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Ch 100: Fundamentals for Chemistry
Chapter 1: Introduction
Lecture Notes
What is Chemistry?
• Chemistry is often described as the “central”
science
• Chemistry is the study of matter
• Matter is the “stuff” that makes up the universe, i.e.
anything that has mass and occupies space
• The fundamental questions of Chemistry are:
1. How can matter be described?
2. How does one type of matter interact with other
types of matter?
3. How does matter transform into other forms of
matter?
Major Developments in Chemistry I
~400 BC: Democritus proposed the concept of the “atom”
st
~300 BC: Aristotle developed 1 comprehensive model of matter
~700 AD: Chinese alchemists invent gunpowder
1661:
Robert Boyle proposed the concept of elements
1770-90: 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
st
1869:
Mendeleev creates the 1 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 Bose-Einstein condensate) in the laboratory
Scientific Method
1. (OBSERVATION) Recognize a problem
– Make observation
– Formulate a question
2. (EXPLANATION) Make an educated guess - a
hypothesis
– Predict the consequences of the
hypothesis
3. (VALIDATION) Perform experiments to test
the predictions
– Does experimental data support or
dispute hypothesis?
4. Formulate the simplest rule that organizes the
3 main ingredients - develop a theory
EXPLANATIONS
Hypothesis
Theory
Tentative Explanation
of
Certain Facts
Explanation of the
General Principles
of Certain Phenomena
Provides a
Basis for
Further Experimentation
Considerable Evidence
or Facts
Support It
Law
Simple Statement
of Natural Phenomena
No Exceptions
Under the Given
Conditions
Bottom Line: The Scientific Attitude
• 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”
The Particulate Nature of Matter
• Matter is the tangible substance of nature, anything
with mass that occupies space
• At the most fundamental level, matter is discrete or
particulate in nature
• The smallest, most basic units of matter are called
atoms
• All matter is thus comprised of individual atoms, or
specific combinations of atoms called molecules
• Molecules can be broken apart into their constituent
atoms but atoms cannot be further broken apart and
still retain the properties of matter
• Matter can exist in one or more physical states (or
phases)
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
Matter
Pure Substance
Constant Composition
– Impure (or mixture)
Mixture
Variable Composition
Homogeneous
• 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
Separation of Matter
•
•
•
A pure substance cannot be broken down into its
component substances by physical means only by a
chemical process
1. The breakdown of a pure substance results in
formation of new substances (i.e. chemical change)
2. For a pure substance there is nothing to separate (its
only 1 substance to begin with)
Mixtures can be separated by physical means (and also
by chemical methods, as well)
There are 2 general methods of separation
1. Physical: separation based on physical properties
a. Filtration
b. Distillation
c. Centrifugation
2. Chemical: separation based on chemical properties
Ch 100: Fundamentals for Chemistry
Chapter 2: Measurements & Calculations
Lecture Notes
Types of Observations
• Qualitative
– Descriptive/subjective in nature
– Detail qualities such as color, taste, etc.
Example: “It is really warm outside today”
• Quantitative
– Described by a number and a unit (an accepted reference scale)
– Also known as measurements
• Notes on Measurements:
• 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
Example: “The temperature is 85oF outside today”
Application of Scientific Notation
Writing numbers in Scientific Notation
1 Locate the Decimal Point
2 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 Scientific Notation numbers in Conventional 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 other countries
• 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 sub-set set of metric units
• Used by scientists and most science textbooks
• Not always the most practical unit system for lab work
Measurements & the Metric System
• 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
• When a measurement has a specific metric unit (i.e. 25 cm) it
can be expressed using different metric units without
changing its meaning
Example: 25 cm is the same as 0.25 m or even 250 mm
• The choice of measurement unit is somewhat arbitrary, what
is important is the observation it represents
Measurement, Uncertainty & Significant Figures
• 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
• 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
Examples:
1. The measurement: 25.2 cm uncertainty: 0.1 cm
2. The measurement: 25.20 cm uncertainty: 0.01 cm
3. The measurement: 25.200 cm uncertainty: 0.001 cm
Rules for Counting Significant Figures
• 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
• If the digit to be removed is
1. less than 5, the preceding digit stays the same
2. equal to or greater than 5, the preceding digit is increased by 1
• 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
Exact Numbers are numbers that are assumed to have
unlimited number of significant figures are considered to
be known with “absolute” certainty. You do not need to
consider or count significant figures for exact numbers.
The following are considered exact numbers for CH100:
1. Counting numbers, such as:
•
•
The number of sides on a square
The number of apples on a desktop
2. Defined numbers such as those used for conversion
factors, such as:
•
•
•
•
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
3. Numbers or constants defined in equations, such as:
•
y = 3x + 15 (both the “3” and the “15” are exact numbers)
Converting between Unit Systems
•
•
•
Converting units from one unit system to another (especially
within the Metric system) can appear daunting at first
glance. However, with a little guidance, and a lot of practice,
you can develop the necessary skill set to master this
process.
To begin, here is a simple mnemonic to guide you through
the unit conversion process:
1. Eliminate
2. Replace
3. Relate
All unit conversions, regardless of how complex they appear,
involve these 3 simple steps. In the following sections, you
will be stepped through the unit conversion process using
these 3 words as a guide.
Example: Unit Conversion
1. Convert 25.0 m to cm
2. Convert 1.26 g to kg
Metric Prefixes
Mass
•
•
•
•
Mass is the quantity of matter in a substance
Mass is measured in units of grams
Mass does not reflect how much volume something has
The kilogram (kg) unit is the preferred unit of mass in the SI
system.
– 1 kilogram is equal to the mass of a platinum-iridium cylinder
kept in a vault at Sevres, France.
– 1 kg has the weight equivalent (on Earth) of 2.205 lb
Conservation of Mass: The total quantity of mass is never
created nor destroyed during a chemical process
Distinguishing Mass vs. Weight
• The terms mass and weight are commonly used
interchangeably but they are fundamentally different!
• The following are some important differences between
mass and weight:
• Weight is the effect (or
• Mass is a fundamental
force) of gravity on an
property of matter, it is the
object’s mass
amount of “stuff” in an object
• Mass represents an object’s • Weight depends on
location (& local gravity)
inertia (tendency to resist
• Weight is not a
change in motion)
fundamental property of
• Mass is the same everywhere matter
in the universe
• SI units of weight are
• SI Units of mass are kilograms newtons (N)
(kg)
• USCS units are pounds (lb)
Volume
• Volume is the 3-dimensional space that an object occupies
• Volume Units:
– The SI unit for volume is the cubic meter, or m3 (meters x meters x meters)
– The more common metric unit of volume is the Liter (L)
1 m3 = 103 L
– In the laboratory, the milliliter (mL) is often more convenient
1 mL = 10-3 L
area
height
height
width
length
Note: mass and volume are not the same thing (try not to confuse them…).
Two objects with the same volume (e.g. a pillow & a sack of potatoes can
have different masses and vice versa)
Density
Density is a property of matter representing the mass per unit
volume
• For equal volumes, a denser object has greater mass
• For equal masses, a denser object has smaller volume
Commonly used units:
1. Solids = g/cm3 (Note: 1 cm3 = 1 mL)
Mass
2. Liquids = g/mL
Density 
Volume
3. Gases = g/L
Useful Notes on Density:
• Volume of a solid can be determined by water displacement
• Density of matter in various states: solids > liquids >>>
gases (exception: water)
– In a heterogeneous mixture, the denser matter will tend
to sink to the bottom
Manipulating the Density Equation
Mass
Density 
Volume
mass
density
volume
Mass
Volume 
Density
Mass  Density  Volume
Ch 100: Fundamentals for Chemistry
Chapter 3: Elements & Compounds
Lecture Notes
Chemical Symbols & Formulas
• Each element has a unique chemical symbol
• Examples of chemical symbols:
– Hydrogen: H
– Oxygen: O
– Aluminum: Al
• Each molecule has a unique chemical formula
• The chemical formula of a molecule indicates
1. the chemical symbol for each of the elements present
2. 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)
The Periodic Table
• All of the known elements are arranged in a chart called
the Periodic Table
• Each element in the Periodic Table is identified by both its
chemical symbol and its Atomic Number
– The elements are organized left-to-right and top-tobottom according to their Atomic Number
– The elements are arranged by similarity of chemical
properties
• The columns are called Groups
– Elements of each group typically have similar properties
• The rows are called Periods, and reflect the periodicity of
chemical properties as atomic number increases
The Periodic Table of Elements
Elements and the Periodic Table
The elements can be categorized as:
1. Metals
• The leftmost elements in the periodic table
• Roughly 70% of all of the elements are metals
2. Nonmetals
• The rightmost elements of the periodic table
3. Semimetals (metalloids)
• The elements that reside along the “stair step”
between the metals and nonmetals in the Periodic
Table
• The properties of semimetals are not quite metallic or
non-metallic, but rather somewhere in between
Dmitri Mendeleev (1834-1907)
• Russian born chemist
• Considered one of the greatest
science teachers of his era
• He organized the known
elements of his time into the first
“periodic table”
– Elements were organized by
chemical properties (& by
weight) -> called periodic
properties
– Surprisingly, his periodic table
predicted the existence of 3
new elements (which were
subsequently discovered)
Ch 100: Fundamentals for Chemistry
Chapter 4: Properties of Matter
Lecture Notes
Physical & Chemical Properties
•
Physical Properties are the characteristics of matter that
can be changed without changing its composition
– These characteristics are directly observable or
measurable
– Types of Physical Properties:
1. Extrinsic Physical Properties are unique to objects (i.e. size,
shape, mass, etc.)
2. Intrinsic Physical Properties are unique to substances (i.e.
density, conductivity, color, etc.)
•
Chemical Properties are the characteristics of a
substance that determine the tendency of the matter to
transform in composition as a result of the interaction with
other substances, the influence of energy or both
– These are characteristics that describe the behavior of
matter
Physical & Chemical Changes
•
Physical Changes are changes that do not result in a
change the fundamental composition of the substance
Typical Examples:
1. Physical State Changes: boiling, melting, condensing,
etc.
2. Shape, Size or Texture Changes
• Chemical Changes involve a change in the fundamental
composition of the matter
Notes on Chemical Change:
1. Production of a new substance(s)
2. Referred to as chemical reactions
3. The basic representation: Reactants  Products
Note: Both physical and chemical changes will likely produce
an alteration of appearance, the key is to discern the type of
change that has occurred
Energy
Energy is loosely described as the capacity of something to do
work (or alter the physical or chemical state of an object or
system)
• Common Forms of Energy
– mechanical, chemical, thermal, electrical, radiant,
nuclear
• The SI unit of energy is the Joule (J)
– Other commonly used units are Calories (cal) and
Kilowatt-hours (kW.hr)
• Types of energy:
1. Potential: stored energy
2. Kinetic: energy associated with motion and vibration
3. Heat: energy that flows from high to low temperature
Principle of Energy Conservation: energy is never created
nor destroyed (but it does change from one type to another!)
Distinguishing Heat Energy & Temperature
Temperature is _____.
1.
2.
How hot or cold something is (an extrinsic physical property), it represents a
particular thermal state
Related to the average (kinetic) energy of the substance (not the total energy
but the average energy)
3.
Measured in units of:
•
•
•
Degrees Fahrenheit (oF)
Degrees Celsius (oC)
Kelvin (K)
Heat is _____.
1.
2.
3.
Energy that flows from hot objects to cold objects. Heat is not a physical
property.
Energy absorbed or released by an object resulting in its temperature
change
Measured in units of:
•
•
•
Joules (J)
Calories (Cal)
Kilowatt Hours (kW.hr)
Bottom Line: Heat energy absorbed or released is measured by changes in
temperature but do not confuse heat energy for temperature
Temperature Scales
The 2 traditional temperature scales, Fahrenheit and Celsius,
were originally defined in terms of the physical states of
water at sea level:
1. Fahrenheit Scale, °F
– For water: freezing point = 32°F, boiling point = 212°F
2. Celsius Scale, °C
– For water: freezing point = 0°C, boiling point = 100°C
– 1 Celsius temperature unit is larger than 1 Fahrenheit unit
The SI unit for temperature is a variant of the Celsius scale
3. Kelvin Scale, K
– For water: freezing point = 273 K, boiling point = 373 K
– The Kelvin temperature unit is the same size as the Celsius unit
Temperature of ice water and boiling water.
Heat (Energy)
•
•
•
•
Heat is energy that flows due to a temperature difference
– Heat energy flows from higher temperature to lower
temperature
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: burning paper is an exothermic process because
energy is produced as heat (the temperature rises!).
2. Endothermic = A process that absorbs energy.
•
Example: melting ice to form liquid water is an endothermic
process because heat energy must be absorbed to change the
physical state (in this case the temperature does not change!).
•
•
Heat (cont.)
When something absorbs or loses heat energy, 1 of 2 things
can occur:
1. Its temperature will change (e.g. hot coffee will cool
down)
2. Its physical state will change (e.g. ice will melt)
For the former case above, the heat energy absorbed or lost
by an object is proportional to:
1. The mass of the object (m)
2. The change in temperature the object undergoes (DT)
3. The specific heat capacity (s) (a physical property unique to the
substance)
Q
To calculate heat gained (Q):
Q = s  m  DT
s
m
DT
Specific Heat Capacity (s)
•
Specific heat capacity reflects how absorbed heat energy
relates to the corresponding increase in temperature for a
given amount of mass, i.e. energy per unit mass per unit
temperature change or
Q
s=
m  DT
•
•
Specific Heat Capacity is commonly measured in units of:
1. J/goC (SI)
2. cal/goC (metric & more useful in the lab)
Specific Heat Capacity is a unique intrinsic physical
property of matter. Typically, __________.
1. Metals have low specific heat capacity
2. Non-metals have higher specific heat capacity than
metals
3. Water has an unusually large specific heat capacity
0.900
0.473
Table of Specific Heat for Various Substances
Substance
J/g.K
cal/g.K
J/mol.K
Aluminum
0.900
0.215
24.3
Iron
0.473
0.113
26.4
Copper
0.385
0.0921
24.5
Brass
0.380
0.092
...
Gold
0.131
0.0312
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
Alcohol (ethyl)
2.138
0.511
111
Water
4.184
1.000
75.2
Ice (-10 C)
2.059
0.492
36.9
Granite
.790
0.19
...
Glass
.84
0.20
...
Ch 100: Fundamentals for Chemistry
Chapter 5: Early Atomic Theory & Structure
Lecture Notes
Early Model of Matter: Aristotle (384-322 BC)
• Introduced observation as an important step in
understanding the natural world
• According to his model of nature, all forms 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 any other substance by altering
the relative proportion of these elements and
qualities (i.e. lead to gold)
Dalton’s Atomic Theory
1. Each element consists of individual
particles called atoms
2. Atoms can neither be created nor
destroyed
3. All atoms of a given element are
identical
4. Atoms combined chemically in definite
whole-number ratios to form
compounds
5. Atoms of different elements have
different masses
The Modern Atomic Model
According to our modern model of the matter, 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
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
Electric Charge
•
•
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 (qproton) or electron (qelectron) is the
smallest amount that occurs in nature, it is called the
quantum of charge:
1. qproton = +1.602 x 10-19 Coulombs (or 1+)
2. qelectron = -1.602 x 10-19 Coulombs (or 1-)
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 are electrically charged atoms and thus carry electric charge:
• The electric charge of an ion is due to the imbalance of electrons
and protons
• When an atom has lost one or more of its electrons it carries a
positive charge
“1+” for each electron that is lost
• When an atom has gained one or more of its electrons it carries a
positive charge
“1-” for each excess electron that is gained
• When an atom/molecule is an ion, its charge must be specified:
– Sodium ion:
Na+
– Chloride ion:
Cl– Hydroxide ion:
OH• Notes on Electric Charge:
– Opposite charges attract
-
+
– Like charges repel
+
+
-
-
Atomic Bookkeeping
• Atomic number (Z)
– The number of protons in an atom or ion
– The number that defines the identity of the atom
• Mass number (A)
– The number of protons & neutrons in a specific atom or
isotope
– The number that represents the mass of an 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
1. Unique atoms of the same element with different mass numbers
(i.e. they have different numbers of neutrons)
2. Unique isotopes are identified by their mass number
Isotope notation:
Mass #
(Atomic Symbol)
Atomic #
Example:
carbon-12 (12C ) & carbon-14 (14C )
6
6
Atomic mass
1. The average total mass of an element’s various naturally occurring
isotopes
2. The unit of Atomic Mass is the Dalton (or 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 slightly 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
Ch 100: Fundamentals for Chemistry
Chapter 6: Nomenclature of Inorganic
Compounds
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 focus on only “simple compounds”
Types of Compounds (cont.)
For “practical” purposes will separate all simple compounds into
2 general categories:
1. Ionic Compounds
a.
b.
c.
d.
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
2. Non-Ionic (aka: Molecular or Covalent) Compounds
a.
b.
c.
Made up of atoms that share their outer electrons
Electric charge plays no direct role in their formation
There are usually no metals are present in these compounds
Naming Compounds
The easiest way (usually) to identify an ionic compound is to
ask whether or not there is a metal present in the chemical
formula (or the name):
Is a metal present?
– Yes -> it is an Ionic Compound (e.g. CaCl2)
– No -> it is a Non-Ionic Compound (e.g. CCl4) or an Acid
Notes:
1. Ionic compounds do not use the Greek prefixes and are
named according to the identity of the ions present
2. Non-Ionic compounds require the use of Greek prefixes to
indicate the number of each element present in one
molecule
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
Determining Chemical Formula of an Ionic Compound
To determine the chemical formula of an ionic compound from its chemical
name:
1.
Identify the ions present, both cation(s) and anion(s), from the name.
Example: potassium sulfide
Cation: potassium
Anion: sulfide
2.
Determine the ionic charge of the ions
Example: {from above}
potassium ion, K+
sulfide ion, S23.
Determine the number of each ion needed to obtain a neutral
compound
Example: {from above}  2 K+ ions are needed for every S23.
Combine the chemical sysmbols of the ions to get the final chemical
formula
Example: {from above}  K2S is the formula for potassium sulfide
Ionic Charges & the Periodic Table
The position of an element in the Periodic Table is a useful indicator of the
type of ion an element is capable of forming:
1. Group 1 metals form 1+ cations (Na+ sodium ion)
2. Group 2 metals form 2+ cations (Ca2+ calcium ion)
3. Group 13 metals form 3+ cations (Al3+ aluminum ion)
4. Group 3-12 Metals (plus Sn, Pb, & Bi) can form more than one type
of cation
Roman numerals are used to indicate the charge of the cation
Example:
Fe3+ is called iron(III)
FeCl3 is called iron(III) chloride
Notable Exceptions:
Ag+, Cd2+ & Zn2+
5.
6.
7.
8.
Group 15 nonmetals form 3- anions (e.g. N3- nitride ion)
Group 16 nonmetals form 2- anions (e.g. O2- oxide ion)
Group 17 nonmetals form 1- anions (e.g. Cl- chloride ion)
Group 18 elements do not form ions
Greek Prefixes for Compound Names
1)
2)
3)
4)
5)
MonoDiTriTetraPenta-
CCl4 is carbon tetrachloride
6) Hexa7) Hepta8) Octa9) Nona10) DecaC3H8 is tricarbon octahydride
Notes:
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
• Polyatomic ions are usually ions formed from non-ionic
molecules
e.g. The sulfate ion, SO42-, is essentially a molecular
compound containing S and O with 2 additional electrons
• When you encounter polyatomic ions in compounds, do not
freak out!!
• Become familiar with the common polyatomic ions on the
handout
Example: The nitrate ion (NO3-)
• Fortunately, the naming of ionic compounds containing
polyatomic ions 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 or release hydrogen
cations, H+, (usually when dissolved in water)
– The chemical formula for acids usually begins with H
Example: hydrochloric acid (HCl)
HCl(aq)  H+ + Cl- (aq)
Bases
– Taste bitter (Note: it is not advised to taste strong
bases…)
– Usually metal containing hydroxides
– Substances that accept hydrogen cations (H+) when
dissolved in water
Example: potassium hydroxide (KOH)
KOH(aq) + H+  K+(aq) + H2O (l)
Naming Acids
Lets separate acids into 2 types:
1. Acids that contain oxygen
2. Acids that do not contain oxygen
Naming acids containing oxygen:
1. For acids containing “-ate” anions:
a. Use root of the anion (for sulfate, SO42-, use sulfur)
b. Add “-ic” suffix then end with “acid”
Example:
H2SO4 is sulfuric acid
2. For acids with “-ite” anions:
a. Use root of the anion (for sulfite, SO32-, use sulfur)
b. Add “-ous” suffix then end with “acid”
Example:
H2SO3 is sulfurous acid
Naming Acids (cont.)
Naming acids not containing oxygen:
1. Add “hydro-” prefix to beginning
2. Use root of the anion (i.e. Cl- use chlor)
3. Add “-ic” suffix then end with “acid”
Example:
HCl is hydrochloric acid
Name the following acids:
HF
HNO2
HCN
H3PO4
Antoine Lavoisier (1743-1794)
•
•
Considered by many to be the “Father
of Modern Chemistry”
Major contributions included
1. Demonstrated that water cannot be
transmuted to earth
2. Established the Law of Conservation
of Mass
3. Developed a method of producing
better gunpowder
4. Observed that oxygen and hydrogen
combined to produce water (dew)
5. Invented a system of chemical
nomenclature (still used in part today!)
6. Wrote the 1st modern chemical
textbook
Ch 100: Fundamentals for Chemistry
Ch 7: Quantitative Composition of Compounds
Lecture Notes
(Sections 7.1 to 7.3)
The Mole
• The mole is a counting unit (analogous to the dozen unit)
– 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
• Molar mass is the mass in grams of 1 mole of a substance
• Molar mass refers to both atoms & molecules
1. 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
2. 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. To convert from atoms (or molecules) to moles, divide the
# of atoms (or molecules) by Avogadro’s #
Example: How many moles are 1.0x1024 atoms?
1 mol


(1.0×10 atoms) 
= 1.7 mol
23 
 6.022×10 
24
2. To convert from moles to atoms (or molecules), multiply
the # of atoms (or molecules) by Avogadro’s #
Example: How many molecules are in 2.5 moles?
 6.022×1023 
24
(2.5 mol) 
=1.5×10
molecules

1 mol


Mole-Mass Calculations
1. To convert from moles to grams, multiply the # of moles by
atomic mass
Example: How many grams in 2.5 moles of carbon?
 12.01 g 
1
(2.5 mol) 
=
30.
g
(or
3×10
)

 1 mol 
2. To convert from grams to moles, divide the mass in grams
by atomic mass
Example: How many moles are in 2.5 g of lithium?
 1 mol 
1
(2.5 g) 
=
0.36
mol
(or
3.6×10
)

 6.941 g 
Percent Composition
•
•
Percent composition is the percentage of each element in
a compound (by mass)
Percent composition can be determined from either:
1. the formula of the compound
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:
1. Determine the molar mass of compound
2. Multiply the molar mass of the element of interest by the number of
atoms per molecule then
3. 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.
2.
3.
The molar mass of NaCl is 58.44 g/mol
There is 1 atom of Na in each NaCl molecule
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
Amadeo Avogadro
(1743-1794)
•
•
Italian lawyer turned chemist
Major contributions included:
1. Established difference between atoms & molecules:
•
Oxygen & nitrogen exist as molecules O2 & N2
2. Reconciled the work of Dalton & Guy-Lussac
3. Establishing Avogadro’s Principle: equal volumes of all gases at the
same temperature and pressure contain the same number of molecules.
•
Note: Avogadro did not determine Avogadro’s number nor the mole
(these concepts came later)
1. Avogadro is honored because the molar volume of all gases should
be the same
2. Much of Avogadro’s work was acknowledged after he died, by
Stanislao Cannizarro
Ch 100: Fundamentals for Chemistry
Chapter 8: Chemical Equations
Lecture Notes
Chemical Equations (Intro)
• Chemical equations are used to symbolically describe
chemical reactions
• In a chemical equation (or reaction for that matter) the
substances that undergo chemical change(s) are called the
reactants
• The resulting substances formed are called the products
• The standard representation of a chemical equation:
Reactant(s)  Product(s)
Example: The production of water
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 Equations (cont.)
Balanced chemical equations indicate the ____
1. identity of each reactant & product involved in the
reaction
2. phase of each reactant and product involved in the
reaction (i.e. solid (s), liquid (l) or gas (g))
3. relative quantity of each reactant and product involved in
the reaction (the coefficients!)
4. relative molar quantity of each reactant and product
involved in the reaction (the coefficients!)
Balancing Chemical Equations
•
According to the Law of Mass Conservation (& John Dalton!)
matter is never created nor destroyed during chemical
reactions
–
All of the atoms in the reactants of a chemical reaction must be
accounted for in the products
The Basic Process of Balancing Chemical Equations:
1. Identify all reactants & products in the reaction & write out their
formulas (this is the unbalanced chemical equation)
2. 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!)
4. 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)
5. Repeat until all atoms are balanced for the reaction
6. Now you should have a balanced chemical equation!
Balancing Chemical Equations (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)
1. Note the nitrate ion (NO3-) gets swapped between the Ag+
and the Ca2+ ions in this reaction
2. So NO3- can be treated as a whole unit when balancing this
reaction
3. Balance it!
Common Classifications for Chemical Reactions
1.
4.
Combination (or Synthesis): reactions in which reactants combine to make one
product
Decomposition: reactions in which one reactant breaks down into smaller
products
Single Displacement: reactions where a part of one reactant is displaced and
combined with another reactant
2Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g)
Double Displacement: reactions where a part of two reactants is displaced and
exchanged
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
Examples:
a. Acid-base neutralization
b. Formation of insoluble products (Precipitation reactions)
c. Metal oxide + acid
d. Gas formation
5.
Oxidation-Reduction Reactions: reactions involving the transfer or
rearrangement of electrons
2.
3.
Combination & Decomposition Reactions
1. Reactions in which chemicals combine to make one
product are called Combination or Synthesis Reactions
a. Metal + Nonmetal reactions can be classified as Combination
Reactions
2 Na(s) + Cl2(g)  2 NaCl(s)
b. Reactions between Metals or Nonmetals with O2 can be classified
as Combination Reactions
N2(g) + O2(g)  2 NO(g)
Note: these two types of Combination Reactions are also subclasses
of Oxidation-Reduction Reactions
2. Reactions in which one reactant breaks down into smaller
molecules are called Decomposition Reactions
a. Decomposition reactions are generally initiated by the addition of
energy (via electric current or heat)
b. Decomposition reactions are the opposite of Combination
Reactions:
2 NaCl(l)  2 Na(l) + Cl2(g)
Single Displacement Reactions
Single displacement reactions involve one part of a reactant
being transferred to another
The basic pattern of the single displacement reaction:
XY + A  X + AY
Example 1: Metal + Acid  Salt + Hydrogen
Zn(s) + 2 HCl(aq)  ZnCl2(aq) + H2(g)
Example 2: Metal + Water  Hydrogen + Metal Oxide (or metal hydroxide)
3 Fe(s) + 4 H2O(l)  4 H2(g) + Fe2O3(s)
Example 3: Metal + Salt  Metal + Salt
2 Al(s) + Fe2O3(s)  2 Fe(s) + Al2O3(s)
Example 4: Halogen + Halide Salt  Halogen + Halide Salt
Cl2(g) + 2 NaBr(s)  Br2(g) + 2 NaCl(s)
Double Displacement Reactions
Double Displacement Reactions involve the double exchange of a
component (such as ions) between two reactants
The basic form of double displacement reactions is:
XY + AB  XB + AY
where X, Y, A, and B are the components of the reactants
Example 1: Acid Base Neutralization
H2SO4(aq) + Ca(OH)2(aq)  CaSO4(aq) + 2 H2O(l)
or 2HOH
Example 2: Metal Oxide + Acid
CaO(s) + 2HCl(aq)  CaCl2(aq) + H2O(l)
or HOH
Example 3: Formation of an Insoluble Precipitate (Precipitation)
KCl(aq) + AgNO3(aq)  KNO3(aq) + AgCl(s)
Example 4: Formation of a Gas
HCl (aq) + ZnS (s)  ZnCl2(aq) + H2S (g)
Solubility & Precipitation Reactions
•
When 2 solutions are combined and result in the formation of an
insoluble product:
1.
2.
•
Solubility is an intrinsic physical property and a measure of how well a
substance (solute) will dissolve in another substance (solvent)
1.
2.
3.
•
•
•
The product will not dissolve in the solvent
The product will form a precipitate
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
Eg. NaCl is soluble in water
A solute is insoluble if no appreciable amount of it will dissolve in solvent
Eg. AgCl is insoluble in water
Precipitation (formation of an insoluble solid) is one indication that a
chemical change has occurred!
General Rules for Solubility
1. Most compounds that contain NO3- ions are soluble
2. Most compounds that contain Na+, K+, or NH4+
ions are soluble
3. Most compounds that contain Cl- ions are soluble,
except AgCl, PbCl2, and Hg2Cl2
4. Most compounds that contain SO42- ions are
soluble, except BaSO4, PbSO4, CaSO4
5. 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
6. Most compounds that contain S2-, CO32-, or PO43ions are slightly soluble (will precipitate)
Oxidation-Reduction Reactions
•
Reactions that involve transfer or rearrangement of electrons
are called oxidation-reduction reactions.
Examples of oxidation-reduction reactions:
1. Metal + Nonmetal: 2Na(s) + Cl2(g)  2NaCl(s)
a. The metal loses an electron(s) and becomes a cation (oxidation 
metal gets oxidized: Na  Na+ + e-)
b. The nonmetal gains an electron(s) and becomes an anion (reduction
 nonmetal gets reduced: Cl + e-  Cl-)
c. In this reaction, electrons are transferred from the metal to the
nonmetal
2. O2 as a reactant or product: CH4(s) + O2(g)CO2(g) + H2O(g)
a. In this reaction, it is not obvious that electron transfer has taken
place. In this case, oxidation states are altered.
b. Often this type of reaction involves the release of large amounts of
energy, even combustion
Rates of Chemical Reactions
How quickly a chemical reaction occurs is indicated by its
reaction rate
1. How quickly the concentration of products increases
2. How quickly the concentration of reactants decreases
The Factors that influence reaction rates:
1. Reactants must be in contact
•
•
Reactions occur due to collisions
Without contact between reactants there can be no reaction
2. Concentration of reactants
•
The more reactant molecules packed into a given space the
more likely a collision (& reaction) will occur
3. 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
The Role of Energy in Chemical Reactions
Energy transformations always accompany chemical reactions:
1. Energy is required to break bonds (energy absorbed or activation
energy)
2. Energy is released when bonds are formed
Note: The amount of energy required to break a chemical bond
is the same as the energy released when that type of bond
is formed, this is called the Bond Energy
For a chemical reaction to occur:
1. Energy must be absorbed in order to break chemical bonds in the
reactants
2. Energy is released as new bonds are formed in the products
•
•
Endothermic reactions absorb more energy than they
release
N2(g) + O2(g) + 393 kJ  2NO(g)
Exothermic reactions release more energy than they absorb
H2(g) + Cl2(g)  2NO(g) + 185 kJ
Energy in Chemical Reactions
Exothermic Reactions
Potential
Energy
Activation
Energy (EA)
Reactants
Energy
Released (Q)
Products
Endothermic Reactions
Potential
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
• Participation of catalysts in a chemical 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.
catalyst
Reactants  Products
Example: The breakdown of hydrogen peroxide
2H2O2 (aq)
catalase
 H2O (l) + O2 (g)
Catalysts & Energy in Reactions
Catalysts lower Activation Energy
Activation Energy
without catalyst
Potential
Energy
Reactants
Activation Energy
with catalyst
Products
Ch 100: Fundamentals for Chemistry
Ch 9: Calculations from Chemical Reactions
Lecture Notes (Sections 9.1 to 9.5)
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
a. Molecule to molecule ratios
b. Mole to mole ratios
c. Even mass to mass ratios can be determined (with use of molar
mass values)
Information Given by the Chemical Equation
A balanced chemical equation provides the relationship
between the relative numbers of reacting molecules and
product molecules
Example: The formation of carbon dioxide from carbon
monoxide and oxygen gas
2 CO + O2  2 CO2
In this chemical equation, it is indicated that 2 CO molecules
react with every 1 O2 molecule to produce 2 CO2 molecules
Alternative interpretation: there is a 2:1 (numerical) ratio of
CO to O2 for this completed reaction, 2 CO:1 O2 :2 CO2
Interpretation of the Chemical Equation
•
Since the information given in a balanced chemical reaction is relative:
2 CO + O2  2 CO2
the following are alternative interpretations of the chemical equation:
a.
b.
c.
d.
•
•
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
Note: The coefficients in the balanced chemical equation also 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
Mole and Mass Ratios in Chemical Equations
For the following chemical equation:
2 CO + O2  2 CO2
The following mole relations are implied:
2 moles CO : 1 mole O2 : 2 moles CO2
Note the molar masses of the compounds in this reaction:
a. 1 mole of CO = 28.01 g
b. 1 mole O2 = 32.00 g
c. 1 mole CO2 = 44.01 g
The mass ratio of the compounds in this reaction can be determined
using the molar mass values:
2(28.01) g CO : 1(32.00) g O2 : 2(44.01) g CO2
The mass ratio of the compounds in this reaction are:
56.02 g CO : 32.00 g O2 : 88.02 g CO2
Example
Determine the Number of Moles of Carbon Monoxide required to
react with 3.2 moles Oxygen, and the moles of Carbon Dioxide
produced
1.
2.
3.
Write the balanced equation
2 CO + O2  2 CO2
Use the coefficients to find the mole relationship
2 moles CO : 1 mol O2 : 2 moles CO2
Use dimensional analysis to obtain the # of moles
a. The # mol of CO:
 2 moles CO 
3.2 moles O2 x 
  6.4 moles CO
 1 mole O2 
b. The # mol of CO2:
 2 moles CO2 
3.2 moles O2 x 
  6.4 moles CO2
 1 mole O2 