Ch 100: Fundamentals for
Chemistry
Chapter 1: Introduction
Lecture Notes
What is Chemistry?
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•
•
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Chemistry is considered to be the central science
Chemistry is the study of matter
Matter is the “stuff” that makes up the universe
The fundamental questions of Chemistry are:
• How can matter be described?
• How does one type of matter interact with other types of
matter?
• How does matter transform into other forms of matter?
Scientific Method
1. Recognize a problem
 Make observation
 Ask a question
2. Make an educated guess - a hypothesis

Predict the consequences of the hypothesis
3. Perform experiments to test the predictions
 Does experiment support or dispute hypothesis?
4. Formulate the simplest rule that organizes the 3
main ingredients - develop a theory
The Scientific Attitude
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All hypotheses must be testable (i.e. there
must be a way to prove them wrong!!)
Scientific: “Matter is made up of tiny
particles called atoms”
Non-Scientific: “There are tiny particles of
matter in the universe that will never be
detected”
Major Developments in Chemistry I
~400 BC: Democritus proposed the concept of the “atom”
~300 BC: Aristotle developed 1st comprehensive model of matter
~700 AD: Chinese alchemists invent gunpowder
1661: Robert Boyle proposed the concept of elements
1770-1790: Lavoisier proposed the concept of compounds & the Law of
Mass Conservation
1774: Priestly isolates oxygen
1797: Proust proposed the Law of Definite Proportions
1803: Dalton re-introduces the concept of the atom and establishes
Dalton’s Laws
1869: Mendeleev creates the 1st Periodic Table
1910: Rutherford proposes the “nuclear” model of the atom
1915: Bohr proposes a “planetary” model of the hydrogen atom
1920: Schroedinger publishes his wave equation for hydrogen
1969: Murray Gell-Mann proposes the theory of QCD (proposing the
existence of quarks)
Major Developments in Chemistry II
Discovery of subatomic particles:
1886: Proton (first observed by Eugene Goldstein)
1897: Electron (JJ Thompson)
1920: Proton (named by Ernest Rutherford)
1932: Neutron (James Chadwick)
Other Important Discoveries:
1896: Antoine Henri Becquerel discovers radioactivity
1911: H. Kamerlingh Onnes discovers superconductivity in low temperature
mercury
1947: William Shockley and colleagues invent the first transistor
1996: Cornell, Wieman, and Ketterle observe the 5th state of matter (the BoseEinstein condensate) in the laboratory
Ch 100: Fundamentals for
Chemistry
Chapter 2: Measurements & Calculations
Lecture Notes
Types of Observations
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Qualitative
Descriptive/subjective in nature
Detail qualities such as color, taste, etc.
Example: “It is really warm outside today”
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Quantitative
Described by a number and a unit (an accepted
reference scale)
Also known as measurements
Example: “The temperature is 85oF outside
today”
Measurements
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Described with a value (number) & a unit
(reference scale)
Both the value and unit are of equal
importance!!
The value indicates a measurement’s size
(based on its unit)
The unit indicates a measurement’s
relationship to other physical quantities
Scientific Notation
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Technique Used to Express Very Large or
Very Small Numbers
Based on Powers of 10
To Compare Numbers Written in Scientific
Notation
First Compare Exponents of 10 (order of
magnitude)
Then Compare Numbers
Writing Numbers in Scientific Notation
1
2
Locate the Decimal Point
Move the decimal point to the right of the
non-zero digit in the largest place
The new number is now between 1 and 10
3
Multiply the new number by 10n
where n is the number of places you moved the
decimal point
4
Determine the sign on the exponent, n
If the decimal point was moved left, n is +
If the decimal point was moved right, n is –
If the decimal point was not moved, n is 0
Writing Numbers in Standard Form
1
Determine the sign of n of 10n
If n is + the decimal point will move to the right
If n is – the decimal point will move to the left
2
Determine the value of the exponent of 10
Tells the number of places to move the decimal
point
3
Move the decimal point and rewrite the
number
Measurement Systems
There are 3 standard unit systems we will focus
on:
1. United States Customary System (USCS)
 formerly the British system of measurement
 Used in US, Albania, and a couple others
 Base units are defined but seem arbitrary (e.g. there are
12 inches in 1 foot)
2. Metric
 Used by most countries
 Developed in France during Napoleon’s reign
 Units are related by powers of 10 (e.g. there are 1000
meters in 1 kilometer)
3. SI (L’Systeme Internationale)
 a special set of metric units
 Used by scientists and most science textbooks
 Not always the most practical unit system for lab work
Related Units in the Metric System
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All units in the metric system are related to
the fundamental unit by a power of 10
The power of 10 is indicated by a prefix
The prefixes are always the same,
regardless of the fundamental unit
Units & Measurement
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When a measurement has a specific unit (i.e.
25 cm) it can can be expressed using
different units without changing its meaning
Example:
» 25 cm is the same as 0.25 m or even 250 mm
•
The choice of unit is somewhat arbitrary,
what is important is the observation it
represents
Measurement & Uncertainty
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A measurement always has some
amount of uncertainty
Uncertainty comes from limitations of
the techniques used for comparison
To understand how reliable a
measurement is, we need to understand
the limitations of the measurement
Measurements & Significant Figures
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To indicate the uncertainty of a single
measurement scientists use a system
called significant figures
The last digit written in a measurement
is the number that is considered to be
uncertain
Unless stated otherwise, the uncertainty
in the last digit is ±1
Rules for Counting Significant Figures
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Nonzero integers are always significant
Zeros
 Leading zeros never count as significant figures
 Captive zeros are always significant
 Trailing zeros are significant if the number has a
decimal point
•
Exact numbers have an unlimited number of
significant figures
Rules for Rounding Off
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If the digit to be removed
• is less than 5, the preceding digit stays the same
• is equal to or greater than 5, the preceding digit
is increased by 1
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In a series of calculations, carry the extra
digits to the final result and then round off
Don’t forget to add place-holding zeros if
necessary to keep value the same!!
Exact Numbers
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Exact Numbers are numbers known with
certainty
Unlimited number of significant figures
They are either
 counting numbers
 number of sides on a square
 or defined
 100 cm = 1 m, 12 in = 1 ft, 1 in = 2.54 cm
 1 kg = 1000 g, 1 LB = 16 oz
 1000 mL = 1 L; 1 gal = 4 qts.
 1 minute = 60 seconds
Converting between Unit Systems
To convert from one unit to another:
Identify the relationship between the units (e.g.
100 cm = 1 m)
Write out the starting measurement and multiply
it by a quantity that will yield the desired value:
25 cm (
) = _____ m
The number in the “( )” is called the “conversion
factor”
Metric Prefixes
Weight vs. Mass
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Mass is the amount of
“stuff” in an object
Mass is inertia
Mass is the same
everywhere in the
universe
SI Units of mass are
kilograms (kg)
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Weight is the effect
of gravity on an
object’s mass
Weight is a force
Weight depends on
location
SI units of weight
are newtons (N)
USCS units are
pounds (lb)
Volume
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The 3-D space an object occupies
The SI unit is m3 (meters x meters x meters)
The common metric unit is the Liter (L)
Mass and volume are not the same thing
Do not confuse mass & volume
Density
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Density is a property of matter representing the mass per
unit volume
For equal volumes, denser object has larger mass
For equal masses, denser object has small volume
Solids = g/cm3
Mass
 1 cm3 = 1 mL
Density 
Volume
Liquids = g/mL
Gases = g/L
Volume of a solid can be determined by water
displacement
Density : solids > liquids >>> gases
In a heterogeneous mixture, denser object sinks
Using Density in Calculations
Mass
Density 
Volume
Mass
Volume 
Density
Mass  Density  Volume
Ch 100: Fundamentals for
Chemistry
Chapter 3: Matter & Energy
Lecture Notes
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Aristotle (384-322 BC)
Introduced observation as an important
step in understanding the natural world
All types of matter are mixtures of one of
4 basic “elements”:
1) Earth
3) Air
2) Water
4) Fire
All matter has one or more of 4 basic
“qualities”:
1) Cold
3) Hot
2) Moist
4) Dry
According to Aristotle:
 Any substance could be transformed into
another substance by altering the relative
proportion of these qualities (i.e. lead to gold)
Physical & Chemical Properties
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Physical Properties are the characteristics of
matter that can be changed without changing
its composition
 Characteristics that are directly observable
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Chemical Properties are the characteristics
that determine how the composition of matter
changes as a result of contact with other
matter or the influence of energy
 Characteristics that describe the behavior of matter
Physical & Chemical Changes
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Physical Changes are changes to matter
that do not result in a change the
fundamental components that make that
substance
State Changes : boiling, melting, condensing
•
Chemical Changes involve a change in the
fundamental components of the substance
Produce a new substance
Chemical reaction
Reactants  Products
States of Matter
Solid → Liquid → Gas
+Energy
State
Solid
Liquid
Gas
Shape
Keeps
Shape
Takes
Shape of
Container
Takes
Shape of
Container
+Energy
Volume
Compress
Flow
Keeps
Volume
Keeps
Volume
No
No
No
Yes
Takes
Volume of
Container
Yes
Yes
Solid ← Liquid ← Gas
+Energy
+Energy
Classification of Matter
Matter can be classified as either Pure or Impure:
 Pure
 Element: composed of only one type of atom

Composed of either individual atoms or molecules (e.g. O2)
 Compound: composed of more than one type of atom

Consists of molecules
 Impure (or mixture)
 Homogeneous: uniform throughout, appears to be one thing



pure substances
solutions (single phase homogeneous mixtures)
Suspensions (multi-phase homogeneous mixtures)
 Heterogeneous: non-uniform, contains regions with different properties
than other regions
Matter
Pure Substance
Constant Composition
Homogeneous
Mixture
Variable Composition
Separation of Mixtures
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A pure substance cannot be broken down into its
component substances by physical means only by a
chemical process
 The breakdown of a pure substance results in formation of
new substances (i.e. chemical change)
 For a pure substance there is nothing to separate (its only 1
substance to begin with)
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Mixtures can be separated by physical means (and
also by chemical methods, as well)
There are 2 general methods of separation
 Physical separation
 Chemical separation
Methods of Separation
•
There are 2 ways of separating various substances:
1) Physical separation: separation of substances by their physical
properties (such as size, solubility, etc.)
 Mixtures can be separated by physical separation
 There are several methods of separating mixtures


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Filtration (solids from liquids)
Distillation (liquids from liquids)
Centrifugation (liquids from liquids)
2) Chemical separation: separation of substances by their chemical
properties
 Usages:

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Compounds can be separated into their individual elements
Mixtures can be separated by chemical separation as well
 There are several methods of chemical separation

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Ion exchange (such as water purification systems)
Chemical affinity (using antibodies to isolate specific proteins)
Various Chemical reactions
Energy
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The capacity of something to do work
chemical, mechanical, thermal, electrical,
radiant, sound, nuclear
The SI unit of energy is the Joule (J)
 Other common units are
 Calories (cal)
 Kilowatt-hour (kW.hr)
•
Types of energy:
 Potential
 Kinetic
 Heat
•
Energy cannot be created nor destroyed (but it
does change from one type to another!)
Heat & Temperature
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Temperature is _____.
 how hot or cold something is (a physical property)
 related to the average (kinetic) energy of the substance
(not the total energy)
 Measured in units of
 Degrees Fahrenheit (oF)
 Degrees Celsius (oC)
 Kelvin (K)
•
Heat is energy that _____.
 flows from hot objects to cold objects
 is absorbed/released by an object resulting in its change
in temperature
•
Heat absorbed/released is measured by changes in
temperature
Temperature Scales
•
Fahrenheit Scale, °F
 Water’s freezing point = 32°F, boiling point = 212°F
•
Celsius Scale, °C
 Temperature unit larger than the Fahrenheit
 Water’s freezing point = 0°C, boiling point = 100°C
•
Kelvin Scale, K
 Temperature unit same size as Celsius
 Water’s freezing point = 273 K, boiling point = 373 K
Temperature of ice water and boiling water.
Heat
•
Heat is the flow of energy due to a temperature
difference
 Heat flows from higher temperature to lower
temperature
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Heat is transferred due to “collisions” between
atoms/molecules of different kinetic energy
When produced by friction, heat is mechanical
energy that is irretrievably removed from a
system
Processes involving Heat:
1. Exothermic = A process that releases heat energy.
 Example: when a match is struck, it is an exothermic
process because energy is produced as heat.
2. Endothermic = A process that absorbs energy.
 Example: melting ice to form liquid water is an endothermic
process.
Heat (cont.)
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The heat energy absorbed by an object is
proportional to:
The mass of the object (m)
The change in temperature the object undergoes
(DT)
Specific heat capacity (s) (a physical property unique to
the substance)
•
To calculate heat (Q):
Q = s . m . DT
Specific Heat Capacity (s)
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The amount of heat energy (in J or Cal) required to
increase the temperature of 1 gram of a
substance by 1oC (or 1K)
The Units of Specific Heat Capacity:
1. J/goC (SI)
2. cal/goC (metric & more useful in the lab)
•
Specific Heat Capacity is a unique physical
property of different substances
 Metals have low specific heat capacity
 Non-metals have higher specific heat capacity
 Water has an unusually large specific heat capacity
s = Q/(mDT)
Table of Specific Heat for various
substances @ 20oC
0.900
c in cal/gm K or
Btu/lb F
0.215
Molar C
J/mol K
24.3
Bismuth
0.123
0.0294
25.7
Copper
0.386
0.0923
24.5
Brass
0.380
0.092
...
Gold
0.126
0.0301
25.6
Lead
0.128
0.0305
26.4
Silver
0.233
0.0558
24.9
Tungsten
0.134
0.0321
24.8
Zinc
0.387
0.0925
25.2
Mercury
0.140
0.033
28.3
2.4
0.58
111
Water
4.186
1.00
75.2
Ice (-10 C)
2.05
0.49
36.9
Granite
.790
0.19
...
Glass
.84
0.20
...
Substance
c in J/gm K
Aluminum
Alcohol(ethyl)
Ch 100: Fundamentals for
Chemistry
Chapter 4: Elements, Ions & Atoms
Lecture Notes
Dmitri Mendeleev (1834-1907)
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Russian born chemist
Considered one of the
greatest teachers of his time
Organized the known
elements into the first
“periodic table”
 Elements organized by
chemical properties (& by
weight) -> called periodic
properties
 Predicted the existence of 3
new elements
Chemical Symbols & Formulas
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Each element has a unique chemical symbol
Examples of chemical symbols:
 Hydrogen: H
 Oxygen: O
 Aluminum: Al
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Each molecule has a chemical formula
The chemical formula indicates
 the chemical symbol for each of the elements present
 The # of atoms of each element present in the molecule
•
Examples of chemical formulas:
 Elemental oxygen: O2 (2 O atoms per molecule)
 Water: H2O (2 H atoms & 1 O atom)
 Aluminum sulfate: Al2(SO4)3 (2 Al, 3 S & 12 O atoms)
Dalton’s Atomic Theory
1.
2.
3.
4.
5.
Each element consists of individual particles
called atoms
Atoms can neither be created nor destroyed
All atoms of a given element are identical
Atoms combined chemically in definite
whole-number ratios to form compounds
Atoms of different elements have different
masses
The Atom
The atom has 2 primary regions of interest:
1) Nucleus
 Contains protons & neutrons (called nucleons, collectively)
 Establishes most of the atom’s mass
 Mass of 1 neutron = 1.675 x10-27 kg
 Mass of 1 proton = 1.673 x10-27 kg
 Small, dense region at the center of the atom
 The radius of the nucleus ~ 10-15 m (1 femtometer)
2) The Electron Cloud
 Contains electrons
 Mass of 1 electron = 9.109 x10-31 kg
 Establishes the effective volume of the atom
 The radius of the electron cloud ~ 10-10 m (1 Angstrom)
 Determines the chemical properties of the atom
 During chemical processes, interactions occur between the outermost
electrons of each atom
 The electron properties of the atom will define the type(s) of interaction
that will take place
Structure of the Atom
Electric Charge
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Electric charge is a fundamental property of matter
We don’t really know what electric charge is but we do know
that there are 2 kinds:
 Positive charge (+)
 Negative charge (-)
•
Opposite charge polarity is attractive:
+ attracts -
•
Same charge polarity is repulsive:
+ repels + and
•
•
– repels –
The magnitude of electric charge (q) is the same for protons
and electrons:
The charge of a proton or electron is the smallest amount
that occurs in nature, it is called the quantum of charge:
 qproton = +1.602 x 10-19 Coulombs
 qelectron = -1.602 x 10-19 Coulombs
What holds the atom together?
•
Electromagnetic interaction (a.k.a. electric force) holds
the electrons to the nucleus
 The negative charge (-) of the electrons are attracted to
the positive charge (+) of the nucleus
•
Strong interaction (a.k.a. strong force) holds the nucleons
together within the nucleus
 The positive charge of the protons repel each other
 All nucleons, protons and neutrons, possess a STRONG
attraction to each other that overcomes the protons’
mutual repulsion
Atomic Bookkeeping
•
Atomic number (Z)
 The number of protons in an atom
 The number of protons in an uncharged atom
 Determines the identity of the atom
•
Mass number (A)
 The number of protons & neutrons in an element
 Determines the weight of the atom
•
To determine number of neutrons in an atom:
# of neutrons = (Mass #) – (Atomic #)
Or
# of neutrons = A - Z
Mass # vs. Atomic Mass
•
Isotopes are the equivalent of sibling members of an
element
 Unique atoms of the same element with different mass numbers (i.e.
they have different numbers of neutrons)
 Unique isotopes are identified by their mass number
•
•
•
Isotope notation:
Mass #
Atomic #
(Atomic Symbol)
14
12
Example: carbon-12 ( C ) & carbon-14 ( C )
6
6
Atomic mass
 The average total mass of an element’s various naturally occuring
isotopes
 The unit of Atomic Mass is the Dalton (formerly called the amu)
 1 Dalton = one twelfth mass of one 12C atom = 1.661x10-27 kg
 Note: There 6 protons & 6 neutrons in a 12C atom but the mass of a 12C
atom is actually less than the combined mass of all of the nucleons
individually.
 Where is this lost mass? It’s released as energy when the nucleons
combine (bind) to form the nucleus of the atom.
Examples of Isotopes
The Periodic Table
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•
•
All of the known elements are arranged in a
chart called the Periodic Table
The elements are arranged by similarity of
chemical properties
Each element is identified by its Atomic
Number
The elements are organized left-to-right and
top-to-bottom by their Atomic Number
The columns are called Groups
 Elements of each group have similar properties
•
The rows are called Periods
Elements and the Periodic Table
The elements can be categorized as
Metals
The leftmost elements of the periodic table
Roughly 70% of all of the elements
Nonmetals
The rightmost elements of the periodic table
Semimetals (metalloids)
The elements between the metals and nonmetals
Properties are not quite metal or non-metal
Ions
•
•
•
•
Atoms (or molecules) that have gained or
lost one or more electrons
Ions that have lost electrons are called
cations
Ions that have gained extra electrons are
called anions
Ionic compounds have both cations and
anions (so that their net charge is zero)
•
Ions (cont.)
Ions have electric charge:
“+” when 1 or more electrons are lost
“-” when 1 or more electrons are gained
•
When an atom/molecule is an ion, its charge
must be specified:
 Sodium ion:
 Chloride ion:
 Hydroxide ion:
•
Na+
ClOH-
Notes on Electric Charge:
 Opposite charges attract
+
-
 Like charges repel
+
+
-
-
Ch 100: Fundamentals for
Chemistry
CH 100: Chemical Nomenclature
(a.k.a. naming compounds)
Antoine Lavoisier (1743-1794)
•
•
Considered by many to be the
“Father of Modern Chemistry”
Major contributions included
 Demonstrated that water cannot be
transmuted to earth
 Established the Law of Conservation
of Mass
 Developed a method of producing
better gunpowder
 Observed that oxygen and hydrogen
combined to produce water (dew)
 Invented a system of chemical
nomenclature (still used in part today!)
 Wrote the 1st modern chemical
textbook
Types of Compounds
•
•
•
When compounds are formed they are held
together by the association of electrons
This association is called a chemical bond
There are 3 general types of chemical bonds:
1. Ionic
2. Covalent (or molecular)
3. Polar covalent
•
Simple compounds are classified (and thus
named) according to the type of chemical bond(s)
that hold together its atoms
Note: many compounds have more than one type of
chemical bond present, but we will only work with
“simple compounds”
Types of Compounds (cont.)
For “practical” purposes will separate compounds into
2 general categories:
• Ionic
 Made up of ions (both positive and negative charge)
 Must have no net charge (i.e. combined charge of zero)
 Depend on the attraction between positive and negative
charges of the ions
 Usually a metal is present as a cation and a nonmetal is
present as an anion
•
Molecular (or covalent)
 Made up of atoms that share their outer electrons
 Charge plays no direct role in their formation
 Usually no metals are present
Naming Compounds
•
Easiest way to identify an ionic compound is
to ask whether or not it has a metal present:
Yes -> ionic (e.g. CaCl2)
No -> covalent (e.g. CCl4)
•
•
Covalent compounds require the use of
Greek prefixes to indicate the number of
each element present in one molecule
Ionic compounds do not use the Greek
prefixes
Naming Simple Compounds
A “simple” or binary compound is a compound made
of only 2 types of elements
•
When the first element is a metal:
• The first element (metal) keeps its full name
• The non-metal goes by its root with the suffix “-ide”
added to the end
Example: NaCl is sodium chloride
•
When there are no metals present
• Same as above except
• Greek prefixes must be used to identify the number of
each element present in the compound
Example: CO2 is carbon dioxide
Ionic Charges & the Periodic Table
 Group 1 metals form 1+ cations (Na+ sodium ion)
 Group 2 metals form 2+ cations (Ca2+ calcium ion)
 Group 13 metals form 3+ cations (Al3+ aluminum ion)
 All other metals (i.e. the transitional metals, Pb, etc.) form
more than one type of cation
 Roman numerals must be used to indicate the charge of the
cation
 Example:
Fe3+ is called iron(III)
FeCl3 is called iron(III) chloride
 Exceptions:
Ag+, Cd2+ & Zn2+
 Group 15 nonmetals form 3- anions (N3- nitride ion)
 Group 16 nonmetals form 2- anions (O2- oxide ion)
 Group 17 nonmetals form 1- anions (Cl- chloride ion)
 Group 18 nonmetals do not form ions
Greek Prefixes for Compound Names
1)
2)
3)
4)
5)
MonoDiTri
TetraPenta-
CCl4 is carbon tetrachloride
Notes:
6) Hexa7) Hepta8) Octa9) Nona10) DecaC3H8 is tricarbon octahydride
1) Prefixes are used when the compound does not have a metal
present (or when H is the first element in the formula)
2) Prefixes must be used for every element present in the compound
3) mono- is not used for the first element in a compound name (e.g.
carbon dioxide)
Ionic Compounds containing
Polyatomic ions
Some ionic compounds are made up of
polyatomic ions
• When you encounter this, do not freak out!!
• Become familiar with the polyatomic ions on
the handout
Example: the nitrate ion (NO3-)
• The naming of this type of compound is
similar to that for ionic compounds
•
Acids
 From the Latin term for “sour”{Acids are sour to the taste}
 Acids are substances that donate protons (H+) (usually
when dissolved in water)
 Chemical formula usually begins with H
Example: hydrochloric acid
HCl(aq) + H2O(l)  H3O+ + Cl- (aq)
Bases
 Taste bitter
 Usually metal containing hydroxides
 Substances that accept protons (H+) when dissolved in
water
Example: potassium hydroxide
KOH(aq) + H3O+  K+(aq) + H2O (l)
Naming Acids
Lets separate acids into 2 types:
 Acids that contain oxygen
 Acids that do not contain oxygen
Naming acids containing oxygen:
 For acids containing “-ate” anions:
1. Use root of the anion (for sulfate, SO42-, use sulfur)
2. Add “-ic” suffix then end with “acid”
Example:
H2SO4 is sulfuric acid
 For acids with “-ite” anions:
1. Use root of the anion (for sulfite, SO32-, use sulfur)
2. Add “-ous” suffix then end with “acid”
Example:
H2SO3 is sulfurous acid
Naming Acids (cont.)
Naming acids not containing oxygen:
 Add “hydro-” prefix to beginning
 Use root of the anion (i.e. Cl- use chlor)
 Add “-ic” suffix then end with “acid”
Example:
HCl is hydrochloric acid
Name the following acids:
HF
HNO2
HCN
H3PO4
Ch 100: Fundamentals for
Chemistry
Chapter 6: Chemical Reactions
Chemical Reactions (Intro)
When matter undergoes chemical changes these
processes are called chemical reactions
• Substances that undergo the change(s) are called
the reactants
• The resulting substances are called the products
• Standard form of a chemical reaction:
Reactant(s)  Product(s)
Example:
2H2 (g) + 1O2 (g)  2H2O (g)
•
•
The underlined numbers are called coefficients.
 The number of each molecule for each reactant &
product in the chemical reaction
 They are always whole numbers
Chemical Reactions (cont.)
Balanced chemical reactions indicate the ____
identity of each reactant & product involved in
the reaction
phase of each reactant and product involved in
the reaction (i.e. solid (s), liquid (l) or gas (g))
relative quantity of each reactant and product
involved in the reaction (the coefficients!)
relative molar quantity of each reactant and
product involved in the reaction (the
coefficients!)
Rates of Chemical Reactions
•
How quickly a chemical reaction occurs is
indicated by its reaction rate
 How quickly the concentration of products increases
 How quickly the concentration of reactants decreases
•
Factors that influence reaction rates:
 Reactants must be in contact
 Reactions occur due to collisions
 Without contact between reactants there can be no reaction
 Concentration of reactants
 The more reactant molecules packed into a given space the
more likely a collision (& reaction) will occur
 Temperature
 the average KE of each reactant affects how much energy will
be transferred between reactants during a molecular collision
 Molecules must transfer enough KE to break the existing bonds
Energy in Chemical Reactions
Exothermic Reactions
Internal
Energy
Activation
Energy (EA)
Reactants
Energy
Released (Q)
Products
Endothermic Reactions
Internal
Energy
Activation
Energy (EA)
Products
Energy
Absorbed (Q)
Reactants
Energy in Reactions (cont.)
Example: Sodium Water Reaction
Internal
Energy
Low Activation
Energy (EA)
2Na(s) + 2H2O(l)
Large amount of
Energy Released
(Q)
2NaOH(aq) + H2(g)
Catalysts
•
Catalysts are substances that speed up chemical
reactions
 Allow reactions to occur that might not otherwise take
place (due to low temperature for example)
 Lower activation energy for a chemical reaction
 Do not participate in the reaction
 They may undergo a chemical change as a reactant but they are
always recycled as a product (so there is no net change in the
catalyst molecule)
Catalysts are indicated in a chemical reaction by
placing the chemical formula over/under the
reaction arrow.
Example:
catalyst
Reactants  Products
•
Catalysts & Energy in Reactions
Catalysts lower Activation Energy
Activation Energy
without catalyst
Internal
Energy
Reactants
Activation Energy
with catalyst
Products
Endothermic or Exothermic?
(that is the question…)
In chemical reactions:
 Energy is required to break bonds (energy absorbed)
 Energy is released when bonds are formed
• The amount of energy required to break a chemical bond is
the same as the energy released when the bond is formed,
this is called Bond Energy
• During a chemical reaction:
 Energy is absorbed equal to the bond energies for all
bonds broken in the reactants
 Energy is released equal to the bond energies for all
bonds formed in the products
• Endothermic reactions absorb more energy than they
release
• Exothermic reactions release more energy than they absorb
Balancing Chemical Reactions
•
According to the Law of Mass Conservation (& John
Dalton!) matter is never created nor destroyed
 All atoms in the reactants of a chemical reaction must be
accounted for in the products
•
The Basic Process:
 Identify all reactants & products in the reaction & write out
their formulas (this is the unbalanced chemical equation)
 Count the number of each atom for each compound for each
reactant & product
(these values must be the same for both reactants & products when the reaction
is balanced!)
 Starting with the most “complicated” molecule,
systematically adjust the coefficients to balance # of the
atoms on each side of the reaction (balance one atom at a
time)
 Repeat until all atoms are balanced for the reaction
 Now you have a balanced chemical equation!
Balancing Chemical Reactions
(example)
When sodium metal is added to water a violent
reaction takes place producing aqueous
sodium hydroxide and releasing hydrogen
gas.
1. Write out the unbalanced chemical reaction:
2.
Now, balance the chemical reaction:
Balancing Chemical Reactions (Hint)
•
•
When a polyatomic ion(s) appears on both
the reactant & product side of the reaction
unchanged, treat the whole ion as a “unit”
when balancing the reaction
Example:
AgNO3(aq) + CaCl2 (aq) AgCl(s) +
•
•
•
Ca(NO3)2(aq)
Note the nitrate ion (NO3-) gets swapped
between the Ag + and the Ca2+ ions in this
reaction
So NO3- can be treated as a whole unit when
balancing this reaction
Balance it!
Ch 100: Fundamentals for
Chemistry
Chapter 7: Chemical Reactions in
Aqueous Solutions
Driving Forces & Chemical Reactions
•
•
The tendency for reactants to undergo chemical
changes (reactions) to form products are called
“driving forces”
There are 4 common “driving forces”:
1.
2.
3.
4.
•
•
Formation of a solid (precipitation reaction)
Formation of water (acid-base reaction)
Transfer of electrons (oxidation-reduction reaction)
Formation of a gas (bad taco reaction )
When 2 or more chemicals are brought together,
if any of these things can happen, a chemical
change is likely to occur
When one of these processes occurs, we
describe the resulting chemical reaction based on
the driving force
Solubility
•
A measure of how much of a solute will dissolve in a solvent
is called its solubility
 Solubility is temperature dependent
 Solid solubility increases with increased temperature (i.e. you can
dissolve more sugar in hot water than in cold water)
 Gas solubility increases with decreased temperature (i.e. you can
dissolve more CO2 in cold water than hot water)
 A solute is soluble if any of it will dissolve in a solvent
 NaCl is soluble in water
•
A solute is insoluble if no appreciable amount of it will
dissolve in solvent
 AgCl is insoluble in water
•
When 2 solutions are combined and result in the formation of
an insoluble product:
 The product will not dissolve in the solvent
 The product will form a precipitate
Precipitation (formation of a solid) is one indication that a
chemical change has occurred!
Precipitation Reactions
•
•
in all precipitation reactions, the ions of one
substance are exchanged with the ions of another
substance when their aqueous solutions are
mixed
At least one of the products formed is insoluble in
water
KI(aq) + AgNO3(aq)  KNO3(aq) + AgIs
K+
Ag+
K+
Ag
I-
NO3-
NO3-
I
Dissociation
•
ionic compounds
 metal + nonmetal (Type I & II)
 metal + polyatomic anion
 polyatomic cation + anion
•
•
when ionic compounds dissolve in water the
anions and cations are separated from each
other; this is called dissociation
we know that ionic compounds dissociate when
they dissolve in water because the solution
conducts electricity
•
Dissociation (examples)
potassium chloride dissociates in water into
potassium cations and chloride anions
KCl(aq) = K+ (aq) + Cl- (aq)
K
•
Cl
K+
Cl-
copper(II) sulfate dissociates in water into
copper(II) cations and sulfate anions
CuSO4(aq) = Cu+2(aq) + SO42-(aq)
Cu SO4
Cu+2
SO42-
Dissociation (example)
•
potassium sulfate dissociates in water into
potassium cations and sulfate anions
K2SO4(aq) = 2 K+ (aq) + SO42-(aq)
K
SO4 K
K+
SO42K+
Process for Predicting the Products of
a Precipitation Reaction
1)
2)
Determine what ions each aqueous reactant has
Exchange Ions
 (+) ion from one reactant with (-) ion from other
3)
4)
Balance Charges of combined ions to get formula
of each product
Balance the Equation
 count atoms
5)
Determine Solubility of Each Product in Water
 solubility rules
 if product is insoluble or slightly soluble, it will
precipitate
Solubility Rules
1.
2.
3.
4.
5.
6.
Most compounds that contain NO3- ions are soluble
Most compounds that contain Na+, K+, or NH4+
ions are soluble
Most compounds that contain Cl- ions are soluble,
except AgCl, PbCl2, and Hg2Cl2
Most compounds that contain SO42- ions are
soluble, except BaSO4, PbSO4, CaSO4
Most compounds that contain OH- ions are slightly
soluble (will precipitate), except NaOH, KOH, are
soluble and Ba(OH)2, Ca(OH)2 are moderately
soluble
Most compounds that contain S2-, CO32-, or PO43ions are slightly soluble (will precipitate)
Ionic Equations
•
•
equations which describe the chemicals put into the water and
the product molecules are called molecular equations
KCl(aq) + AgNO3(aq)  KNO3(aq) + AgCl(s)
equations which describe the actual ions and molecules in the
solutions as well as the molecules of solid, liquid and gas not
dissolved are called ionic equations
K+ (aq) + Cl- (aq) + Ag+ (aq) + NO3- (aq) K+ (aq) + NO3- (aq) + AgCl(s)
•
ions that are both reactants and products are called spectator
ions
K+ (aq) + Cl- (aq) + Ag+ (aq) + NO3- (aq) K+ (aq) + NO3- (aq) + AgCl(s)
•
an ionic equation in which the spectator ions are dropped is
called a net ionic equation
Cl- (aq) + Ag+ (aq) AgCl(s)
Electrolytes
•
•
•
•
•
electrolytes are substances whose aqueous
solution is a conductor of electricity
all electrolytes have ions dissolved in water
in strong electrolytes, virtually all the molecules
are dissociated into ions
in nonelectrolytes, none of the molecules are
dissociated into ions
in weak electrolytes, a small percentage of the
molecules are dissociated into ions
Reactions that Form Water:
Acids + Bases
•
•
•
•
Acids all contain H+ cations and an anion
Bases all contain OH- anions and a cation
when acids dissociate in water they release
H+ ions and their anions
when bases dissociate in water they release
OH- ions and their cations
Acid-Base Reactions
•
•
in the reaction of an acid with a base, the H+ from
the acid combines with the OH- from the base to
make water
the cation from the base combines with the anion
from the acid to make the salt
acid + base salt + water
•
H2SO4(aq) + Ca(OH)2(aq)  CaSO4(aq) + 2 H2O(l)
the net ionic equation for an Acid-Base reaction is
always
H+ (aq) + OH- (aq)  H2O(l)
Reactions of Metals with Nonmetals
(Oxidation-Reduction)
•
The metal loses electrons and becomes a
cation
We call this process oxidation
•
The nonmetal gains electrons and becomes
an anion
We call this process reduction
•
In the reaction, electrons are transferred
from the metal to the nonmetal
Oxidation-Reduction Reactions
•
•
All reactions that involve a transfer of one or
more electrons are called oxidationreduction reactions
We say that the substance that loses
electrons in the reaction is oxidized and the
substance that gains electrons in the reaction
is reduced.
Predicting Products of
Metal + Nonmetal Reactions
•
metal + nonmetal  ionic compound
 ionic compounds are always solids unless dissolved in
water
•
•
•
in the ionic compound the metal is now a cation
in the ionic compound the nonmetal is now an
anion
to predict direct synthesis of metal + nonmetal
1) determine the charges on the cation and anion
(from their position on the Periodic Table)
2) determine numbers of cations and anions needed to
have charges cancel
3) balance the equation
Another Kind of
Oxidation-Reduction Reaction
•
•
•
Some reactions between two non-metals are also
oxidation-reduction reaction
Any reaction in which O2 is a reactant or a product
will be an oxidation-reduction reaction
Examples:
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g)
2 SO3(g)  2 SO2(g) + O2(g)
Ways to Classify Reactions
•
•
•
Reactions that involve solid formation are
called precipitation reactions
Reactions that involve water formation are
called acid-base reactions
Both precipitation reactions and acid-base
reactions involve compounds exchanging
ions, ion exchange reactions are called
double displacement reactions
Double Displacement Reactions
•
•
•
two ionic compounds exchange ions
X Y (aq) + AB (aq)  XB + AY
reaction will not occur unless one of the
products either (1) precipitates or (2) is water
Ways to Classify Reactions
•
Reactions that involve electron transfer are called
oxidation-reduction reactions
 Metals + Nonmetal
 O2 as a reactant or product
Reactions that occur in aqueous solution where one
of the products is a gas are called gas forming
reactions
NaHCO3(aq) + HCl(aq)  NaCl(aq) + CO2(g) + H2O(l)
•
Ways to Classify Reactions
•
•
Reactions that involve one ion being
transferred from one cation to another are
called single replacement reaction
X Y + A  X + AY
Examples:
Zn(s) + 2 HCl(aq)  ZnCl2(aq) + H2(g)
Fe2O3(s) + 2 Al(s)  2 Fe(s) + Al2O3(s)
Other Ways to Classify Reactions
•
•
•
•
•
Reactions in which O2(g) is reacted with a
carbon compound are called Combustion
Reactions
Combustion reactions release a lot of energy
Combustion reactions are a subclass of
Oxidation-Reduction reactions
Combustion of carbon compounds produces
CO2(g)
Combustion of compounds that contain
hydrogen produces H2O(g)
C3H8(g) + 5 O2(g)  3 CO2(g) + 4 H2O(g)
Other Ways to Classify Reactions
•
•
•
•
Reactions in which chemicals combine to
make one product are called Synthesis
Reactions
Metal + Nonmetal reactions can be classified
as Synthesis Reactions
2 Na(s) + Cl2(g)  2 NaCl(s)
Reactions of Metals or Nonmetals with O2
can be classified as Synthesis Reactions
N2(g) + O2(g)  2 NO(g)
These two types of Synthesis Reactions are
also subclasses of Oxidation-Reduction
Reactions
Other Ways to Classify Reactions
•
•
Reactions in which one reactant breaks
down into smaller molecules are called
Decomposition Reactions
Generally initiated by addition of energy
Addition of electric current or heat
•
Opposite of a Synthesis Reaction
2 NaCl(l)  2 Na(l) + Cl2(g)
electric
current
Ch 100: Fundamentals for
Chemistry
Chapter 8 Lecture Notes
(Sections 8.1 to 8.5)
Amadeo Avogadro
(1743-1794)
•
•
Italian lawyer turned chemist
Major contributions included:
 Established difference between atoms & molecules:
 Oxygen & nitrogen exist as molecules O2 & N2
 Reconciled the work of Dalton & Guy-Lussac
 Establishing Avogadro’s Principle: equal volumes of all gases
at the same temperature and pressure contain the same number of
molecules.
•
Did not determine Avogadro’s number nor the mole
(these concepts came later)
 Avogadro is honored because the molar volume of all
gases should be the same
 Much of Avogadro’s work was acknowledged after he
died, by Stanislao Cannizarro
The Mole
•
A counting unit (similar to the dozen)
 A large unit used to describe large quantities such as
number of atoms
•
•
1 mole = 6.022 x 1023 units
6.022 x 1023 is known as Avogadro’s number (NA)
Relationship between the mole & the Periodic
Table
 The atomic mass is the quantity (in grams) of 1 mole of
that element
 The units of atomic mass are grams/mole
 Mass is used by chemists as a way of “counting”
number of atoms/molecules of a substance
•
Mole calculations
Got mole problems?
Call Avogadro at 602-1023.
What do you get if you have
Avogadro's number of
donkeys?
Answer: molasses (a mole of asses)
Molar Mass
•
•
•
Mass in grams of 1 mole of a substance
Refers to both atoms & molecules
Elements (atoms)
Examples:
1 mole of Na has a mass of 22.99 g
1 mole of Cl has a mass of 35.45
1 mole of Cl2 has a mass of 70.90 g
•
Compounds (molecules)
Examples:
1 mole of NaCl has a mass of 58.44 g
 Mass of Na (22.99 g) + Mass of Cl (35.45 g)
1 mole of CO2 has a mass of 44.01 g
 Mass of C (12.01 g) + 2 x Mass of O (16.00 g)
Mole Calculations (1)
•
Atoms/Molecules to Moles
 Divide # of atoms (or molecules) by Avogadro’s #
Example: How many moles are 1.0x1024 atoms?
 1 mole 
(1.0 10 atoms) 
 1.7moles
23 
 6.022 10 
24
•
Moles to Atoms/Molecules
 Multiply # of atoms (or molecules) by Avogadro’s #
Example: How many molecules are in 2.5 moles?
 6.022 1023 
24
(2.5 moles) 

1.5

10
molecules

 1 mole 
Mole Calculations (2)
•
Moles to Grams
 Multiply the # of moles by atomic mass
Example: How many grams in 2.5 moles of carbon?
 12.01 grams 
1
(2.5 moles) 

30.
grams
(
or
3

10
)

 1 mole 
•
Grams to Moles
 Divide the mass in grams by atomic mass
Example: How many moles are in 2.5 grams of lithium?
 1 mole 
1
(2.5 grams) 

0.36
moles
(
or
3.6

10
)

 6.941 grams 
Percent Composition
•
•
Percentage of each element in a compound (by
mass)
Can be determined from:
1. the formula of the compound or
2. the experimental mass analysis of the compound
 part 
% Composition  
 100%
 whole 
Note: The percentages may not always total to 100%
due to rounding
•
Percent Composition Calculations
To determine % Composition from the chemical
formula:
 Determine the molar mass of compound
 Multiply the molar mass of the element of interest by the
number of atoms per molecule then
 Divide this value by the molar mass of the compound
 (# atoms of A)(atomic mass of A) 
% Composition of A  
 100%
molar mass of compound


Example: The % Composition of sodium in table salt
1. The molar mass of NaCl is 58.44 g/mol
2. There is 1 atom of Na in each NaCl molecule
3. The atomic mass of Na is 22.99
 1 22.99 
% Composition of Na  
 100%  39.33%
 58.44 
Percent Composition Calculations
Perform the following % Composition
calculations:
1.The % composition of carbon in carbon
monoxide
2.The % composition of oxygen in water
3.The % composition of chlorine in sodium
hypochlorite
Ch 100: Fundamentals for
Chemistry
Ch 9: More on Chemical Reactions
Lecture Notes (Sections 9.1 to 9.2)
Chemical Equations:
What do they tell us?
•
A properly written chemical equation will
provide the following information:
1. All reactants & products involved in the
reaction
2. The physical state of all reactants & products
3. The presence of any catalysts involved in the
chemical reaction
4. The relative quantity of all reactants &
products
Information Given by the
Chemical Equation
•
Balanced equation provides the relationship
between the relative numbers of reacting
molecules and product molecules
2 CO + O2  2 CO2
2 CO molecules react with 1 O2 molecules to
produce 2 CO2 molecules
•
Information Given by the
Chemical Equation
Since the information given is relative:
2 CO + O2  2 CO2
200 CO molecules react with 100 O2 molecules to produce
200 CO2 molecules
2 billion CO molecules react with 1 billion O2 molecules to
produce 20 billion CO2 molecules
2 moles CO molecules react with 1 mole O2 molecules to
produce 2 moles CO2 molecules
12 moles CO molecules react with 6 moles O2 molecules to
produce 12 moles CO2 molecules
Information Given by the
Chemical Equation
•
•
The coefficients in the balanced chemical
equation shows the molecules and mole
ratio of the reactants and products
Since moles can be converted to masses, we
can determine the mass ratio of the reactants
and products as well
Information Given by the
Chemical Equation
2 CO + O2  2 CO2
2 moles CO = 1mole O2 = 2 moles CO2
Since 1 mole of CO = 28.01 g, 1 mole O2 = 32.00
g, and 1 mole CO2 = 44.01 g
2(28.01) g CO = 1(32.00) g O2 = 2(44.01) g CO2
Example:
Determine the Number of Moles of Carbon Monoxide
required to react with 3.2 moles Oxygen, and determine the
moles of Carbon Dioxide produced
1.
Write the balanced equation
2 CO + O2  2 CO2
2.
Use the coefficients to find the mole
relationship
2 moles CO = 1 mol O2 = 2 moles CO2
Example (cont.)
Determine the Number of Moles of Carbon Monoxide
required to react with 3.2 moles Oxygen,
and determine the moles of Carbon Dioxide produced
3.
Use dimensional analysis
2 moles CO
3.2 moles O2 x
 6.4 moles CO
1 mole O2
2 moles CO2
3.2 moles O2 x
 6.4 moles CO2
1 mole O2
Ch 100: Fundamentals for
Chemistry
Ch 14: Solutions & Concentration
Lecture Notes
Solutions
•
•
Solutions are single phase homogenous mixtures
Solutions consist of:
 Solvent: the component in largest quantity
 Solute(s): the other components
•
•
•
The solute is considered to be dissolved in the
solvent
When a solution has not reached its limit of
dissolved solute it is an unsaturated solution
When a solution has reached its limit of dissolved
solute and any added solute will not dissolve, it is
a saturated solution
Concentration
•
•
A measure of how much of a substance (solute)
is dissolved in another substance (solvent)
To calculate [concentration]:
conc. of
•

amount of solute 
solute 
amount of solvent 
Common usages of concentration:
 Mass (m/v) conc. (units are grams/L, grams/mL, etc.)
 Volume (v/v) conc. (unit-less, often % is used)
 Molarity (units are moles/L or M)
Mass Percent (%)
•
•
Concentration of a solute dissolved in a solvent
(in grams per unit gram of solution)
To determine mass percent
 Divide mass of solute (in grams) by the total mass of
solution (in grams) and multiply this ration by 100%
 mass of solute 
 100%
Mass %  
 mass of solution 
Example: What is the mass percent of 30.0 grams of NaCl
in a 150.0 gram solution?
 30.0 grams NaCl 
 100%  20.0%
Mass %  
 150.0 grams solution 
Questions: (a) How much CaCl2 is in 250.0 grams of solution
where the mass percent of CaCl2 is 30.0%? (b) How much Clis in this solution?
Molarity
•
•
•
Concentration of a solute (in moles per unit
volume) dissolved in a solvent
SI units are moles/liter, or M (molarity or molar
concentration)
To determine molarity from mass concentration
 Simply a unit conversion from grams to moles (using
atomic or molar mass as the unit conversion)
Example: What is the molarity of a NaCl solution with
concentration of 30.0 grams/L?
The molar mass of NaCl is 22.99 + 35.45 = 58.44 grams/mol
 30.0 grams 

1 mol



  0.513 M
NaCl  


1L

 58.44 grams 
Question: What is the molarity of Na+ and Cl- in this
solution?
The pH Scale
•
•
•
The acidity (or concentration of H+) of a solution is
often measured using the pH scale
The pH scale is based on the molarity of H+ ions in
solution
The pH scale ranges from 0 (acidic) to 14 (basic)
 When pH=7.0 the solution is neutral acidity, there is
equal concentration of H+ and OH- in the bulk liquid
•
To calculate pH from [H+] (in mol/L): pH=-log10[H+]
Example: a solution with [H+]=1.0x10-5 M
[H+]=1.0x10-5 M then pH = -log10(1.0x10-5) = 5.0
•
To calculate [H+] from pH:
[H+]=10-pH
Example: a solution with pH = 9.0
pH = 9.0 then [H+]=10-9.0 M = 1.0x10-9 M
pH Concept Questions
What is the [H+] for a 0.5 M HCl solution?
 0.5 M
• What is the [H+] for a 0.5 M H3PO4 solution?
 Less than 0.5 M
• Do the 2 solutions above have the same pH?
 No, pH depends on [H+] not [acid]
• Why or why not?
 HCl is a strong acid but H3PO4 is a weak acid
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How does a strong acid differ from a weak acid?
 Strong acids dissociate all of their H+ ions when in
water whereas weak acids do not!
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