. or s.f.) of
Fig. 1: Types of zeroes in a certain number.
Rules in Finding Significant Figures in Calculations
What is the sum of 20.5, 200.5,
9.756 and 9.00009?
If our final answer should be 239.5,
what rule/s regarding significant
figures can we infer from the
calculation?
1. Rule for Addition and Subtraction Calculations: Round
the calculated answer so that it contains the same number
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of decimal places as the measurement with the least
number of decimal places (LNDP)
3. In converting back from scientific notation to decimal
notation/number, you need to reverse back the movement
of decimal places.
What is the area of the rectangle with 14 units width and
25.056 length?
Solution:
𝐴𝑟𝑒𝑎=𝑙 ×𝑤
𝐴𝑟𝑒𝑎=14 𝑢 ×25.056 𝑢
𝐴𝑟𝑒𝑎=350.784 𝑠𝑞𝑢𝑎𝑟𝑒 𝑢𝑛𝑖𝑡𝑠
≈350 𝑠𝑞𝑢𝑎𝑟𝑒 𝑢𝑛𝑖𝑡𝑠 (2 sig fig)
When rounding off numbers:
· If the digit following the last significant digit is 0 to 4,
drop the non-significant digit/s on the right and leave
the remaining digit unchanged.
· If the digit following the last significant digit is 5 to 9,
round up the last significant digit and drop the nonsignificant digit/s on the right.
2. Rule for Multiplication and Division Calculations:
Round the calculated answer so that it contains the same
number of significant figures as the measurement with
the least number of significant figures (LNSF).
Note: In considering significant figures always round up
numbers after calculations, not in between calculations.
Scientific Notation
In science, some values can be too large to be written out.
For example, one mole of carbon atom is equal to 602 000
000 000 000 000 000 000 carbon atoms. Such long values
need to be shortened with the use of exponential/
scientific notation.
Guidelines:
1. Move decimal point until there is a single digit to its left.
Include all significant digits based on the number so
significant values
Post Test:
How many significant figures are there in the following
values or final answers for calculations?
1. 7.5000 m
– 5 sig. fig.
2. 0.0040 km
– 2 sig. fig.
3. 10.0340 g + 0.003874g – 6 sig. fig.
4. 23.567 moles ÷ 1.26 L
– 3 sig. fig
Convert the following from numbers/decimal to scientific
notation and vice-versa.
1. 2,400,000 ug
– 2.4 x 106 ug
2. 0.00256 kg
– 2.56 x 10-3 kg
-5
3. 7 x 10 km
– 0.00007 km
4. 6.2 x 104 mm
– 62,000 mm
GENERAL CHEMISTRY CM3: THE LAWS OF MATTER AND
THE WORLD OF ATOMS
Laws of Matter and the Atomic Theory
John Dalton first introduces the Atomic Theory in 18031807. This theory will give us insight on how atoms and
elements interact in the chemical reaction.
Natural Laws exist already in the universe. It’s for us
scientist to observe these laws.
Dalton’s Atomic Theory
2. Add “x 10n” where n will represent how many time you
move the decimal places until you reach a number with
single digit on the left. Take note that the value of n will be
positive if you move the decimal point to the left, while
value of n will be negative if you move the decimal to the
right.
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From here, we can identify different laws of matter, which
helps Dalton formulate the atomic theory:
1. Law of Conservation of Mass: states that “matter can
be neither created or destroyed by chemical means.” In a
chemical reaction, the total mass of the substances before
reaction is equal to the total mass of the substances
formed after the reaction.
2. Law of Definite Proportion: states that in a given
compound, there’s always be the same proportions/ratio
by mass of its elemental components.
atoms), and mass number (A) is the sum of protons and
neutrons in a certain element
Fig. 4: Atomic
notation and how
they were
determined.
Isotopes of a certain element have the same atomic
number, but different mass number- basically, different
number of neutrons.
Most elements in the periodic table have naturally
occurring isotope. For example, carbon has three naturally
occurring isotopes, Carbon-12, Carbon-13 and Carbon-14.
Isotopes almost have an identical chemical behavior.
Fig. 2: Image of John Dalton appeared at the down left corner.
3. Law of Multiple Proportion: “If two elements A and B
combine to form more than one compound, the masses of
B that can combine with a given mass of A are in the ratio
of small whole numbers.” This law is best shown with
compounds that contain same elements, but different
subscripts. For example, water (H2O) and hydrogen
peroxide (H2O2).
View of the Atomic Structure
Dalton first describes atoms as indestructible part of
matter. But later advancement in the studies of atom prove
Dalton’s notion of indestructible atom wrong due to the
discovery of the sub-atomic particles.
To accommodate the naturally occurring isotopes, we need
to determine the relative abundance of each isotopes in the
element’s mass. This derived mass is called atomic
weight/mass (amu), which is stated in our periodic table.
Example: Carbon has two stable isotopes. Compute for the
average atomic mass of carbon.
Answer: 12.0110 amu
This value can be also seen in your periodic table.
Solve: Determine the Average Atomic Mass (in amu) of
Nitrogen with the given isotopes and their relative
abundance.
Two of the sub-atomic particles, protons and neutron, is
located in the atom’s nucleus. The mass of the atom is
concentrated to its nucleus (mass number in the periodic
table is the sum of number of protons and neutrons).
Electrons, on the other hand, are circling around the
nucleus of the atom. Electrons maintain their movement
around the nucleus with the constant attraction from the
protons in the nucleus.
Usually, elements have no charge (neutral) due to the
equal number of protons to electrons.
Elements and their Isotopes
In the periodic table, each element is arranged in
increasing atomic number. Based on figure 4, atomic
number (Z) pertains to the number of protons of a certain
element (and number of protons if neutral charged
GENERAL CHEMISTRY CM5: ELEMENTS AND PERIODIC
TRENDS
From protons and neutron, we will focus more to the
negative charged subatomic particle, electron. As you will
observe in our discussion, the difference in the electron
arrangement in atoms can be correlated to the different
physical and chemical characteristics of the elements.
We can determine the maximum number of electrons per
principal energy level using the “2n2” rule.
Ex. If n=4, the maximum electron for 4th energy level is
2(4)2= 32 electrons
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Electron Configuration
1. Principal Energy Level
(n): pertains to the
distance of electron to the
nucleus. The higher the
principal energy level, the
farther the electron to the
nucleus. In the periodic
table, this refers to the
period number.
2. Valence Electrons:
electrons that are in the
outermost shell and
involved in the formation
of bonds in a compound.
3. Sublevel orbitals (l): this describes the shape of the
electron orbital (denoted as “s”, “p”, “d”, “f”). There are
limited number of electrons occupying in a sublevel.
s-orbital: 2 ed-orbital:10 ep-orbital: 6 ef-orbital:14 eElectron configuration of different elements can be
represented in three ways. Two of these are given in Table 1:
1. Full Electron Configuration: configuration with fully
written energy levels (n), sublevels (l) and number of
electrons (represented in superscript).
2. Condensed Electron/ Noble Gas Configuration:
shortened electron configuration, which some part of the
configuration is replaced by the respective noble gas. Take
note that the replaced configuration should be
corresponding to the given noble gas.
Atomic number of phosphorous is 15, therefore number of
electrons are also 15 (neutral atom).
2. Second, take note of the sequence of the sublevel based
on the diagonal rule and the number of electrons that
each sublevel can accommodate. Electrons are placed as a
superscript of orbitals. The total value of superscripts
must be equal to the number of electrons of the given
element.
Maximum number of electrons per orbital:
s-orbital: 2 ed-orbital:10 ep-orbital: 6 ef-orbital:14 eExample: Phosphorous (Atomic number:15)
Following the diagonal rule (Check the total of
superscripts.) 1s2 2s2 2p6 3s2 3p3
3. You may shorten the configuration using noble gas
configuration with the corresponding noble gas before the
element. Replaced the corresponding noble gas to its
electron configuration.
Example: Phosphorous (Atomic number:15)
Full electron configuration: [1s2 2s2 2p6]3s2 3p3
Noble gas configuration: [Ne]3s2 3p3
Atomic Properties Trends in Periodic Table
Atoms in the periodic table may manifest different
characteristics due to the attraction and repulsion
(electrostatic interaction) of different sub-atomic particles
within the atom). These attraction and repulsion effects are
the following:
1. Effective nuclear charge attraction: attraction of
electron towards the nucleus due to the protons. The higher
the number of protons, the greater the effect of the
attraction.
2. Electron shielding effect: this describes the weakening of
the effective nuclear charge farther the nucleus, which will
heighten the repulsion of the core (inner) and
valence/outer electrons. Atomic Properties and Trends:
Writing Electron Configuration
The writing for the full electron configuration is based on
the diagonal line rule (see figure 1).
Fig. 1: Diagonal rule for writing electron configuration.
Example: Write the full electron configuration of
Phosphorous (neutral atom).
1. First, take note of the number of electrons of
phosphorous.
1. Atomic Size (Atomic Radius): atomic size is measured
based by the distance of two bonded molecule with the
same element or two adjacent atoms.
Trend: · Increasing from top to bottom of periodic table due
to the increasing energy levels. Increasing energy level will
result to larger atomic radius (increased energy shell).
2. Ionization Energy: pertain to the energy needed to
remove an electron from the atom. The greater the
ionization energy, the harder to remove electron from its
atom (Brown et. al.,2012).
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Trend (for first ionization energy):
· Increasing from left to right of periodic table due to the
increasing effective nuclear charge, making the electrons
harder to remove (requires high energy to remove).
· Decreasing from top to bottom of periodic table due to the
weak effective nuclear charge and stronger electron
shielding in increasing energy levels. This makes the outer
electron easier to remove, which means lower ionization
energy.
3. Electron Affinity: this characteristic refers to the energy
release/absorbed when accepting electron to form
negative ion. The more energy is released, the more stable
the ion formed).
Trend: · Increasing magnitude from left to right of periodic
table due to the increasing effective nuclear charge. Atoms
with higher effective nuclear charge can make a stable
negative ion when electron was added.
· Decreasing magnitude from top to bottom of periodic
table because larger atoms have weaker effective nuclear
charge, which decreases its tendency to accept electrons.
For the reported values of electron affinity:
· negative value means that energy was released.
· positive value when energy was absorbed/needed
during the reaction. Take note that in electron affinity,
electron is added to atom. While in ionization energy,
electron is being removed.
4. Electronegativity: measures the ability of the element to
attract two electrons when bonded to another atom. The
higher the value, the more the atom can attract pair of
electrons. This property is relevant in bonding formation.
Trend: · Increasing from left to right of periodic table due to
the increasing effective nuclear charge. High effective
nuclear charge can highly attract the pair of electrons being
shared.
· Decreasing from top to bottom of periodic table due to the
weak effective nuclear charge and increasing electron
shielding effect. This will cause less attraction to the
shared/bonded electron
Post Test:
Arrange the following elements base on the following:
a. Mg, F, Hg (decreasing atomic size)
b. Sr, Ge, I (increasing electron affinity)
c. K, Au, Po (from high to low ionization energy)
d. Cl, Tc, Fr (from low to high electronegativity).
Answers:
a. Hg, Mg, F
b. Ge, Sr, I
c. K, Po, Au
d. Fr, Tc, C
INTERVENTION TO CHEMISTRY WEEK 2
READING MATERIALS
LESSON: Atomic Theory and Structure
• Fundamental Chemistry Laws
Law of Conservation of Mass
Mass is neither created nor destroyed in a chemical
reaction.
Law of Definite Proportion
Different pure samples of a compound always contain the
same elements in the same proportion by mass; this
corresponds to atoms of these elements combined in fixed
numerical ratios.
Law of Multiple Proportions
Different compounds made up of the same elements differ
in the number of atoms of each kind that combine.
Dalton’s Law of Atomic Theory
- The postulates....
- All matter consists of atoms.
- Atoms of one element cannot be converted into atoms of
another element.
- Atoms of an element are identical in mass and other
properties and are different from atoms of any other
element.
- Compounds result from the chemical combination of a
specific ratio of atoms of different elements.
ATOM - is the basic unit of
an element that can enter
into chemical
combination.
Components of Atom
• Electron (e-) – negatively charged particle
• Mass of electron = 9.11x10-31 kilogram
• Nucleus – contains all the positive charge and essentially
all the mass of the atom.
• Nucleus - consists of protons and neutrons
• Protons (p+) – positively charged particle
located inside the nucleus
• Neutrons (n0) – uncharged dense particles located inside
the nucleus
Components of an Atom
𝑨
𝒁𝑿 − 𝑠𝑦𝑚𝑏𝑜𝑙 𝑜𝑓 𝑎𝑡𝑜𝑚 𝑜𝑟 𝑖𝑠𝑜𝑡𝑜𝑝𝑒𝑠
X – atomic symbol of the element Z – atomic number
Z – number of protons
A – mass number
THE BOOK LOUNGE PH| 8
𝐴 = 𝑍 + 𝑛0
• An atom is an electrically neutral, spherical entity
composed of a positively charged central nucleus
surrounded by one or more negatively charged electrons.
• The magnitude of charge possessed by a proton is equal
to that of an electron, but the signs of the charges are
opposite.
Components of an Atom
𝑨
𝒁𝑿 − 𝒔𝒚𝒎𝒃𝒐𝒍 𝒐𝒇 𝒂𝒕𝒐𝒎 𝒐𝒓 𝒊𝒔𝒐𝒕𝒐𝒑𝒆𝒔
• Z = number of protons (p+)
• A = mass number
Atomic Mass (Atomic Weight) – average of the masses of
its naturally occurring isotopes weighted according to their
abundances.
Silicon (Si) is essential to the computer industry as a major
component of semiconductor chips. It has three naturally
occurring isotopes: 28Si, 29Si, and 30Si. Determine the
number of protons, neutrons, and electrons in each silicon
isotope.
The atomic number of silicon is 14. Therefore
28
Si has 14p+, 14e- and 14n0 (28 – 14 = 14)
Si has 14p+, 14e- and 15n0 (29 – 14 = 15)
30Si has 14p+, 14e- and 16n0 (30 – 14 = 16)
𝒏𝟎
𝑨=𝒁+
𝑨 = 𝒑+ + 𝒏𝟎
29
• For atom (#protons = #electrons)
𝒑+ = 𝒆−
Example
• For Na (sodium), it has mass # of 23; it contains 11
protons, 11 electrons, and 12 neutrons
• For He (helium), it has mass # of 4; it contains 2 protons,
2 electrons, and 2 neutrons
• For Al (aluminum), it has mass # of 27; it contains 13
protons, 13 electrons, and 14 neutrons
Sample Problems
𝟏𝟗
𝟗𝑿
𝟐𝟑
𝟏𝟏𝑵𝒂
𝑋=𝐹
𝐴 = 19
𝑍=9
𝑒− = 9
𝑝+ = 9
𝑛0 = 10
𝑋 = Na
𝐴 = 23
𝑍 = 11
𝑒− = 11
𝑝+ = 11
𝑛0 = 12
𝟖𝟎
𝟑𝟓𝑩𝒓
𝟕𝟓
𝟑𝟑𝑨𝒔
𝑋 = Br
𝐴 = 80
𝑍 = 35
𝑒− = 35
𝑝+ = 35
𝑛0 = 45
𝑋 = As
𝐴 = 75
𝑍 = 33
𝑒− = 33
𝑝+ = 33
𝑛0 = 42
𝟑𝟗
______
𝟐𝟕
_______
𝑋 = __
𝐴 = 39
𝑍 = __
𝑒− = __
𝑝+ = 19
𝑛0 = __
𝑋 = __
𝐴 = 27
𝑍 = __
𝑒− = __
𝑝+ = __
𝑛0 = 14
Try these:
Isotopes
-atoms of an element with the same number of protons,
but a different number of neutrons
Silver (Ag: Z = 47) has 46 known isotopes, but only two occur
naturally, 107Ag and 109Ag. Given the following mass
spectrometric data, calculate the atomic mass of Ag:
𝐴𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 𝐴𝑔 = 55.4195+52.4485
𝑨𝒕𝒐𝒎𝒊𝒄 𝒎𝒂𝒔𝒔 𝑨𝒈 = 𝟏𝟎𝟕.𝟖𝟔𝟖 𝒂𝒎𝒖
Cations
IONS – contains charge. Either positive or negative charge
CATION – ion that contains positive charge. It tends to lose
electron. Example:
• 𝑵𝒂+𝟏 = sodium ion loses 1 electron due to +1 as its charge.
• 𝑴𝒈+𝟐 = magnesium ion loses 2 electrons due to +2 as its charge.
• 𝑨𝒍+𝟑 = aluminum ion loses 3 electrons due to +3 as its charge
Anions
IONS – contains charge. Either positive or negative charge
ANION - ion that contains negative charge. It tends to gain
electron. Example
• 𝑭−𝟏 = fluorine ion gains 1 electron due to -1 as its charge.
• 𝑶−𝟐 = oxygen ion gains 2 electrons due to -2 as its charge.
• 𝑵−𝟑 = nitrogen ion gains 3 electrons due to -3 as its charge
Note: The mass #, # protons, and # neutrons will not change
for the ions. The only thing that will change for the ion is the
# of electrons.
For Ions:
• Atomic number (Z) = # protons (#p+)
• Mass # (A) = # protons (#p+) + # neutrons (#n0)
• # neutrons (#n0) = Mass # (A) - # protons (#p+)
THE BOOK LOUNGE PH| 9
• # protons (#p+) = Mass # (A) - # neutrons (#n0)
•
Ex. CI-1 ion
𝟑𝟓 -1
𝟏𝟕𝑪𝒍
Number of electrons for ions
#𝒆𝒍𝒆𝒄𝒕𝒓𝒐𝒏𝒔 = 𝒂𝒕𝒐𝒎𝒊𝒄 𝒏𝒖𝒎𝒃𝒆𝒓 𝒁 −𝒄𝒉𝒂𝒓𝒈𝒆
OR #𝒆𝒍𝒆𝒄𝒕𝒓𝒐𝒏𝒔 = #𝒑𝒓𝒐𝒕𝒐𝒏𝒔 (#𝒑+) – 𝒄𝒉𝒂𝒓𝒈𝒆
Examples
𝑋
𝐴
𝑍
# 𝑝+
# 𝑒−
# 𝑛0
Charge
Cation/Anion/Atom
CI
35
17
17
18
18
-1
Anion
LESSON: Electron Configuration and Effect of
Nuclear Charge
Electron configuration – distribution of electron of an
atom in atomic orbitals.
Examples:
• Ex. O-2 ion
𝟏𝟔 -2
𝟖𝑶
𝑋
𝐴
𝑍
# 𝑝+
# 𝑒−
# 𝑛0
Charge
Cation/Anion/Atom
•
O
16
8
8
10
8
-2
Anion
Ex. Na+1 ion
𝟐𝟑
+1
𝟏𝟏𝑵𝒂
𝑋
𝐴
𝑍
# 𝑝+
# 𝑒−
# 𝑛0
Charge
Cation/Anion/Atom
•
Na
23
11
11
10
12
+1
Cation
Ex. Ba+2 ion
𝟏𝟑𝟕
+2
𝟓𝟔𝑩𝒂
𝑋
𝐴
𝑍
# 𝑝+
# 𝑒−
# 𝑛0
Charge
Cation/Anion/Atom
Ba
137
56
56
54
81
+2
Cation
Categories of Electrons
• Inner (core) electrons - are those seen in the previous
noble gas and any completed transitions series. They fill all
the lower energy levels of an atom.
• Outer electrons - are those in the highest energy level
(highest n value). They spend most of their time farthest
from the nucleus.
• Valence electrons - are those involved in forming
compounds. Among the main group elements, the valence
electrons are the outer electrons.
Nuclear Charge Effect
•Effect of Nuclear Charge
- a higher nuclear charge (more protons in the
nucleus) increases nucleus electron attractions
and, thus, lowers sublevel energy (stabilizes the
atom).
• Shielding
ability of inner electrons to lessen nuclear
attraction for an outer electron
• Effective Nuclear Charge
- the nuclear charge an electron actually experiences
with shielding, and this lower nuclear charge makes
the electron easier to remove.
THE BOOK LOUNGE PH| 10
Penetration
the process by which an outer electron moves through the
region occupied by the core electrons to spend more of its
time near the nucleus; increases average effective nuclear
charge for that electron.
Key Points
• Greater nuclear charge lowers orbital energy and make
electrons harder to remove
• Electron - electron repulsion raise orbital energy and
make electrons easier to remove. Repulsions have the
effect of shielding electrons from the full nuclear charge,
reducing it to an effective nuclear charge. Inner electrons
shield outer electrons most effectively.
• Greater penetration makes an electron harder to remove
because it is attracted more strongly and shielded less
effectively. As a result, an energy level (shell) is split into
sublevels (subshells) with the energy order s<p<d<f.
ENERGY LEVEL AND SUBSHELLS
Orbital Diagram
A notation that shows how many electrons an atom has in
each of its occupied electron orbitals.
HOW TO WRITE ELECTRON CONFIGURATION
1st Determine the number of electrons of an atom using
the atomic number. Use the diagram of electron
configuration. Follow the arrow in the said diagram.
2nd Always start with 1s? followed by 252 then 2p53s?
and so on...
3rd NOTE: Add all the superscripts in the electron
configuration, make sure that the sum should be equal to
the number of electrons or atomic number of a given
element/atom.
ELECTRON CONFIGURATION
Isoelectronic Series
Write the electron configuration of Na+, F-, and Ne
THE BOOK LOUNGE PH| 11
Metals, Nonmetals, and Metalloids
•Group 1A (H, Li, Na, K, Rb, Cs, Fr) has 1 valence
electron (v.e-. = 1e-)
•Group 2A (Be, Mg, Ca, Sr, Ba, Ra) has 2 valence
electrons (v.e-. = 2e-)
•Group 3A (B, Al, Ga, In, Tl) has 3 valence electrons
(v.e-. = 3e-)
•Group 4A (C, Si, Ge, Sn, Pb) has 4 valence electrons
(v.e-. = 4e-)
•Group 5A (N, P, As, Sb, Bi) has 5 valence electrons
(v.e-. = 5e-)
•Group 6A (O, S, Se, Te, Po) has 6 valence electrons
(v.e-. = 6e-)
•Group 7A (F, Cl, Br, I, At) has 7 valence electrons
(v.e-. = 7e-)
•Group 8A (Ne, Ar, Kr, Xe, Rn, except He) has 8
valence electrons (v.e-. = 8e-)
LESSON: Periodic Table and Periodic Trends
Periodic Trends
Atomic Radius – size of the atom
• Ionization Energy - Energy required to remove an
electron from a gaseous atom or ion.
• Electron Affinity - Energy change associated with the
addition of an electron to a gaseous atom.
• Electronegativity - The ability of an atom in a
molecule to attract shared electrons to itself.
Atomic Radius
Ionization Energy
Periodic Table
Mendeleev is given the most credit for the current
version of the periodic table.
Originally constructed to represent the patterns
observed in the chemical properties of the elements
Electron Affinity
THE BOOK LOUNGE PH| 12
Electronegativity
Magnetic Quantum Number (ml)
• It is an integer from -l through 0 to +l
• It prescribes the orientation of the orbital in the space
around the nucleus and is sometimes called the orbitalorientation quantum number.
• An orbital with l = 0 can have only ml = 0. However, an
orbital with l=1can have any one of three ml values, -1,
0, or +1; thus, there are three possible orbitals with l =
1, each with its own orientation.
Hierarchy of Quantum Numbers for Atomic Orbitals
Problem
• Arrange the following elements in increasing atomic
radius, ionization energy, and electron affinity.
B, Cs, K, F, O, Na
𝐴𝑅: 𝐹 < O < B < Na < K < Cs
𝐼𝐸: 𝐶𝑠 < K < Na < B < O < F
𝐸𝐴: 𝐶𝑠 < K < Na < B < O < F
Problem
Arrange the following elements in decreasing atomic
radius, ionization energy, and electron affinity.
N, F, C, Fr, Rb, Be
𝐴𝑅: Fr > Rb > Be > C > N > F
𝐼𝐸: F > N > C > Be > Rb > Fr 𝐸
𝐴: F > N > C > Be > Rb > Fr
Problem
• Which has the lowest electronegativity? F, Rb, B, K, N
𝑨𝒏𝒔𝒘𝒆𝒓: 𝑹𝒃
LESSON: Quantum Numbers and Orbitals
Quantum Numbers
Principal Quantum Number (n)
• It is a positive integer (1, 2, 3, 4, 5, 6, 7, ...)
• It indicates the relative size of the orbital, and it
specifies the energy level
Angular Momentum Quantum Number (l)
• other term is Azimuthal Quantum Number
• It is an integer from 0 to n-1.
• It is related to the shape of the orbital and is
sometimes called the orbital-shape quantum number.
• Note that the principal quantum number; that is, n
limits l.
• For an orbital with n = 1, l can have a value of only 0.
For orbitals with n = 2, l can have a value of 0 or 1; for
those with n = 3, l can be 0, 1, or 2
Level and Sublevel
The energy states and orbitals of the atom are
described with specific terms and associated with one
or more quantum numbers:
1. Level - the atom's energy levels, or shells, are given
by the n value: the smaller the n value, the lower the
energy level and the greater the probability of the
electron being closer to the nucleus.
2. Sublevel - the atoms' levels contain sublevels, or
subshells, which designate the orbital shape. Each
sublevel or subshell has a letter designation
l = 0 is a s sublevel (s = sharp)
l = 1 is a p sublevel (p = principal)
l = 2 is a d sub level (d= diffuse)
l = 3 is a f sublevel (f = fundamental)
Orbitals
3. Orbital - each allowed combination of n, l, ml values
specify one of the atom's orbitals. Thus, the three
quantum numbers that describe an orbital express its
size (energy), shape, and spatial orientation.
• A 2s sublevel has only one orbital, and its quantum
numbers are n = 2, l = 0.
• A 3p sublevel has three orbitals, and its quantum
numbers are n = 3, l = 1
• A 3d sublevel has five orbitals, and its quantum
numbers are n = 3, l = 2
THE BOOK LOUNGE PH| 13
Quantum
Numbers
n
l
ml
Energies
1 < 2 < 3 < 4....
s<p<d<f
Orbitals in the same subshell have
equal energies
Subshell
S
P
D
F
Maximum Electrons Per
Subshell
2
6
10
14
Quantum Numbers
Electron-Spin Quantum Number (ms)
• has values of either +1/2 (arrow up) or -1/2 (arrow
down)
• each electron in an atom is described completely by a set
of four quantum numbers: the first three describe its
orbital, and the fourth describes its spin.
• Note: Each orbital can hold a maximum of two electrons
(+1/2 if one electron or unpaired; -1/2 if two electrons or
paired).
• Pauli Exclusion Principle - in a given atom no two
electrons can have the same set of four quantum
numbers.
• Aufbau Principle - The orbitals of an atom must be filled
up in increasing energy levels.
• Hund’s Rule - The most stable arrangement of electrons
in subshells is the one with more parallel spins.
INTERVENTION TO CHEMISTRY WEEK 3
READING MATERIALS
• In forming chemical bonds, main group elements gain,
lose, or share electrons to achieve configuration in
which they are surrounded by eight valence electrons
Types of Chemical Bond
Types of Chemical Bonds
1. Metal with nonmetal: electron transfer and ionic
bonding
2. Nonmetal with nonmetal: electron sharing and
covalent bonding
3. Metal with metal: electron pooling and metallic
bonding
Ionic Bond
- Chemical bond resulting from the transfer of
electrons from one bonding atom to another.
- In an ionic bond, the positively charged ion
(cation) is attracted to negatively charged ion
(anion)
- Static electrical attraction is the basis for ionic
bonds.
LESSON: CHEMICAL BONDING
Chemical Bond
• Attractive force between two atoms holding them
together to form a molecule or a chemical compound
• Forces that link together atoms to form different kinds
of matter
Duet Rule
• Hydrogen forms stable
molecules where it
shares two electrons.
Octet Rule
• Elements form stable molecules when surrounded by
eight electrons.
THE BOOK LOUNGE PH| 14
Ionic Bond
Ionic Compounds
• compound of positive and negative ions combined so
that the charges are neutralized
• NaCl, Fe2O3
Covalent Bond
- Formed by a shared pair of electrons between two
atoms
- Chemical bond formed when valence electrons are
shared by nonmetal elements
Ionic Bond
• Ionic bonding occurs
when ions assemble into
an extended array called a
lattice and are held
together by the attraction
between oppositely
charged ions.
Types of Covalent Bond based on Bond Polarity
Nonpolar covalent bond
Polar covalent bond
Nonpolar Covalent Bond
A bond that has an even distribution of charge due to an
equal sharing of bonding electrons
Insert pic
Polar Covalent Bond
• A bond that has uneven distribution of charge due to
unequal sharing or bonding electrons
• A chemical bond between two atoms that have
different electronegativities, such that one end of the
bond takes on a partial positive charge and the other end
takes on a partial negative charge and constitute a dipole
Ionic Bond
• The valence electrons for Na and O which are located
in a partial orbital diagram are shown below.
Electrostatic forces and the reason ionic compounds
crack.
Electrical Conductance
Bond Polarity
• due to difference in electronegativity of atoms
• If the electronegativity difference is zero, the
bond is classified as nonpolar covalent.
• The greater the electronegativity difference, the more
polar the bond.
• When the electronegativity difference greater than or
equal to 2.0, the bond is classified as ionic.
Electronegativity
Ionic Character
ΔEN
≥ 2.0
0.4 – 1.9
< 0.4
0
IONIC CHARACTER
Mostly ionic
Polar Covalent
Mostly covalent
Non-polar Covalent
Examples
C – O:
• EC = 2.5; EO = 3.5 ΔEN = |3.5 – 2.5| = 1.0 (Polar)
Na – Cl
• ENa = 0.9; ECl = 3.0 ΔEN = |3.0 – 0.9| = 2.1 (Ionic)
THE BOOK LOUNGE PH| 15
F–F
• EF = 4.0 ΔEN = |4.0 – 4.0| = 0.0 (Non-Polar)
P – Br
• EP = 2.1; EBr = 2.8 ΔEN = |2.8 – 2.1| = 0.7 (Polar)
TRY THESE!
N-O
• E = 3.0; E, = 3.5
AEN = 13.5 -3.0| = 0.5 (Polar)
Bond Energy
• The energy released when isolated atoms form a
covalent bond
• The amount of energy required to break a bond (Bond
dissociation energy)
• A measure of the strength of chemical bond
• Triple bond is stronger than double bond
• Double bond is stronger than single bond
Ba - F
• EBa = 0.9; Ef = 4.0
DEN = (4.0-0.9) = 3.1 (lonic)
N-N
• E== 3.0
AEN = 13.0 -3.0| = 0.0 (Non-Polar)
Types of Bonds
Based on the number of electron pair shared
• Single bond
• Double bond
• Triple bond
Single Bond
• A covalent bond in which two atoms share one pair of
electrons.
Double Bond
• A covalent bond in which two atoms share two pairs of
electrons.
Triple Bond
• A covalent bond in which two atoms share three pairs
of electrons.
N=N
Bond Length
• The distance between the nuclei of the bonded atoms
• Single bond is longer than double bond
• Double bond is longer than triple bond
Bond Order
the number of electron pairs shared between two atoms
in the formation of the bond
PROBLEM:
Rank the bonds in each set-in order of decreasing bond
length and bond strength or bond energy:
C = O, C─O, C O
SOLUTION:
Bond length: C─O > C = O > C O
Bond strength: C O > C = O > C─O
Covalent Compounds
• Compound that has atoms held together by covalent
bond
• Always involve two nonmetals
• Examples:
CO2, O2, CH4
Single Bond • Bond Order = 1
Double Bond • Bond Order = 2
Triple Bond • Bond Order = 3
THE BOOK LOUNGE PH| 16
Metallic Bonding
Metallic: sharing by forming a mobile “sea of electrons”
Positively charged metal nuclei arranged in a lattice.
Electrons move, more or less, freely throughout the
whole lattice. Free movement allows metals to conduct
electricity
LESSON: EXAMPLES OF NAMING AND WRITING OF
INORGANIC COMPOUNDS
Topic Learning Outcomes
•Write the chemical formula of an inorganic compound
•Name the inorganic compound based on its chemical formula
THE BOOK LOUNGE PH| 17
LESSON: NAMING AND WRITING OF CHEMICAL
FORMULA
The formation of an ionic compound.
Transferring electrons from the atoms of one element to
those of other results in an ionic compound.
A.
B.
C.
D.
The Elements (Lab View)
The Elements (Atomic View)
Electron Transfer
The compound (atomic view): Na+ and CI- in the
crystal
E. The compound (lab view) sodium chloride
crystal
Predicting the Ion and Element Forms
PROBLEM: What monatomic ions do the following
elements form?
(a) Iodine (Z = 53) (b) Calcium (Z = 20) (c) Aluminum (Z = 13)
PLAN: Use Z to find the element. Find its relationship to
the nearest noble gas. Elements occurring before the
noble gas gain electrons and elements following lose
electrons.
SOLUTION:
¡- Iodine is a nonmetal in Group 7A(17). It gains one
electron to have the same number of electrons as 5Xe.
Ca2+ Calcium is a metal in Group 2A(2). It loses two
electrons to have the same number of electrons as 18Ar.
Al 3+ Aluminum is a metal in Group 3A(13). It loses three
electrons to have the same number of electrons as 10Ne.
THE BOOK LOUNGE PH| 18
Formation of a covalent bond between two H atoms.
Covalent bonds form when elements share electrons,
which usually occurs between nonmetals.
Naming binary ionic compounds
The name of the cation is written first, followed by that
of the anion. The name of the cation is the same as the
name of the metal.
Elements that are polyatomic > A polyatomic ion
Types of Chemical Formulas
A chemical formula is comprised of element symbols and
numerical subscripts that show the type and number of
each atom present in the smallest unit of the substance.
An empirical formula indicates the relative number of
atoms of each element in the compound. It is the
simplest type of formula.
The empirical formula for hydrogen peroxide is HO.
A molecular formula shows the actual number of atoms
of each element in a molecule of the compound.
The molecular formula for hydrogen peroxide is H2O2.
A structural formula shows the number of atoms and the
bonds between them, that is, the relative placement and
connections of atoms in the molecule.
The structural formula for hydrogen peroxide is H-O-O-H.
Many metal names end in -ium.
The name of the anion takes the root of the nonmetal
name and adds the suffix -ide.
Calcium and bromine form calcium bromide.
Naming Binary Ionic Compounds
PROBLEM: Name the ionic compound formed from the
following pairs of elements:
(a) magnesium and nitrogen (b) iodine and cadmium
(c) strontium and fluorine (d) sulfur and cesium
PLAN: Use the periodic table to decide which element is
the metal and which the nonmetal. The metal (cation) is
named first and we use the -ide suffix on the nonmetal
name root.
SOLUTION:
(a) magnesium nitride
(b) cadmium iodide
(c) strontium fluoride
(d) cesium sulfide
Determining Formulas of Binary Ionic Compounds
PROBLEM: Write chemical formulas for the compounds
named in Sample
Problem 2.5.
PLAN: Compounds are neutral. We find the smallest
number of each ion which will produce a neutral
formula. Use subscripts to the right of the element
symbol.
THE BOOK LOUNGE PH| 19
SOLUTION:
(a) Mg2+ and N3-; three Mg2+(6+) and two N3-(6-);
Mg3N2
(b) Cd2+ and I-; one Cd2+(2+) and two I-(2-); CdI2
(c) Sr2+ and F-; one Sr2+(2+) and two F-(2-); SrF2
(d) Cs+ and S2-; two Cs+(2+) and one S2- (2-); Cs2S
Metals With Several Oxidation States
Determining Names and Formulas of Ionic Compounds
of Elements That Form More Than One Ion
PROBLEM: Give the systematic names for the formulas
or the formulas for the names of the following
compounds:
(a) tin(II) fluoride (c) ferric oxide
(b) CrI3
(d) CoS
Determining Names and Formulas of Ionic Compounds
Containing Polyatomic Ions
PROBLEM: Give the systematic names or the formula or
the formulas for the names of the following compounds:
(a) Fe(ClO4)2
(b)sodium sulfite
(c)Ba(OH)2 8H2O
PLAN: Compounds are neutral. We find the smallest
number of each ion which will produce a neutral formula.
Use subscripts to the right of the element symbol.
PLAN: Note that polyatomic ions have an overall charge
so when writing a formula with more than one
polyatomic unit, place the ion in a set of parentheses.
SOLUTION:
(a) Tin (II) is Sn2+; fluoride is F-; so the formula is SnF2.
(b) The anion I is iodide(I-); 3I- means that Cr(chromium)
is +3. CrI3 is chromium (III) iodide
(c) Ferric is a common name for Fe3+; oxide is O2-,
therefore the formula is Fe2O3.
(d) Co is cobalt; the anion S is sulfide (2-); the
compound is cobalt (II) sulfide.
SOLUTION:
(a) ClO4- is perchlorate; iron must have a 2+ charge.
This is iron (II) perchlorate.
(b) The anion sulfite is SO32- therefore you need 2
sodium’s per sulfite. The formula is Na2SO3.
(c) Hydroxide is OH- and barium is a 2+ ion. When water
is included in the formula, we use the term “hydrate”
and a prefix which indicates the number of waters. So, it
is barium hydroxide octahydrate.
Some Common Polyatomic Ions
Recognizing Incorrect Names and Formulas of Ionic
Compounds
PROBLEM: Something is wrong with the second part of
each statement.
Provide the correct name or formula.
(a) Ba(C2H3O2)2 is called barium diacetate.
(b) Sodium sulfide has the formula (Na)2SO3.
(c) Iron (II) sulfate has the formula Fe2(SO4)3.
(d) Cesium carbonate has the formula Cs2(CO3).
THE BOOK LOUNGE PH| 20
SOLUTION:
(a) Barium is always a +2 ion and acetate is -1.
The “di-” is unnecessary.
(b) An ion of a single element does not needs
parentheses. Sulfide is S2-, not SO32-. The correct
formula is Na2S.
(c) Since sulfate has a 2- charge, only 1 Fe2+ is needed.
The formula should be FeSO4.
(d) The parentheses are unnecessary. The correct
formula is Cs2CO3.
NAMING ACIDS
1) Binary acids solutions form when certain gaseous
compounds dissolve in water.
For example, when gaseous hydrogen chloride (HCI)
dissolves in water, it forms a solution called
hydrochloric acid. Prefix hydroanion nonmetal root + suffix -ic + the word acid hydrochloric acid.
2) Oxoacid names are similar to those of the oxoanions,
except for two suffix changes:
Anion "-ate" suffix becomes an "-ic" suffix in the acid.
Anion "-ite "suffix becomes an "-ous" suffix in the acid.
The oxoanion prefixes "hypo-" and "per-" are retained.
Thus, BrO; is perbromate, and HBrO, is perbromic acid;
10 is iodite, and HIO is iodous acid.
Determining Names and Formulas of Anions and Acids
PROBLEM: Name the following anions and give the
names and formulas of the acids derived from them:
(a) Br
(b) 103
(c) CN
(d) SO,2(e) NO,
SOLUTION:
(a) The anion is bromide; the acid is hydrobromic acid,
HBr.
(b) The anion is iodate; the acid is iodic acid, HIO.
(c) The anion is cyanide; the acid is hydrocyanic acid,
HCN.
(d) The anion is sulfate; the acid is sulfuric acid, H,SO,
(e) The anion is nitrite; the acid is nitrous acid, HNO.
Determining Names and Formulas of Binary Covalent
Compounds
PROBLEM:
(a) What is the formula of carbon disulfide?
(b) What is the name of PCly?
(c) Give the name and formula of the compound whose
molecules each consist of two N atoms and four O
atoms.
SOLUTION:
(a) Carbon is C, sulfide is sulfur S and di-means 2 - CS2.
(b) P is phosphorous, Cl is chloride, the prefix for 5 is
penta-. Phosphorous pentachloride.
(c) N is nitrogen and is in a lower group number than O
(oxygen). Therefore, the formula is N2O4 - dinitrogen
tetraoxide.
Recognizing Incorrect Names and Formulas of Binary
Covalent Compounds
PROBLEM: Explain what is wrong with the name of
formula in the second part of each statement and
correct it:
(a) SF4 is monosulfur pentafluoride.
(b) Dichlorine heptaoxide is Cl,06(c) Nag is dinitrotrioxide.
SOLUTION:
(a) The prefix mono- is not needed for one atom; the
prefix for four is tetra-. So the name is sulfur
tetrafluoride.
(b) Hepta- means 7; the formula should be Cl,O.
(c) The first element is given its elemental name so this
is dinitrogen trioxide.
INTERVENTION TO CHEMISTRY WEEK 4
READING MATERIALS
LESSON: CHEMICAL REACTIONS
Balancing of Chemical Reaction
• Number of atoms of each element must be the same
before and after chemical reactions.
Number of atoms in reactants = Number of atoms
in products
Law of Conservation of Mass
• Mass can neither be gained nor lost through a
chemical reaction
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total mass of reactants = total mass of products
Steps in Balancing Chemical Reactions
-Write a formula equation with correct symbols and
formulas.
-Count the number of atoms of each element on each
side of the arrow.
-Balance atoms by using coefficients.
-Check your work by counting atoms of each element.
Example A
Balance the ff. reaction:
1. CO + O2 → CO2
Balanced reaction: 2CO + O2 → 2CO2
2. Ca + O2 → CaO
Balanced reaction: 2Ca +O2→2CaO
Example B
Balance the ff. reaction:
KClO3→KCl + O2
Balanced reaction: 2KClO3 → 2KCl + 3O2
Cl2 +KBr → KCl +Br2
Balanced reaction: Cl2 + 2KBr → 2KCl + Br2
Example C
Balance the ff. reaction:
H2SO4, + NaOH> Na2SO4+ H2O
Balanced equation: H2SO4, + 2NaOH > Na2SO4 + 2H20
Try this!
Balance the ff. reaction:
N2O5→NO2+ O2
Balanced reaction: 2N2O5→4NO2+ O2
Types of Reaction
•Synthesis or Combination
• Single Displacement Reaction or Single Replacement
Reaction
𝐴+𝐵𝑋→𝐴𝑋+𝐵
Examples:
𝑍𝑛+ 2𝐻𝐶𝑙 → 𝑍𝑛𝐶𝑙2 + 𝐻2
𝐹𝑒 + 𝐶𝑢𝑆𝑂4 → 𝐹𝑒𝑆𝑂4 + 𝐶𝑢
•Double Displacement Reaction or Double
Replacement Reaction or Metathesis
𝐴𝐵+𝐶𝐷 →𝐴𝐷+𝐶𝐵
Examples:
𝐶𝑎𝐶𝑂3+2𝐻𝐶𝑙→𝐶𝑎𝐶𝑙2+𝐻2𝐶𝑂3
𝐵𝑎𝐶𝑙2+𝑁𝑎2𝑆𝑂4→𝐵𝑎𝑆𝑂4+2𝑁𝑎𝐶𝑙
Steps in Balancing Combustion Reaction
- Balance first carbon
- Next, balance hydrogen
- Count the number of oxygens in the product
side, then divide it by “2”. This will be the
coefficient of O2. The “2” is the subscript in the
O2
Note: If there is an oxygen present in the compound
(which is to be burned/oxidized), take the difference
first of the number of oxygen in the product with the
subscript of oxygen in the compound, then divide it by
“2”. This will be the coefficient of O2.
COMBUSTION REACTION
Balance the ff. reaction:
1. CH4 + O2 → CO2 + H2O
Balanced reaction:
CH4 + 2O2 →CO2 + 2H2O
2. C3H8 + O2 → CO2 + H2O
𝐴+𝐵→𝐶
Examples:
2H2(g) + 02(g) - 2H2O(g)
4Fe + 302 - 2Fe2O3
Balanced reaction:
C3H8 + 5O2 → 3CO2 + 4H2O
Balance the ff. reaction:
3. C2H6 + O2 →CO2 + H2O
• Decomposition
𝐴→𝐵+𝐶
Examples:
2H2O(g) → 2H2(g) + + 02(g)
2KCIO3 -> 2KCI + 302
Balanced reaction:
C2H6 + 7/2O2 → 2CO2 + 3H2O
To eliminate the fraction, multiply the whole
reaction to 2
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(C2H6 + 7/2O2 → 2CO2 + 3H2O) x 2
2C2H6 + 7O2 → 4CO2+ 6H2O
Balance the ff. reaction:
4. C4H10 + O2 → CO2 + H2O
Balanced reaction:
C4H10+ 13/2O2→4CO2+ 5H2O
To eliminate the fraction, multiply the whole
reaction to 2
(C2H6+ 13/2O2→4CO2+ 5H2O) x 2
2C2H6+ 13O2→8CO2+ 10H2O
Balance the ff. reaction:
1.C2H6O+O2→CO2+H2O
Balanced reaction:
C2H6O + 3O2→2CO2+ 3H2O
2.C4H10O +O2 → CO2+H2O
Balanced reaction: C4H10O + 6 O2→ 4 CO2+ 5H2O
LESSON: MOLE CONCEPTS
Chemical Stoichiometry
Mole Concept
Molecular Weight (MW)
• 1 molecule H2O = 18 amu
• 1 molecule O2 = 32 amu
• 1 molecule NaF = 42 amu
Molar Mass (MM) for a Molecule/Compound
• 1 mol H2O= 18 grams of H2O
• 1molO2 =32gramsofO2
• 1molNaF=42gramsofNaF
We use the unit “amu” or atomic mass unit for every 1
molecule
We use the unit grams for every 1 mol of a compound
or molecule
Note: In chemistry, the term molecular weight (MW) is
used interchangeably with molar mass of a
molecule/compound
The unit for the molar mass of a compound/molecule is
g/mol
How to compute the molar mass for a
molecule/compound or molecular weight
1. Determine the atomic weights / molar masses of the
elements present in the compound / molecule
2. Multiply the subscript of the element by its respective
atomic weight or molar mass.
3. Add all the values.
Stoichiometry – The study of quantities of materials
consumed and produced in chemical reactions.
Mole (mol) – basic unit of a substance
Mole Concept
Atomic Weight (AW)
• 1 atom C = 12 amu
• 1 atom F = 19 amu
• 1 atom Na = 23 amu
Molar Mass (MM) for Element/atom
• 1 mol C = 12 grams of C
• 1 mol F = 19 grams of F
• 1 mol Na = 23 grams of Na
•We use the unit “amu” or atomic mass unit for every 1
atom
•We use the unit grams for every 1 mol of an element
•Note: In chemistry, the term atomic weight (AW) is
used interchangeably with molar mass of an element.
•The unit for the molar mass of an element is g/mol
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MASS TO MOLE
Formula for Mass to Mole
𝑋=𝑐𝑎𝑛 𝑏𝑒 𝑎𝑛 𝑒𝑙𝑒𝑚𝑒𝑛𝑡 𝑜𝑟 𝑎 𝑐𝑜𝑚𝑝𝑜𝑢𝑛𝑑/𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑒
TRY THESE:
1.Convert 50grams of C to mol of C
(Molar mass: C=12g/mol)
2.Convert 150 grams of C6H6 to mol of C6H6
(Molar mass: C=12g/mol; H=1g/mol)
MOLE TO MASS
•Formula for Mole to Mass
𝑋=𝑐𝑎𝑛 𝑏𝑒 𝑎𝑛 𝑒𝑙𝑒𝑚𝑒𝑛𝑡 𝑜𝑟 𝑎 𝑐𝑜𝑚𝑝𝑜𝑢𝑛𝑑/𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑒
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3.Convert 2.3 mol of NH4Cl to grams of NH4Cl
(Molar mass: N=14g/mol; H=1g/mol; Cl=35.45g/mol)
Avogadro’s Number
Example:
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2.Determine the percent composition of Al(OH)3
TRY THESE....
1. Convert 35 grams of C to C atoms
(Molar mass: C=12g/mol)
3.Determine the percent composition of (NH4)3PO4
2. Convert 0.05 mol of Cu to Cu atoms
3. Convert 55 grams of H2SO4 to H2SO4
molecules
(Molar mass: H = 1 g/mol; S = 32 g/mol; O = 16 g/mol)
LESSON: STOICHIOMETRY
Stoichiometric Calculations: Amounts of Reactants and
Products
Percent Composition
𝑋 = 𝑒𝑙𝑒𝑚𝑒𝑛𝑡
Calculating Masses of Reactants and Products in Reactions
• Balance the equation for the reaction
• Convert the known mass of the reactant or product to
moles of that substance.
•Use the balanced equation to set up the appropriate
mole ratios.
• Use the appropriate mole ratios to calculate the number
of moles of desired reactant or product.
•Convert from moles back to grams if required by the
problem.
Calculating Masses of Reactants and Products in Reactions
Ex: % Composition of element in a compound
1.Determine the percent composition of Na2SO4
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Mole Ratio
Balanced Chemical Reaction
4𝐹𝑒 + 3𝑂 → 2𝐹𝑒 𝑂 223
4 mol Fe reacts with 3 mols of O2
4 mol Fe produces 2 mols Fe2O3
3 mols of O2 reacts with 4 mols of Fe
3 mols of O2 produces 2 mol of Fe2O3
Mole Ratio
• Balanced Chemical Reaction
4 𝐹𝑒 + 3𝑂2 → 2𝐹𝑒 2𝑂3
Mole Ratio–it is based on the coefficients from a
balanced chemical reaction
𝑃4+5𝑂2 → 2𝑃2𝑂5
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• Hydration – the process in which an ion is surrounded by
water molecules arranged in specific manner
Solubility
- The maximum amount of solute that dissolves in a fixed
quantity of a particular solvent at a specified temperature
when excess solute is present.
- Property which allows solute particles to form uniform
mixture with solvent particles.
INTERVENTION TO CHEMISTRY WEEK 4
READING MATERIALS
LESSON: SOLUTIONS
Solution
Homogeneous mixture of two or more substances, the
relative proportion of which may vary within certain limit
• A homogeneous mixture of two or more substances in a
single physical state.
Properties of Solution
• Particles in a solution are very small.
• Particles in solution are evenly distributed or
intermingled uniformly on a molecular level.
• The particles in solution will not separate no matter how
long the solution is allowed to stand under constant
conditions.
Components of Solution
Solute
• substance that is dissolved
• smaller amount than solvent
Solvent
• dissolving medium
• larger amount than solute
Water – “universal solvent”
Types of Solution
• Saturated - contains maximum amount of solute that a
solvent can dissolve at a given temperature
• Unsaturated - there is less solute that can be dissolved
at a given temperature
• Supersaturated – unstable condition in which there is
more solute in solution than can normally exist at a given
temperature.
Solute can be...
Soluble - A given solute can easily dissolve in a given
amount of solvent.
Slightly Soluble - the solute is partially dissolved in a given
amount of solvent.
Insoluble - a given solute does not dissolve in a given
amount of solvent.
Miscibility
- Solubility of liquid with another liquid
Miscible
• if two liquids dissolve in each other in any proportion.
Partially Miscible
• when two liquid components form a single phase when
mixed in certain proportions but form two phases when
mixed in different proportions.
Immiscible
• two components are insoluble in each other
Oil is immiscible to water
Aqueous Solution
A solution in which the solvent is water
A solute that dissolved in water can be:
Electrolyte
•a substance that when dissolved in water, results in a
solution that can conduct electricity.
Non-electrolyte
• a substance that does not conduct electricity when
dissolved in water
Aqueous Solution
Solvation and Hydration
• Solvation – the process in which an ion or a molecule is
surrounded by solvent molecules arranged in specific
manner
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Factors Affecting Solubility
Percent by Mass (Prob 1)
- NATURE OF SOLUTE AND SOLVENT
- TEMPERATURE
- PRESSURE (USUALLY FOR SOLUBILITY OF GASES)
Nature of Solute and Solvent
• When two substances are similar in nature, they can
dissolve each other ("like dissolves like").
• Polar solute – polar solvent
• Non-polar solute – non – polar solvent
Percent by Mass (Prob 2)
Polar and Non-Polar Substances
Concentration of Solution
• The amount of a solute in a given amount of solvent or
solution
• A complete description of a solution
Ways of Expressing Concentrations
• Percent by Mass
• Percent by Volume
• Molarity
• Molality
• Mole Fraction
• Normality
Percent by Volume
Volume of solution = volume of solute + volume of solvent
Percent by Mass
Note: the unit of numerator must be the same with unit of
the denominator
Mass of solution = mass of solute + mass of solvent
Percent by Volume (Prob 1)
Note: the unit of numerator must be the same with unit of
the denominator
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Molarity (Prob 2)
Percent by Volume (Prob 2)
Molality (Molal Concentration)
Molarity (Molar Concentration)
Note: The unit of molality is “m” or mol/kg.
Molality (Prob 1)
Note: The unit of molarity is “M” or mol/L.
Molarity (Prob 1)
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Molality (Prob 2)
Mass % to molarity (given density in g/mL)
ADDITIONAL PROBLEM
Mole Fraction (X)
Mole Fraction (Prob 1)
ADDITIONAL PROBLEM
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Normality (NORMAL SOLUTION)
•Unit of Normality is eq/L or “N”
Normality (Prob 1)
Normality (Normal Solution)
1. “f” for acids
• f = number of H in the compound
2. f” for bases
• f = number of OH in the compound
3. “f” for ionic compound / salt
• f = charge of the cation times the numerical value of the
subscript of the cation in the compound.
Normality (Prob 2)
Normality
Normality (Prob 3)
FORMULA OF NORMALITY
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LESSON: GASEOUS STATE
GENERAL PROPERTIES OF GASES
➢ Indefinite shape
➢ Indefinite volume
➢ Take the shape and volume of container
➢ Particles are far apart
➢ Particles move fast
➢ High Kinetic Energy - particles can separate and
move throughout container
GENERAL PROPERTIES OF GASES
Solid Liquid
Behavior of Gases
• IDEAL GAS
– gas described in kinetic molecular theory and strictly
follow gas laws
• REAL OR NON-IDEAL GAS
- deviates from ideal gas behavior at high pressure and
temperature
- the intermolecular attractions of these gases hold their
molecules close to one another and allow gases to be
liquefied
Gas
GENERAL PROPERTIES OF GASES
- the most compressible of the states of matter
- mix evenly and completely when confined to the same
container
- Easily effuse (to flow through small holes) and diffuse (to
spread to occupy available space)
- have much lower densities than liquids and solids
- Exert pressure on the containing vessel
Some substances found as Gases under normal
atmospheric condition (1 atm and 25oC)
Variables Affecting Behavior of Gases
• Temperature (T)
• Volume (V)
• Amount or number of moles (n)
• Pressure (P)
Temperature
• A measure of the average kinetic energy of a gas sample
measured in Kelvin
• The motion of the molecules is dependent on the
temperature
• The molecular motion of molecules measured in terms of
the average kinetic energy increases as temperature is
increased
• K = oC+273.15
Volume (V) or Capacity
• Space occupied by the sample of gas which is equal to
the volume of its container
• Measured in liters (L) due to low densities of gases
Kinetic Molecular Theory postulated by: Daniel Bernoulli
(1738)
• Explains the regularity of the behavior of all gases
Postulates:
1. Gases consist of small molecules that are in constant
random motion
2. The volumes of all molecules of a gas are small
compared to the space between molecules (A gas is mostly
empty space).
3. Intermolecular forces between particles are negligible
4. Collisions between molecules and with their container
are perfectly elastic (no energy is lost due to friction and
the pressure in the container does not vary with time at
any given temperature)
5. Average kinetic energy of the molecules is proportional
to absolute Temperature.
Moles of Gas
• The amount or the number of particles present in a gas
sample
• interms of mole of gas(n)
n = moles of gas
Pressure
• Force exerted by gas molecules on the wall of the
container
• Gases exert pressure on any surface with which they
come in contact because
gas molecules are constantly in motion
Atmospheric Pressure
• Pressure exerted by earth’s atmosphere
• Actual value depends on location, temperature and
weather conditions
• Standard Atmospheric Pressure = 1 atm
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Conversion of Units for Pressure
• Combined Gas Law (combination of Boyle’s and Charles’
Law)
Examples (Convert the ff:)
Sample Problem
A sample of argon gas occupies 105 mL at 0.871 atm. If the
temperature and number of moles of argon remains
constant, what is the volume (in L) at 0.259 atm?
GAS LAW
• Boyle’s Law (volume is inversely proportional to the
pressure of gas, at constant temperature and number of
moles)
• Charles’ Law (Volume of gas is directly proportional to
absolute temperature of gas at constant pressure and
number of moles)
Sample Problem
A balloon is filled with 1.95 L of air at 25oC and then placed
in a car in the sun. What is the volume of the balloon (in L)
when the temperature in the car reaches 90oC?
• Avogadro’s Law (volume of gas is directly proportional to
number of moles of gas, at constant temperature and
pressure)
• Gay-Lussac’s Law (Pressure of gas is directly proportional
to absolute temperature of gas at constant volume and
number of moles)
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IDEAL GAS LAW
DENSITY OF A GAS
Example 1
An automobile tire at 23°C with an internal volume of 25.0
L is filled with air to a total pressure of 3.18 atm.
Determine the number of moles of air in the tire.
Solution:
Unit of density of gas is g/L
STANDARD TEMPERATURE AND PRESSURE (STP)
Example 2
What is the pressure in a 19-L tank that contains 5.67 g of
helium at 25°C?
Solution:
Example 4
Example 3
Sulfur hexafluoride (SF6) is a gas used as a long-term
tamponade (plug) for a retinal hole to repair detached
retinas in the eye. If this compound is introduced into an
evacuated 500.0-mL container at 83°C with a pressure of
760 torr, what is the mass (in grams) of the compound?
Molar mass :S=32g/mol; F=19g/mol
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Example 5
GAS MIXTURE
Example 6
TRY THESE
Mixture of Gases
• Gases mix homogeneously in any proportions
• Each gas in a mixture behaves as if it were the only gas
present (assuming no chemical interactions).
Dalton’s Law of Partial Pressures
• In a mixture of unreacting gases, the total pressure is the
sum of the partial pressures of the individual gases.
Partial Pressure
• It is the pressure exerted by each gas in a mixture
• In a mixture of unreacting gases, the total pressure is the
sum of the partial pressures of the individual gases.
Partial Pressures and Mole Fraction
• Each component in a mixture contributes a fraction of
the total number of moles in the mixture.
• Mole Fraction (X)
Mole Fraction (X)
Since the total pressure is due to the total number of moles,
the partial pressure of a gas in gas mixture is equal to the
total pressure multiplied by the mole fraction of a gas.
Sample Problem
Dry air contains the following mole fraction: Calculate the
partial pressure (in kPa) for N2, O2, and Ar. The total pressure
is equal to the atmospheric pressure (101.325 kPa)
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Sample Problem
INTERVENTION TO CHEMISTRY WEEK 5
READING MATERIALS
LESSON: ACIDS AND BASES
Arrhenius Acid – Base Definition
EXAMPLE: A mixture of consisting of 7.0 g of CO and 10.0 g of
SO2 has a total pressure of 0.33 atm when placed in a sealed
container. What are the partial pressures (in atm) of CO and
SO2? Molar mass: CO = 28 g/mol; SO2 = 64 g/mol
Acid – a substance that has H in its formula and dissociates in
water to yield H3O+ or H+.
• Example: HCl, HNO3, HCN
• Base – a substance that has OH in its formula and
dissociates in water to yield OH-.
• Example: NaOH, KOH, Ba(OH)2
• Neutralization – reaction between acid and base.
Bronsted – Lowry Acid – Base Definition
• An acid is a proton donor, any species that donates an H+.
An acid must contain H in its formula: HNO3 and H2PO4-. All
Arrhenius acids are Bronsted-Lowry acids.
• A base is a proton acceptor, any species that accepts an H+
ion. A base must contain a lone pair of electrons to bind the H+
ion. A few examples are NH3, CO3-2, and F-, as well as OH-.
Bronsted- Lowry bases are not Arrhenius bases, but all
Arrhenius bases contains the Bronsted-Lowry Base OH-.
• An acid and a base always work together in the transfer of a
proton. One species behaves as an acid only if another species
simultaneously behaves as a base, and vice versa.
Acid – Base dissolve in water...
• When an acid or a base merely dissolves in water, an acidbase reaction occurs because water acts as the other partner.
1. Acid donates a proton to water
2. Base accepts a proton from water
TRY THIS
A chemical engineer places a mixture of noble gases
consisting of 5.50 g of He, 15.0 g Ne, and 35.0 g of Kr in
piston-cylinder assembly at STP. Calculate the partial pressure
(in atm) of each gas. Molar mass: He = 4 g/mol ; Ne = 20.18
g/mol ; Kr = 83.79 g/mol
Answers:
Conjugate Acid – Base Pair
• Using the reversible reaction below:
H2S acts as an acid by donating an H+ to NH3 which acts as a
base. Notice that the acid, H2S, becomes a base, HS-, and the
base, NH3, becomes an acid, NH4+.
• H2S and HS- are a conjugate acid – base pair
• NH3 and NH4+ are a conjugate acid – base pair
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Conjugate Acid – Base Pair
Using the reversible reaction below:
Acid –Dissociation Constant, Ka
Using the dissociation reaction of water
• HS- is the conjugate base of the acid H2S.
• NH4+ is the conjugate acid of the base NH3
Conjugate Acid – Base Pair
Every acid has a conjugate base, and every base has
a conjugate acid.
The conjugate base has one fewer H and one more
minus charge than the acid.
The conjugate acid has one more H and one fewer
minus charge than the base.
Example
1. 𝐻𝐹+𝐻𝑂 ↔ 𝐹−+𝐻𝑂+ 23
2. 𝐻𝐶𝑂𝑂𝐻 + 𝐶𝑁− ↔ 𝐻𝐶𝑂𝑂− + 𝐻𝐶𝑁
3. 𝑁𝐻++𝐶𝑂−2 ↔ 𝑁𝐻 +𝐻𝐶𝑂− 4333
4. 𝐻𝑃𝑂−+𝑂𝐻− ↔ 𝐻𝑃𝑂−2+𝐻𝑂
• Strong acid → higher [H3O+] → larger Ka
• Strong acid dissociates completely into ions in water.
[H3O+] = initial concentration of the acid.
Base –Ionization Constant, K
pH and pOH
Example
What is the conjugate base of the acid, HPO4-2?
•Answer: PO4-3
What is the conjugate acid of the base SO3-2?
•Answer: HSO3-
pH measures the acidity of the solution.
•pH < 7 acidic solution
•pH = 7 neutral solution
•pH > 7 basic or alkaline solution
Relative Acid – Base Strength
Relating pH and pOH using Kw
Take the negative log of both sides of equation
pH for acid and base solution
•For acidic solution
• For Basic / Alkaline Solution
Autoionization of Water (The Ion-Product Constant for
Water, Kw)
•Using the dissociation reaction of water
𝐻+ and [𝑂𝐻−] concentration
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EXAMPLES
1.What is the pH of 10-3M H3O+? Assume the acid is a
strong acid.
EXAMPLE 7
2. What is the pH of 0.55 M HCl? HCl is a strong acid.
LESSON: COLLIGATIVE PROPERTIES
3.What is the pH of 0.85M NaOH? NaOH is a strong base.
EXAMPLE 4
A property of a solution that depends on the number, not
the identity, of solute particles.
• Colligative properties were measured to explore the
nature of a solute in aqueous solution and its extent of
dissociation into ions (electrolyte solutions).
Colligative Properties of Non-volatile and non- electrolyte
Solution
Vapor Pressure Lowering
Boiling Point Elevation
Freezing Point Depression
Osmotic Pressure
Vapor Pressure Lowering (∆P)
• The vapor pressure of a solution of a non-volatile nonelectrolyte is always lower than the vapor pressure of the
pure solvent.
• Using Raoult’s Law:
EXAMPLE 5
EXAMPLE 5
Derivation of Vapor Pressure Lowering (∆P)
EXAMPLE 6
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Boiling Point Elevation (ΔTb)
• Boiling Point → it is the temperature at which the vapor
pressure is equal to the external pressure.
• External Pressure → atmospheric pressure
• Boiling Point Elevation (ΔTb) →A solution boils at a
higher temperature than the pure solvent.
Example 1 (Vapor Pressure Lowering)
Calculate the vapor pressure lowering, ∆P (in torr) when
0.137 mol of glycerol is added to 27.4 mol of H2O at 50oC.
At this temperature, the vapor pressure of pure water is
92.5 torr.
Boiling Point Elevation (ΔTb)
Example 2 (Vapor Pressure Lowering)
Calculate the vapor pressure lowering (in torr) of a
solution of 2.00 grams of aspirin (molar mass = 180.15
g/mol) in 50.0 grams of methanol (molar mass = 32 g/mol)
at 21.2oC. Also, calculate the vapor pressure (in torr) of
the solution. Pure methanol has a vapor pressure of 101
torr at this temperature.
Example 1
Example 2
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EXAMPLE 1
EXAMPLE 2
Freezing Point Depression (ΔTf) → the solution freezes at
a lower temperature than the solvent
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EXAMPLE 2
A 5 grams of glucose (Molar mass glucose = 180 g/mol) is
dissolved in water to make 1200 mL solution. What is the
osmotic pressure of the solution at 30oC?
Osmotic Pressure (∏)
• Osmosis – when two solutions of different concentrations
are separated by semi-permeable membrane, one that
allows solvent, but not the solute, molecules to pass
through.
• Osmotic Pressure – defined as the applied pressure
required to prevent the net movement of water from
solvent to solution.
EXAMPLE 1
A 0.30 M solution of sucrose that is at 37oC has
approximately the same osmotic pressure as blood does.
What is the osmotic pressure of blood (in atm)?
𝑇 =37+273.15=310.15𝐾
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