Chapter 1
Introduction to Chemistry
I. Chemistry
A. What is matter?
1. Chemistry is the study of the composition of matter and the
changes that it undergoes.
a) Matter is anything that has mass and occupies space.
b) Because living and nonliving things are made of matter,
chemistry affects all aspects of life and most natural
events.
B. Areas of Study
1. Five traditional areas of study are organic chemistry,
inorganic chemistry, biochemistry, analytical chemistry, and
physical chemistry.
a) Organic chemistry is defined as the study of all
chemicals containing carbon.
b) Inorganic chemistry is the study of chemicals that, in
general do not contain carbon.
c) The study of processes that take place in organisms is
biochemistry.
d) Analytical chemistry is the area of study that focuses on
the composition of matter.
e) Physical chemistry is the area that deals with the
mechanism, the rate, and the energy transfer that occurs
when matter undergoes a change.
C. Pure and Applied Chemistry
1. Pure chemistry is the pursuit of chemical knowledge for its
own sake.
2. Applied chemistry is research that is directed toward a
practical goal or application.
3. Pure research can lead directly to an application, but an
application can exist before research is done to explain how it
works.
a) Nylon and aspirin provide examples of these two
approaches.
4. Technology is the means by which society provides its
members with those things needed and desired.
D. Why Study Chemistry?
1. Chemistry can be useful in explaining the natural world,
preparing people for career opportunities, and producing
informed citizens.
a) read pages 10-11
II. Chemistry Far and Wide
A. Materials
1. Chemists design materials to fit specific needs.
2. Two different ways of looking at the world are the
microscopic view and the macroscopic view.
B. Energy
1. Energy is necessary to meet the needs of a modern
society.
a) There are two ways to meet the demand for energy –
conserve energy recourses and produce more energy.
b) Chemists play an essential role in finding ways to
conserve energy, produce energy, and store energy. Read
page 13
C. Medicine and Biotechnology Chemistry
1. Chemistry supplies the medicines, materials, and
technology that doctors use to treat their patients.
a) Medicines – There are over 2000 prescription drugs.
Many drugs are effective because they inte4ract in a
specific way with chemicals in cells. Knowledge of the
structure and function of these target chemicals helps a
chemist design safe and effective drugs.
b) Materials – chemistry can supply materials to repair or
replace body parts.
c) Biotechnology applies science to the production of
biological products or processes.
D. Agriculture
1. Chemists help develop more productive crops and safer,
more effective ways to protect crops.
E. The Environment
1. Chemists help to identify pollutants and prevent pollution.
a) A pollutant is a material found in air, water, or soil that
is harmful to humans or other organisms.
F. The Universe
1. To study the universe, chemists gather data from afar and
analyze matter that is brought back to Earth.
III.Thinking Like a Scientist
A. Alchemy
1. The word chemistry comes from alchemy.
2. Alchemists developed the tools and techniques for working
with chemicals.
B. An Experimental Approach to Science
1. By the 1500s in Europe, there was a shift from alchemy to
science.
2. Lavoisier helped to transform chemistry from a science of
observation to the science of measurement that it is today.
C. The Scientific Method
1. An organized plan for gathering, organizing, and
communicating information is called a scientific method.
2. The goal of any scientific method is to solve a problem or
to better understand an observed event.
a) Making Observations – An observation is information
that you obtain through your senses.
(1) Repeatable observations are known as facts.
b) Forming a Hypothesis – A hypothesis is a proposed
answer to a question.
(1) A hypothesis must be testable.
c) Testing a Hypothesis – Scientist perform experiments
to test their hypothesis.
(1) Manipulated variable (independent variable) – the
variable that causes a change in another variable.
(2) Responding variable (dependent variable) – the
variable that changes in response to the manipulated
variable.
(3) A controlled experiment is an experiment in which
only one variable, the manipulated variable, is deliberately
changed at a time. While the responding variable is
observed for changes, all other variables are kept constant,
or controlled.
d) Drawing Conclusions – A scientist must decide if the
data from tests supports his hypothesis. If the data does
not support his hypothesis he must revise the hypothesis
or propose a new one.
e) Developing a Theory – Once a hypothesis has been
supported in repeated experiments, scientist can begin to
develop a theory.
(1) A scientific theory is a well-tested explanation for a set
of observations or experimental results.
3. Scientific Laws
a) A scientific law is a statement that summarizes a
pattern found in nature.
b) A scientific law describes an observed pattern in nature
without attempting to explain it. The explanation of such
a pattern is provided by a scientific theory.
D. Collaboration and Communication
1. When scientists collaborate and communicate, they
increase the likelihood of a successful outcome.
2. Read pages 24-25
IV.
Problem Solving in Chemistry
A. Skills Used in Solving Problems
1. Effective problem solving Always involves developing a
plan and then implementing that plan.
B. Solving Numeric Problems
1. The steps for solving a numeric word problem are analyze,
calculate, and evaluate.
a) Analyze – To solve a word problem, you must first
determine where you are starting from (identify what is
known) and where you are going (identify the unknown).
b) Calculate – If you make an effective plan, doing the
calculations is usually the easiest part of the problem.
c) Evaluate – After you calculate an answer, you should
evaluate it. Is the answer reasonable?
C. Solving Conceptual Problems
1. The steps for solving a conceptual problem are analyze
and solve.
Assignment: Read pages 28-32
1.4 section assessment; page 32; 30 - 33
Workbook section 1.4
Chapter 2
Matter and Change
I. Properties of Matter
A. Describing Matter
1. Properties used to describe matter can be classified as
extensive or intensive.
a) Extensive properties are properties that depend on the
amount of matter in a sample.
(1) The volume of an object is a measure of the space
occupied by the object.
(2) Mass and volume are examples of extensive properties.
b) Intensive properties are properties that depend on the
type of matter in a sample, not the amount of matter.
(1) Hardness is an example of an intensive property.
B. Identifying Substances
1. Matter that has a uniform and definite composition is called
a substance.
2. Every sample of a given substance has identical intensive
and extensive properties.
3. A physical property is a quality or condition of a substance
that can be observed or measured without changing the
substance’s composition.
C. States of Matter
1. Three states of matter are solid, liquid, and gas.
a) A solid is a form of matter that has a definite shape and
volume.
b) A liquid is a form of matter that has an indefinite
shape, flows, and yet has a fixed volume.
(1) Liquids are almost incompressible, but they tend to
expand slightly when heated.
c) A gas is a form of matter that takes both the shape and
volume of its container.
(1) Vapor describes the gaseous state of a substance that is
generally a liquid or solid at room temperature.
d) page 41; figure 2.3
assignment: page 42; 1-8
Workbook section 2.1
II. Mixtures
A. Classifying Mixtures
1. A mixture is a physical blend of two or more components.
2. Based on the distribution of their components, mixtures
can be classified as heterogeneous mixtures or homogeneous
mixtures.
a) A mixture in which the composition is not unifrom
throughout is a heterogeneous mixture.
b) A homogeneous mixture is a mixture in which the
composition is uniform throughout.
(1) Another name for a homogeneous mixture is a
solution.
(a) Most solutions are liquids, but some are gases.
c) The term phase is used to describe any part of a sample
with uniform composition and properties.
B. Separating Mixtures
1. Differences inn physical properties can be used to separate
mixtures.
2. The process that separates a solid from the liquid in a
heterogeneous mixture is called filtration.
3. During a distillation, a liquid is boiled to produce a vapor
that is then condensed into a liquid.
assignment: page 47; 11-17
workbook section 2.2
III.Elements and Compounds
A. An element is the simplest form of matter that has a unique
set of properties.
B. Distinguishing Elements and Compounds
1. A compound is a substance that contains two or more
elements chemically combined in a fixed proportion.
a) Compounds can be broken down into simpler
substances by chemical means, but elements cannot.
b) A chemical change is a change that produces matter
with a different composition than the original matter.
c) The properties of compounds are quite different from
those of their component elements.
C. Distinguishing Substances and Mixtures
1. If the composition of a material is fixed, the material is a
substance. If the composition of a material may vary, the
material is a mixture.
D. Symbols and Formulas
1. Chemists use chemical symbols to represent elements,
and chemical formulas to represent compounds.
a) Each element is represented by a one- or two-letter
chemical symbol.
(1) The first letter of a chemical symbol is always
capitalized. When a second letter is used, it is lowercase.
2. Chemical symbols provide a shorthand wat to write the
chemical formulas of compounds.
Assignment: page 52; 20-27
Workbook section 2.3
IV.
Chemical Reactions
A. Chemical Changes
1. The ability of a substance to undergo a specific chemical
change is called a chemical property.
2. During a chemical change, the composition of matter
always changes.
3. A chemical change is also called a chemical reaction.
a) One or more substances change into one or more new
substances during a chemical reaction.
b) A substance present at the start of a reaction is a
reactant.
c) A substance produced in the reaction is a product.
B. Recognizing Chemical Changes
1. Possible clues to chemical change include a transfer of
energy, a change in color, the production of a gas, or the
formation of a precipitate.
a) Evidence of a chemical change.
(1) A change in color.
(2) Production of a gas.
(3) Formation of a precipitate.
(a) Any solid that forms and separates from a liquid
mixture is called a precipitate.
(4) Production of heat or light.
C. Conservation of Mass
1. During any chemical reaction, the mass of the products is
always equal to the mass of the reactants.
2. The law of conservation of mass states that in any physical
change or chemical reaction, mass is conserved.
Assignment: page 55; 28-34
Workbook section 2.4
Chapter 3
Scientific Measurement
I. Measurements and Their Uncertainty
A. Using and Expressing Measurements
1. Measurements are fundamental to the experimental
sciences. For that reason, it is impourtant to be able to make
measurements and to decide whether a measurement is
correct.
2. In scientific notation, a given number is written as the
product of two numbers: a coeffieient and 10 raised to a
power.
B. Accuracy, Precision, and Error
1. Accuracy is a measure of how close a measurement comes
to the actual or true value of whatever is measured.
2. Precision is a measure of how close a series of
measurements are to one another.
3. To evaluate the accuracy of a measurement, the measured
value must be compared to the correct value. To evaluate the
precision of a measurement, you must comparethe values of
two or more repeated measurements.
4. Determining error
a) Accepted value is the correct value based on reliable
references.
b) Experimental value is the valued measured in lab.
c) The difference between the experimental value and the
accepted value is called error.
d) The percent error is the absolute value of the error
divided by the accepted value, multiplied by 100%.
5. Significant Figures in Measurements
a) The significant figures in a measurement include all of
the digits that are known, plus a last digit that is
estimated.
b) Measurements must always be reported to the correct
number of significant figures because calculated answers
often depend on the n umber of significant figures in the
values used in the calculation.
(1) Rules for determining whether a digit in a measured
value is significant; pages 66-67
assignment: page 68; practice problems 1-2
6. Significant Figures in Calculations
a) In general, a calculated answer cannot be more precise
than the least precise measurment from which it was
calculated.
assignment: page 69; practice problems 3-4
b) the answer to an addition or subtraction calculation
should be rounded to the same number of decimal places.
assignment: page 70; practice problems 5-6
c) In calculations involving multiplication and division,
you need to round the answer to the same number of
signficant figures as the measurement with the least
number of significant figures.
assignment: page 71; practice problems 7-8
page 72; 9-15
II. The International System of Units
A. Measuring with SI Units
1. The International Systerm of Units is a revised version of
the metric system.
2. The five SI base units commonly used by chemists are the
meter, the kilogram, the kelvin, the second, and the mole.
a) Page 73; table 3.1
B. Units and Quantities
1. Common metric units of length include the centimeter,
meter, and the kilometer.
2. Common units of volume include the liter, milliter, cubic
centimeter, and microliter.
3. Common metric units of mass include the kilogram, gram,
milligram, and microgram.
4. Scientist commonly use two equivalent units of
temperature, the degree celsius and the kelvin.
a) Absolute zero (zero kelvin) is the temperature at which
all molecular motion ceases.
5. Energy is the capacity to do work or produce heat.
a) The joule and the calorie are common units of energy.
Assignment: page 79; 18-27
Workbook section 3.2
III.Conversion Problems
A. Conversion Factors
1. A conversion factor is a ratio of equivalent measurements.
2. When a measurement is multiplied by a conversion factor,
the numerical value is generally changed, but the actual size
of the quantity measured remains the same.
B. Dimensional Analysis
1. Dimensional analysis is a way to analyze and solve
problems using the units, or dimensions, of the
measurements.
2. Dimensional analysis provides you with an alternative
approach to problem solving.
Assignment: page 82; 28-29
page 83; 30-31
C. Converting Between Units
1. Problems in which a measurement with one unit is
converted to an eqivalent measurement with another unit are
easily solved using dimensional analysis.
Assignment: page 84; 32-33
page 85; 34-35
page 86; 36-37
page 87; 38-45
workbook section 3.3
IV.
Density
A. Determining Density
1. Density is the ratio of the mass of an object to its volume.
a) Density 
mass
volume
2. Density is an intensive property that depends only on the
composition of a substance, not on the size of the sample.
3. The density of a substance generally decreases as its
temperature increases.
Assignment: page 91; 46-47
page 92; 48-49
page 93; 50-56
workbook section 3.4
Chapter 4
Atomic Structure
I. Defining the Atom
B. Early Models of the Atom
1. An atom is the smallest particle of an element that retains
its identity in a chemical reaction.
2. Democritus’ Atomic Philosophy
a) The Greek philosopher Democritus (460 B.C. – 370
B.C.) was among the first to suggest the existence of
atoms.
b) Democritus believed that atoms were indivisible and
indestructible.
c) Democritus’ ideas lacked experimental support and his
approach was not based on the scientific method.
3. Dalton’s Atomic Theory
a) The modern process of discovery regarding atoms
began with John Dalton (1766 – 1844), an English chemist
and schoolteacher.
b) By using experimental methods, Dalton transformed
Democritus’ ideas on atoms into a scientific theory.
c) Dalton’s atomic theory includes the following ideas.
(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.
(3) Atoms of different elements can physically mix
together or can chemically combine in simple wholenumber 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.
C. Sizing up the Atom
1. The radii of most atoms fall within the range of
to 2x10-10m.
5x10-11m
2. Despite their small size, individual atoms are observable
with instruments such as scanning tunneling microscopes.
Assignment: Page 103; 1-7
workbook section 4.1
V. Structure of the Nuclear Atom
A. Subatomic Particles
1. Electrons
a) Discovered in 1897 by English Physicist J.J. Thomson.
b) Electrons are negatively charged subatomic particles.
c) Thomson’s Experiments
(1) Thomson built a sealed tube with gas at low pressure
and a metal discs at each end.
(2) When an electric current was passed through the tube
a glowing beam or cathode ray would appear.
(3) When charged plates were placed on either side of the
tube, the beam was bent away from the negative plate and
towards the positive plate.
d) In 1916 U.S. Physicist Robert Milikan carried out
experiments to find the quantity of charge carried by an
electron.
(1) An electron carries exactly one unit of negative charge
and its mass is 1/1840 the mass of a hydrogen atom.
2. Protons and Neutrons
a) In 1886 Eugen Goldstein discovered positively charge
subatomic particles called protons.
b) In 1932 James Chadwick confirmed the existence of the
neutron.
(1) Neutrons are subatomic particles with no charge but
with a mass nearly equal to that of a proton.
B. The Atomic Nucleus
1. Rutherford’s Atomic Theory
a) In 1899, Ernest Rutherford discovered that uranium
emits fast-moving particles that have a positive charge.
He named them alpha particles.
b) The gold foil experiment
(1) Rutherford hypothesized that the mass and charge at
any location in the gold would be too small to change the
path of an alpha particle.
(2) When Rutherford’s student, Ernest Marsden, aimed a
narrow beam of alpha particles at the gold the screen
around the gold produced a flash of light when struck by
the alpha particles.
(3) This allowed Marsden to figure out the path of an
alpha particle after it passed though the gold.
(a) Most of the particles passed straight through.
(b) Some particles, about one out of every 2000, were
deflected.
(c) A few particles even bounced straight back.
(4) Rutherford concluded that the positive charge of an
atom is not evenly spread throughout the atom. It is
concentrated in a very small, central area that Rutherford
called the nucleus.
(a) The nucleus is a dense, positively charged mass located
in the center of the atom.
(5) According to Rutherford’s model, all of an atom’s
positive charge is concentrated in its nucleus.
2. In the nuclear atom, the protons and neutrons are located
in the nucleus. The electrons are distributed around the
nucleus and occupy almost all the volume of the atom.
a) page 106; table 4.1
Assignment: Page 108; 8-14
workbook section 4.2
VI.
Distinguishing Among Atoms
A. Atomic Number
1. Elements are different because they contain different
number of protons.
a) The atomic number of an element is the number of
protons in the nucleus of an atom of that element.
B. Mass Number
1. The total number of protons and neutrons in an atom is
called the mass number.
2. The number of neutrons in an atom is the difference
between the mass number and atomic number.
a) number of neutrons = mass number – atomic number
Assignment: page 112; 17-18
page 113; 19-20
C. Atomic Mass
1. An atomic mass unit (amu) is defined as one twelfth of the
mass of a carbon-12 atom.
2. The atomic mass of an element is a weighted average mass
of the atoms in a naturally occurring sample of the element.
a) To calculate the atomic mass of an element, multiply
the mass of each isotope by its natural abundance,
expressed as a decimal, and then add the products.
Assignment: page 116; 21-22
page 117; 23-24
D. The Periodic Table – A Preview
1. A periodic table is an arrangement of the elements in which
the elements are separated in groups based on a set of
repeating properties.
2. A periodic table allows you to easily compare the
properties of one element (or group of elements) to another
element (or group of elements).
a) Each horizontal row of the periodic table is called a
period.
b) Each vertical column of the periodic table is called a
group, or family.
http://www.chemicool.com/
(1) Assignment: page 119; 25-33
workbook section 4.3
Chapter 5
Electrons In Atoms
I. Models of the Atom
E. The Development of Atomic Models
1. Rutherford’s atomic model could not explain the chemical
properties of elements.
F. The Bohr Model
1. Niels Bohr (1885-1962) believed Rutherford’s model needed
improvement.
2. Bohr proposed that an electron is found only in specific
circular paths, or orbits, around the nucleus.
3. The fixed energies an electron can have are called energy
levels.
4. A quantum of energy is the amount of energy required to
move an electron from one energy level to another energy
level.
a) The amount of energy an electron gains or loses in an
atom is not always the same. The higher energy levels are
closer together.
5. The Quantum Mechanical Model
a) Erwin Schrödinger wrote a mathematical equation to
describe the behavior of electrons in an atom.
(1) The quantum mechanical model determines the
allowed energies an electron can have and how likely it is
to find the electron in various locations around the
nucleus.
(2) http://www.phobe.com/s_cat/s_cat.html
6. Atomic Orbitals
a) An atomic orbital is often thought of as a region of
space in which there is a high probability of finding an
electron.
b) Each energy sublevel corresponds to an orbital of a
different shape which describes where the electron is likely
to be found.
c) Read pages 131-132; know tables 5.1 and 5.2
assignment: page 132; 1-7
workbook section 5.1
II. Electron Arrangement in Atoms
G. Electron Configurations
1. The ways in which electrons are arranged in various
orbitals around the nuclei of atoms are called electron
configurations.
2. Three rules – the aufbau principle, the Pauli exclusion
principle. and Hund’s rule – tell you how to find the electron
configurations of atoms.
a) Aufbau principle – electrons occupy the orbitals of
lowest energy first.
b) Pauli exclusion principle – an atomic orbital may
describe at most two electrons.
c) Hund’s rule – electrons occupy orbitals of the same
energy in a way that makes the number of electrons with
the same spin directions large as possible.
assignment: page 135; 8-9
H. Exceptional Electron Configurations
1. If you would write the electron configurations for copper
and chromium you would get
Cr 1s22s22p63s23d44s2
Cu 1s22s22p63s23d94s2
these would be incorrect.
2. The correct electron configurations are as follows:
2
2
Cr 1s 2s 2p63s23d54s1
Cu 1s22s22p63s23d104s1
3. Some actual electron configurations differ from those
assigned using the aufbau principle because half filled
sublevels are not as stable as filled sublevels, but they are
more stable than other configurations.
Assignment: Page 136; 10-13
Workbook section 5.2
III.Physics and the Quantum Mechanical Model
A. Light
1. The amplitude of a wave is the wave’s height froms zero to
the crest.
2. The wavelength, represented by λ, is the distance between
the crests.
3. The frequency, represented by ν, is the number of wave
cycles to pass a given point per unit of time.
a) The units of frequency are usually cycles per second.
The SI unit of cycles per second is called a hertz.
4. The product of frequency and wavelength always equals a
constant ©, the speed of light
a) c=λν
(1) c = 2.998 x 108 m/s
b) The wavelength and frequency of light are inversely
proportional to each other.
5. According to the light model, light consist of
electromagnetic waves.
a) Electromagnetic radiation includes radio waves,
microwaves, infrared waves, visible light. ultraviolet
waves, X-rays, and gamma rays.
b) When sunlight pass through a prism, the different
frequencies separate into a spectrum of colors.
B. Atomic Spectra
1. When atoms absorb energy, electrons move to higher
energy levels, and these electrons lose energy by emitting
light when they return to lower energy levels.
a) The frequencies of light emitted by an element separate
into discrete lines to give the atomic emission spectrum of
the element.
(1) The emission spectrum of each element is like a
person’s fingerprint.
C. An Explanation of Atomic Spectra
1. The lowest possible energy of the electron is its ground
state.
2. The light emitted by an electron moving from a higher to a
lower energy level has a frequency directly proportional to the
energy change of the electron.
3. page 143; figure 5.14
D. Quantum Mechanics
1. Einstein
a) Quanta – tiny bundles of energy.
b) The quanta behave as if they were particles.
c) Light quanta are called photons.
2. De Broglie
a) Wavelike behavior of particles – matter waves.
b) Classical mechanics adequately describes the motions
of bodies much larger than atoms, while quantum
mechanics describes the motions of subatomic particles
and atoms as waves.
3. Heisenberg
a) 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.
Assignment: page 146; 16-21
workbook section 5.3
Chapter 6
The Periodic Table
I. Organizing the Elements
A. Searching For an Organizing Principle
1. Chemists used the properties of elements to sort them into
groups.
a) Dobereiner used the properties of elements to organize
them into triads.
b) Mendeleev arranged the elements in his periodic table
in order of increasing atomic mass.
B. The Periodic Law
1. In the modern periodic table, elements are arranged in
order of increasing atomic number.
2. periodic law: When elements are arranged in order of
increasing atomic number, there is a periodic repetition of
their physical and chemical properties.
C. Metals, Nonmetals, and Metalloids
1. Three classes of elements are metals, nonmetals, and
metalloids.
a) Metals
(1) are good conductors of heat and electric current.
(2) have metallic luster
(3) Malleable – can be hammered into thin sheets
(4) Ductile – can be drawn into wires.
b) Nonmetals
(1) are poor conductors of electricity.
(2) tend to be brittle as solids.
(3) commonly will be found in gas and liquid states.
c) Metalloids
(1) A metalloid will generally have properties that are
similar to those of metals and nonmetals.
(a) Under some conditions, a metalloid may behave like a
metal, Under other conditions, it may behave like a
nonmetal.
(i) For example, pure silicon is a poor conductor
of electric current, like most nonmetals. But if a
small amount of boron is mixed with the silicon,
the mixture is a good conductor of electric
current.
Assignment: page 160; 1-7
workbook section 6.1
II. Classifying the Elements
A. Squares in the Periodic Table
1. The periodic table displays the symbols and names of the
elements along with information about the structure.
2. Background colors are used to distinguish groups of
elements.
a) The Group 1A elements are called alkali metals.
b) The Group 2A elements are called alkaline earth
metals.
c) The nonmetals of Group 7A are called the halogens.
B. Electron Configurations in Groups
1. Elements can be sorted into noble gases, representative
elements transition metals, or inner transition metals based
on their electron configurations.
2. The noble gases are the elements in Group 8A of the
periodic table.
a) Inert gases because they rarely take part in a reaction.
b) The s and p sublevels are completely filled.
3. Elements in Groups 1A through 7A are often referred to as
representative elements because they display a wide range of
physical and chemical properties.
a) Some are metals, some are nonmetals, and some are
metalloids.
b) Most are solids at room temperature.
c) for any representative element, its group number
equals the number of electrons in the highest occupied
energy level.
C. Transition Elements
1. There are two types of transition elements – transition
metals and inner transition metals.
a) In atoms of a transition metal, the highest occupied s
sublevel and nearby d sublevel contain electrons.
(1) Commonly called d – block elements.
b) In atoms of an inner transition metal, the highest
occupied s sublevel and nearby f sublevel generally
contain electrons.
(1) commonly called f block elements.
2. Blocks of elements
a) page 166; figure 6.12
Assignment: page 167; 8-15
workbook section 6.2
III.Periodic Trends
A. Trends in Atomic Size
1. The atomic radius is one half of the distance between the
nuclei of two atoms of the same element when the atoms are
joined.
2. In general atomic size increases from top to bottom within
a group and decreases from left to right across a period.
3. Group trends in Atomic Size
a) As the atomic number increases within a group, the
charge on the nucleus increases and the number of
occupied energy levels increases.
(1) The increase in positive charge draws the electrons
closer to the nucleus.
(2) The increase in the number of occupied orbitals shields
electrons in the highest occupied energy level from the
attraction of protons in the nucleus.
(a) the shielding effect is greater than the effect of the
increase in nuclear charge therefore the atomic size
increases.
4. Periodic Trends in Atomic Size
a) In general, atomic size decreases across a period from
left to right.
(1) Each element has one more proton and one more
electron than the preceding element.
(2) The shielding effect is constant for all the elements in a
period.
(3) The increasing nuclear charge pulls the electrons in
the highest occupied energy level closer to the nucleus and
the atomic size decreases.
B. Ions
1. Positive and negative ions form when electrons are
transferred between atoms.
a) An ion with a positive charge is called a cation.
b) An ion with a negative charge is called an anion.
2. Trends in Ionization Energy
a) The energy required to remove an electron from an
atom is called ionization energy.
(1) First ionization energy tends to decrease from top to
bottom within a group and increase from left to right
across a period.
(a) page 173; figure 6.1
b) Group Trends in Ionization Energy
(1) As the size of the atom increases, nuclear charge has a
smaller effect on the electrons in the highest occupied
energy level.
(2) less energy is required to remove an electron from this
energy level and the first ionization energy is lower.
c) Periodic Trends in Ionization Energy
(1) The nuclear charge increases across the period, but the
shielding effect remains constant.
(2) There is an increase in the attraction of the nucleus for
an electron. Therefore it takes more energy to remove an
electron from an atom.
3. Trends in Ionic Size
a) Cations are always smaller than the atoms from which
they form.
b) Anions are always larger than the atoms from which
they form.
c) page 176; figure 6.21
4. Trends in Electronegativity
a) Electronegativity is the ability of an atom of an element
to attract electrons when the atom is in a compound.
b) 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.
5. Summary of Trends
a) the trends that exist among these properties can be
explained by variations in atomic structure.
b) page 178; figure6.22
Assignment: Page 178; 16-23
workbook section 6.3
Chapter 7
Ionic and Metallic Bonding
I. Ions
A. Valence electrons
1. Valence electrons are the electrons in the highest occupied
energy level of an element’s atoms.
2. To find the number of valence electrons in an atom of a
representative element, simply look at the group number.
B. Octet Rule
1. 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
electrons or to share electrons with another nonmetallic
element to achieve a complete octet.
a) Gilbert Lewis (1916) – octet rule
C. Formation of Cations
1. An atom’s loss of valence electrons produces a cation, or a
positively charged ion.
a) read pages 188-190
D. Formation of Anions
1. the gain of negatively charged electrons by a neutral atom
produces an anion.
a) The ions that are produced when atoms of chlorine and
other halogens gain electrons are called halide ions.
b) read pages 191 – 192
assignment: page 193; practice problems 1-2
section assessment 3-11
workbook section 7.1
II. Ionic Bonds and Ionic Compounds
A. Formation of Ionic Compounds
1. Compounds composed of cations and anions are called
ionic compounds.
a) Although they are composed of ions, ionic compounds
are electrically neutral.
2. Ionic bonds
a) The electrostatic forces that hold ions together in ionic
compounds are called ionic bonds.
3. Formula Units
a) A chemical formula shows the kinds and numbers of
atoms in the smallest representative unit of a substance.
b) A formula unit is the lowest whole-number ratio of ions
in an ionic compound.
Assignment: Page 196; 12-13
B. Properties of Ionic Compounds
1. Most ionic compounds are crystalline solids at room
temperature.
2. Ionic compounds generally have high melting points.
3. The coordination number of an ion is the number of ions
of opposite charge that surround the ion in a crystal.
4. Ionic compounds can conduct an electric current when
melted or dissolved.
Assignment: page 199; 14-22
workbook section 7.2
III.Bonding in Metals
A. Metallic Bonds and Metallic Properties
1. The valence electrons of metal atoms can be modeled as a
sea of electrons.
2. Metallic bonds consist of the attraction of the free-floating
valence electrons for the positively charged metal ions.
a) page 201; figure 7.12
B. Crystalline Structure of Metals
1. Metal atoms are arranged in very compact and orderly
patterns.
a) page 202; figure 7.14
C. Alloys
1. Alloys are mixtures composed of two or more elements, at
least one of which is a metal.
2. Alloys are important because their properties are often
superior to those of their component elements.
Assignment: page 203; 23-29
workbook section 7.3
Chapter 8
Covalent Bonding
I. Molecular Compounds
A. Molecules and Molecular Compounds
1. The atoms held together by sharing electrons are joined by
a covalent bond.
2. A molecule is a neutral group of atoms joined together by
covalent bonds.
a) A diatomic molecule is a molecule consisting of two
atoms.
(1) Hydrogen, Oxygen, Nitrogen, Fluorine, Chlorine,
Bromine, and Iodine commonly form diatomic molecules.
3. A compound composed of molecules is called a molecular
compound.
a) Molecular compounds tend to have relatively lower
melting and boiling points than ionic compounds.
B. Molecular Formulas
1. A molecular formula is the chemical formula of a molecular
compound.
a) A molecular formula shows how many atoms of each
element a molecule contains.
Assignment: page 216; 1-6
workbook section 8.1
II. The Nature of Covalent Bonding
A. The Octet Rule in Covalent Bonding
1. In forming covalent bonds, electron sharing usually occurs
so that atoms attain the electron configurations of the noble
gases.
B. Single Covalent Bonds
1. Two atoms held together by sharing a pair of electrons are
joined by a single covalent bond.
a) An electron dot structure such as H:H represents the
shared pair of electrons of the covalent bond by two dots.
b) A structural formula represents the covalent bonds by
dashes and shows the arrangement of covalently bonded
atoms.
c) A pair of valence electrons that is not shared between
atoms is called an unshared pair, also known as a lone pair
or a nonbonding pair.
2. read pages 218- 220
Assignment page 220 7-8
C. Double and Triple Covalent Bonds
1. Atoms form double or triple covalent bonds if they can
attain a noble gas structure by sharing two pairs of three pairs
of electrons.
a) A bond that involves two shared pairs of electrons is a
double covalent bond.
b) A bond formed by sharing three pairs of electrons is a
triple covalent bond.
D. Coordinate Covalent Bonds
1. A coordinate covalent bond is a covalent bond in which
one atom contributes both bonding electrons.
a) In a coordinate covalent bond, the shared electron pair
comes from one of the bonding atoms.
(1) A polyatomic ion, such as NH4+, is a tightly bound
group of atoms that has a positive or negative charge and
behaves as a unit.
(2) Page 224; table 8.2
Assignment: page 225; 9-11
E. Bond Dissociation Energies
1. The energy required to break the bond between two
covalently bonded atoms is known as the bond dissociation
energy.
a) A large bond dissociation energy corresponds to a
strong covalent bond.
F. Resonance
1. The actual bonding of oxygen atoms in ozone is a hybrid,
or mixture, of the extremes represented by the resonance
forms.
2. A resonance structure is a structure that occurs when it is
possible to draw two or more valid electron dot structures that
have the same number of electrons pairs for a molecule or
ion.
G. Exceptions to the Octet Rule
1. The octet rule cannot be satisfied in molecules whose total
number of valence electrons is an odd number. there are also
molecules in which an atom has fewer, or more, than a
complete octet of valence electrons.
Assignment: page 229; 13-22
workbook section 8.2
III.Bonding Theories
A. Molecular Orbitals
1. When two atoms combine, this model assumes that their
atomic orbitals overlap to produce molecular orbitals, or
orbitals that apply to the entire molecule.
2. Just as an atomic orbital belongs to a particular atom, a
molecular orbital belongs to a molecule as a whole.
3. A molecular orbital that can be occupied by two electrons
of a covalent bond is called a bonding orbital.
a) When two atomic orbitals combine to form a molecular
orbital that is symmetrical around the axis connecting two
atomic nuclei, a sigma bond is formed.
(1) the symbol for this bond is the Greek letter sigma (σ).
(2) Page 230; figure 8.13
b) In a pi bond (symbolized by the Greek letter π), the
bonding electrons are most likely to be found in sausage
shaped regions above and below the bond axis of the
bonded atoms.
(1) Page 231; figure 8.15
B. VSEPR Theory
1. Valence-Shell Electron-Pair Repulsion Theory
a) According to VSEPR theory, the repulsion between
electron pairs causes molecular shapes to adjust so that
the valence-electron pairs stay as far apart as possible.
(1) Page 232; figure 8.16
(2) page 233; figure 8.17
(3) page 233; figure 8.18
C. Hybrid Orbitals
1. Orbital hybridization provides information about both
molecular bonding and molecular shape.
a) In hybridization, several atomic orbitals mix to form
the same total number of equivalent hybrid orbitals.
(1) Page 234; figure 8.19
(2) page 235; figure 8.20
2. Hybridization involving double and triple bonds.
a) read pages 235-236
assignment: page 236; 23-29
workbook section 8.3
IV.
Polar Bonds and Molecules
A. Bond Polarity
1. When the atoms in the bond pull equally the bonding
electrons are share equally, and the bond is a nonpolar
covalent bond.
2. A polar covalent bond, known also as a polar bond, is a
covalent bond between atoms in which the electrons are
shared unequally.
a) The more electronegative an atom attracts electrons
more strongly and gains a slightly negative charge. The
less electronegative atom has a slightly positive charge.
(1) Page 238; table 8.3(know for test)
assignment: page 239; 30-31
B. Polar Molecules
1. In a polar molecule, one end of the molecule is slightly
negative and the other end is slightly positive.
2. A molecule that has two poles is called a dipolar molecule
or a dipole.
a) When polar molecules are placed between oppositely
charged plates, they tend to become oriented with respect
to the positive and negative plates.
C. Attraction Between Molecules
1. Intermolecular attractions are weaker than either ionic or
covalent bonds.
a) The two weakest attractions between molecules are
collectively called van der Waals forces.
(1) Dipole interactions occur when polar molecules are
attracted to one another.
(2) Dispersion forces, the weakest of all molecular
interactions, are caused by the motion of electrons.
b) Hydrogen bonds
(1) Hydrogen bonds are attractive forces in which a
hydrogen covalently bonded to a very electronegative atom
is also weakly bonded to and unshared electron pair of
another electronegative atom.
c) Intermolecular Attractions and Molecular Properties
(1) Melting a network solid would require breaking
covalent bonds throughout the solid.
(a) Page 244; table 8.4
assignment: page 244; 32-38
workbook section 8.4
Chapter 10
Chemical Quantities
I. The Mole: A Measurement of Matter
A. Measuring Matter
1. You often measure the amount of something by one of
three different methods – by count, by mass, and by
volume.
Assignment: page 289; 1-2
B. What is a Mole?
1. A mole of a substance is 6.02 x 1023 representative
particles of that substance and is the SI unit for measuring
the amount of a substance.
a) The number of representative particles in a mole,
6.02 x 1023, is called Avogadro’s number.
b) The term representative particle refers to the
species present in a substance: usually atoms,
molecules, or formula units.
c) A mole of any substance contains Avogadro’s
number of representative particles, or 6.02 x 1023
representative particles.
Assignment: page 291; 3-4
page 292; 5-6
C. The Mass of a Mole of an Element
1. The atomic mass of an element expressed in grams is the
mass of a mole of the element.
a) The mass of a mole of an element is its molar mass.
2. to calculate the molar mass of a compound, find the
number of grams of each element in one mole of the
compound. Then add the masses of the elements in the
compound.
Assignment: page 296; 7-8
page 296; 9-15
workbook section 10.1
II. Mole-Mass and Mole-Volume Relationship
A. The Mole-Mass Relationship
1. Use the molar mass of an element or compound to
convert between the mass of a substance and the moles of a
substance.
Assignment: page 298; 16-17
page 299; 18-19
B. The Mole-Volume Relationship
1. Avogadro’s hypothesis states that equal volumes of gases
at the same temperature and pressure contain equal numbers
of particles.
2. Standard temperature and pressure (STP) means a
temperature of 0˚C and a pressure of 101.3 kPa, or 1
atmosphere (atm).
3. At STP, 1 mol or 6.02 x 1023 representative particles, of
any gas occupies a volume of 22.4 L.
Assignment: page 301; 20-21
page 302; 22-23
page 303; 24-31
workbook section 10.2
III. Percent Composition and Chemical Formulas
A. The Percent Composition of a Compound
1. The relative amounts of the elements in a compound are
expressed as the percent composition.
2. 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%.
Assignment: page 306; 32-33
page 307; 34-35
B. Empirical Formulas
1. The basic ratio, called the empirical formula, gives the
lowest whole-number ratio of the atoms of the elements in a
compound.
2. The empirical formula of a compound shows the smallest
whole-number ratio of the atoms in the compound.
Assignment: page 310; 36-37
C. Molecular Formulas
1. The molecular formula of a compound is either the same
as its experimentally determined empirical formula, or it is
a simple whole-number multiple of its empirical formula.
Assignment: page 312; 38-39
page 312; 40-46
workbook section 10.3
Chemical Reactions
Chapter 11
I. Describing Chemical Reactions
A. Writing Chemical Equations
1. Word Equations
a) To write a word equation, write the names of the
reactants to the left of the arrow separated by plus
signs; write the names of the products to the right of
the arrow, also separated by plus signs.
(1) Example: “Hydrogen peroxide decomposes
to form water and oxygen gas.”
hydrogen peroxide → water + oxygen
2. Chemical Equations
a) A chemical equation is a representation of a
chemical reaction; the formulas of the reactants are
connected by an arrow with the formulas of the
products.
b) A skeleton equation is a chemical equation that
does not indicate the relative amounts of the reactants
and products.
c) page 323; figure 11.1
B. Balancing Chemical Equations
1. To write a balanced chemical equation, first write the
skeleton equation. Then use coefficients to balance the
equation so that it obeys the law of conservation of mass.
a) Coefficients are small whole numbers that are
placed in front of the formulas in an equation in order
to balance it.
Assignment: page 324; 1-2
page 327; 3-4
page 328; 5-6
page 329; 7-12
workbook section 11.1
II. Types of Chemical Reactions
A. Classifying Reactions
1. The five general types of reaction are combination,
decomposition, single-replacement, double replacement,
and combustion.
a) A combination reaction is a chemical change in
which two or more substances react to form a single
new substance.
(1) Magnesium metal and oxygen gas combine
to form the compound magnesium oxide.
2Mg(s) + O2(g) → 2MgO(s)
Assignment: page 331; 13-14
b) A decomposition reaction is a chemical change in
which a single compound breaks down into two or
more simpler products.
(1) When mercury (II) oxide is heated, it
decomposes into mercury and oxygen.
2HgO(s) → 2Hg(l) + O2(g)
Assignment: page 332: 15-16
c) Single-Replacement Reactions
(1) A single-replacement reaction is a
chemical change in which one element replaces
a second element in a compound.
(a)Whether one metal will displace another
metal from a compound depends upon the
relative reactivities of the two metals.
(b) The activity series of metals lists
metals in order of decreasing reactivity.
(i) page 333; table 11.2
assignment: page 334; 17
d) Double-Replacement Reactions
(1) A double-replacement reaction is a
chemical change involving an exchange of
positive ions between two compounds.
(a)For a double-replacement reaction to
occur, one of the following is usually true.
(i) One of the products is only
slightly soluble and precipitates from
solution.
(a) Na2S(aq) + H2SO4(aq) → CdS(s) + 2NaNO3(aq)
(ii)
One of the products is a gas.
(a)
2NaCN(aq) + H2SO4(aq) → 2HCN(g) + Na2SO4(aq)
(iii) One product is a molecular
compound such as water.
(a) Ca(OH)2(aq) + 2HCl(aq) → CaCl2(aq) + 2H2O(l)
Assignment: page 335; 18-19
e) A combustion reaction is a chemical change in
which an element or compound reacts with oxygen.
(1) The complete combustion of a
hydrocarbon produces carbon dioxide and water.
(2) The complete combustion of gasoline in a
car engine is shown by the equation:
(a)2C8H18(l) + 25O2(g) → 16CO2(g) +
18H2O(l)
(b) The reactions between oxygen and
some elements other than carbon are also
examples of combustion reactions. For
example, both magnesium and sulfur will
burn in the presence of oxygen.
Assignment: page 337; 20-21
B. Predicting the Products of a Chemical Reaction
1. The number of elements and/or compounds reacting is a
good indicator of possible reaction type and thus possible
products.
a) read pages 338-339
Assignment: page 339; 22-27
workbook section 11.2
III. Reactions in Aqueous Solution
A. Net Ionic Equations
1. A complete ionic equation is an equation that shows
dissolved ionic compounds as dissociated free ions.
a) Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) →
AgCl(s) + Na+(aq) + NO3-(aq)
b) An ion that appears on both sides of an equation
and is not directly involved in the reaction is called a
spectator ion.
c) The net ionic equation is an equation for a reaction
in solution that shows only those particles that are
directly involved in the chemical change.
d) A net ionic equations show only those particles
involved in the reaction and is balanced with respect
to both mass and charge.
Assignment: Page 343; 28-29
2. Predicting the Formation of a Precipitate
a) You can predict the formation of a precipitate by
using the general rules for solubility of ionic
compounds.
(1) Page 334; table 11.3
Assignment: page 344; 30-35
workbook section 11.3