Chemistry
EOC Review
Key Concepts by TEK
Must Knows!!!
Adapted from Stemscopes, © 2012, Rice University
Physical and Chemical Changes and Properties
Student Expectation:
to differentiate between physical and chemical
changes and properties
identify extensive and intensive properties
classify matter as pure substances or mixtures
through investigation of their properties.
Physical and Chemical Changes and Properties
Key Concepts
A physical change to matter will not alter the
composition or identity of a substance.
A chemical change to matter will always result in
the formation of a new substance.
The physical properties of matter include
properties that describe the substance such as
color, smell, boiling point, density or others
Physical and Chemical Changes and Properties
Extensive properties are dependent on the amount
of a substance present, such as mass, number of
particles, or energy.
Intensive properties are physical properties of
matter that are not dependent on the amount of a
substance present such as density, ductility, and
boiling point.
Physical and Chemical Changes and Properties
All matter on Earth can be classified as either a
pure substance or a mixture.
A pure substance consists of a single substance
with its own definite composition and properties.
A mixture consists of a combination of two or more
pure substances with variable composition and
properties.
States of Matter
Student Expectation:
compare solids, liquids, and gases in terms of
compressibility, structure, shape, and volume.
States of Matter
Key Concepts
Kinetic Molecular Theory states that all matter is
composed of particles that are in constant motion.
Solid particles have low energy and vibrate in fixed
position.
Liquid particles have higher energy and are able to move
freely past each other.
Gas particles are widely spaced with high energy and
little attraction to each other
States of Matter
Solids are relatively incompressible, have a fixed
shape, and a definite volume
Liquids are more compressible than solids, but
much less compressible than gases. Liquids have a
definite volume but will take the shape of the
container
Gases are very compressible and do not have a
definite shape. Gases will fill the volume of a closed
container. If the container is opened, gases will not
have a definite volume.
States of Matter
Phase
Solid
Energy –
move
Little
Liquid
Some
Gas
Lot
Particle
spacing
Close
Rigid
Close
Slide
Far
Volume
Shape
Compressible
Definite
Definite
No
Definite
Not
Definite
Not
Definite
No
Not
Definite
Yes
The Periodic Table
Student Expectation:
to explain the use of chemical and physical
properties in the historical development of the
Periodic Table
use the Periodic Table to identify and explain the
properties of chemical families, including alkali
metals, alkaline earth metals, halogens, noble
gases, and transition metals.
The Periodic Table
Key Concepts
Early scientists, such as Dmitri Mendeleev and
Lothar Meyer, observed that the chemical
properties of the known elements repeated when
placed in order of increasing mass
These patterns led to the development of the
modern Periodic Table. Group A elements (the
representative elements) within the same Group
(column) have similar chemical properties.
The Periodic Table
Group 1A Alkali metals
1 valence electron is lost easily, forms a cation with
a 1+ charge.
The group has similar physical properties
Metallic appearance.
The Periodic Table
Group 2A Alkaline earth metals
2 valence electrons that can be readily lost, forms a
cation with a 2+ charge.
The group has similar physical properties
Metallic appearance.
The Periodic Table
Group 7A Halogens (nonmetals)
highly reactive with 7 valence electrons, forming an
anion with a 1- charge
Group 8A noble gases
have a complete octet (8) of valence electrons
they have little tendency to gain or lose electrons
and are non-reactive.
The Periodic Table
Group B transition metals
in the middle of the Periodic Table
Number of outer electrons varies, leading to a
variety of charges for each element.
they share many of their chemical and physical
properties
The Lanthanide series and the Actinide series,
located at the bottom of the periodic table, are the
inner transition metals (the rare earth metals).
Periodic Trends
Student Expectation:
to use the Periodic Table to identify and explain
periodic trends, including atomic and ionic radii,
electronegativity, and ionization energy
Periodic Trends
Key Concepts
Periodic trends
patterns that occur across a row (from left to right) or
down a column (family)
can be used to predict certain properties of elements in
their atomic or ionic form.
Periodic Trends
Atomic radii
Decreases when moving from left to right on the
Periodic Table due to the increasing number of
positive protons within the nuclei pulling on the
valence electrons.
Increases down a Group due to additional electron
shells between the nucleus and repulsion among
electrons.
Periodic Trends
Ionic radii - compared based on their numbers
of protons and electrons
Across a period, With a greater number of protons,
ions will be smaller due to attractive forces
between the nucleus and the valence electrons
Down a Group, atoms will have an increase in ionic
radius for both + and - ions due to the addition of
an electron shell.
Down a group, a greater number of electrons will
cause an ion to become larger due to electron
repulsion
Periodic Trends
electronegativity - ability to attract electrons.
increases moving from left to right
Decreases moving down a Group, making fluorine
the most electronegative element.
Periodic Trends
Ionization energy - is the amount of energy
required to remove an electron from a neutral
atom
increases moving from left to right
Within the same family decreases with increasing
atomic number
Atomic Theory
Student Expectation:
to understand the experimental design and
conclusions used in the development of modern
atomic theory, including Dalton's Postulates,
Thomson's discovery of electron properties,
Rutherford's nuclear atom, and Bohr's nuclear
atom.
Atomic Theory
Key Concepts
Concepts of the atom and the nature of matter
originated with Greek philosophers more than
2000 years ago.
These ideas, though not scientifically tested, formed the
basis for later scientists to build on and develop modern
atomic theory.
Atomic Theory
John Dalton (1800’s)
investigated the nature of gases in order to gain
a better understanding of the laws of
conservation of mass and of multiple
proportions.
His five postulates of atomic theory helped to
define the structure and nature of the atom.
The scientific community accepted his postulates
due to his sound experimental evidence.
Atomic Theory
J.J. Thomson (late 1800’s)
Using cathode rays, he discovered that the rays
were actually negatively charged particles with
a charge of 1-, and that they were much smaller
than atoms.
Also studied the relationship between electric
charge and matter.
Thomson developed the “plum pudding model,” in
which electrons were embedded in a positively
charged sphere.
Atomic Theory
Ernest Rutherford (early 1900s)
developed the nuclear model of the atom
Rutherford’s scattering experiment, (Gold foil
experiment)
found that atoms contained an extremely small,
dense, and positively charged nucleus
the area around the nucleus was mostly empty
space with a few negative electrons.
Atomic Theory
Niels Bohr
refined the findings of Rutherford
Used spectral light emissions to conclude that
electrons had specific energy levels
His atomic model consisted of spherical shells of
electrons on various states surrounding the
positively charged nucleus.
Atomic Equations
Student Expectation:
to understand the electromagnetic spectrum and
the mathematical relationships between energy,
frequency, and the wavelength of light
c=λxf
To calculate the wavelength, frequency, and energy
of light using Planck's constant and the speed of
light.
E = hf or E = hc/ λ
Atomic Equations
Key Concepts
The electromagnetic spectrum displays the full
range of electromagnetic energy based on wave
properties, from high-energy gamma rays to lowenergy radio waves.
Electromagnetic waves are characterized by energy,
frequency, and wavelength
Long wavelengths = low frequency and low energy
Short wavelengths = high frequency and high energy
Atomic Equations
c=λxf
All energy waves travel at the same velocity known
as the speed of light (c), which equals 3.0 x 108
m/s.
Wavelength (λ) is defined as the distance between
two peaks or two troughs on a wave.
Frequency (f) is defined as the number of waves
passing a given point per second.
Since the speed of light is a constant, frequency
and wavelength are inversely proportional
Atomic Equations
E = hf or E = hc/ λ
Planck’s constant equals 6.63 x 10-34 Js
the energy of an electromagnetic wave is
directly proportional to its frequency
Atomic Structure
Student Expectation:
to use isotopic composition to calculate average
atomic mass of an element
express the arrangement of electrons in atoms
through electron configurations and Lewis valence
electron dot structures.
Atomic Structure
Key Concepts
isotopes - elements with the same number of
protons but a different number of neutrons
The average atomic mass equals the average of the
masses of all the naturally occurring atoms and
isotopes for an element.
Atomic Structure
Electron configuration shows the location and
number of electrons in an atom.
Energy levels are divided into four sublevels; s, p, d, f.
The sublevels are filled with the lowest energy available
orbitals filled first.
Lewis valence electron dot structures can be used
to represent the outer electrons of an atom.
Uses the chemical symbol of an element surrounded by
dots, each dot represents an electron found in the
valence shell.
Nuclear Chemistry
Student Expectation:
to describe the characteristics of alpha, beta, and
gamma radiation
describe radioactive decay process in terms of
balanced nuclear equations
compare fission and fusion reaction.
Nuclear Chemistry
Key concepts
Radiation emitted by an element can be
characterized as alpha, beta, or gamma radiation
Alpha radiation
helium nuclei with a +2 positive charge.
Beta radiation
electron particles with a -1 negative charge
gamma radiation
high-energy photons with a neutral charge.
Nuclear Chemistry
Balanced radioactive decay is written to show
the conservation of mass number and atomic
number during the transmutation of one
element into another occurs.
Nuclear fission
large, unstable atoms split into smaller atoms to
achieve a more stable state.
Nuclear fusion is the opposite of fission
occurs when smaller atoms bind together to form
more a larger, more stable atom.
Chemical Formulas
Student Expectation:
to name ionic compounds, covalent compounds,
acids, and bases, using International Union of Pure
and Applied Chemistry (IUPAC) nomenclature rules
write the chemical formulas of common polyatomic
ions, ionic compounds, covalent compounds, acids,
and bases
construct electron dot formulas to illustrate ionic
and covalent bonds.
Chemical Formulas
Key Concepts
Ionic compounds
name the positive ion (cation) first followed by the
negative ion (anion).
If the anion is one atom the suffix -ide is added to the
anion name.
Use a roman numeral for transition metals to indicate the
correct charge of the metal cation.
The electrons lost by the cation must equal the electrons
gained by the anion to form a neutral ionic compound.
Chemical Formulas
Covalent compounds
Use prefixes to show the number of atoms.
Then second element ends in -ide.
Non-metal atoms share electrons, forming a more stable
compound so each atom can achieves a full octet of
electrons.
Chemical Formulas
Acids
Binary acids contain just hydrogen and a nonmetal.
The prefix hydro- is used followed by the root name of
the anion and then suffix -ic. (hydro – ic)
Oxyacids (acids that contain hydrogen and a polyatomic
ion containing oxygen)
ite – ous acid
ate – ic acid
Chemical Bonding
Student Expectation:
to construct electron dot formulas to illustrate ionic
and covalent bonds
describe the nature of metallic bonding
apply the theory to explain metallic properties such
as thermal and electrical conductivity, malleability
and ductility.
Chemical Bonding
Key Concepts
Electron dot formulas are used to show how bonds
are formed
Ionic bonds, (formula units)
Anions and cations are shown in brackets with their
respective charges.
Lithium oxide
Chemical Bonding
Covalent bonds, (molecules)
the valence electrons are shared between two atoms.
Starting with the central atom, the electron pairs are
placed around each atom in order to fulfill the octet rule
for each atom.
Carbon dioxide
Chemical Bonding
Metallic bonding
electrons are delocalized, they do not remain close to
any one atom.
solid state, the valence electrons of the metal atoms
move freely in atom, forming what is known as the
“electron sea”.
Metallic bonding is formed due to the attraction of the
electrons for the metal cations.
metallic bonding are more flexible than ionic or
covalent bonds
metals are ductile and malleable
metals are good thermal and electrical conductors.
Molecular Structure
Student Expectation:
to predict molecular structure
Linear (180o), trigonal planar (120o), or tetrahedral (109.5o),
using Valance Shell Electron Pair Repulsion (VSEPR) theory.
Molecular Structure
Key Concepts
Linear (180o), trigonal planar (120o), or tetrahedral
(109.5o)
- or bent
The Mole
Student Expectation:
to define and use the concept of a mole
To use the mole concept to calculate the number of
atoms, ions, or molecules in a sample of material
The Mole
Key Concepts
The mole is to describe an amount of a substance.
The mole is equal to 6.02 × 1023 atoms, molecules, or
formula units of a substance.
Is defined by the number of atoms in exactly 12
grams of carbon-12.
The Mole
One mole = Molar mass
The molar mass of a substance (g/mol), can be found by
adding the atomic masses of the atoms on the Periodic
Table.
1 mole of a gas = 22.4 Liters of a gas
(for gases only at STP, standard temperature and
pressure)
Percent Composition Calculations
Student Expectation:
to calculate percent composition and empirical and
molecular formulas.
Percent Composition Calculations
Key Concepts
Percent composition of a compound represents the
percent of each element in a compound by mass.
Equals the molar mass of the whole compound divided by
the mass of a single element. Then, multiplying by 100 to
make a percent.
Empirical formula is the lowest whole-number ratio
of elements in a compound.
Calculated by
1st: Convert the masses of each element in the
compound into mole ratios
2nd: Divide all by the smallest mole
3rd: Write the smallest whole number ratios with each
element
Percent Composition Calculations
A molecular formula is the actual chemical formula of
a compound
Determined based on the molar mass and the empirical
formula of the compound.
1st: The molar mass of the empirical formula is
calculated
2nd: The molar mass of the compound is divided by the
molar mass of the empirical formula to find a whole
number integer.
3rd: The empirical formula is then multiplied by this
integer to calculate the molecular formula.
Chemical Equations
Student Expectation:
to use the law of conservation of mass to write and
balance chemical equations.
Chemical Equations
Key Concepts
Chemical equations must follow the Law of
Conservation of Mass - mass is neither created nor
destroyed in a non-nuclear change.
The total mass of the reactants must equal the total
mass of the products.
Reactants  products
Chemical Equations
How to balance
1st: count the number of atoms of each element on the
reactant side and compare that to the number of atoms of
the same elements on the product side
2nd: Use coefficients in front of the chemical formulas to
make the number of atoms on each side of the arrow equal.
Stoichiometry
Student Expectation:
to perform stoichiometric calculations
determination of mass relationships between reactants and
products
calculation of limiting reagents
percent yield.
Stoichiometry
Key Concepts
Mole ratios are used to determine the relationships
among moles in a reaction
these ratios come from the coefficients of a balanced
chemical equation.
The molar masses of reactants and products are used
as conversion factors to calculate mass relationships.
Stoichiometry
A limiting reagent determines the amount of product
formed in a reaction, as it is the reactant that is
completely consumed first.
When one of the reactants is consumed in a chemical
reaction, the reaction stops and no further products may be
formed.
Use stoichiometry to calculate the limiting reactant
determine which reactant is consumed first, and then
use the amount of this reactant to find the moles or
mass of product formed.
Stoichiometry
Percent yield of a reaction is found by dividing the
actual amount of a product by the theoretical
amount of the product, and then multiplying by 100
to make a percent.
% yield = actual x 100
theoretical
Stoichiometry is used to find the theoretical (ideal)
yield of a chemical
The actual yield is a measurement of the actual
amount of product made during a reaction.
Gas Laws
Student Expectation:
to describe and calculate the relations between
volume, pressure, number of moles, and temperature
for an ideal gas as described by Boyles's law, Charles'
law, Avogadro's law, Dalton's law of partial pressure,
and the ideal gas law
describe the postulates of kinetic molecular theory.
Gas Laws
Key Concepts
Kinetic Molecular theory is used to describe the
behavior of ideal gases.
ideal gas particles are of negligible size compared to the
space between them
the particles are in continuous, rapid, random motion
their collisions are elastic
there are no significant interactions among the particles of
a gas
When temperature increases for an ideal gas, the kinetic
energy of the particles increases proportionally.
Gas Laws
Boyle’s law, P1V1 = P2V2
the pressure exerted by gas particles is inversely
proportional to the volume occupied by the gas.
At constant temperature, an increase in pressure will result
in a decrease in volume.
Gas Laws
Charles’s law, V1/T1 = V2/T2
The temperature of a gas is directly proportional to the
volume occupied by the gas. When the pressure is held
constant, a decrease in temperature results in a decrease in
volume.
temperature must be in Kelvin.
Gas Laws
Avogadro’s law, V1/n1 = V2/n2
relationship between the number of moles (n) of a gas, and
the volume occupied by the gas at a constant temperature
and pressure, where the volume of a gas is directly
proportional to the number of moles of a gas.
Avogadro’s principle states that equal volumes of gases
contain the same number of atoms or particles.
Gas Laws
Dalton’s law of partial pressures,
Ptotal = P1 + P2 + P3 +…Pn
The total pressure exerted by a mixture of gases is equal to
the sum of the individual pressures of all the gases in the
mixture.
Gas Laws
The ideal gas law, PV = nRT
describes the relationship between the temperature,
pressure, volume, and the number of moles of a gas under
specific conditions.
This law represents a combination of the relationships
described in Boyle’s law, Charles’ law, and Avogadro’s law.
The ideal gas constant R is a physical constant and is
specific to the units used for the pressure and volume.
Temperature must be in Kelvin.
Gas Stoichiometry
Student Expectation:
to perform stoichiometric calculations
including determination of mass and volume relationships
between reactants and products for reactions involving
gases.
Gas Stoichiometry
Key Concepts
Using a balanced chemical equation, stoichiometric
ratios, and molar masses to determine the mole and
mass relationships between the amounts of products
and reactants in a chemical equation involving gases.
For reactions not at STP, stoichiometric ratios and the
Ideal Gas Law can be used to find mole and volume
relationships between products and reactants.
Aqueous Solutions
Student Expectation:
to describe the unique role of water in chemical and
biological systems
develop and use general rules regarding solubility
through investigations with aqueous solutions.
Aqueous Solutions
Key Concepts
Water has several unique properties.
amphoteric, (can be an acid or a base)
is highly polar,
has a higher boiling and melting point than other
compounds of the same structure due to the hydrogen
bonding between water molecules.
The strong attraction between water molecules (hydrogen
bonding) accounts for properties such as cohesion and
surface tension.
Aqueous Solutions
Life on Earth is highly dependent on the unique
properties of water
most of the metabolic processes in biological
organisms take place in aqueous solutions.
water is an essential component of the energy
transformation processes of photosynthesis and cellular
respiration.
water’s high specific heat plays a critical role in
Earth’s relatively moderate temperature variations.
Use STAAR Chart
to determine
Which Ionic
compounds are
soluble (aq) and
insoluble (s)
Molarity
Student Expectation:
to calculate the concentration of solutions in units of
molarity
use molarity to calculate the dilutions of solutions
Molarity
Key Concepts
Molarity (M) is used to express the concentration of a
solution
moles of a solute dissolved per liter of solution, or mol/L.
This unit of mol/L is also called “molar.”
Diluted solutions may be created from concentrated
solutions by merely adding more solvent.
M1V1 = M2V2,
Solutions and Solubility
Student Expectation:
to distinguish between types of solutions such as
electrolytes and nonelectrolytes and unsaturated,
saturated, and supersaturated solutions
investigate factors that influence solubilities and rates
of dissolution such as temperature, agitation, and
surface area.
Solutions and Solubility
Key Concepts
When ionic compounds dissolve in a solvent, the
charged ions in the solution can conduct an electric
current and are called electrolytes.
When molecular compounds dissolve in a solvent,
the molecules in solution do not conduct electric
current and are called nonelectrolytes.
Solutions and Solubility
saturated solution - a solution in which the maximum
amount of solute is dissolved in the solvent
supersaturated - the solution contains more than the
maximum amount of solute
unsaturated - it contains less than the maximum
amount.
Solutions and Solubility
Factors that affect the solubility of a solute in a
solvent
temperature of the solvent
intermolecular forces
(for gases) the partial pressure of a gas over a liquid solute.
Rates of dissolution of a solute in a solvent depend
on
Temperature
Agitation (stirring)
surface area of the solute
Acids and Bases
Student Expectation:
to define acids and bases and distinguish between
Arrhenius and Brønsted-Lowry definitions
Understand, differentiate, and predict products in
acid base reactions, precipitation reactions, and
oxidation-reduction reactions
define pH and use the hydrogen or hydroxide ions
concentrations to calculate the pH of a solution
distinguish between degrees of dissociation for strong
and weak acids and bases.
Acids and Bases
Key Concepts
The Arrhenius definition of acids and bases in
aqueous solutions,
an acid forms more hydrogen ions (H+)
a base forms more hydroxide ions (OH-)
The Brønsted-Lowry definition of acids and bases
an acid can donate a proton (H+)
a base can accept a proton.
Acids and Bases
Acid base reactions involve the transfer of protons. In
these neutralization reactions, the acid and the base
normally react to form water and salt.
A precipitation reaction occurs when an insoluble
precipitate is formed.
Both precipitation and acid-base reactions are types
of double replacement reactions.
Oxidation-reduction (redox) reactions involve the
transfer of electrons
Acids and Bases
The pH scale is used to determine the strength of an
acid or base
pH = – log [H+]
The ranges on the pH scale
below 7 for acids
above 7 for bases
= 7 for neutral solutions
[H+] x [OH- ] = 1 x 10-14
Acids and Bases
Complete dissociation of ions occurs with strong
acids and strong bases,
partial dissociation of ions occurs with weak acids
and weak bases.
The conjugate base of a strong acid is a weak base.
The conjugate acid of a strong a base is a weak acid.
Energy
Student Expectation:
to understand energy and its forms, including kinetic,
potential, chemical, and thermal energies
understand the law of conservation of energy
the processes of heat transfer.
Energy
Key Concepts
Matter may contain one or more types of energy,
where energy is defined as the ability to do work.
Kinetic energy (energy of motion)
based on the movement of an object or a substance.
Potential energy (energy of position)
stored energy based on the composition of a substance
the position of an object in space.
Energy
Chemical energy (form of potential energy)
the energy stored in the bonds between the atoms of
molecules and ionic crystals.
the potential of a chemical substance to undergo a
transformation through a chemical reaction
the ability of a system (the chemical reaction) to do work
during chemical reactions.
During chemical reactions, energy is transferred,
either producing thermal energy or requiring thermal
energy.
Energy
Thermal energy is the total internal energy that a
substance possesses as well as the total KE of the
particles in a system.
It depends on the temperature of that system
relates to both heat and temperature.
Temperature is a measure of the average KE within the
system
heat is the movement of thermal energy from one
substance to another.
Energy
The law of conservation of energy
states that energy can neither be created nor destroyed,
but instead is transferred between a system and its
surroundings.
This energy transfer often occurs in the form of heat
flow. Heat, or the movement of thermal energy from
a warmer substance to a cooler substance, flows
between a system and its surroundings
Thermochemical Equations
Student Expectation:
to use thermochemical equations to calculate energy
changes that occur in chemical reactions
classify reactions as exothermic or endothermic
perform calculations involving heat, mass,
temperature change, and specific heat
Thermochemical Equations
Key Concepts
A thermochemical equation is a balanced chemical
equation that includes the amount of energy
absorbed or released during a chemical reaction.
The total heat energy of the system is known as the
enthalpy (H) of the system.
ΔH is often expressed in kilojoules (kJ).
Thermochemical Equations
Exothermic reaction, ΔH is negative
heat flows from the system to the surroundings.
The energy released during the reaction is represented as a
product in the thermochemical equation.
Endothermic reaction, ΔH is positive
heat flows into the system from the surroundings.
The energy absorbed during the reaction is represented as
a reactant in the equation.
Thermochemical Equations
The ΔH (change in enthalpy) of a system is equivalent
to the heat of the system (q)
q = mc ΔT
q = m (mass in grams) x c (specific heat) x ΔT (change in
temperature).
The specific heat (c) of a substance is defined as the
amount of heat required to raise the temperature of
1 gram of the substance by 1 degree Celsius.
Calorimetry
Student Expectation:
to use calorimetry to calculate the heat of a chemical
process.
Calorimetry
Key Concepts
Calorimetry includes the complete set of equations
and experimental procedures used to measure the
heat flow for physical and chemical processes.
A calorimeter is the device used to measure the heat
absorbed or released during a chemical process.