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Chapter 1
Introduction to Chemistry
What is Chemistry?
Matter – anything that has mass
and occupies space.
Chemistry – study of the composition of matter
and the changes that matter
undergoes.
Because living and nonliving things are made of matter,
chemistry affects all aspects of life
Areas of Chemistry
Organic – study of all chemicals containing carbon
Inorganic – study of chemicals that, in general, do
not contain carbon. (found mainly in nonliving things)
Biochemistry – study of processes that take place
in organisms. (digestion, muscle contraction)
Analytical – focuses on the composition of matter.
(measuring lead in drinking water)
Physical – area of study that deals with the
mechanism, the rate, and the energy transfer that
occurs when matter undergoes a change.
See page 8 for examples
Pure & Applied Chemistry
Pure Chemistry – pursuit of chemical knowledge for
its own sake
•
Chemists does not expect there to be any immediate
practical use for the knowledge
Applied – research that is directed toward a
practical goal or application
In practice, pure chemistry & applied chemistry are linked.
The Scientific Method
Logical, systematic approach to the solution of a
scientific problem.
1. Making Observations – using your senses to obtain
information. An observation can lead to a question.
2. Making a Hypothesis – a proposed explanation for an
observation. A hypothesis is only useful if it accounts for
what is actually observed.
3. Experiment – a procedure that is used to test a
hypothesis.
a) Independent variable – a variable that you change
during an experiment
The Scientific Method
b) Dependent variable – a variable that is observed
during the experiment.
c) For the results of an experiment to be accepted,
the experiment must produce the same result no
matter how many times it is repeated or by whom.
4. Developing a Theory – a well-tested explanation for a
broad set of observations.
5. Scientific Law – concise statement that summarized
the results of many observations and experiments.
Ex. Gas Laws
The Scientific Method
Observations
Hypothesis
Experiments
A hypothesis may be
revised based on
experimental data
A theory is tested by
more experiments &
modified if necessary
Scientific
Law
Steps do not have to occur
in the order shown
Theory
Summarizes the results of
many observations and
experiments
Solving Numeric Problems
1. Analyze – identify what is known
and what is unknown.
2. Calculate – make the calculations. You may need to
convert a measurement or rearrange an equation
before you can solve.
3. Evaluate – after you calculate, evaluate your
answer. Is the answer reasonable? Does it make
sense?
Chapter 2
Matter
Matter – anything that has mass and takes up space
Mass – measure of the amount of matter that an
object contains
Volume – measure of the space occupied by the
object
Extensive & Intensive Properties
What you observe when you examine a sample of
matter is its properties.
1. Extensive Property – a property that depends
on the amount of matter in a sample
Ex. Mass, volume, weight, length
2. Intensive Property – a property that depends
on the type of matter in a sample (prefix–in means
within)
Ex. Hardness, color, odor, luster, conductivity,
malleability, ductility, freezing point, boiling point,
melting point, density
Substances
Substance – Matter that has a uniform and definite
composition
•
Either an element or a compound
•
Also called pure substance
•
Rarely found in nature
•
Fixed proportions to each other
Examples
 Diamond
 Water
 Gold
 Copper
 Sugar
 Nitrogen
Mixtures
Mixture – a physical blend of two or more
substances that are not chemically combined
•
Do not exist in fixed proportions to each other
•
Most natural substances are mixtures
•
Can usually be separated back into its original
components
Examples
 Concrete
 Soil
 Salt water
 Milk
 Coke
 Gasoline
 Fruit salad
 Atmosphere
Two Types of Mixtures
Homogeneous Mixture (solution) – a mixture in
which the composition is uniform throughout.
•
Consists of a single phase
•
Can’t see them separately or separate them
physically
Examples
 stainless steel
 air
 olive oil
 vinegar
Two Types of Mixtures
Heterogeneous Mixture – a mixture in which the
composition is not uniform throughout.
•
Consists of a two or more phases
Examples
 chicken soup
 oil & vinegar mixed
 milk
 rice crispy treats
Separating Mixtures
Differences in physical properties can be used to separate
mixtures
Filtration – process that separates a solid from a
liquid
Examples
 coffee filters
 draining pasta
Separating Mixtures
Distillation – process of boiling a liquid to produce a
vapor and then condensing the vapor into a liquid
Example
 separating water from other
substances in the water
States of Matter
1. Solid
3. Gas
2. Liquid
States of Matter
Solid
Definite shape
Definite volume
Not easily compressed
•
•
•
•
Characteristics
Does not take the shape of the container
Particles packed tightly together, and often in orderly
arrangement
Almost incompressible
Expands only slightly when heated
States of Matter
Liquid
Indefinite shape
Definite volume
Not easily compressed
•
•
•
•
Characteristics
Take the shape of the container in which it is placed
Particles in close contact, but arrangement of particles
is not orderly (can flow past each other)
Almost incompressible
Expands slightly when heated
States of Matter
Gas
Indefinite shape
indefinite volume
Easily compressed
•
•
•
•
Characteristics
Take the shape of the container in which it is placed
Can expand to fill any volume
Particles are much farther apart
Easily compressed into a smaller volume
Physical Change
Physical Change
Some properties of a material
change, but the composition of
the material does not change
Examples

Changes of state such as boiling water,
condensation (boil, freeze, melt, condense)

Physical deformation such as cutting, denting,
stretching, breaking, crushing
Chemical Change
Chemical Change
A change that produces matter
with a different composition than
the original matter
Examples

Silver spoon tarnishes

Metal rusts
 Methane burns

Methane burns
 Sugar ferments

Burn, rot, rust, decompose, ferment, explode,
corrode usually mean a chemical change
Elements
Element – simplest form of matter that has a
unique set of properties.
• cannot be broken down into simpler substances
Examples
 Hydrogen
 Nitrogen
 Oxygen
Compounds
Compound – substance that contains two or more
elements chemically combined in a fixed
proportion.
• Compounds can be broken down into simpler
substances by chemical means
Examples
 Sugar (C12H22O11)
 Salt (NaCl)
 Water (H2O)
Classifying Matter
Any sample of matter is
either an element, a
compound, or a mixture
Matter
Can be separated
physically
Substance
Definite composition
Mixture
Variable composition
Can be separated
chemically
Element
Simplest form
Silver
Compound
Salt
Homogeneous
Mixture
Heterogeneous
Mixture
Uniform; also called
a solution
Nonuniform;
Distinct phases
Stainless
Steel
Cement
Symbols Derived From Latin
Sodium
Na
Potassium
Antimony
Copper
Gold
Silver
Iron
Lead
Tin
K
Sb
Cu
Au
Ag
Fe
Pb
Sn
Physical Properties
Physical Property – a quality or
condition of a substance that can be
observed or measured without
changing the substance’s
composition
Examples
 Appearance
 Density
 Texture
 Malleability
 Color
 Boiling Point
 Odor
 Melting Point
 Conductivity
 Hardness
Chemical Property
Chemical Property
Ability of a substance to undergo a
specific chemical change
• Chemical properties can be observed
only when a substance undergoes a
chemical change.
Examples
Gasoline -- burns in air
Iron -- rusts
Baking Soda -- reacts with vinegar
Copper -- rusts in water
Table salt -- does not react with vinegar
Recognizing Chemical Changes
Words such as burn, rot, rust,
decompose, ferment, explode, and
corrode usually signify a chemical
change.
During a chemical change, the composition
of matter always changes.
Examples
Gasoline -- burns in air
Iron -- rusts
Baking Soda -- reacts with vinegar
Copper -- rusts in water
Table salt -- does not react with vinegar
Recognizing Chemical Changes
Possible Clues
• Transfer of energy
• A change in color
• The production of gas
• The formation of a precipitate
Precipitate – solid that forms and settles out of a
liquid mixture
Ex. – ring of soap scum in your bathtub
The only way to be sure a chemical change has
occurred is to test the composition of a sample
before and after the change
Law of Conservation of Mass
During any chemical reaction, the mass of the products is
always equal to the mass of the reactants.
Example
2H2 + O2
2g
2g
reactants
2H20
=
=
4g
product
Chapter 3
Observation,
Measurement
and Calculations
Precision and Accuracy
Accuracy – measure of how close a measurement comes to the
actual or true value of whatever is being measured.
Precision – measure of how close a series of measurements are to
one another.
Neither accurate
nor precise
Precise but not
accurate
Precise AND
accurate
Determining Error
Accepted Value – the correct
value based on reliable references
Experimental Value – the value
measured in the lab
Error(can be +or-)=experimental value – accepted value
Percent error = absolute value of error x 100%
accepted value
Rules for Counting Significant
Figures
Nonzero integers always count as
significant figures.
3456 has
4 sig figs.
Rules for Counting Significant
Figures
Leading zeros do not count as
significant figures.
0.0486 has
3 sig figs.
Rules for Counting Significant
Figures
Zeros at the end of a number and
to the right of a decimal point are
always significant.
9.000 has
4 sig figs
1.010 has
4 sig figs
Rules for Counting Significant
Figures
Captive zeros always count as
significant figures.
16.07 has
4 sig figs.
Rules for Counting Significant
Figures
Zeros at the rightmost end that
lie at the left of an understood
decimal point are not significant.
7000 has
1 sig fig
27210 has
4 sig figs
Rules for Counting Significant
Figures
Exact numbers have an infinite
number of significant figures.
1 inch = 2.54 cm, exactly
Rules for Significant Figures in
Mathematical Operations
Multiplication and Division: # sig figs in
the result equals the number in the least
precise measurement used in the
calculation.
6.38 x 2.0 =
12.76  13 (2 sig figs)
Rules for Significant Figures in
Mathematical Operations
Addition and Subtraction: The number
of decimal places in the result equals the
number of decimal places in the least
precise measurement.
6.8 + 11.934 =
18.734  18.7 (3 sig figs)
Sig Fig Practice #1
How many significant figures in each of the following?
1.0070 m 
5 sig figs
17.10 kg 
4 sig figs
100,890 L 
5 sig figs
3.29 x 103 s 
3 sig figs
0.0054 cm 
2 sig figs
3,200,000 
2 sig figs
International Systems of Units
There are seven SI base units
Quantity
SI Base Units
SI base unit
Symbol
Length
Mass
Temperature
Time
Amount
Luminous intensity
Electric current
Meter
kilogram
kelvin
second
mole
candela
ampere
m
kg
K
s
mol
cd
A
Metric Prefixes
Mega (M)
Kilo (k) 103
Hecto (hm) 102
Deka(da) 101
Meter (m)
left
Deci (d) 10-1
Centi (c) 10-2
Milli (m) 10-3
right
Micro (µ) 10-6
Nano (nm) 10-9
Pico (pm) 10-12
Other Common Conversions
1 cm3 = 1ml
1dm3 = 1L
1 inch = 2.54 cm
1kg = 2.21 lb
454 g = 1 lb
4.18 J = 1 cal
1 mol = 6.02 x 1023 pieces
1 GA = 3.79 L
Units of Length
meter – the basic SI unit of length or linear measure
Common metric units of length include the centimeter
(cm), meter (m), and kilometer (km)
Units of Volume
Volume -the space occupied by any sample of matter
Volume (cube or rectangle) = length x width x height
The SI unit of volume is the amount of space occupied by a
cube that is 1m along each edge. (m3)
Liter (L) – non SI unit – the volume of a cube that is 10cm
along each edge (1000cm3)
The units milliliter and cubic centimeter are used
interchangeably.
1 cm3 = 1ml
1dm3 = 1L
Units of Mass
Common metric units of mass include the kilogram,
gram, milligram and microgram.
Weight – is a force that measures the pull on a given
mass by gravity.
Weight is a measure of force and is different than mass.
Mass – measure of the quantity of matter.
Although, the weight of an object can change with its
location, its mass remains constant regardless of its
location.
Objects can become weightless, but not massless
Units of Temperature
Temperature – measure of how hot or cold an object is.
The objects temperature determines the direction of
heat transfer.
When two objects at different temperatures are in
contact, heat moves from the object at the higher
temperature to the object at the lower temperature.
Scientist use two equivalent units of temperature, the
degree Celsius and the Kelvin.
Units of Temperature
A change of 1 º on the Celsius scale is equivalent to one
kelvin on the Kelvin scale.
The zero point on the Kelvin scale, 0K, or absolute zero, is
equal to -273.15º C.
K = ºC + 273
ºC = K - 273
.
Units of Energy
Energy – the capacity to do work or to produce heat.
The joule and the calorie are common units of energy.
The joule (J) is the SI unit of energy named after the
English physicist James Prescott Joule.
1 calorie (cal) - is the quanity of heat that raises the
temperature of 1 g of pure water by 1ºC.
1 J = 0.2390 cal
1 cal = 4.184 J
Dimensional Analysis
Dimensional analysis – a way to analyze and solve
problems using the units, or dimensions, of the
measurements.
How many minutes are there in exactly one week?
60 minutes = 1 hour
24 hours = 1 day
7 days = 1 week
1 week 7 days 24 hours 60 minutes = 10,080 min
1 week 1 day
1 hour
1.0080 x 104 min
Dimensional Analysis
How many seconds are in exactly a 40-hr work week?
60 minutes = 1 hour
7 days = 1 week
24 hours = 1 day
60 seconds = 1 minute
40 hr 60 min 60 sec = 144,000 s
1 hr 1 min
1.44000 x 105 s
Dimensional Analysis
Gold has a density of 19.3 g/cm3. What is the density in kg/m3
19.3 g 1 kg
1 x 106 cm3 = 1.93 x 104 kg / m3
1000 g m3
cm3
There are 7.0 x 106 red blood cell (RBC) in 1.0 mm3 of
blood. How many red blood cells are in 1.0 L of blood?
7.0 x 106 RBC 1 x 106 mm3 1 dm3 = 7.0 x 1012
1.0 mm3
dm3
1L
Density
If a piece of led and a feather of the same volume are
weighted, the lead would have a greater mass than the
feather.
It would take a much larger volume of feather to equal the
mass of a given volume of lead.
Density = mass / volume
D=m/v
Mass is a extensive property (a property that depends on the
size of the sample)
Density is an intensive property (depends on the composition
of a substance, not on the size of the sample)
Questions
A student finds a shiny piece of metal that she thinks is
aluminum. In the lab, she determines that the metal has
a volume of 245cm3 and a mass of 612g. Was is the
density? Is it aluminum?
D = 612g / 245cm3 = 2.50g/cm3
D of aluminum is 2.70 g/cm3; no it is not aluminum
A bar of silver has a mass of 68.0 g and a volume of 6.48
cm3. What is the density?
D = 68.0g / 6.48 cm3 = 10.5 g/cm3
Chapter 4
Atomic Structure
The Atom
You cannot see the tiny fundamental particles that make
up matter.
Yet, all matter is composed of such particles, called atoms
Atom – the smallest particles of an element that retains
its identity in a chemical reaction
Several early philosophers and scientists could not
observe individual atoms, but still were able to propose
ideas on the structure of atoms.
Democritus’s Atomic Philosophy
Greek philospher Democritus (460B.C – 370 B.C.) was
among the first to suggest the existence of atoms.
Democritus believed that matter consisted of tiny,
indivisible and indestructible.
• Democritus’s ideas did not explain chemical behavior.
• Lacked experimental support, because his approach
was not based on scientific method.
Dalton’s Atomic Theory
According to Dalton’s atomic theory, and element is composed of
only one kind of atom, and a compound is composed of particles
that are chemical combinations of different kinds of atoms.
1. All elements are composed of tiny indivisible particles called
atoms
2. Atoms of the same element are identical. The atoms of any one
element are different from those of any other element.
Dalton’s Atomic Theory
3. Atoms of different elements can physically mix together or can
chemically combine in simple whole-number ratios to form
compounds.
4. Chemical reactions occur when atoms are separated, joined, or
rearranged. Atoms of one element, however, are never changed
into atoms of another element as a result of a chemical reaction.
Subatomic Particles
Most of Dalton’s atomic theory is accepted today. Except,
we now know atoms to be divisible.
Atoms can be broken down into smaller particles, called
subatomic particles.
There are 3 kinds of subatomic particles.
1. electrons
2. Protons
3. neutrons
Electrons
In 1897, English physicist J.J. Thomson discovered the
electron.
Electrons – negatively charged subatomic particles.
Dalton performed experiments that involved passing
electric current through gases at low pressure.
Protons and Neutrons
After a hydrogen atom loses an electron, what is left?
A particle with one unit of positive charge should remain
when a typical hydrogen atom loses an electron.
In 1886, Eugene Goldstein observed a cathode-ray tube
and found rays traveling in the direction opposite to
that of the cathode rays.
He concluded they were positive particles.
Protons – positively charged subatomic particles.
Protons and Neutrons
English physicist James Chadwick confirmed the existence
of another subatomic particle.
Neutron – subatomic particles with no charge but with a
mass nearly equal to that of a proton.
Particle Symbol Relative Relative
Charge
Mass
electron
e11/1840
Actual mass
(g)
9.11 x 10-28
proton
p+
1+
1
1.67 x 10-24
neutron
n0
0
1
1.67 x 10-24
Rutherford’s Gold-foil Experiment
However, the great majority of alpha particles passed
straight through the gold atoms, without deflection.
Also, a small fraction of the
alpha particles bounced off
the gold foil at very large
angles.
Rutherford’s Gold-foil Experiment
Based on his experimental results, Rutherford suggested a
new theory of the atom.
He proposed that the atom is mostly empty space, thus
explaining the lack of deflections of most of the alpha
particles.
He concluded that all the positive charge and almost all
the mass are concentrated in a small region that has
enough positive charge to account for the great
deflection .
Nucleus – the tiny central core of an atom and is
composed of protons and neutrons.
Questions
Describe Thomson’s and Millikan’s contributions to atomic
theory.
Thomson – Cathode ray experiments which concluded
that electrons must be parts of the atoms of all
elements. Millikan determined the charge and mass of
the electron.
What experimental evidence led Rutherford to conclude
that an atom is mostly empty space?
The great majority of the alpha particles passed straight
through the gold foil
Questions
Compare Rutherford’s expected outcome of the gold-foil
experiment with the actual outcome.
Expected all alpha particles to pass straight through with
little deflection. Found that most passed straight
through, but some particles were deflected at large
angles and some bounced back.
Distinguishing Among Atoms
How are atoms of hydrogen different from atoms of
oxygen?
Elements are different because they contain different
number of protons.
Atomic number – of an element is the number of protons
in the nucleus of an atom of that element.
Example – all hydrogen atoms have 1 proton and the
atomic number of hydrogen is 1.
The atomic number identifies an element.
Distinguishing Among Atoms
Most of the mass of an atom is concentrated in its nucleus
and depends on the number of protons and neutrons.
Mass number – the total number of protons and neutrons
in an atom
Example: Helium atom contains 2 protons and two
neutrons, so its mass number is 4
If you know the atomic number and mass number of an
atom of any element, you can determine the atom’s
composition.
Distinguishing Among Atoms
Example: Oxygen
Atomic number is 8 = number of p+ = e- (So oxygen has 8
electron s and 8 protons.)
Mass number is 16 = number of p+ plus the number of n0.
(So oxygen has 8 neutrons)
Number of neutron = mass number – atomic number
Mass number
Atomic number
197
79
Au
Isotopes
There are some elements that have different kinds of
atoms of the same element
Example – there are three different kinds of Neon atoms
Isotopes – are atoms that have the same number of
protons, but different numbers of neutrons.
Because isotopes of an element have different numbers of
neutrons, they also have different mass numbers.
Isotopes are chemically alike because they have identical
numbers of protons and electrons, which are the subatomic
particles responsible for chemical behavior.
Chemical Symbols of Isotopes
Write the chemical symbols for three isotopes of oxygen.
Oxygen 16, oxygen 17, and oxygen 18.
Mass Number
(# protons + # neutrons)
16
17
O
8
18
O
8
O
8
Atomic number
(# proton = # electrons)
Atomic Mass
The slight difference takes into account the larger masses,
but smaller amounts of the other two isotopes of
hydrogen.
Atomic mass – of an element is a weighted average mass
of the atoms in a naturally occurring sample of the
element.
The atomic mass of copper is 63.546 amu. Which of
copper’s two isotopes is more abundant: copper -63 or
copper-65?
Atomic mass of 63.546 is closer to 63 than 65, thus
copper-63 must be more abundant.
Atomic Mass
Atomic mass = multiply the mass of each isotope by its
natural abundance, expresses as a decimal, and then add
the products.
Element X has two natural isotopes. The isotope with a mass of
10.012 amu has a relative abundance of 19.91%. The isotope with
a mass of 11.009 amu has a relative abundance of 80.09%.
Calculate the atomic mass of this element.
(10.012 amu x 0.1991) + (11.009 amu x 0.8009)
(1.993 amu)
+
(8.817 amu)
Atomic mass = 10.810
Question
Copper – 63 has a mass of 62.93 amu and 69.2%
abundance. Copper-65 has a mass of 64.93 amu and
30.8% abundance. What is copper’s average atomic
mass?
(62.93 amu x 0.692) + (64.93 amu x 0.308)
(43.548 amu)
+
(19.998 amu)
Atomic mass = 63.55
Periodic Table
Each element is identified by its symbol place in a square.
The atomic number of the element is shown centered above the
symbol. Elements are listed in order of increasing atomic number,
from left to right and from top to bottom.
Period - each horizontal row of the periodic table. Within a given
period, the properties of the elements vary as you move across it
from element to element.
Group – each vertical column of the periodic table. Elements within
a group have similar chemical and physical properties. Each
group is identified by a number and the letter A or B.
Chapter 5
Models of the Atom
Atomic Models
Rutherford used existing ideas bout the atom and
proposed an atomic model in which the electrons move
around the nucleus.
However, Rutherford’s atomic model could not explain the
chemical properties of element.
Niels Bohr, a student of Rutherford’s, changed
Rutherford’s model to include how the energy of an atom
changes when it absorbs or emits light.
The Bohr Model – he proposed that an electron is found
only in specific circular paths, or orbits, around the
nucleus.
The Bohr Model
Each possible electron orbit in Bohr’s model has a fixed
energy. The fixed energies an electron can have are called
energy levels.
The fixed energy levels of electrons are somewhat like the
rungs of the ladder in which the lowest rung of the ladder
corresponds to the lowest energy level.
An electron can jump from one energy level to another.
Electrons in an atom cannot be between energy levels.
The Bohr Model
To move from one energy level to another, an electron must
gain or lose jus the right amount of energy.
In general, the higher an electron is on the energy ladder,
the farther it is from the nucleus.
A quantum of energy is the amount of energy required to
move and electron from one energy level to another energy
level.
The energy of an electron is said to be quantized.
The term quantum leap originates from the ideas found in
the Bohr model of the atom.
The Quantum Mechanical Model
The Quantum Mechanical Model is the modern description of
the electrons in atoms comes from the mathematical solution to
the Schrodinger equation.
Like the Bohr model, the quantum mechanical model restricts
the energy of electrons to certain values.
Unlike the Bohr model, the quantum mechanical model does
not involve an exact path the electron takes around the nucleus.
The quantum mechanical model determines the allowed
energies an electron can have an how likely it is to find the
electron in various locations around the nucleus
The Quantum Mechanical Model
How likely it is to find the electron in a particular location is
described by probability.
The quantum mechanical model
describes of how the electron
moving around the nucleus is
similar to the motion of a rotating
propeller blade.
The propeller blade has the same probability of being
anywhere in the blurry regions it produces, but you cannot
tells its precise location at any instant.
The Quantum Mechanical Model
The probability of finding an electron within a certain volume of
space surrounding the nucleus can be represented as a fuzzy
cloud.
The cloud is more dense where the probability of finding the
electron is high. The cloud is less dense where the probability of
finding the electron is low.
It is unclear where the cloud ends, there is at least a slight
chance of finding the electron at a considerable distance form
the nucleus.
Number Number
Energy
of
of
Number
Energy
Sublevel Orbitals Orbitals of e- per
Level
( # = n)
per
per
Sublevel
Type
Level
n=1
1s
1
n=2
2s
2p
1
3
n=3
3s
3p
3d
1
3
5
n=4
4s
4p
4d
4f
1
3
5
7
Max e- in
Sublevel
Maximum
e- in
Energy
Level
(2n2)
1
2e-
2e-
2 e-
4
2e2e-
2e6e-
8 e-
9
2e2e2e-
2e6e10e-
18 e-
16
2e2e2e2e-
2e6e10e14e-
32 e-
Electron Configuration
In most natural phenomena, change proceeds toward the
lowest possible energy.
In the atom, electrons and the nucleus interact to make the
most stable arrangement possible.
The way in which electrons are arranged into various orbitals
around the nuclei of atoms are called electron configuration.
Three rules tell you how to find the electron configurations of
atoms.
•The aufbau principle
•The Pauli exclusion principle
•Hund’s rule
Electron Configuration Rules
aufbau Principle
Electrons occupy the orbitals of lowest energy first.
Pauli Exclusion Principle
• An orbital can hold a maximum of 2 electrons.
• 2 electrons in the same orbital must have opposite
spins.
• An electron is "paired" if it is sharing an orbital with
another electron with an opposite spin.
• An electron is "unpaired" if it is alone in an orbital
Paired
unpaired
Electron Configuration Rules
Hund’s Rule
•Electrons occupy orbitals of the same energy in a way that
makes the number of electrons with the same spin direction
as large as possible.
•One electron enters each orbital until all the orbitals
contain one electron with the same spin direction
•For example, three electron would occupy three orbitals of
equal energy as follows:
•Second electrons then occupy each orbital so that their
spins are paired with the first electron in the orbital. Thus
each orbital can eventually have two electrons with paired
spins.
Electron Configuration Practice
Write the electron configuration for each atom. How many
unpaired electrons does each atom have?
Carbon (atomic number 6 so 6 protons = 6 electrons)
1s22s22p2
2 unpaired electrons
Argon
1s22s22p63s23p6
no unpaired electrons
Silicon
1s22s22p63s23p2
2 unpaired electrons
Exceptional Electron Configurations
Some actual electron configurations differ from those assigned
using the aufbau principle because half-filled sublevels are not
as stable as filled sublevels.
You can obtain correct electron configurations for the elements
up to vanadium (atomic number 23) by following the aufbau
diagram for orbital filling.
Cr 1s22s22p63s23p64s23d4 using aufbau
Cr 1s22s22p63s23p64s13d5
correct
Exceptional Electron Configurations
Transition elements are some exceptions to the filling rules.
These exceptions can be explained by the atom’s tendency to
keep its energy as low as possible.
These exceptions help explain the unexpected chemical
behavior of transition elements.
Shorthand Electron Configurations
Electron configurations are often abbreviated by naming the last
element with a filled shell (halogens) in brackets and listing only
the orbitals after the filled shell.
Na: 1s22s22p63s1
shorthand
Al:
Na: [Ne] 3s1
1s22s22p63s23p1
shorthand
Al: [Ne] 3s23p1
V: 1s22s22p63s23p6 4s23d3
shorthand
V: [Ar] 4s23d3
Waves
Each complete wave cycle
starts at zero, increases to
its highest value, passes
through zero to reach its
lowest value, and returns
to zero again.
Amplitude of a wave is
the wave’s height from
zero to the crest.
Wavelength (λ) is the
distance between the
crests.
Waves
Frequency (ν) is the number of wave cycles to pass a given point
per unit of time.
The units of frequency are usually cycles per second. The SI unit
of cycles per second is called a hertz (Hz)
A hertz can also be expressed as a reciprocal seconds (s-1)
Hz = s-1
Light
The product of frequency and wavelength always equal a
constant (c) = the speed of light
c = λν
The wavelength and frequency of light are inversely
proportional to each other. As the wavelength increases, the
frequency decreases.
According to the wave model, light consists of electromagnetic
waves.
Electromagnetic radiation includes radio waves, microwaves,
infrared waves, visible light, ultraviolet waves, X-rays, and
gamma rays.
Light
All electromagnetic waves travel in a vacuum at a speed of 2.998
x 108 m/s
c = 2.998 x 108 m/s
Sunlight consists of light with a continuous range of wavelengths
and frequencies. The color of light depends on its frequency.
When sunlight passes through a prism, the different frequencies
separate into a spectrum of color.
A rainbow is an example of this phenomenon.
Electromagnetic
Spectrum
Each color of the spectrum blends into the next in the order red,
orange, yellow green, blue and violet.
In the visible spectrum, red light has the longest wavelength and
the lowest frequency.
Sample Problems
What is the wavelength of radiation with a frequency of 1.50 x
1013 Hz? Does this radiation have a longer or shorter wavelength
than red light?
c = λν
or
λ=c/ν
λ = (2.998 x 108 m/s) / (1.50 x 1013 s-1)
λ = 2.00 x 10-5 m (longer wavelength than red light)
What frequency is radiation with a wavelength of 5.00 x 10-8m?
In what regions of th e electromagnetic spectrum is this
radiation?
c = λν
or
ν=c/λ
ν = (2.998 x 108 m/s) / (5.00 x 10-8 m)
ν = 6.00 x 1015 s-1 (ultraviolet)
When light passes through a prism, the frequencies of light
emitted by an element separate into discrete lines to give the
atomic emission spectrum of the element.
Explanation of Atomic Spectra
Atomic line spectra were known before Bohr proposed his
model of the H atom. However, Bohr’s model explained why the
emission spectrum of H consists of specific frequencies of light.
In the Bohr model, the lone electron in the H atom can have
only certain specific energies.
The lowest possible energy of the electron is its ground state.
In the ground state, the electron’s principal quantum number is
1 (n=1)
Excitation of the electron by absorbing energy raises it from the
ground state to an excited state with n = 2,3,4,5…
Explanation of Atomic Spectra
A quantum of energy in the form of light is emitted when the
electron drops back to a lower energy level.
The emission occurs in a single abrupt step, called an electronic
transition.
Bohr knew from earlier work that the quantum of energy (E) is
related to the frequency (ν) of the emitted light by the equation
E=hxν
h is the fundamental constant of nature, the “Planck constant”
and is equal to 6.626 x 10-34 J·s
Explanation of Atomic Spectra
(Transition to n = 3 energy level,
infrared range of spectra)
(Transition to n = 2 energy level
Visible end of the spectra)
(Transition to the n = 1 energy level
electron
moving
from
a higher
Ultraviolet
part of the
spectra)
The light emitted by an
to a
lower energy level has a frequency directly proportional to the
energy change of the electrons.
Each transition produces a line of a specific frequency in the
spectrum.
Quantum Mechanics
Albert Einstein successfully explained experimental data by
proposing that light could be described as quanta of energy.
The quanta behave as if they were particles.
Light quanta are called photons.
Although the wave nature of light was well known, the dual waveparticle behavior of light was difficult for scientists to accept.
Louis de Broglie a French graduate student, asked an important
question: Given that light behaves as waves and particles, can
particles of matter behave as waves?
The proposal that matter moves in a wavelike way would not be
accepted unless experiments confirmed its validity.
Quantum Mechanics
German physicist Werner Heisenberg examined another feature of
quantum mechanics that is absent is classical mechanics.
The Heisenberg uncertainly principle states that it is impossible to
know exactly both the velocity and the position of a particle at the
same time.
This limitation is critical in dealing with small particles such as
electrons.
The Heisenberg uncertainty principle does not matter, however, for
ordinary-sized objects such as cars or airplanes.
Recap
The frequency and wavelength of light waves are inversely related.
As the wavelength increases, the frequency decreases. (c = λν)
The electromagnetic spectrum consists of radiation over a broad
band of wavelengths. The visible light portion is very small. It is in
the 10-7 m wavelength rand 1015 Hz (s-1) frequency range.
When atoms absorb energy, electrons move into higher energy
levels, and these electrons lose energy by emitting light when they
return to lower energy levels.
Recap
A prism separates light into the colors it contains. For white light
this produces a rainbow of colors. Light from a helium lamp
produces discrete lines.
An electron microscope can produce sharp images of a very small
object, because of the small wavelength of a moving electron
compared with that of light.
The Heisenberg uncertainty principle states that it is impossible to
know exactly both the velocity and the position of a particle at the
same time.
Chapter 6
The Periodic Table
The Periodic Law
Mendeleev developed his table before scientists knew
about the structure of atoms. He did not know that the
atoms of each element contain a unique number of
protons.
A British physicist, Henry Moseley, determined an
atomic number for each known element.
In the modern periodic table, elements are arranged in
order of increasing atomic number.
The Periodic Law
There are seven rows, or periods in
the table.
Period 1 has 2 elements, Period 2
has 8 elements, Period 4 has 18
elements & Period 6 has 32
elements.
Each period corresponds to a principal energy level.
There are more elements in higher numbered
periods because there are more orbitals in higher
energy levels.
The Periodic Law
The elements within a column or
group in the periodic table have
similar properties.
The properties of the elements
within a period change as you
move across a period from left to
right.
The pattern of properties within a period repeats as you
move from one period to the next.
The Periodic Law
Periodic Law – when elements are arranged in order
of increasing atomic number, there is a periodic
repetition of their physical and chemical properties.
Group 1 – (alkali metals) are all highly reactive and are
rarely found in elemental form in nature
Group 2 – (alkaline earth metals) are silvery colored,
soft metals
Group 17- (halogens) the only group which contains
elements in all three familiar states of matter at
standard temperature and pressure.
Metal, Nonmetals, and Metalloids
The International Union of Pure and Applied Chemistry
(IUPAC) set the standard for labeling groups in the periodic
table.
They numbered the groups from left to right 1 – 18,
The elements can be grouped into three broad classes
based on their general properties.
• Metals
• Nonmetals
• Metalloids
Across the period, the properties of elements become
less metallic and more nonmetallic.
Metals
About 80 % of the elements are metals.
Properties of Metals
• Good conductors of heat and electric current.
• Have a high luster or sheen caused by the ability
to reflect light
• Solids at room temperature (except Hg)
• Many metals are ductile (can be drawn into wires)
• Most metals are malleable (they can be
hammered into thin sheets without breaking)
Nonmetals
Nonmetals are in the upper-right corner of the periodic
table.
There is a greater variation in physical properties
among nonmetal than among metals.
Properties of Nonmetals
• Most are gases at room temperature. S and P are
solids, Br is a liquid.
• Nonmetals tend to have properties that are
opposite to those of metals.
• In general, nonmetals are poor conductors of heat
and electric current. Solid nonmetals tend to be
brittle.
Metalloids
There is a heavy stair-step lines that separates the
metals from the nonmetals.
Most of the elements that border this line are
metalloids.
Properties of Metalloids
• Generally has properties that are similar to metals
and nonmetals.
• Under some conditions they behave like a metal.
Under other conditions they behave like a nonmetal.
Questions
How did chemists begin the process of organizing elements?
Used the properties of elements to sort them into groups.
What property did Mendeleev use to organize his periodic table?
In order of increasing atomic mass
How are elements arranged in the modern periodic table?
In order of increasing atomic number
Name the three broad classes of elements.
Metals, nonmetals, and metalloids
Squares in the Periodic Table
The periodic table displays the symbols and names of the
elements along with information about the structure of their
atoms.
The symbol for the element is located in the center of the
square.
The atomic number is above the symbol.
The element name and average atomic mass are below the
symbol.
Squares in the Periodic Table
The background colors in the squares are used to
distinguish groups of elements. (Ex:2 shades of gold are
used for the metals in Groups IA and 2A)
Group IA elements are called alkali metals. Group 2A
elements are called alkaline earth metals.
The nonmetals of Group 7A are called halogens.
Group 8A elements are called Noble Gases
Groups 1B – 8B are called transition metals
The two periods usually located at the bottom of the
periodic table separate from the main table are called inner
transition elements. Period 8 is called the Lanthanide
Series and Period 9 is called the Actinide Series
Electron Configuration in Groups
Electrons play a key role in determining the properties of
elements.
So there is a connection between an element’s electron
configuration and its location in the periodic table.
Elements can be sorted into noble gases, representative
elements, transition metals, or inner transition metals based
on their electron configurations.
The Noble Gases are in Group 8A and are sometimes
called inert gases because they rarely take part in a
reaction.
Electron Configuration in Groups
Helium (He)
1s2
Neon (Ne)
Argon (Ar)
Krypton (Kr)
1s22s22p6
1s22s22p63s23p6
1s22s22p63s23p63d104s24p6
The highest occupied energy level for each element, (the s & p
sublevels) are completely filled with electrons.
s sublevel
p sublevel
The Representative Elements
Elements in groups 1A through 7A are often referred to as
representative elements because they display a wide
range of physical and chemical properties.
In atoms of representative elements, the s and p sublevels
of the highest occupied energy level are not filled.
Lithium(L)
1s22s1
Sodium (Na) 1s22s22p63s1
Potassium (K) 1s22s22p63s23p64s1
s sublevel
The Representative Elements
Carbon (C)
1s22s22p2
Silicon (Si)
1s22s22p63s23p2
Germanium (Ge) 1s22s22p63s23p64s23d104p2
In atoms of carbon, silicon, and germanium, in Group 4A, there
are four electrons in the highest occupied energy level
For any representative elements, its group number equals the
number of electrons in the highest occupied energy level.
s sublevel
p sublevel
Transition Metals
Elements in the B groups are referred to as transition
elements.
There are two types of transitions elements: transition
metals and inner transition metals
In atoms of a transition metal, the highest occupied s
sublevel and a nearby d sublevel contain electrons.
These elements are characterized by the presence of
electrons in d orbitals.
Ions
Some compounds are composed of particles called ions.
An ion is an atoms or group of atoms that has a positive
or negative charge.
An atom is electrically neutral because it has equal
numbers of protons and electrons.
Positive and negative ions from when electrons are
transferred between atoms.
Atoms of metallic elements tend to form ions by losing one
or more electrons from their highest occupied energy
levels.
A sodium atom tend to lose one electron.
Cations
In the sodium ion, the number of electrons (10) is no
longer equal to the number of protons (11).
Because there is more positively charged protons than
negatively charged electrons, the sodium ion has a net
positive charge.
An ion with a positive charge is called a cation.
The charge for a cation is written as a number followed by
a plus sign. (Example: 1+ )
If the charge is 1+, the number 1 is usually omitted from
the complete symbol for the ions. (Na+)
Anions
Atoms of nonmetallic elements, such as chlorine, tend to
form ions by gaining one or more electrons.
A chlorine atom tend to gain one electron.
In a chlorine ion, the number of electrons (18) is no longer
equal to the number of protons (17).
Because there are more negatively charged electrons
than positively charged protons, the chloride ion has a net
negative charge.
An ion with a negative charge is called an anion.
Examples: Cl-, S2-
Trends in Ionization Energy
Recall that electrons can move to higher energy levels
when atoms absorb energy.
Sometimes there is enough energy to overcome the
attraction of the protons in the nucleus.
The energy required to remove an electron from an atom
is called ionization energy.
The energy to remove the first electron from an atom is
called the first ionization energy.
The cation produced has a 1+ charge.
Ionization Energy
The energy to remove the first electron from an atom is
called the first ionization energy. The cation produced
has a 1+ charge.
The second ionization energy is the energy required to
remove an electron from an ion with a 1+ charge. The ion
produced has a 2+ charge.
The third ionization energy is the energy required to
remove an electron from an ion with a 2+ charge. The ion
produced has a 3+ charge.
Trends in Electronegativity
There is a property that can be used to predict the type of
bond that will form during a reaction.
This property is electronegativity, which is the ability of
an atom of an element to attract electrons when the atom
is in a compound.
In general, electronegativity values decrease from top to
bottom within a group.
For representative elements, the values tend to increase
from left to right across a period.
Trends for Groups 1A
Atomic size decreases
Through 8A
• Can be explained by variations
Ionization energy increases
in atomic structure
• Increase in nuclear charge
Electronegativity increases
within groups & across
periods, also shielding within
Nuclear charge increases
groups
Shielding increases
Nuclear charge increases
Electronegativity decreases
Ionic size increases
Atomic size increases
Ionization Energy decreases
Shielding is constant
Size of cation decreases
Size of anions decreases
Chapter 7
Ionic and Metallic Bonding
valence Electrons
Scientists learned that all of the elements within each group of
the periodic table behave similarly because they have the
same number of valence electrons.
valence electrons are the electrons in the highest occupied
energy level of an element’s atom.
The number of valence electrons largely determines the
chemical properties of an element.
To find the number of valence electrons in an atom of a
representative elements, simply look at its group number
Elements of Group IA have one valence electron. Elements in
Group 4A have four valence electrons, and so forth
valence Electrons
The noble gases, Group 8A, are the only exceptions to the
group-number rule.
Helium has two valence electrons, and all of the other noble
gases have eight.
valence electrons are usually the only electrons used in
chemical bonds.
As a general rule, only the valence electrons are shown in
electron dot structures.
Electron dot structures are diagrams that show valence
electrons as dots.
Electron Dot Structures
The Octet Rule
Noble gases, such as neon and argon, are unreactive in chemical
reactions. (They are stable)
Gilbert Lewis explained why atoms form certain kinds of ions and
molecules in the octet rule
The Octet Rule - in forming compounds, atoms tend to achieve the
electron configuration of a noble gas. An octet is a set of eight. (each
noble gas except helium has eight electrons in its highest energy
level)
Atoms of the metallic elements tend to lose their valence electrons,
leaving a complete octet in the next-lowest energy level. Atoms of
some nonmetallic elements tend to gain electron or to share
electrons with another nonmetallic element to achieve a complete
octet.
Formation of Cations
Using electron dot structures, you can show the ionization of
some elements more simply.
Na·
Na+
Sodium atom
Sodium ion
neutral
1 unit of + charge
·Mg·
Mg2+
+
e-
electron
1 unit of - charge
+
2e-
Magnesium atom Magnesium ion
electron
neutral
2 unit of + charge 2 units of - charge
Transition Metals
For transition metals, the charges of cations may vary.
An atom of iron (Fe) may lose two, or three electrons forming
either Fe2+ or Fe3+ ions.
Some ions formed by transition metals do not have noble gas
electron configurations and are therefore exceptions to the
octet rule.
Ag is an example - 1s22s22p63s23p63d104s24p64d105s1
To achieve the structure of krypton, which is the preceding
noble gas, a silver atom would have to lose eleven electrons.
Transition Metals
Ions with charges of three or greater are uncommon, and
losing eleven electrons is highly unlikely.
If Ag loses its 5s1 electron, the configuration that results,
(4s24p64d10) with 18 electrons in the outer energy level and all
of the orbitals filled, is relatively favorable in compounds.
Such a configuration is known as pseudo noble-gas electron
configuration.
Ag forms a positive ion (Ag+) in this way.
Formation of Anions
The gain of negatively charge electrons by a neutral atom
produces an anion.
The name of an anion of a nonmetallic element is not the
same as the element name. The name of the ion typically ends
in -ide.
Chlorine atom (Cl) forms a chloride ion (Cl-)
Oxygen atom (O) forms an oxide ion (O2-)
Because they have relatively full valence shells, atoms of
nonmetallic elements attain noble-gas electron configurations
more easily by gaining electrons than by losing them.
Formation of Anions
Chlorine belongs to Group 7A and has seven valence electrons.
A gain of one electron gives chlorine an octet and converts a
chlorine atom into a chloride ion.
Atoms of nonmetallic elements form anions by gaining enough
valence electrons so as to attain the electron configuration of
the nearest noble gas.
The chloride ion has the same electron configuration as the
noble gas argon.
Chloride ion (Cl-) 1s22s22p63s23p6
Argon (Ar) 1s22s22p63s23p6
Formation of Ionic Compounds
Compounds composed of cations and anions are called ionic
compounds.
Ionic compounds are usually composed of metal cations and
nonmetal anions. Ex: NaCl is formed from Na+ + ClAlthough they are composed of ions, ionic compounds are
electrically neutral. The total + charge of the cations equals
the total – charge of the anions.
Anions and cations have opposite charges and attract one
another by means of electrostatic forces.
The electrostatic forces that hold ions together in ionic
compounds are called ionic bonds.
Formation of Ionic Compounds
Look at the reaction of a Na atom and a chlorine atom.
Na has 1 valence electron that it can easily lose. (Na is in group
1A of the representative elements, thus has 1 valence electron)
Cl has seven valence electrons and can easily gain one
electron. (Cl is in group 7A of the representative elements, thus
has 7 valence electrons)
If Na loses its valence electron it achieves the stable electron
configuration of neon. If Cl gains a valence electron, it achieves
the stable electron configuration of argon. (Remember the
Octet Rule)
Formation of Ionic Compounds
When Na and Cl react, the Na atom gives its one valence
electron to a Cl atom. They react in a 1:1 ratio and both ions
have stable octets.
+
Na+
Cl-
1s22s22p6
1s22s22p63s23p6
Formula Units
Chemists represent the composition of substances by writing
chemical formulas. A chemical formula shows the kinds and
numbers of atoms in the smallest representative unit of a
substance.
NaCl is the chemical formula for sodium chloride.
A Formula unit is the lowest whole-number ratio of ions in an
ionic compound. One Na+ to each Cl-, thus the formula unit for
sodium chloride is NaCl.
Even though ionic charges are used to derive the correct
formulas, they are not shown when you write the formula unit
of the compound
Formula Units
The ionic compound Magnesium chloride (MgCl2) contains
magnesium cations (Mg2+) and chloride anions (Cl-)
In MgCl2, the ratios of Mg2+ to Cl- is 1:2 (One Mg2+ to two Cl). Its formula unit is MgCl2
Because there are twice as many Cl- (each with a 1- charge)
as Mg2+ (each with a 2+ charge), the compound is
electrically neutral.
Another example: Al3+ + Br- combine to form AlBr3.
Metallic Bonds & Properties
Metals are made up of closely packed cations rather than neutral
atoms.
The valence electrons of metal atoms can be modeled as a sea of
electrons. (they are mobile and can drift freely from one part
of the metal to another).
Metallic bonds consists of the attraction of the free-floating
valence electrons from the positively charged metal ion.
The sea-of-electrons model explains many physical properties of
metals.
– Good conductors of electrical current because electrons can
flow freely.
– Ductile – they can be drawn into wires.
– Malleable – they can be hammered or forced into shapes.
Crystalline Structure of Metals
There are several closely packed arrangements that are
possible.
• body-centered cubic arrangement
• face-centered cubic arrangement
• hexagonal close-packed arrangement
Body-centered cubic
Every atom (except those on the
Surface) has eight neighbors.
Crystalline Structure of Metals
Face-centered cubic arrangement
• every atom has twelve neighbors.
Crystalline Structure of Metals
Hexagonal close-packed arrangement
• every atom also have twelve neighbors. Because of the
hexagonal shape, the pattern is different from the facecentered.
Alloys
Very few of the metallic items that you use every day are pure
metals. Ex: spoons.
Most of the metals you encounter are alloys.
Alloys are mixtures composed of two or more elements., at
least on of which is a metal. Ex: Brass (Cu & Zn)
Alloys properties are often superior to those of their
component elements.
Sterling silver (92.5% silver & 7.5% copper) is harder and more
durable than pure silver, but still soft enough to be made into
jewelry and tableware.
Laws Governing Formulas & Names
Law of Definite Proportions
A chemical formula tells you (by subscripts) the ratio of
atoms of each element in the compound.
Ratios of atoms can also be expressed as ratios of
masses.
100 g of MgS breaks down into 43.12g Mg and 56.88g of
sulfur.
100g MgS 1 mol MgS 1 mol Mg
24.305g Mg = 43.12g Mg
56.4g MgS 1 mol MgS 1 mol Mg
100g MgS 1 mol MgS 1 mol S
32.06g S = 56.88g S
56.4g MgS 1 mol MgS 1 mol S
Chapter 10
Chemical Quantities
Measuring Matter
Avogadro’s number is the number of representative particles
in a mole, 6.02 x 1023.
The term representative particles refers to the species present
in a substance: usually atoms, molecules or formula units.
Representative particles for ionic compounds is the formula
unit : CaCl2 , NaCl
Representative particles for molecular compounds is the
molecule: H2O , H2
Representative particles for most elements is the atom: Fe, Li
Measuring Matter
A mole of any substance contains Avogadro’s number of
representative particles or 6.02 x 1023 representative
particles.
The relationship, 1 mole = 6.02 x 1023 representative particles,
is the basis for a conversion factor to convert numbers of
representative particles to moles.
How many moles of Mg is 1.25 x 1023 atoms of Mg?
1.25 x 1023 atoms Mg (1 mol Mg / 6.02 x 1023 atoms Mg)
Measuring Matter
How many atoms are in 2.12 mol of propane (C6H8)?
In the formula of a molecule of C3H8 , the subscripts show that
propane is composed of 14 atoms: 3 atoms of C and 8 atoms
of H.
2.12 mol C6H8 6.02 x 1023 molecules C6H8 11 atoms
1 mol C6H8
1 molecule of C6H8
1.40 x 1025 atoms
Mass of a Mole
The atomic mass of an element (mass of a single atom) is
expressed in atomic mass units (amu)
The atomic masses are relative values based on the mass of
the most common isotope of carbon 12.
The atomic mass of an element expressed in grams is the mass
of a mole of the element.
The mass of a mole of an element is its molar mass.
Molar mass of C is 12.0 g. H – 1.0 g, S – 32.1g
Molar mass is the atomic mass of an element rounded off to
the first decimal place.
Molar Mass
If you were to compare 12.0g of C atoms with 16.0g of O
atoms, you would find they contain the same number of
atoms.
The molar mass of any element contains 1 mole or 6.02 x 1023
atoms of that element.
12.0g of C is 1 mol of C atoms
1.0 g of H is 1 mol of H atoms
Molar mass is the mass of 1 mole of atoms of any element.
Mass of a Mole of a Compound
To find the mass of a mole of a compound, you must know the
formula of the compound.
A molecule sulfur trioxide, SO3, is composed of one atom of
sulfur and three atoms of oxygen.
Calculate the mass of a molecule of SO3 by adding the atomic
masses of the atoms making up the molecule.
The atomic mass of Sulfur is 32.1g and the mass of three
Oxygen atoms is 48.0g (3 x 16.0), so the molecular mass of
SO3 is 80.1g (32.1 + 48.0)
The molar mass of any compound is the mass of 1 mole of
that compound.
Mass of a Mole of a Compound
1 mole of SO3 has a mass of 80.1g and is the mass of 6.02 x
1023 molecules of SO3
To calculate the molar mass of a compound, find the number
of grams of each element in one mole of the compound and
then add the masses of the elements.
The method for calculating molar mass applies to any
compound, molecular or ionic.
Mole/Mass Relationship
You need 3.00 mol of NaCl. How do you measure this amount?
What mass in grams is 3.00 mol of NaCl?
3.00 mol NaCl 58.5 g NaCl = 176g NaCl
1 mol NaCl
(use the molar mass)
When you measure 176g of NaCl on a balance, you are
measuring 3.00 mol of NaCl.
What is the mass of 9.45 mol of aluminum oxide? (Al2O3)
9.45 mol Al2O3
102.0g Al2O3 = 964 g Al2O3
1 mol Al2O3
Mole/Mass Relationship
How many moles of sodium sulfate (Na2SO4) is in 10 g of
Na2SO4?
10.0 g Na2SO4 1 mol Na2SO4 = 7.04 x 10-2 mol Na2SO4
142.1 g Na2SO4
How many moles of iron(III) oxide are contained in 92.2 g of
pure Fe2O3?
92.2 g Fe2O3
1 mol Fe2O3 = 0.578 mol Fe2O3
159.6 g Fe2O3
Mole/Volume Relationship
The volume of one mole of different solid and liquid
substances are not the same. However, the volumes of moles
of gases measured under standard condition are much more
predictable.
Avogadro’s hypothesis states that equal volumes of gases at
the same temperature and pressure contain equal numbers of
particles.
If you buy a party balloon filled with helium and take it home
on a cold day, you might notice that the balloon shrinks while
it is outside.
The volume of a gas varies with a change in temperature.
Mole/Volume Relationship
The volume of a gas also varies with a change in pressure. An
increase in pressure causes the volume of the gas to
decrease.
Because of these variation due to temperature and pressure,
the volume of a gas is usually measured at standard
temperature and pressure.
Standard temperature and pressure (STP) means a
temperature of 0ºC and a pressure of 101.3 kPa (1atm)
At STP, 1 mole or 6.02 x 1023 representative particles of any
gas occupies 22.4L.
22.4 L is called the molar volume of gas.
Mole/Volume Relationship
If you have 0.375 mol of O2 gas, what volume at STP will this
gas occupy?
0.375 mol O2 22.4L O2 = 8.40 L O2
1 mol O2
Determine the volume in liters of 0.60 mole of SO2 gas at STP.
0.60 mol SO2
22.4L SO2
1 mol SO2
= 13 L SO2
How many moles of H2 are in 0.200 L at STP?
0.200 L H2
1 mol H2
22.4 L H2
= 8.93 x 10-3 mol H2
Molar Mass From Density
Different gases have different densities.
Usually the density of a gas is measured in grams per liter (g/L)
The density of a gas at STP and the molar volume at STP can be
used to calculate the molar mass of the gas.
The density of a gaseous compound containing C and O is
1.964 g/L at STP. What is the molar mass of the compound?
1.964 g
L
22.4 L
1 mol
= 44.0 g/mol
Percent Composition
The relative amounts of the elements in a compound are
expressed as the percent composition or the percent by mass
of each element in the compound.
The percent composition of a compound consists of a percent
value for each different element in the compound.
The percent composition of K2CrO4 is K = 40.3%, Cr = 26.8%, O
= 32.9%. (They must total 100%)
The percent by mass of an element in a compound is the
number of grams of the element divided by the mass in grams
of the compound, multiplied by 100%.
Percent Composition
% mass of element = mass of element
mass of compound
x 100%
When a 13.60 g sample of a compound containing only Mg
and O is decomposed, 5.40g of O is obtained. What is the
percent composition of this compound?
% O = 5.40 g / 13.60g x 100% = 39.7%
% Mg = 13.60 g – 5.40 g / 13.60g x 100% = 60.3%
Percent Composition by Formula
% mass = mass of element in 1 mol compound x 100%
molar mass of compound
Calculate the percent composition of propane C3H8
% C = 36.0 g / 44.0 g x 100% = 81.8%
% H = 8.0 g / 44.0 g x 100% = 18.0%
Percent Composition as a Conversion
Factor
How much C and H are contained in 82.0 g of propane? (C3H8)
Calculate the percent composition of propane C3H8
% C = 36.0 g / 44.0 g x 100% = 81.8%
% H = 8.0 g / 44.0 g x 100% = 18.0%
In a 100 g sample of propane you would have 81.8 g of C and
18 g of O.
(82.0 g propane)(81.8 g C / 100 g propane) = 67.1 g C
(82.0 g propane)(18 g O / 100 g propane) = 15 g H