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Wester Midterm Review
Info: This should cover everything on the study guide. Copy the notes from
previous note packets to fill this out. Any unnecessary notes have been taken out.
* Replacement reactions are still in the notes even though we didn’t learn them as
they are included in the study guide
Chapter 1: Matter and Change
MAKE STATES OF MATTER ITS OWN SECTION RIGHT BELOW THE MATTER TREE
DIAGRAM
Topics:
 Definition of chemistry
 Classification of matter
 Properties of matter
 Energy
Definition of chemistry
Chemistry is a physical science, and is the study of…



chemical - any substance that has a definite composition (gold, glucose, water, CO)
Classification of matter
- the amount of three-dimensional space an object occupies.
- a measure of the amount of matter.
- anything that has mass and takes up space.
Basic Building Blocks of matter
pure substance – has a fixed composition; every sample has exactly the same characteristics and
composition
element – a pure substance that cannot be broken down into simpler, stable substances and is
made of one type of atom.
atom – the smallest unit of an element that maintains the chemical identity of that element.
compound – a substance that can be broken down into simple stable substances. Each compound
is made from the atoms of two or more elements that are chemically bonded.
molecule – the smallest unit of a compound that maintains the chemical identity of that
compound. (ionic ~ crystal/formula unit)
mixture – a blend of two or more kinds of matter, each of which retains its own identity and
properties
homogeneous mixture – (solutions) uniform in composition (salt-water solution)
heterogeneous mixture – not uniform throughout (clay-water mixture)
Separation techniques for mixtures:

___________________ - one substance caught, one goes through

___________________ - one substance vaporized, collected, and condensed

___________________ - one substance travels more slowly along a surface (paper)
Properties of matter
Physical Properties and Physical Changes



a characteristic that can be observed or measured without changing the identity of the
substance, like melting point and boiling point.
a change in a substance that does not involve a change in the identity of the substance,
like cutting, melting, and boiling.
mass, weight, length, volume, density, solubility (dissolving), state/phase, brittleness,
flexibility, conductivity, energy, specific heat
States of matter
Description
particles touch each other
Solid
Liquid
Gas
definite volume
definite shape
particles vibrate
particles have translational motion
(place-to-place)
particle arrangement
____________________ - a high-temperature physical state of matter in which atoms lose most
of their electrons, particles that make up atoms.
Chemical Properties and Chemical Changes

relates to a substance’s ability to undergo changes that turn it into different substances

chemical change = chemical reaction
carbon

+ oxygen

carbon dioxide
possible evidence of a chemical change: color, temperature, odor, precipitate, bubbles
Demonstrations of physical and chemical changes:
Description
Observations
Physical or
chemical?
Energy

Energy is involved when ___________________ or ___________________ changes
occur.

Forms of energy:

Law of Conservation of Energy - energy can be
in a change, it is not destroyed or created
or
___________________ – average kinetic energy of the particles (motion)
___________________ – energy transferred between samples of matter because of a difference
in their temperatures
Chapter 2: Measurements and Calculations
Topics:
1. Units of measurement
2. Significant figures
3. Scientific notation
4. Useful chemistry measurements
5. Analyzing lab data
Units of measurement
Le Système International d’Unités (SI) is a single set of measurements agreed upon by scientists
around the world. It is similar to the metric system, but the “base unit” may be different.
SI units of measurement (Know the first 5)
Measurement
SI unit
Abbrev.
Additional info
length
meter
m
mass
kilogram
kg
time
second
s
temperature
Kelvin
K
Not C
amount of substance
mole
mol
a way of counting particles
electric current
ampere
A
Do not need to know.
luminous intensity
candela
cd
Do not need to know.
Not grams
Significant figures
Significant figures in a measurement are all the digits known with certainty plus the estimated
(uncertain) digit


Tells you how many digits to round to in your answer.
Tells you how precise your measurement tool is.
Atlantic-Pacific Method
Does the number have a decimal point?
Pacific Ocean
Atlantic Ocean
Present



Absent
Start counting digits on the correct side of the number according to the A-P guide.
Move to the first non-zero digit.
Count that digit and every digit that comes after it, including any zeros.
Examples:
1.
40.0
mL = _____ significant figures
2.
982,500
3.
20.540
4.
0.06670
5.
1.34
cm = _____ significant figures
mL = _____ significant figures
kg = _____ significant figures
m = _____ significant figures
Calculations and significant figures

When adding or subtracting decimals, the answer must have the same number of digits
to the right of the decimal point as there are in the measurement having the fewest digits
to the right of the decimal point. (decimal places)
Example: 5.44 m - 2.6103 m =

For multiplication or division, the answer can have no more significant figures than are
in the measurement with the fewest number of significant figures. (digits)
Example: 2.4 g/mL  15.82 mL =
Combining orders of operation

Follow order of operations, rounding the answer for each type of operation (add-sub or
mult-div) along the way. (Depending on the textbook, you will get different instructions
about this, but for this class, follow this method.)
Example: Find the density of a sample of an object that has a mass of 23.56 g. When placed in a
graduated cylinder of 32.0 mL of water, the water level rises to 37.5 mL.
Scientific notation

Use the 2nd function EE combination on your graphing calculator instead of parentheses.
Examples:
1. 856,000 km = ___________________
2. 350. g = ___________________
3. 0.000433 g = ___________________
Useful chemistry measurements
mass – is a measure of the quantity of matter.


The SI standard unit for mass is the kilogram, however in class, grams (g), is used.
Mass does not depend on gravity.
weight – (Derived, not SI) a measure of the gravitational pull on matter (Newtons, N).
volume – (Derived, not SI) the amount of space occupied by an object.




Multiply 3 lengths (length x width x height)
The cubic centimeter, cm3, works well for solids.
The milliliter (mL) or liter (L) works well for liquids.
Displacement method: volume of object = final volume – initial volume
temperature – measure of the average kinetic energy of the particles in a sample


The SI standard unit for temperature is Kelvin, however lab thermometers are in C.
K = C + 273
density – (Derived, not SI) the ratio of mass to volume, or mass divided by volume.
density = mass
volume
D=m
V
Common units: g/mL or g/cm3

Density is a characteristic physical property of a substance that is useful for identification
Analyzing lab data
Accuracy & Precision
Accuracy refers to the closeness of measurements to the correct or accepted value of the quantity
measured. (correct)
Precision refers to the closeness of a set of measurements of the same quantity made in the same
way. (repeatable)
Interpreting data
average – add up the values and divide by the total number; round the answer to the correct
number of significant figures
experimental value theoretical value -
percent yield =
experimental value
theoretical value
x 100
percent error = (experimental value – theoretical value) x 100
theoretical value


positive = too high
negative = too low
Note: Depending on what you need to know about your data, you could make these calculations
with data from each individual trial or you could average your data first and then make these
calculations.
Example: A chemical reaction is run that theoretically releases 755 mL of gas at room
temperature. A lab group repeats the experiment three times and collects the following volumes
of gas: 780. mL, 730. mL, and 754 mL.
1. What is the average volume of gas produced?
2. Is their collection of data accurate according to the average?
3. Is their collection of data precise?
4. Calculate the percent yield based on the average.
5. Calculate the percent error based on the average.
6. Is the lab group’s average volume too high or too low compared with the theoretical
amount of gas that should have been produced?
Chapter 3: Atoms: The Building Blocks of Matter
Topics:
1. Foundations of atomic theory
2. Development of the modern atomic theory
3. The structure of the atom
4. Counting subatomic particles
5. The mole
Foundations of atomic theory
________________________________________________ - mass is neither created nor
destroyed during ordinary chemical reactions or physical changes
S + O2  SO2
H2 + O2 
H2O
________________________________________________ - a chemical compound contains the
same elements in exactly the same proportions by mass regardless of the size of the sample or
source of the compound
________________________________________________ - if two or more different
compounds are composed of the same two elements, then the ratio of the masses of the second
element combined with a certain mass of the first element is always a ratio of small whole
numbers
Development of the Modern Atomic Theory
Dalton’s Atomic Theory
All matter is composed of extremely small
particles called atoms.
Atoms of a given element are identical in size,
mass, and other properties; atoms of different
elements differ in size, mass, and other propertie
Atoms cannot be subdivided, created, or
destroyed.
Modern Atomic Theory
Atoms of different elements combine in simple
whole-number ratios to form chemical
compounds.
In chemical reactions, atoms are combined,
separated, or rearranged.
Discovery of the Electron and Nucleus
Scientist
J.J. Thomson
Time
period
late 1800s
Contribution
Sketch of experiment
Robert Millikan
late 1800s
oil drop
Ernest Rutherford
early 1900s
gold foil
cathode ray tube
The structure of the atom
1. An atom is the smallest particle of an element that retains the properties of that element.
2. The nucleus is a small, dense region located at the center of an atom.
3. The nucleus is made up of at least one positively charged particle called a proton and
usually one or more neutral particles called neutrons.
4. Short-range nuclear forces hold the protons and neutrons together.
5. Surrounding the nucleus is a region with negatively charged particles called electrons.
6. A proton has a positive charge equal in magnitude to the negative charge of an electron.
7. Atoms are electrically neutral because they contain equal numbers of protons and
electrons.
8. A neutron is electrically neutral.
9. Atoms of different elements differ in their number of protons and therefore in the amount
of positive charge they possess.
10. The number of protons determines that atom’s identity.
11. Protons, neutrons, and electrons are often referred to as subatomic particles.
Counting subatomic particles
atomic number (Z) =
mass number =
isotopes =

The isotopes of a particular element all have the same number of protons and electrons
but different numbers of neutrons.

Most of the elements consist of mixtures of isotopes.
nuclide 
notations: uranium-235 or
235
92
U

Example: How many protons and neutrons are in each nuclide?
12
6
7
3
C =
24
12


1
1
Mg =
Li =
H =

Counting subatomic particles

protons (p+) =
neutrons (n0) =
mass number =
electrons (e-) =
Name
Symbol
iron
40
20
Ca

Atomic
number (Z)
Mass
number
(A)
56
Protons
Neutrons
28
31
40
Electrons
argon-40
18
Atomic mass
relative atomic mass – the masses of elements are reported relative to the mass of carbon, which
has been arbitrarily assigned as the standard
atomic mass unit = amu = exactly 1/12 the mass of a carbon-12 atom
average atomic mass - weighted average of the atomic masses of the naturally occurring isotopes
of an element

The average atomic mass of an element depends on mass and relative abundance.
Example 1: Copper consists of 69.15% copper-63, which has an atomic mass of 62.929 601 amu,
and 30.85% copper-65, which has an atomic mass of 64.927 794 amu.
The Mole
mole - amount of a substance that contains as many particles as there are atoms in exactly 12 g of
carbon-12.
Avogadro’s number = 6.02  1023 = number of particles in exactly one mole of a pure substance.
molar mass - mass of one mole of a pure substance (g/mol) (= atomic mass for an element)
Practice:
1. How many moles of silver, Ag, are in 3.01  1023 atoms of silver?
2. What is the mass in grams of 3.50 mol of the element copper, Cu?
3. What is the mass in grams of 1.20  108 atoms of copper, Cu?
Chapter 4: Arrangement of Electrons in Atoms
Topics:
1. Properties of light
2. Line-emission spectrum
3. Bohr model
4. Quantum model of the atom
5. Atomic orbitals and quantum numbers
6. Electron configurations and orbital diagrams
7. Interpreting electron configurations
Properties of light
__________________________________________ – a form of energy that exhibits wavelike
behavior as it travels through space; moves at c = 3.00 x 108 m/s (light)
____________________ – () the distance between corresponding points on adjacent waves (m,
nm, cm); __________ (n) = 10-9
____________________ – () the number of waves that pass a given point in a specific time
(usually one second) (waves/second, hertz – Hz)
Equation connecting wavelength to frequency:
c = 
high frequency  short wavelength
low frequency  long wavelength
__________________________________________ – electromagnetic radiation causes ejection
of electrons from the surface of a metal

For a given metal, only specific frequencies eject electrons.

Max Planck proposed that energy is emitted in specific packets called quanta.

Albert Einstein proposed that light could be thought of as a stream of particles: photons.

Matter absorbs energy only in whole numbers of photons.
Planck’s constant = h = 6.626 x 10-34 Js
E  h 
hc

Line-emission spectrum



Since late 1800s, scientists knew that the light emitted by energized hydrogen was not
continuous (all the colors), but limited to a few, specific frequencies.
Johan Balmer and Johannes Rydberg expressed the line spectrum for hydrogen
mathematically.
line-emission spectrum – a series of specific wavelengths of light emitted from an energized
sample of atoms (also called an atomic spectrum)
ground state – the lowest energy state of the
electrons in an atom
excited state – the electrons of an atom have a
higher potential energy than in the ground state
The energy diagram for hydrogen shows possible
energy transitions for its electron as it emits
energy from a higher energy level to a lower
energy level.
Rydberg equation:  
RH  1
1 



2
2 
h 
n f n i 
Bohr model
 interpreted the data as electrons traveling in one of a limited number of specific
Niels Bohr
orbits, or energy levels.



Since the electron in hydrogen would only emit specific frequencies of light, only
specific amounts of energy were available to it.
The Bohr model represents the basic idea of principal energy levels very well and works
for hydrogen
The Bohr model does not work for multi-electron atoms.
Quantum model of the atom
p  mv

In classical mechanics, momentum = mass x velocity:

Louis de Broglie built on Einstein’s work and proposed that an electron behaves as a
particle and a wave.
deBroglie equations:

momentum  p 
h


wavelength   
h
mv
quantum theory – describes mathematically the wave properties of electrons and other
very small particles


Werner Heisenberg explained the difficulty in locating an electron that is both a particle
and a wave: the photons (light) we use to detect particles have about the same energy as
an electron so they knock an electron off its course.

Heisenberg uncertainty principle – it is impossible to determine simultaneously both the
position and velocity of an electron or any other particle.

Erwin Schrodinger developed equations to describe the most probable location of an
electron: H n 

1  me4 


n 2 8 02 h 2  n
wave function – mathematical equation that gives the probability of finding an electron
(90% of the time)

Matter AND radiation both have properties of waves AND particles
Atomic orbitals and quantum numbers
When solved, the Schrodinger equation provides quantum numbers that describe the electron and
its orbital.
orbital – 3-D region around the nucleus that indicates the probable location of an electron.
quantum numbers – specify the properties of an atomic orbital and the electrons in the orbital
Information about quantum numbers
Name of quantum
number
Principal quantum
number
Possible values
Description
Provides a general sense of an electron’s
overall energy
Also:
 principal energy level
 shell
Higher values of n = further distance
from the nucleus = more energy
Angular momentum
quantum number
Indicates the shape of the orbital: s, p, d, f
Also:
 sublevel
 subshell
Magnetic quantum
number
Indicates the orientation of an orbital
around the nucleus.
Spin quantum number
Accounts for the two types of magnetic
fields created by electrons.

Using n + l describes the orbital.

Adding m and ms provides information about individual electrons.

Each orbital can hold 2 electrons.
Quantum numbers and number of electrons
Principal
quantum number
(n)
Sublevels in
principal energy
level
(l = n-1)
Number of
orbitals per
sublevel
(-l thru +l)
Number of
electrons per
sublevel
Number of
electrons in the
principal energy
level
1
2
3
4
Electron configurations and orbital diagrams
An atom’s electron configuration shows the arrangement of its electrons, and there are three
types of notation:
1. electron configuration notation – shows principal energy level, sublevel, and number of
electrons in the sublevel (eg. 1s22s22p6)
2. orbital notation – shows all of the orbitals of a sublevel and the pairing of the electrons
(uses arrows instead of superscripts)
3. noble gas configuration – abbreviates the electron configuration by using the prior noble
gas ([Ar]3d54s1)
Rules of Electron Configurations and Orbital Diagrams
Rule 1:

Electrons fill orbitals of lower energy first.

The “s” sublevel is always the lowest energy sublevel within a principal energy level.
Rule 2:

Each orbital describes at most 2 electrons.

Electrons must have opposite spin (shown with up or down arrows).
Rule 3:

When electrons occupy a sublevel, one electron enters each orbital until all the orbitals
contain one electron with parallel spins, then they begin to pair up.
Aufbau diagram:
Practice electron configurations:
hydrogen, H (1 e-)
beryllium, Be (4 e-)
neon, Ne (10 e-)
Interpreting electron configurations

To identify an element by its electron configuration, add the superscripts, which equals
the number of electrons. This number equals the atomic number for an atom.
Categorizing electrons:
______________________________ – electrons in the highest principal energy level; the
outermost electrons; involved in bonding.
 Using the electron configuration, add up the superscripts for the highest principal energy
level only (1-8)
______________________________ – inner electrons held close to the nucleus; match the noble
gas configuration; not usually involved in bonding
Practice: How many valence electrons do these elements have?
hydrogen
helium
carbon
chlorine
Connecting electron configuration to the periodic table
Practice: Identify these elements:
1. Electron configuration ends in 3s2.
_____________________________________
2. Electron configuration ends in 5p5.
_____________________________________
3. Electron configuration ends in 2p6.
_____________________________________
Limits of the Aufbau diagram






The Aufbau diagram is a pattern that scientists came up with to model observations.
The Aufbau diagram works best for atoms up to calcium (Z = 20) and for many after that.
Transition metals have d orbitals that are very close to the valence s orbital, so for some
of them the Aufbau diagram does not correctly describe their arrangement.
Electrons want the lowest energy possible with the most stability.
They are affected by repulsions of nearby electrons and shielding.
Examples:
chromium (Cr)
[Ar]3d54s1
copper (Cu)
[Ar]3d104s1
molybdenum (Mo) [Kr]4d55s1
silver (Ag)
[Kr]4d105s1
gold (Au)
[Xe]4f145d106s1
Chapter 5: The Periodic Law
Topics:
1. History of the Periodic Table
2. Sections of the Periodic Table
3. Periodic trends
History of the Periodic Table
Dmitri Mendeleev

Noticed that when the elements were arranged in order of
increasing ___________________________, similarities in their
properties appeared at regular intervals.

Created a table in which elements with similar properties
were grouped together—a periodic table of the elements.

Predicted the existence and properties of
___________________ to fill empty spaces.
Henry Moseley

Discovered that the elements fit into patterns better when they
were arranged according to _____________________________, rather
than atomic mass.
_____________________________ - states that the physical and chemical properties of the
elements are periodic, or systematic, functions of their atomic numbers.
The Modern Periodic Table

An arrangement of the elements by _____________________________ so that elements
with similar properties fall in the same column, or group.

Repeating patterns are referred to as _____________________________.
Sections of the Periodic Table

_______________ = column = up and down = family = similar chemical properties

_______________ = row = across = gradual changes in properties (trends)
Sections of the Periodic Table
Metals, nonmetals, and metalloids
s-block elements

Group 1 - _____________________________ – highly reactive

Group 2 - _____________________________ – a little less reactive, but do not exist
freely in nature

_____________________________ (1s1) does not share the same properties as the
elements of Group 1.

_____________________________ (2s2) is part of Group 18 because its highest
occupied energy level is filled by two electrons, which gives it helium chemical stability.
p-block elements

All of the elements of Groups 13–18 except _____________________________.

p-block and s-block elements are called the ___________________________________.

The properties of elements of the p block vary greatly.

________________________ – to the right, plus hydrogen

________________________ – boron, silicon, germanium, arsenic, antimony, tellurium

________________________ – to the left

Group 17 – halogens - most reactive nonmetals

Group 18 – noble gases – sublevels completely filled; extremely low reactivity
d-block elements

The d-block elements are metals with typical metallic properties and are often referred to
as
.
f-block elements

Fit between Groups 3 and 4 in the sixth and seventh periods.

____________________ – top row - shiny metals similar in reactivity to the Group 2
alkaline metals

____________________ - bottom row - are all radioactive
Periodic Properties
All of these properties are most noticeable in the main group elements (blocks s and p).
Atomic radius
atomic radius -
Z = _____________________________ = _____________________________

Top to bottom: get larger  more energy levels

Left to right: get smaller  higher Z pulls electrons closer
Ionization Energy
______________ - an atom or group of bonded atoms that has a positive or negative charge.
______________ - any process that results in the formation of an ion
ionization energy, IE -

Top to bottom: decreases  valence electrons further away from protons

Left to right: increases  valence electrons closer to protons
Ionization energies are numbered:

IE1 = first ionization energy, IE2 = second ionization energy, etc.

Each successive IE is higher because the ion feels a stronger nuclear charge

_______________________________________ – the electrons between the nucleus and
the outer electrons reduce the attractive force felt by the outer electrons
Valence Electrons
valence electrons -
Electronegativity
electronegativity -
*Valence electrons hold atoms together in bonds and tend to concentrate near one of the atoms in
the bond due to Z and atomic radius.
 Top to bottom: decreases  shared electrons are further from the pull of the protons

Left to right: increases  more protons in the same principal energy level
Summary of periodic trends
Chapter 6: Chemical Bonding
Topics:
1. Introduction to chemical bonding
2. Covalent bonding and molecular compounds
3. Molecular geometry
4. Ionic bonding and ionic compounds
Section 1: Introduction to chemical bonding

Independent atoms have higher
due to
electrostatic forces between protons and electrons.

Bonding

Most atoms exist as part of a
potential energy for most atoms and creates more
.
chemical bond -
ionic bond –
covalent bond –
Subtract electronegativities to predict bond type:
Bond type
ionic
polar
covalent
Electronegativity
DIFFERENCE
Description
nonpolar
covalent
Generalizations:
ionic = M + NM (metal + nonmetal)
covalent = 2 NM (two nonmetals)
Section 2: Covalent bonding and molecular compounds
molecule –
molecular compound –
chemical formula –
Remember: covalent = sharing = molecular
Formation of a Covalent Bond

The electrons of one atom and protons of a second atom

The two nuclei and the electrons
each other.
one another.

These two forces cancel out to form a
at a length where the
potential energy is at a minimum.

When two atoms form a covalent bond, their shared electrons form
, this achieves a noble-gas configuration.
The Octet Rule

___________________ atoms generally do not react because their electron
configurations are stable.

Their outer s and p orbitals are completely _______________.

Other atoms can fill their outermost s and p orbitals by _______________ electrons.
octet rule –
Exceptions to the Octet Rule:

Some atoms cannot fit 8 electrons (_______________ is complete with 2 electrons;
_______________ has only 3 valence electrons and shares up to 6 electrons)

Some larger atoms can fit more than 8 electrons (_______________________________)
electron-dot notation –
Valence
electrons



Electron-dot
notation
Example
Valence
electrons
Electron-dot
notation
Example
A pair of dots between two atoms represents a bond.
Usually the pair of dots is turned into a straight line.
A single bond is one line, a double bond is two lines, and a triple bond is three lines.
Lewis structures –
1. Write the electron-dot notation for each type of atom in the molecule.
2. Draw a skeleton structure. If C is present, it is the central atom. H and F are always
terminal atoms. The element with the lowest ionization energy/electronegativity goes in
the middle (with some exceptions).
3. Count the total number of available valence electrons. If it is a negative ion, add the
electrons; if it is a positive ion, subtract the electrons.
4. Count the total number of electrons needed for each atom to have a full valence shell.
5. Find the number of bonding electrons: bonding = full – valence.
6. Connect the atoms in the skeleton structure with pairs of bonding electrons. Usually a
straight line is used to represent the bond and it equals two electrons.
7. Add the remaining electrons as lone pairs to each nonmetal atom (except hydrogen) to
complete the octet. lone pair electrons = valence – bonding
8. Count the electrons in the structure to be sure that the number of valence electrons used
equals the number available.
resonance structures –
Notes:
 Polyatomic ions: take into account the different numbers of electrons
 Hydrogen, H, has 2 valence electrons. Boron, B, has 6 valence electrons.
Some covalent bonds do NOT form molecules
network covalent solids –

Can be elements or compounds: carbon silicon, silicon dioxide
Section 3: Molecular Geometry

The properties of molecules depend on the bonding and the

Molecular geometry and the polarity of individual bonds determines the overall
_______________ of the molecule, which in turn helps predict its properties.
hybridization –
Example: Carbon is 1s22s22p2. Based on configuration, it seems as if carbon should have one
unshared pair of electrons and two lone electrons for bonding, but instead it has 4 bonding sites.
VSEPR theory – (“VES-pur” = “valence-shell electron-pair repulsion”)

The name of the shape of a molecule refers to the positions of _______________ only.
Molecular geometries
A = central atom
B = bonded atom of any kind
E = electron pairs
Shape
ABE notation
Sketch
AB2
AB2E or AB2E2
AB3
AB3E
AB4
Intermolecular Forces
intermolecular forces –

The boiling point of a liquid is a good measure of the intermolecular forces between its
molecules: the _______________ the boiling point, the _______________ the forces
between the molecules.

Intermolecular forces vary in strength but are generally _______________ than bonds
between atoms in molecules, ions in ionic compounds, or metal atoms in solid metals.
dipole –
dipole-dipole forces –
hydrogen bonding –
(hydrogen bonds have FON)
London dispersion forces –
Contributing to strength:

Hydrogen bonding

Increases with more electrons, larger molecular weight, and greater surface area
Section 4: Ionic bonding and ionic compounds
ionic compound –
formula unit –

Most ionic compounds exist as ______________________________.

A crystal of any ionic compound is a
of positive and negative ions mutually attracted to each other.

The chemical formula of an ionic compound represents not molecules, but the
(formula unit)
Formation of Ionic Compounds

The _______________ electronegativity of nonmetals means they will gain an electron
and take on a _______________ charge.

The _______________ electronegativity of metals means they will lose an electron and
take on a _______________ charge.
crystal lattice –
lattice energy –
polyatomic ion –
Chapter 7: Formulas and Compounds
Topics
 Section 1 - Ionic names and formulas
 Section 2 - Molecular names and formulas
 Section 4 - Mole & mass conversions
 Section 5 - Percent composition
 Section 6 - Determining chemical formulas
Section 1 - Ionic names and formulas
IONIC COMPOUNDS – the chemical formula represents a “formula unit,” which is the lowest
whole-number ratio of ions in the lattice.
Monatomic ions – formed when atoms gain or lose electrons
Groups:
1+
2+
3+
3-
2-
1-
Metals with unpredictable charges (form cations)

Some d-block elements and metals below the staircase can form more than one ion. You
will be able to figure out the charge from the name or formula that it is part of. However,
it is important to recognize when these are being used so that they are named correctly:
o See the elements on your reference sheet that have a dotted line in their boxes.

It is difficult to predict the charges that some of these metals take on. MEMORIZE
THESE TWO:
zinc: Zn2+ AND silver: Ag+
Polyatomic ions – Re-memorize the list from the reference sheet.
Ionic formulas:
1. Write the symbol with the charge.
2. Criss-cross the charges to find subscripts.
3. Drop the + or – symbol.
4. Do not write the number 1.
5. Reduce the subscripts, if necessary.
6. Use parentheses around polyatomic ions with a subscript larger than 1.
Naming:
7. Write the regular name of the positive (metallic) ion.
8. Use the Roman numeral designation after the name of metals with multiple charges.
9. Change the ending of the negative (nonmetallic) ion to –ide.
10. Use the name of polyatomic ions as is.
Hydrates – water makes up part of the crystal structure
barium hydroxide octahydrate
Ba(OH)28H2O
lead (II) perchlorate trihydrate
Pb(ClO4)23H2O
(You will not have to name these. They are here so that you are familiar with them.)
Section 2 - Molecular names and formulas
MOLECULAR FORMULAS – provide the exact number and kind of atoms in the molecule.
There can be multiple combinations between the same two elements, so you have to be given
either the name or the formula. Do not reduce the subscripts.
1.
2.
3.
4.
5.
If given the name, use the prefixes to write the subscripts.
If given the formula, use the subscripts to write the prefixes.
Change the ending of the second element to –ide.
Do not write the prefix mono for the first element.
The o or a at the end of a prefix is usually dropped if the name starts with a vowel.
Prefixes
1
mono
2
di
3
tri
4
tetra
5
penta
6
7
8
9
10
hexa
hepta
octa
nona
deca
Special cases. Memorize:

water = H2O (not dihydrogen monoxide)

ammonia = NH3 (not nitrogen trihydride)
Diatomic elements - Seven elements exist as bonded as pairs. Memorize:
hydrogen – H2
nitrogen – N2
oxygen – O2
fluorine – F2
chlorine – Cl2
bromine – Br2 (liquid)
iodine – I2 (solid)
Acids and salts. Memorize:
Binary Acids
Oxyacids
(hydro – “root” – ic acid)
HF – hydrofluoric acid
HNO3 – nitric acid
HCl – hydrochloric acid
HNO2 – nitrous acid
HBr – hydrobromic acid
H2SO4 – sulfuric acid
HI – hydroiodic acid
H2SO3 – sulfurous acid
H2S – hydrosulfuric acid
H3PO4 – phosphoric acid
HCN – hydrocyanic acid
CH3COOH – acetic acid
Section 4 - Mole and mass conversions

Chemical formulas are a connection between the lab and chemistry at the particle level.

mole = Avogadro’s number = _______________________________________
molar mass –


The formula mass adds up to the same number, but is in units of amu.
For this section, practice calculating molar mass, mole – mass, and mole – particle
conversions.
Section 5 - Percent composition
percent composition –
% of element in the compound =


Calculate directly from lab measurements by dividing for percentage.
Calculate from a chemical formula starting with mole conversions. This is based on the
assumption that you have one mole of the substance.
Section 6 - Determining chemical formulas
empirical formula –



Two molecular formulas: ethane: C2H4 or butene: C4H8
Both have an empirical formula of CH2.
Starting measurements are either lab measurements of mass or percent composition.
Example: The percent composition of ammonia is 82.4% nitrogen and 17.6%
hydrogen. Calculate the empirical formula of ammonia.
Step 1: Assume the sample is 100 g and make each percentage become a mass.
Step 2: Convert mass to moles.
Step 3: Divide each mole value by the smallest mole value available.
Step 4: Round each division problem to the nearest whole number and use them as subscripts
molecular formula –
Example: Fructose has an empirical formula of CH2O and a molar mass of 180.00 g/mol. What is
its molecular formula?
Step 1: Calculate the formula mass of the empirical formula.
Step 2: Find the multiplying factor, x.
x = molecular formula mass
empirical formula mass
Step 3: Multiply each subscript by the whole number multiplying factor.
Chapter 8: Chemical Reactions
Topics
1. Describing chemical reactions
2. Skeleton equations
3. Balanced chemical equations
4. Synthesis reactions
5. Decomposition reactions
6. Single replacement reactions
7. Double replacement reactions
8. Combustion reactions
Describing chemical reactions
Possible indications of a chemical reaction:
1. ______________________________________________ released
2. ____________ produced
3. ________________________ formed
4. ____________ change
chemical equation -
Characteristics of chemical equations:
1. All
and
must be identified.
2. The reactants and products must be written with the correct
3. The
must be
satisfied (balancing).
Levels of writing a chemical equation:
WORD:
methane + oxygen  carbon dioxide + water
SKELETON:
CH4(g) + O2(g)  CO2(g) + H2O(g)
BALANCED:
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
(not balanced)
Skeleton equations
skeleton equation -
Memorize this list:

Diatomic elements:

S8, P4 (solids)

_________________________________________________ are generally solids.

Liquid elements:

Gases:

Ionic compounds are usually

Acids are usually
and all
.
.
Symbol
Explanation
separates two reactants or two products
“yields” – separates reactants from products
reversible reaction
solid
liquid
gas
aqueous, dissolved in water
gas product
precipitate formed
Heat is supplied to the reaction
The formula of a catalyst is written above the yield symbol (in this
example, platinum, Pt)
Balanced chemical equations
balanced equation –

Balancing equations demonstrates the
(mass is neither created nor destroyed in a closed system)

Use coefficients to balance the equation (not subscripts).
6 H2 O
“6” (coefficient) refers to total number of particles “2” and “1” (subscripts) refer to the number
of atoms in the substance and are part of its
identity
Guidelines for balancing equations:
1. This is done by inspection and trial-by-error, but mini charts can help.
2. Balance the different types of atoms one at a time.
3. First balance the atoms of elements that are combined and that appear only once on each
side of the equation.
4. Balance polyatomic ions that appear on both sides of the equation as single units.
5. Balance H and O atoms last.
6. Count to make sure each type of atom is balanced.
Synthesis Reactions
synthesis reactions -
A + B  AB
Metal + nonmetal: one ionic product
Na(s) + O2(g) 
Nonmetal + oxygen: one molecular product
C(s) + O2(g) 
H2(g) + O2(g) 
Metal oxides + water: one ionic product ending with hydroxide (OH)
CaO(s) + H2O(l) 
Decomposition Reactions
decomposition reactions -
CD  C + D
Binary compounds: two elements
H2O(l) 
Metal carbonates: metal oxide + carbon dioxide
CaCO3(s) 
Metal hydroxides: metal oxide + water
Ca(OH)2(s) 
Metal chlorates: metal chloride + oxygen
KClO3(s) 
Single Replacement/Displacement
single replacement -
E + FG  F + EG
or
E + FG  G + FE
**Use the Activity Series to predict whether or not a reaction will happen.
Metal plus ionic compound
Al(s) + Pb(NO3)2(aq) 
Metal plus acid: ionic compound + hydrogen
Mg(s) + 2HCl(aq) 
Halogen plus ionic compound
Cl2(g) + KBr(aq) 
Double Replacement/Displacement
double replacement -
HI + JK  HK + JI
**The reaction will happen if a precipitate, gas, or water is formed.
Formation of a precipitate
PbI2(s) + KNO3 
Formation of a gas
FeS(s) + HCl(aq) 
Formation of water
HCl(aq) + NaOH(aq) 
Combustion
combustion -
CxHy + O2(g)  CO2(g) + H2O(g/l)
combustion of methane
CH4(g) + O2(g) 