Chemistry EOC Review Key Concepts by TEK Must Knows!!! Adapted
from Stemscopes, 2012, Rice University
Slide 2
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.
Slide 3
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
Slide 4
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.
Slide 5
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.
Slide 6
States of Matter Student Expectation: compare solids, liquids,
and gases in terms of compressibility, structure, shape, and
volume.
Slide 7
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
Slide 8
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.
Slide 9
States of Matter PhaseEnergy move Particle spacing
VolumeShapeCompressible SolidLittleClose Rigid Definite No
LiquidSomeClose Slide DefiniteNot Definite No GasLotFarNot Definite
Not Definite Yes
Slide 10
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.
Slide 11
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.
Slide 12
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.
Slide 13
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.
Slide 14
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.
Slide 15
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).
Slide 16
Periodic Trends Student Expectation: to use the Periodic Table
to identify and explain periodic trends, including atomic and ionic
radii, electronegativity, and ionization energy
Slide 17
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.
Slide 18
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.
Slide 19
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
Slide 20
Periodic Trends electronegativity - ability to attract
electrons. increases moving from left to right Decreases moving
down a Group, making fluorine the most electronegative
element.
Slide 21
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
Slide 22
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.
Slide 23
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.
Slide 24
Atomic Theory John Dalton (1800s) 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.
Slide 25
Atomic Theory J.J. Thomson (late 1800s) 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.
Slide 26
Atomic Theory Ernest Rutherford (early 1900s) developed the
nuclear model of the atom Rutherfords 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.
Slide 27
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.
Slide 28
Atomic Equations Student Expectation: to understand the
electromagnetic spectrum and the mathematical relationships between
energy, frequency, and the wavelength of light c = x f To calculate
the wavelength, frequency, and energy of light using Planck's
constant and the speed of light. E = hf or E = hc/
Slide 29
Atomic Equations Key Concepts The electromagnetic spectrum
displays the full range of electromagnetic energy based on wave
properties, from high-energy gamma rays to low- energy 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
Slide 30
Atomic Equations c = x f All energy waves travel at the same
velocity known as the speed of light (c), which equals 3.0 x 10 8
m/s. Wavelength () 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
Slide 31
Atomic Equations E = hf or E = hc/ Plancks constant equals 6.63
x 10 -34 Js the energy of an electromagnetic wave is directly
proportional to its frequency
Slide 32
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.
Slide 33
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.
Slide 34
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.
Slide 35
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.
Slide 36
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.
Slide 37
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.
Slide 38
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.
Slide 39
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.
Slide 40
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.
Slide 41
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
Slide 42
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.
Slide 43
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
Slide 44
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
Slide 45
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.
Slide 46
Molecular Structure Student Expectation: to predict molecular
structure Linear (180 o ), trigonal planar (120 o ), or tetrahedral
(109.5 o ), using Valance Shell Electron Pair Repulsion (VSEPR)
theory.
Slide 47
Molecular Structure Key Concepts Linear (180 o ), trigonal
planar (120 o ), or tetrahedral (109.5 o ) - o r bent
Slide 48
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
Slide 49
The Mole Key Concepts The mole is to describe an amount of a
substance. The mole is equal to 6.02 10 23 atoms, molecules, or
formula units of a substance. Is defined by the number of atoms in
exactly 12 grams of carbon-12.
Slide 50
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)
Slide 51
Percent Composition Calculations Student Expectation: to
calculate percent composition and empirical and molecular
formulas.
Slide 52
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 1 st : Convert the
masses of each element in the compound into mole ratios 2 nd :
Divide all by the smallest mole 3 rd : Write the smallest whole
number ratios with each element
Slide 53
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. 1 st : The molar
mass of the empirical formula is calculated 2 nd : The molar mass
of the compound is divided by the molar mass of the empirical
formula to find a whole number integer. 3 rd : The empirical
formula is then multiplied by this integer to calculate the
molecular formula.
Slide 54
Chemical Equations Student Expectation: to use the law of
conservation of mass to write and balance chemical equations.
Slide 55
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
Slide 56
Chemical Equations How to balance 1 st : 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 2 nd : Use
coefficients in front of the chemical formulas to make the number
of atoms on each side of the arrow equal.
Slide 57
Stoichiometry Student Expectation: to perform stoichiometric
calculations determination of mass relationships between reactants
and products calculation of limiting reagents percent yield.
Slide 58
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.
Slide 59
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.
Slide 60
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 =
actualx 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.
Slide 61
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.
Slide 62
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.
Slide 63
Gas Laws Boyles law, P 1 V 1 = P 2 V 2 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.
Slide 64
Gas Laws Charless law, V 1 /T 1 = V 2 /T 2 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.
Slide 65
Gas Laws Avogadros law, V 1 /n 1 = V 2 /n 2 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.
Avogadros principle states that equal volumes of gases contain the
same number of atoms or particles.
Slide 66
Gas Laws Daltons law of partial pressures, P total = P 1 + P 2
+ P 3 +P n 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.
Slide 67
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 Boyles law, Charles
law, and Avogadros 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.
Slide 68
Gas Stoichiometry Student Expectation: to perform
stoichiometric calculations including determination of mass and
volume relationships between reactants and products for reactions
involving gases.
Slide 69
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.
Slide 70
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.
Slide 71
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.
Slide 72
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. waters high specific heat
plays a critical role in Earths relatively moderate temperature
variations.
Slide 73
Use STAAR Chart to determine Which Ionic compounds are soluble
(aq) and insoluble (s)
Slide 74
Molarity Student Expectation: to calculate the concentration of
solutions in units of molarity use molarity to calculate the
dilutions of solutions
Slide 75
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. M 1 V 1 = M 2 V 2,
Slide 76
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.
Slide 77
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.
Slide 78
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.
Slide 79
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
Slide 80
Acids and Bases Student Expectation: to define acids and bases
and distinguish between Arrhenius and Brnsted-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.
Slide 81
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 Brnsted-Lowry
definition of acids and bases an acid can donate a proton (H + ) a
base can accept a proton.
Slide 82
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
Slide 83
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
Slide 84
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.
Slide 85
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.
Slide 86
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.
Slide 87
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.
Slide 88
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.
Slide 89
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
Slide 90
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
Slide 91
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. H
is often expressed in kilojoules (kJ).
Slide 92
Thermochemical Equations Exothermic reaction, H 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, H is positive heat
flows into the system from the surroundings. The energy absorbed
during the reaction is represented as a reactant in the
equation.
Slide 93
Thermochemical Equations The H (change in enthalpy) of a system
is equivalent to the heat of the system (q) q = mc T q = m (mass in
grams) x c (specific heat) x T (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.
Slide 94
Calorimetry Student Expectation: to use calorimetry to
calculate the heat of a chemical process.
Slide 95
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.
