NAME:______________________________________________
Period ____
CORE CURRICULUM CHECK-OFF LIST
REGENTS CHEMISTRY
The following information is everything you need to know and be able to do to attain Mastery of the
Regents Chemistry Curriculum. The bold, underlined words are important vocabulary words that you
should be able to define and use properly in explanations. This is a study guide for what you will be tested
on throughout the year. The objectives are divided into categories of “Knowledge” (what you have to
know) and “Application” (what you have to be able to do). Check off each objective when you fully
understand it. If you do not understand an objective, ask questions before you are tested on it.
I. MATHEMATICS ANALYSIS & GRAPHING
o Identify independent and dependent variables in an experiment and correctly plot them on an
axis
Example hypothesis: Chemistry students who do their homework will have higher test scores than students who do not
do their homework.
1.
 X-axis (horizontal): the independent variable is the one that is manipulated by the
experimenter. (“The one I change.” – do homework/not do homework)
 Y-axis (vertical): the dependent variable is the one that changes based on the independent
variable. (The data you collect – test scores)
o Express uncertainty in measurement by properly using significant figures
o Identify the number of sig. figs. in data
o Round to the correct number of sig. figs. on calculations:
2.
 Addition and Subtraction (round to least precise [farthest left] place value)
 Multiplication and Division (round to the lowest # of digits in data)
 Combined Rules (Add/Subtract first, then Multiply/Divide)
3.
o Identify relationships between variables from data tables and graphs (direct or inverse
relationships)
4.
o Understand what is meant by conditions of Standard Temperature and Pressure (STP)
(Table A)
o Recognize and convert between units on various scales of measurement (Tables C, D, and T)
5.
 Temperature: Celsius ↔ Kelvin
 Mass: grams ↔ kilograms
 Thermal Energy: joules ↔ kilojoules
 Length: meters ↔ centimeters ↔ millimeters
 Pressure: kilopascals ↔ atmospheres
 Amount of Substance:
GFM = 1 mole = 6.02x1023 particles
o grams ↔ moles ↔ atoms or molecules
o Use the density equation on Table T to solve for density, mass, or volume, given the other two
values
7. o Calculate percent error (Table T)
6.
II. ATOMIC CONCEPTS
Knowledge
o
1.
o
o
2.
o
o
o
3.
o
Application
o Explain what happened during the goldfoil experiment and what it showed.
Atoms are the basic unit (building block) of
Gold foil was bombarded (hit) with positively
matter.
charged alpha particles. Most alpha particles passed
Atoms of the same kind are called elements.
through the gold foil, but some were deflected. This
The modern model of the atom has developed
showed that:
over a long period of time through the work of
1. the atom is mostly empty space
many scientists.
2. the nucleus is small, positively charged, and
located in the center of the atom
The three subatomic particles that make up an atom are protons, neutrons, and electrons.
The proton is positively charged, the neutron has no charge, and the electron is negatively charged.
(This is also referenced on Table O. Check it out!)
Each atom has a nucleus with an overall
positive charge, made up of protons and
o Determine the nuclear charge of an atom.
neutrons.
(equal to the number of protons in the
The nucleus is surrounded by negatively charged
nucleus)
electrons.
4.
o Atoms are electrically neutral, which means that
they have no charge (# protons = # electrons)
o Ions are atoms that have either lost or gained
electrons and are either positively or negatively
charged
o When an atom gains one or more electrons, it
becomes a negative ion and its radius increases.
o When an atom loses one or more electrons, it
becomes a positive ion and its radius decreases.
o Determine the number of protons or
electrons in an atom or ion when given one
of these values. (Periodic Table)
o Compare the atomic radius and ionic
radius of any given element
Ex: A chloride ion has a larger radius than a
chlorine atom because the ion has an extra electron.
A sodium ion has a smaller radius than a sodium
atom because the ion has lost an electron.
5.
o The mass of each proton and each neutron is
approximately equal to one atomic mass unit
(AMU).
o An electron is much less massive (has almost no
mass) compared to a proton or neutron.
o Calculate the mass of an atom given the
number of protons and neutrons
o Calculate the number of neutrons or
protons, given the other value
6.
o In the modern model of the atom, the WAVE-MECHANICAL MODEL (electron cloud), the
electrons are in orbitals (clouds), which are defined as regions of most probable electron location.
7.
8.
o Each electron in an atom has a specific amount of energy.
o Electrons closest to the nucleus have the lowest energy. As an electron moves away from the
nucleus, it has higher energy.
o Electron configurations show how many
electrons are in each orbital.
o When all of an atom’s electrons are in the
o Distinguish between ground state and
orbitals closest to the nucleus, the electrons are
excited state electron configurations.
in their lowest possible energy states. This is
(Be careful to keep the same number of
called the ground state.
electrons when writing the excited state.)
o When an electron in an atom gains a specific
amount of energy, the electron is at a higher
energy state (excited state)
o
9.
o
o
10.
o
11.
o
12.
o
1.
o
o
o
2.
o
3.
II. ATOMIC CONCEPTS (CONTINUED)
Knowledge
Application
When an electron returns from a higher energy
(excited) state to a lower energy (ground) state, a
specific amount of energy is emitted. This
emitted energy can be used to identify an
o Identify an element by comparing its brightelement.
line spectrum to given spectra
The flame test is an example of the bright-line
spectrum visible to the naked eye. The color
can determine the identity of a positive ion in a
compound.
o Draw a Lewis electron-dot structure of an
The outermost electrons in an atom are called
atom.
the valence electrons. In general, the number
o Distinguish between valence and nonof valence electrons affects the chemical
valence electrons, given an electron
properties of an element.
configuration
Atoms of an element that contain the same
o Calculate the number of neutrons in an
number of protons but a different number of
isotope of an element given the isotope’s
neutrons are called isotopes of that element.
mass
o Given an atomic mass, determine the most
abundant
The average atomic mass of an element is the
isotope
weighted average of the masses of its naturally
o Calculate the atomic mass of an element,
occurring isotopes.
given the masses and abundance of
naturally occurring isotopes
III. THE PERIODIC TABLE
Knowledge
Application
The placement or location of an element on the Periodic
Table gives an indication of physical and chemical
o Explain the placement of an
properties of that element.
unknown element in the periodic
The elements on the Periodic Table are arranged in order of
table based on its properties
increasing atomic number.
The number of protons in an atom (atomic number)
identifies the element. This goes at the bottom left corner
of the symbol for that element.
o Interpret and write isotopic
The sum of the protons and neutrons in an atom (mass
notations.
number) identifies an isotope. The mass number is placed
Ex:
at the top left corner of the symbol for an element OR is
C-12, C-13, and C-14 are isotopes of
placed after the element symbol or name and a dash.
the element carbon
Ex: Three different ways to write carbon with a mass number of 14:
C, carbon-14, or C-14
o Identify the properties of metals,
metalloids, nonmetals and noble
Elements are classified by their properties, and located on
gases
the periodic table as metals, metalloids, nonmetals, or
o Classify elements as metals,
noble gases.
metalloids, nonmetals, or noble
gases by their properties
o
4.
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5.
o
o
6.
o
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7.
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8.
o
o
9.
o
10.
III. THE PERIODIC TABLE (CONTINUED)
Knowledge
Application
An element’s atomic radius, first ionization energy, and electronegativity determine its
physical and chemical properties
Substances can be differentiated by their physical
o Identify and give examples of
properties.
physical properties
Physical properties of substances include melting point,
o Describe the states of the
boiling point, density, conductivity, malleability,
elements at STP (solid, liquid, or
solubility, and hardness.
gas). (Table S)
Substances can be differentiated by chemical
o Identify and give examples of
properties.
chemical properties
Chemical properties describe how an element behaves
o Describe the difference between
during a chemical reaction and include reactivity,
physical and chemical properties
flammability, and toxicity.
of substances
Some elements exist as two or more forms in the same phase. These forms differ in their
molecular or crystal structure and therefore in their properties. The word to describe this
phenomenon is ALLOTROPE.
Ozone and oxygen gases are allotropes of each other. Ozone is O3 and it is very dangerous to our
health. Oxygen gas is O2 and we need it to survive.
Diamonds and graphite (better known as pencil lead) are both forms of the element carbon. They
have different molecular structures and very different properties.
o Determine the group of an
For Groups (also called families) 1, 2, and 13-18 on the
element, given the chemical
Periodic Table, elements within the same group have the
formula of a compound
same number of valence electrons (helium is the exception)
Ex: A compound has the formula
and therefore similar chemical properties.
XCl2, element X is in Group 2
Elements in the same Period (row) have the same number
o Determine the number of
of principal energy levels (shells) which contain
energy levels containing
electrons.
electrons given an element’s
Period and vice versa
The succession of elements within the same GROUP (top
o Compare and contrast properties
to bottom) demonstrates characteristic trends: differences
of elements within a group or a
in atomic radius, ionic radius, electronegativity, first
period for groups 1, 2, and 13-18
ionization energy, and metallic/nonmetallic properties.
on the periodic table
The succession of elements across the same PERIOD (left o Understand and be able to
to right) demonstrates characteristic trends: differences in
explain the trends in terms of
atomic radius, ionic radius, electronegativity, first
nuclear charge and electron
ionization energy, and metallic/nonmetallic properties.
shielding
IV. MATTER, ENERGY, & CHANGE
Knowledge
o Matter is anything that has mass and volume (takes up
space).
o Matter cannot be created nor destroyed, only
1.
transformed.
o Matter is classified as a pure substance (element or
compound) or as a mixture of substances.
Application
o Identify specific examples of matter
as an element, compound, or
mixture
o
o
2.
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4.
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5.
o
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6.
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7.
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8.
IV. MATTER, ENERGY, & CHANGE (CONTINUED)
Knowledge
Application
Energy is not matter (does not have volume)
Energy can exist in different forms, such as kinetic,
o Distinguish between matter and
potential, thermal (heat), sound, chemical, electrical, and
energy
electromagnetic.
Energy cannot be created or destroyed, only transformed.
During a physical change, particles of matter are
rearranged. Examples of physical changes include
freezing, melting, boiling, condensing, dissolving,
o Differentiate between physical and
crystallizing, and crushing into a powder
chemical changes in matter
During a chemical change, NEW substances are
o Identify and give examples of
formed with new properties. Examples of chemical
physical changes and chemical
changes include combustion (burning), rusting, and
changes in matter
neutralizing an acid or base.
Energy can be absorbed or released during physical and
chemical changes.
A pure substance (element or compound) has a uniform
composition and constant properties throughout a given
o Draw and interpret particle
sample, and from sample to sample.
diagrams for elements,
Mixtures are composed of two or more different
compounds, and mixtures
substances that can be separated by physical means.
Elements are substances that are composed of atoms that have the same atomic number. Elements
cannot be broken down by chemical change.
A compound is a substance composed of two or more
o Describe differences in ionic and
different elements that are chemically combined in a fixed
molecular/covalent compounds
proportion.
o Identify a compound as ionic or
A compound can be broken down by chemical means,
molecular/covalent compound
such as during a chemical reaction.
given its properties
Two major categories of compounds are ionic and
o Name compounds based on their
molecular (covalent) compounds.
chemical formulas
A chemical compound can be represented by a specific
o Determine the formula of a
chemical formula and assigned a name based on the
compound given its name
IUPAC system.
When different substances (elements or compounds) are
mixed together and do NOT chemically react, a mixture
is formed.
o Interpret particle diagrams as
The amounts of substances in a mixture can vary. Each
showing homogeneous or
substance in a mixture retains its original properties.
heterogeneous mixtures
The composition of a mixture can vary. If the substances
o Give examples of homogeneous
are uniformly (evenly) distributed throughout the
and heterogeneous mixtures
mixture, it is called a homogenous mixture. If the
substances are unevenly distributed, it is called a
heterogeneous mixture.
Differences in properties such as density, particle size,
o Describe the processes of filtration,
molecular polarity, boiling point, freezing point, and
distillation, and chromatography and
solubility allow physical separation of the components of
the types of mixtures they are used
the mixture.
to separate
o
1.
o
o
2.
o
o
3.
o
4. o
o
o
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5.
o
6. o
V. CHEMICAL BONDING
Knowledge
Application
Atoms bond with other atoms to
gain a stable electron configuration. o Determine the noble gas configuration an atom will
Noble gases are already stable and
achieve when bonding
tend to not bond/react.
When a chemical reaction takes place, existing bonds must be broken in order for new bonds (and
new compounds) to be formed.
When a bond is broken, energy is absorbed. When a bond is formed, energy is released.
Electron-dot diagrams (Lewis
structures) are used to represent
o Draw Lewis dot structures for any given element, ion, or
the valence electron arrangement in
compound
elements, ions, and compounds.
Chemical bonds are formed when
o Identify the type of bonding in a compound (ionic or
valence electrons are:
covalent), given the elements that make it up
 transferred from one atom to
o Demonstrate bonding concepts using Lewis Dot
another (ionic)
Structures (electron-dot diagrams) for ionic and covalent
compounds
 shared between atoms (covalent)
Ionic
compounds – after the transfer of electrons, the positive ion should
 mobile within a metal (metallic)
have
no
dots. The negative ion should have 8 dots around it. Put
Metals tend to react with nonmetals
brackets around the ions. Check the periodic table to find the charge
to form ionic bonds.
Nonmetals tend to react with other associated with the ion and place this charge outside of the brackets. *Be
sure to include coefficients if there are more than one of the same kind of
nonmetals to form molecular
ion.*
(covalent) bonds.
Covalent compounds – each atom in a covalent compound must end up
Ionic compounds containing
with 8 dots around it – except for hydrogen (only 2 dots). *If you use
polyatomic ions have both ionic
lines, remember that one line represents 2 electrons being shared.*
AND covalent bonds.
In a multiple covalent bond, more
o Draw electron-dot diagrams and give examples of
than one pair of electrons is shared
molecules with multiple covalent bonds
between two atoms.
Electronegativity indicates how
o Distinguish between and give examples of nonpolar
strongly an atom of an element
covalent bonds and polar covalent bonds
attracts electrons in a bond.
If two atoms of the same element share electrons, the bond is
The electronegativity difference
nonpolar (ex: H–H)
between two bonded determines the
If atoms of two different elements share electrons, the bond is polar
degree of polarity in the bond.
(ex: H–Cl)
o Molecular polarity is determined
by the distribution of electrons in a
covalent compound.
o Symmetrical distribution of
electrons results in (nonpolar)
7.
molecules (Ex: CO2, CH4 and all
diatomic elements)
o Asymmetrical distribution of
electrons results in (polar)
molecules (Ex: HCl, NH3, H2O)
o Distinguish between bond polarity and molecular polarity
o Draw Lewis Dot Structures for all of the compounds
listed to the left, including the diatomic elements
o Determine whether a molecule is polar or nonpolar,
given its structure
*SNAP*
o
o
o
1. o
o
o
o
2.
o
o
3.
o
o
4.
o
o
5.
1.
VI. CHEMICAL FORMULAS, REACTIONS & STOICHIOMETRY
Knowledge
Application
Chemical formulas are used to represent compounds.
The main types of chemical formulas include: empirical,
molecular, and structural.
o Determine the empirical
An empirical formula is the simplest whole-number ratio of
formula from a molecular
atoms in a compound.
formula
Molecular formulas are chemical formulas that show the actual
o Draw structural formulas for
ratio of atoms in a molecule of that compound.
covalent (molecular)
Structural formulas can also be used to represent covalent
compounds
compounds. These use lines to show covalent bonds between
atoms and also show the geometrical arrangement of atoms in
the compound.
o Calculate the molar mass
One mole of any substance is equal to 6.02 x 1023 pieces of that
(gram-formula mass) of a
substance.
substance
The formula mass of a compound is equal to the sum of the
o Determine the molecular
atomic masses of its atoms (units are atomic mass units)
formula, given the empirical
The molar mass (gram-formula mass) of a substance is equal
formula and the molar mass
to the formula mass in grams – hence “gram-formula mass”.
o Determine the number of
The mass of one mole of any substance is equal to its molar
moles of a substance, given its
mass (gram-formula mass).
mass and vice versa
o Calculate the percent
composition of any element in
The percent composition by mass of each element in a
a given compound
compound can be calculated mathematically.
o Calculate the percent
composition of water in a
given hydrate
o Balance equations, given the
Balanced chemical equations show conservation of matter,
formulas for reactants and
energy, and charge.
products
The coefficients in a balanced equation can be used to
o Calculate simple mole-mole
determine mole ratios in the reaction.
ratios, given balanced
equations
Types of chemical reactions include synthesis,
o Identify the different types of
decomposition, single replacement, and double
chemical reactions, given their
replacement.
chemical equations
VII. PHYSICAL BEHAVIOR OF MATTER
Knowledge
Application
o Physical properties of substances can be explained in
o Predict relative melting and boiling
terms of chemical bonds and intermolecular forces.
points of compounds given
o Intermolecular forces created by an unequal distribution
information about its chemical
of charge result in varying degrees of attraction between
bonds or strength of
molecules. Hydrogen bonding is an example of a
intermolecular forces
strong intermolecular force that occurs in compounds
o Predict relative strength of
containing H bonded to F, O, or N atoms.
intermolecular forces of a
o Physical properties include malleability, solubility,
compound given its melting and
hardness, melting/freezing point, and boiling point.
boiling points
o
2.
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3.
o
o
4.
o
5.
6.
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7.
o
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8.
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9.
10.
o
VII. PHYSICAL BEHAVIOR OF MATTER – SOLIDS, LIQUIDS, AND GASES
Knowledge
Application
o Draw and interpret a particle
The three phases of matter (solids, liquids, and gases)
diagram to differentiate among
have different properties.
solids, liquids, and gases
The structure and arrangement of particles and their
o Explain phase change in terms of
interactions determine the physical state (s, l, or g) of a
the changes in energy and
substance at a given temperature and pressure.
intermolecular distances
Heat is a transfer of energy (thermal energy) from a
body of higher temperature to a body of lower
o Distinguish between heat energy
temperature. Heat (thermal energy) is associated with the
and temperature in terms of
random motion of atoms and molecules.
molecular motion and amount of
(Heat flows from HOT materials to COLD materials)
matter
Temperature is a measure of the average kinetic
o Convert between degrees Celsius
energy of the particles in a sample of matter.
and degrees Kelvin (Table T)
Temperature is NOT energy – it is a measure of heat
energy.
o Identify areas of heating and
The concepts of kinetic and potential energy can be
cooling curves that show changes
used to explain physical processes that include: fusion
in kinetic and potential energy,
(melting), solidification (freezing), vaporization (boiling,
heat of vaporization, heat of
evaporation), condensation, sublimation, and
fusion, and phase changes
deposition.
o Calculate the heat involved in a
The kinetic and potential energy changes involved in
phase or temperature change of a
these physical processes can be illustrated in a heating
sample of matter using Tables B
curve or cooling curve.
& T and/or a given heating or
cooling curve
Physical processes like phase changes can be exothermic o Distinguish between endothermic
or endothermic.
and exothermic phase changes
Entropy is a measure of the randomness or disorder of a o Compare the entropy of different
system. A system with greater disorder has more entropy.
phases of matter
The concept of an ideal gas is a model to explain
o Given a choice of pressure and
behavior of gases.
temperature conditions, identify
A real gas is most like an ideal gas when the real gas is at
those under which gases behave
low pressure and high temperature.
most ideally and/or least ideally
The Kinetic Molecular Theory (KMT) states that, for an IDEAL gas, all gas particles
a. are in random, constant, straight-line motion
b. are separated by great distances relative to their size (have negligible volume)
c. have no attractive forces between them
d. have collisions that may result in a transfer of energy between them, but the total energy of the
system remains constant
o Explain the gas laws in terms of
KMT.
The Kinetic Molecular Theory (KMT) describes the
o Solve problems using the
relationships of pressure, volume, temperature, velocity,
combined gas law (Table T)
frequency, and force of collisions among gas molecules.
o Recognize and draw graphs
showing P vs. T, V vs. T, and P vs.
V
Equal volumes of gases at the same temperature and pressure contain an equal number of
particles.
11.
o
o
o
12.
o
o
13.
o
14.
o
15.
o
VII. PHYSICAL BEHAVIOR OF MATTER – AQUEOUS SOLUTIONS
Knowledge
Application
Physical processes, such as a compound dissolving in a
o Interpret ∆H values for physical
solution, can be exothermic or endothermic.
processes given in Table I
o Identify the solute and the
A solution is a homogeneous mixture of a solute
solvent in a given solution
dissolved in a solvent.
o Give examples of different types
of solutions
The solubility of a solute in a given amount of solvent is
dependent on the temperature, the pressure, and the
o Predict the effect of temperature,
chemical natures of the solute and solvent.
pressure, and nature of solvent on
General rules:
solubility for a given solute
a. solubility of a solid increases as temperature increases o Use a solubility curve to
(direct relationship)
distinguish among unsaturated,
b. solubility of a gas decreases as temperature increases
saturated, and supersaturated
(inverse relationship)
solutions
c. solubility of a gas increases as pressure increases
o Calculate the amount of a specific
(direct)
solute dissolved at different
d. “like dissolves like” – polar solvents dissolve polar
temperatures using Table G
solutes; nonpolar solvents dissolve nonpolar solutes
o Use Table F (Solubility Guidelines)
to determine a compound’s
solubility
Many chemical reactions happen in solution. When
o Determine if a precipitate will
different ionic compounds are mixed together in the same
form when ionic compounds are
solution, a double replacement reaction may occur and a
mixed in solution
stable precipitate (insoluble/solid compound) may form.
o Write and balance chemical
equations for double replacement
reactions
o Calculate solution concentrations
in molarity (M), percent by
volume, percent by mass, or parts
per million (ppm)
The concentration of a solution may be expressed as:
o Describe how you would prepare a
molarity (M), percent by volume (%v/v), percent by
solution from scratch, given the
mass (%m/v), or parts per million (ppm).
desired molarity
o Describe how you would dilute a
solution of known concentration
(must use the equation M1V1 =
M2V2)
The addition of a nonvolatile solute to a solvent causes
the boiling point of the solution to increase and the
o Compare the freezing and boiling
freezing point of the solution to decrease.
points of solutions of different
The greater the concentration of solute particles, the
concentration
greater the increase in b.p. and decrease in f.p.
1.
o
o
2. o
o
o
3.
o
o
o
4.
o
o
5.
o
o
6.
o
7. o
VIII. KINETICS AND EQUILIBRIUM
Knowledge
Application
The Collision Theory states that a chemical reaction is most likely to occur if reactant particles
collide with the proper energy and orientation.
o Use the Collision Theory to explain how
factors such as temperature, surface area, and
The rate (speed) of a chemical reaction depends
concentration influence the rate of reaction
on several factors: temperature, concentration,
Ex: Increasing the temperature, surface area, or
nature of reactants, surface area, and the
concentration all lead to an increase in the rate of a
presence of a catalyst.
reaction because they all increase the number of
Ionic compounds generally react faster than
effective collisions between reactant
covalent (molecular) compounds
particles.
A catalyst provides an alternate reaction
o Explain, in terms of the number of bonds
pathway, which has lower activation energy
broken, why ionic compounds generally react
than an uncatalyzed reaction.
faster than covalent compounds
o Explain how a catalyst speeds up a reaction
o Read and interpret a potential energy diagram
Energy released or absorbed during a chemical
o Draw and label the following parts of a
reaction can be represented by a potential
potential energy diagram for both an
energy diagram.
endothermic and exothermic reaction
The difference in PE of the products and
 PE of reactants and PE of products
reactants is called the heat of reaction (H)
 heat of reaction (H)
H = PE products – PE reactants
 activation energy (for both the forward
H values for many chemical reactions are
and reverse reactions)
listed in Table I
 activation energy with a catalyst present
Chemical and physical changes can reach
equilibrium
o Distinguish between examples of physical
Saturated solutions are examples of systems in
equilibria and chemical equilibria
physical equilibria (aq ↔ s)
o Describe what is happening to the
At equilibrium, the rate of the forward reaction
concentrations or amounts of reactants and
equals the rate of the reverse reaction and the
products in a system at equilibrium
measurable quantities of reactants and products
o Describe the rates of opposing reactions in a
remain constant at equilibrium*CARE*
system at equilibrium
LeChatelier’s principle can be used to predict
the effect of a stress (such as a change in
o Describe, in terms of LeChatelier’s principle,
pressure, volume, concentration, or
the effects of stress on a given system at
temperature) on a system at equilibrium.
equilibrium, including:
According to LeChatelier’s principle, a system at
 Changing the
equilibrium will “shift” to reduce the effects of
temperature/heating/cooling
a stress placed on the system. It will “shift”
AWAY from an INCREASE and will
 Changing the concentration of a reactant
“shift”toward a decrease in concentration or
or product
temperature (“shift” means that either the
 Changing the pressure or volume (this
forward or the reverse reaction will be
affects systems involving gases)
“favored” (go faster) until the rates are again
o Also be able to explain why any shifting
equal and equilibrium is re-established).
occurs in terms of Collision Theory
Changing the pressure or volume only affects
systems that contain gases
Systems in nature tend to undergo changes toward lower energy and higher entropy.
o
1.
o
o
2.
o
o
IX. ORGANIC CHEMISTRY
Knowledge
Application
Organic compounds contain carbon atoms which bond to one another in chains, rings, and
networks to form a variety of structures.
Organic compounds are named using specific IUPAC rules and can be represented using molecular
formulas, structural formulas, or condensed structural formulas.
o Draw structural formulas for alkanes,
Hydrocarbons are organic compounds that
alkenes, alkynes, given their IUPAC names
contain only carbon and hydrogen. *The C–H
o Name a hydrocarbon (IUPAC rules), given
bonds are considered to be nonpolar covalent
its molecular formula, structural formula, or
bonds.*
condensed structural formula
Saturated hydrocarbons contain only single
o Identify whether a hydrocarbon is saturated
C–C bonds.
or unsaturated, given its IUPAC name,
Unsaturated hydrocarbons contain at least
structural formula, condensed structural
one double or triple carbon-carbon bond.
formula, or general formula (see Table Q)
o In a multiple covalent bond, more than one pair
of electrons is shared between two atoms. In a
3.
structural formula, each line represents TWO
shared electrons
o Determine the TOTAL number of electrons
shared in a covalent bond
o Determine the number of electron PAIRS
(lines) shared in a covalent bond
o A functional group is a group of atoms
attached to an organic compound that gives
distinct physical and chemical properties to
organic compounds having that group attached
to it.
o Organic acids, alcohols, esters, aldehydes,
ketones, ethers, halides, amines, amides,
and amino acids are types of organic
4.
compounds that differ in the type of functional
group they have.
o Compounds that have the same functional
group have similar physical and chemical
properties
Ex: all esters have pleasant odors, all organic acids
donate H+ ions in solution, all alcohols have low boiling
points, etc.
o Identify different kinds of functional groups
o Classify an organic compound based on its
structural formula, condensed structural
formula, or IUPAC name
o Draw a structural formula with the functional
group(s) on a straight chain hydrocarbon
backbone, given the correct IUPAC name for
the compound
o Name any of these organic compounds, given
their structural or condensed structural
formulas (see Table R)
o Isomers are organic compounds that have the
5.
same molecular formula, but different structures
and properties.
o Types of organic reactions include:
polymerization, substitution, fermentation,
6.
addition, combustion, esterification, and
saponification. *P.S. FACES*
o Determine if two compounds are isomers,
given their molecular formulas, structural
formulas, condensed structural formulas, or
names
o Identify types of organic reactions, given
balanced chemical equations
o Determine missing reactants or products in a
balanced equation, given the type of reaction.
o
1.
o
o
2.
o
o
o
3.
o
o
4.
o
o
o
5.
o
6.
o
o
7.
X. OXIDATION-REDUCTION REACTIONS
Knowledge
Application
An oxidation-reduction (redox) reaction
involves the transfer of electrons from one
o Determine the number of moles of electrons
species (element or ion) to another.
lost or gained in a redox reaction, given the
The number of electrons lost equals the number
other value
of electrons gained (conservation of charge)
o Assign oxidation states to atoms and ions
Oxidation numbers (states) can be assigned to o Determine if a reaction is a redox reaction
atoms and ions.
given the reaction equation (Hint: any reaction
Changes in oxidation numbers indicate that a redox
in which an element is alone (uncombined with
reaction has occurred.
another element) on one side, but in a compound on
the other side – it’s a redox reaction!)
Losing Electrons is Oxidation (LEO)
o Determine which species undergoes
Gaining Electrons is Reduction (GER)
reduction (oxidation state goes down) and
Oxidized and reduced species are ALWAYS on
which species undergoes oxidation
the LEFT (reactants) side of the equation.
(oxidation state goes up)
o Determine if a given half-reaction is showing
An oxidation half-reaction shows which
oxidation or reduction
species is oxidized and the number of electrons it
o Write and balance oxidation and reduction
loses (electrons go on the right side of the arrow)
half-reactions (*remember conservation of mass
A reduction half-reaction shows which species
and charge – multiply one or both of the halfis reduced and the number of electrons it gains
reactions to make electrons lost = electrons gained if
(electrons go on the left side of the arrow)
they are not equal at first)
An electrochemical cell can either be a voltaic
cell (a battery) or an electrolytic cell.
o Explain, in terms of atoms and ions, why the
In both voltaic and electrolytic cells
anode loses mass and the cathode gains mass
 oxidation occurs at the anode (An Ox)
o Compare and contrast voltaic cells with
 reduction occurs at the cathode (Red Cat)
electrolytic cells
 the anode loses mass
 the cathode gains mass
o Given a diagram of a voltaic cell, identify
and label the cathode, anode, salt bridge, and
A voltaic cell spontaneously converts chemical
the direction of electron flow
energy into electrical energy.
o Explain the function of the salt bridge and
The purpose of the salt bridge is to allow for
the direction of positive and negative ion
the migration of ions between half-cells
migration
o Write balanced oxidation and reduction halfreactions
An electrolytic cell requires electrical energy to
produce a chemical change. Electrolytic cells can o Given a diagram of an electrolytic cell,
be used for electrolysis (splitting a compound
identify and label the cathode, anode, and
into its elements) and for electroplating
direction of electron flow.
(coating something with a metal).
XI. ACIDS, BASES, AND SALTS
Knowledge
o The behavior of many acids and bases can be
explained by the Arrhenius Theory.
1.
o Arrhenius acids produce H+ (hydrogen
ions) as the only positive ions in aqueous
solution. The hydrogen ion may also be
written as H3O+ and called the hydronium
ion.
o Arrhenius bases produce OH– (hydroxide
ions) as the only negative ion is aqueous
solution.
(Table E)
o Arrhenius acids, Arrhenius bases, and salts
(ionic compounds) are all electrolytes. An
electrolyte is a substance which, when
dissolved in water, forms a solution capable of
conducting an electric current (electricity).
2.
Electrolytes can conduct electricity because they
ionize (break apart into ions) in a solution.
o The ability of a solution to conduct an electric
current depends on the concentration of the
ions in it (more ions, more conduction).
o In the process of neutralization, an Arrhenius
acid and an Arrhenius base react to form a salt
3.
and water.
Acid + Base  Salt + Water
o Titration is a laboratory process in which a
volume of solution with a known concentration
4.
is added to another solution of unknown
concentration. Titrations are done to determine
the concentration of the unknown solution.
Application
o Know the definitions of Arrhenius acids and
bases
o If given the properties, chemical formula, or
name, identify a substance as an Arrhenius
acid or Arrhenius base. (Use Tables K, L, and T
to help you remember these definitions. Arrhenius
acids begin with H, Arrhenius bases are metals +
hydroxide ion(s). *Don’t be fooled by alcohols, which
also end in OH, but contain covalent bonds and do not
ionize like bases do in solution. ALCOHOLS
ARE NOT BASES!)
o Given names or chemical formulas, identify
acids, bases, and salts as being electrolytes
o Determine the relative strength (strong or
weak) of an electrolyte given information on
its ability to ionize in solution.
(Strong acids and strong bases are strong electrolytes –
Tables K and L list acids and bases in order from
strongest to weakest.
If a salt is soluble, it is a strong electrolyte – Table F
can be used to determine the solubility of different
salts.)
o Recognize neutralization reactions when
given the reaction equation
o Write neutralization reactions when given the
reactants. (Remember that this is a double
replacement reaction. Just switch the positive ions,
look up their charges and cross down the subscripts if
needed. Then balance the equation.)
o Calculate the concentration or volume of a
solution, using titration data using the
equation
Ma x Va = Mb x Vb
(This equation is on Table T)
5.
o There are alternate acid-base theories. One
such theory states that the acid is a proton
donor (H+ donor) and the base is a proton
acceptor.
o The acidity or alkalinity of an aqueous
solution can be measured using the pH scale.
o The pH scale measures the concentration of
H+/H3O+ in a solution.
[H+] = 10–pH
A pH of 1 means that the [H+] = 10–1 = 0.1M
A pH of 3 means that the [H+] = 10–3 = 0.001M
 Acids have pH values between 0 and 7
(the stronger the acid, the lower the pH and the more
6.
H+)
[H+]  [OH–]
 Neutral solutions have a pH of 7
[H+] = [OH–]
 Bases have pH values between 7 and 14
(the stronger the base, the higher the pH and the
more OH–)
[H+]  [OH–]
o The pH scale is a logarithmic scale, which
means that a change of one pH unit changes the
concentration of H+/H3O+ by a factor of ten
 tenfold = 10 times = 101
 hundredfold = 100 times = 102
7.
 thousandfold = 1000 times = 103
The exponents represent the CHANGE in pH
o If a solution becomes more acidic,
the pH , and the [H+]/[H3O+] 
o If a solution becomes more basic,
the pH , and the [H+]/[H3O+] 
o The pH of a solution can be shown by using
indicators.
8. o An indicator is a substance that changes color
depending on the concentration of
hydrogen/hydronium ions in a solution.
o Give the alternate definitions of acids and
bases
o Use this definition to explain why ammonia is
considered a base
o Identify a solution as acidic, basic
(alkaline), or neutral based upon the pH
value OR the relative concentrations of
H+/H3O+ and OH–
o Describe acidic, basic, and neutral solutions in
terms of pH value and relative H+/H3O+ and
OH– concentrations
o Differentiate between strong acids/bases and
weak acids/bases given pH values or ion
concentrations
o Determine the new pH value of a solution
given the starting pH and the amount of
increase or decrease in [H+]/[H3O+] (such as
tenfold, a hundredfold, or a thousandfold)
Ex: A lake with an initial pH of 6 has been affected by
acid rain. The acid rain has caused a hundredfold change
in the [H+] concentration of the lake. What is the new
pH of the lake?
Answer: pH = 4
o Determine the amount that the [H+]/[H3O+]
would increase or decrease given a certain
change in pH
o Interpret changes in acid-base indicator color
o Explain how different indicators can be used
to distinguish between solutions with
different pH values
o Identify appropriate indicators that can be
used to show changes in pH values, such as
during a titration, given starting and ending
pH values
XII.
NUCLEAR CHEMISTRY
Knowledge
Application
o The stability of an isotope is based on the ratio of the neutrons and protons in the nucleus.
o Usually when the ratio is not 1:1, the nucleus gets a little unstable and starts spitting out particles so
1.
that it will have a more stable 1:1 ratio.
o Although most nuclei are stable, some are unstable and spontaneously emit radiation. We call
these unstable isotopes radioactive isotopes, radioisotopes, or nuclides.
o Spontaneous decay (natural emission of
radiation) by a nuclide (radioactive isotope)
involves the release of particles and/or energy
from the nucleus.
o Each radioactive isotope has a specific decay
mode (the kind of particle or energy it gives off
from its unstable nucleus) (Tables N and O!)
2.
 alpha decay: release of alpha particles
 beta decay: release of beta particles
 positron emission: release of positrons
 gamma radiation : release of gamma rays
o These emissions differ in mass, charge, ionizing
power, and penetrating power.
o Each radioactive isotope has a specific half-life
(rate of decay). The half-life is the time it takes
3.
for half of the radioisotope to decay/transmutate
into something more stable). (Table N)
o Determine decay mode and write nuclear
equations showing alpha decay, beta decay,
positron emission, and gamma radiation
(*Remember to put radioactive emissions on the
RIGHT side of the arrow – if something is
released, it goes on the right)
o Compare and contrast the 4 different types
of radiation in terms of mass, charge,
ionizing power, and penetrating power.
o Calculate the initial amount, the fraction
remaining, time elapsed, or the half-life of a
radioactive isotope, given the other variables
o Nuclear reactions are represented by equations
that include symbols for elements and radioactive o Complete nuclear equations and predict
emissions (with mass number in upper left and
missing particles in nuclear equations
4.
charge/atomic number in lower left)
o Write nuclear equations given word
o These reactions show conservation of mass and
problems
charge
o A change in the nucleus of an atom that changes
it from one element to another is called
5.
transmutation. This can occur naturally or can
be done artificially by bombarding the nucleus
with high-energy particles.
o Distinguish between natural
transmutation (one reactant) and artificial
transmutation (two reactants) given
nuclear equations
o Types of nuclear reactions include fission and
6.
fusion. Fission and fusion can be natural or
artificial transmutations.
o Compare and contrast fission and fusion
reactions.
o Distinguish between fission and fusion
reactions given nuclear equations
o Nuclear changes convert matter into energy
(E = mc2)
7. o Energy released during nuclear reactions is much
greater than the energy released during chemical
reactions.
o Compare and contrast chemical reactions
and nuclear reactions
o Describe benefits of using nuclear fission
o There are risks and problems associated with
radioactivity and the use of radioactive isotopes,
including: biological exposure, long-term storage
8.
and disposal problems, and nuclear accidents
which release radioactive materials into the
environment.
o Describe the risks and problems associated
with using radioactive isotopes
o In addition to using nuclear fission for nuclear power, radioactive isotopes have other beneficial uses
in medicine and industrial chemistry, including:
 radioactive dating (ages of once-living things can be found from the ratio of C-14 to C-12 in the remains; ages of
rocks can be found from the ratio of U-238 to Pb-206)
9.
 tracing chemical and biological processes (radioactive tracers can be injected into the body and then x-rayed.
The radioactive substance will show up on the x-ray and if there are problems, they can be detected easily)
 detecting and treating of disease (Sr-90: diagnosing and treating bone cancer; I-131: diagnosing and
treating thyroid disorders; Co-60: cancer treatment
 radiation can be used to kill bacteria in foods (used with spices, meats, produce)