Chemistry: Content Knowledge (0245 or 5245)
I. Matter and Energy; Heat and Thermodynamics, and Thermochemistry
A. Matter and Energy
a. Organization of matter
i. Pure substances – one with constant composition
1. Elements – substances that cannot be decomposed into simpler substances by
chemical or physical means
2. Compounds – a substance with constant composition that can be broken down into
elements by chemical processes
ii. Mixtures – variable composition
1. Homogenous – having visibly indistinguishable parts
2. Heterogeneous – having visibly distinguishable parts
3. Solutions – a homogeneous mixture
4. Suspensions – heterogeneous mixture containing solid particles; will eventually
settle
iii. States of matter
1. Solid – rigid; fixed volume, shape
2. Liquid – definite volume, no specific shape; assumes shape of container
3. Gas – no fixed volume or shape; assumes shape and volume of container; very
compressible
4. Plasma – similar to gas, with a certain portion of particles ionized
b. Particulate structure of matter
i. Atoms – smallest unit of mass; composed of nucleus and electron(s)
ii. Ions – an atom or group of atoms that has a positive or negative charge
iii. Molecules – unit of atoms held together by equal sharing of electrons (covalent bonds)
c. Differences between chemical and physical properties and chemical and physical changes
i. Chemical vs. physical properties
1. Chemical property – a property that can only be observed by changing the identity
of the substance
a. Changes happen on a molecular level
2. Physical property – can be observed without changing the identity of the substance
a. About energy and states of matter
ii. Chemical vs. physical changes
1. Chemical change – changes that alter the identity of a substance
2. Physical change – changes in matter that do not alter the identity of the substance
iii. Intensive vs. extensive properties
1. Intensive – properties that do not depend on the amount of matter present
a. Examples: color, odor, luster, density, etc.
2. Extensive – properties that do depend on the amount of matter present
a. Examples: mass, weight, volume, length
d. Conservation of energy and the conservation of matter in chemical processes
i. Law of conservation of energy – energy may neither be created nor destroyed; the
sum of all energies in the system is a constant
ii. Law of conservation of matter – mass is neither created nor destroyed
e. Different forms of energy
i. Kinetic and potential
1. Kinetic energy – (½mv2) energy due to the motion of an object
2. Potential energy – energy due to position or composition
ii. Chemical, electrical, electromagnetic, nuclear, and thermal energy
1. Chemical energy – energy stored in the bonds of chemical compounds
2. Electrical energy – the flow of charges along a conductor (electricity)
3. Electromagnetic energy – energy that is reflected or emitted from objects in the
form of electrical and magnetic waves that can travel through space
4. Nuclear energy – energy of an atomic nucleus, which can be released by fusion or
fission or radioactive decay
5. Thermal energy – internal energy present in a system in a state of thermodynamic
equilibrium by virtue of its temperature
iii. Conversions between different forms of energy within chemical systems
1. Battery: chemical to electrical
2. Chemical explosion: chemical to kinetic and thermal
* Transducer – a device that converts one form of energy to another
B. Thermodynamics in Chemistry
a. Temperature, thermal energy, and heat capacity, including temperature scales, units of
energy, and calculations involving those concepts
i. Temperature and temperature scales
1. TK = TC + 273.15 [Kelvins]
2. TC = (TF - 32°F)5°C/9°F [degrees Celcius)
3. TF = (TC × 9°F/5°C) + 32°F [degrees Fahrenheit]
ii. Heat transfer – heat can be transferred from one place to another by three methods:
conduction in solids, convection in fluids (liquids or gases), and radiation through
anything that will allow radiation to pass. It there is a temperature difference in a
system, heat will always move from higher to lower temperatures
iii. Heat capacity and specific heat
1. Heat capacity – (C) measureable physical quantity that characterizes the amount of
heat required to change a substance’s temperature by a given amount (C = heat
absorbed/increase in temperature)
2. Specific heat – the heat capacity per gram of substance [J/°C×g or J/K×g]
iv. Calorimetry – the science of measuring heat flow
1. Heat capacity, C = q/ΔT
2. Specific heat capacity, q = mcΔT
3. ΔE = q + w, where w = -PΔV
b. Concepts and calculations involving phase transitions between the various states of matter
i. Phase transitions and diagrams
1.
Critical point – the point where
the critical temperature (temperature
above which the vapor cannot be
liquefied no matter what pressure is
applied) and the critical pressure
(pressure required to produce
liquefaction at critical temperature)
intersect
2.
Triple point – substance in
equilibrium as solid, liquid, and gas
3.
Supercritical – fluid beyond the
critical point, with characteristics of
both gas and liquid
4. Water has a negative solid/liquid slope because the density of ice is less than that of
liquid water at the melting point; the maximum density of water occurs at 4°C; when
liquid water freezes, its volume increases
ii. Heats of vaporization, fusion, and sublimation
1. Heat of vaporization – energy required to vaporize 1 mole of a liquid at a pressure
of 1 atm; a.k.a. enthalpy of vaporization, ΔHvap (liquid to gas; endothermic because
energy is required to overcome the relatively strong intermolecular forces)
2. Heat of fusion- enthalpy change associated with melting (solid to liquid)
3. Heat of sublimation – enthalpy change associated with sublimation (solid to gas)
iii. Heating curves (*equations for water)
1.
A – q = mciceΔT
2.
B – ΔHfusion
3.
C – q = mcwaterΔT
4.
D –ΔHvaporization
5.
E - q = mcsteamΔT
*Heat added = q
c. Kinetic molecular theory and ideal gas law
i. Assumptions of kinetic molecular theory – simple model that attempts to explain the
properties of an ideal gas
1. Gas molecules have zero volume
2. Gas molecules exert no forces on each other
3. Gas molecule make complete elastic collisions
4. The average kinetic energy is directly proportional to the temperature
ii. Ideal gases and ideal gas laws (e.g., applications, calculations)
1. Ideal gas law: PV = nRT (Pressure, Volume, number of moles, Temperature)
2. (KE)avg = 3/2RT
3. Partial pressure (Pa) = χaPtotal (Χa = mole fraction of particular gas)
iii. Real gas behavior
1. When molecules are close together, the volume becomes significant and the
electrostatic forces increase and become significant
2. High pressure and low temperature push molecules together
3. Vreal > Videal
4. Preal < Pideal
d. Energetics of chemical reactions
i. Exothermic and endothermic reactions
1. Exothermic – refers to a reaction where energy (as heat) flows out of the system
2. Endothermic – refers to a reaction where energy (as heat) flows into the system
ii. Bond energy; Hess’s law
1. Bond energy – the energy required to break a given chemical bond
2. Hess’s law – ΔH is not dependent on the reaction pathway
e. How the laws of thermodynamics relate to chemical reactions and phase changes
i. Laws of thermodynamics
1. The energy of the universe is constant
2. In any spontaneous process there is always an increase in the entropy of the
universe
3. The entropy of a perfect crystal at 0 K is zero
ii. Spontaneous/reversible processes
1. Spontaneous – a process that occurs without outside intervention; may be fast or
slow
2. Reversible – a cyclic process carried out by a hypothetical pathway, which leaves
the universe exactly the same as it was before the process (no real process is
reversible)
iii. Change in enthalpy, entropy, and Gibbs energy in chemical/physical processes
1. Enthalpy – ΔH is independent of pathway; if the reaction is reversed, the sign of ΔH
is also reversed; the magnitude of ΔH is directly proportional to the quantities of
reactants and products in a reaction (if the coefficients of a balanced reaction are
multiplied by an integer, the value of ΔH is multiplied by the same integer)
2. Entropy – the sign of ΔSsurr depends on the direction of the heat flow (+ =
exothermic) and the magnitude of ΔSsurr depends on the temperature (ΔSsurr=-ΔH/T)
ΔSsys
ΔSsurr
ΔSuniv Spontaneous?
Yes
+
+
+
No (opposite direction)
Yes, if ΔSsys > ΔSsurr
+
?
Yes, if ΔSsurr > ΔSsys
+
?
3. Gibbs Free Energy – ΔG=ΔH–TΔS; ΔG tells us the eventual equilibrium position (the
more negative ΔG, the further a reaction will go to the right to reach equilibrium);
ΔG is dependent on pressure [G=G°+RTln(P)]
II. Atomic and Nuclear Structure
A. Current model of atomic structure
a. Description of atomic model (e.g., subatomic particles, orbitals, quantum numbers)
i. Tiny nucleus (protons and neutrons), with electrons moving around it
ii. Protons have a positive charge; neutrons are the same size as protons, but are
uncharged; electrons have a negative charge, and are much smaller
iii. Orbitals – a specific wave function for an electron in an atom
iv. Quantum numbers – each orbital is characterized by a series of numbers
1. Principal quantum number (n) – has integral values: 1,2,3,…; related to the size
and energy of the orbital; as n increases, the orbital becomes larger and the electron
spends more time further from the nucleus; this increases also means higher energy,
because the electron is less tightly bound to the nucleus, and the energy is less
negative
2. Angular momentum quantum number (l) – has integral values from 0 to n-1 for
each value of n; related to the shape of atomic orbitals; l=0, s; l=1, p; l=2, d; l=3, f
3. Magnetic quantum number (ml) – has integral values between l and –l, including
zero; related to the orientation of the orbital in space relative to other orbitals in the
atom
*Number of orbitals per subshell: s=1, p=3, d=5, f=7, g=9
b. Experimental basis (e.g., cathode ray tube, gold foil experiment, spectral lines)
i. Cathode ray tube – J.J. Thomson studied electrical discharges in partially evacuated
tubes called cathode ray tubes; determined the charge to mass ratio of electrons, and
postulated that an atom consisted of a diffuse cloud of positive charge with negative
electrons embedded randomly in it (plum pudding model)
ii. Gold foil experiment – Ernest Rutherford tested the plum pudding model and
disproved it by postulating that the atom was nuclear (with a dense positive center with
electrons around it)
c. Isotopes – atoms of the same element with different numbers of neutrons
i. Mass number – the total number of protons and neutrons in the atomic nucleus of an
atom
ii. Average atomic mass – average total mass of protons, neutrons, and electrons in a
single atom
B. Electron configuration of the elements based on the periodic table
a. Aufbau principle, Hund’s rule, Pauli exclusion principle
i. Aufbau principle – as protons are added one by one to the nucleus to build up the
elements, electrons are similarly added to hydrogenlike orbitals
ii. Hund’s rule – the lowest energy configuration for an atom is the one having the
maximum number of unpaired electrons allowed by the Pauli exclusion principle in a
particular set of degenerate orbitals, with all unpaired electrons having parallel spins
iii. Pauli exclusion principle – in a given atom no two electrons can have the same set of
four quantum numbers
b. Correlation between electron configuration and periodic table
c. Relationship between electron configuration and chemical and physical properties
i. Valence electrons – the electrons in the outermost principal quantum level of an atom;
they are involved in bonding
C. Radioactivity
a. Characteristics of alpha particles, beta particles, and gamma radiation
i. Alpha particle – a helium nucleus
ii. Beta particle – an electron produced in radioactive decay
iii. Gamma particle – a high-energy photon
b. Radioactive decay processes; half life
i. Alpha-particle production – involves a change in mass number for the decaying
nucleus
ii. Beta-particle production – changes a neutron to a proton
iii. Gamma-particle production – release of excess energy by a nucleus
iv. Positron production – changes a proton to a neutron (positron = positive electron)
v. Half-life – the time required for number of nuclides in a radioactive sample to reach
half of the original value
c. Fission, fusion, and other nuclear reactions
i. Fission – the process of using a neutron to split a heavy nucleus into two nuclei with
smaller mass numbers
ii. Fusion – the process of combining two light nuclei to form a heavier, more stable
nucleus
d. Balancing nuclear reactions and identifying products of nuclear reactions
D. How the electronic absorption and emission spectra of elements are related to electron energy
levels
a. Electronic energy transitions in atoms (e.g., ground state, excited states,
emission/absorption of energy)
i. Ground state – the lowest possible energy state of an atom or molecule
ii. Excited state – one or more electrons is excited and in a higher orbital
b. Energy of electronic absorption/emission spectral lines in various regions of the
electromagnetic spectrum
i. Transitions between electronic energy levels, as either emission or absorption of light,
occur at discrete energies or wavelengths
ii. Emission  atom loses energy; energy becomes more negative and ΔE for the atom is
negative
iii. Absorption  atom gains energy; energy increases and ΔE is positive
c. Relationship between energy, frequency, and wavelength
i. Wavelength (λ) – distance between two consecutive peaks or troughs in a wave
ii. Frequency (ν) – number of waves (cycles) per second that pass a given point in space
iii. λ×ν=c (c=speed of light=3.0 × 108 m/s)
iv. ΔE = hν (h=Planck’s constant=6.26 × 10-34 J×s)
III. Nomenclature; the Mole, Chemical Bonding, and Geometry
A. Nomenclature and Chemical Composition
a. Systematic names and chemical formulas of simple inorganic compounds
i. Binary compounds
1. The cation is always named first and the anion second
2. A monatomic cation takes its name from the element
3. A monatomic acnion takes its root and adds –ide
4. Metals that form more than one cation must have the charge specified; the ion with
the higher charge has a name ending in –ic and the lower charged one –ous
5. Prefixes: mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca
ii. Acids, bases, and salts
1. Acids – if the anion does not contain oxygen, the acid is named with the prefix
hydro- and the suffic –ic; if the anion contains oxygen, and ends in –ate, the –ate is
replaced with –ic, and if it ends in –ite, the –ite is replaced by –ous
2. Bases and salts – use normal rules (ex. NaOH – sodium hydroxide)
iii. Hydrates – compounds that attract and bond with water molecules
1. Name the salt
2. Determine the prefix to be used based on the number of water molecules
3. Add the word “hydrate” to the end of the prefix
b. Names of common organic compounds based on functional groups
i. Alkanes, alkenes, and alkynes
1. Alkanes (single)– add the suffix –ane to the Greek root for the number of carbon
atoms (meth, eth, prop, but, pent, hex, hept, oct, non, dec); for a branched chain, use
the longest continuous chain and begin numbering so the largest substituent is
closed to the first position; substituent groups are named by dropping –ane for –yl;
add numbers and prefixes to denote multiple groups of the same kind (2,3Dimethylpentane)
2. Alkenes (double)– use –ene instead of –ane; the location of the double bond is
indicated by the lowest number carbon atom involved in the bond
3. Alkynes (triple)– use –yne instead of –ane
ii. Alcohols, ethers, ketones, aldehydes, amines
1. Alcohols (OH)– drop –e for –ol; give them the lowest carbon number
2. Aldehydes (COH) – drop –e for –al
3. Ketones (CO)– drop –e for –one
4. Carboxylic acid (COOH)– drop –e for –oic acid
5. Esters (COO)– drop –e for –oate
6. Ethers (R-O-R) – use alkoxy = alkyl – ky + oxy (ex. Methoxy)
7. Amine (NR3) – use –amine (Primary: NRH2; secondary: NR2H; tertiary: NR3)
c. Mole concept and how it applies to chemical composition
i. Avogadro’s number, molar mass, and mole conversions
1. Avogadro’s number – one mole of something consist of 6.02 x 1023 units
2. Molar mass – the mass in grams of one mole of the compound
3. In order to perform mole conversions, you must calculate the molar mass of
compounds and use it as a ratio with 1 mole
ii. Calculation of empirical and molecular formulas
1. Empirical formula – the base formula (not necessarily exact)
a. Start with the number of grams; if given percentages, convert to moles using 100
grams as your total
b. Convert the mass of each element to moles using molar mass
c. Divide each mole value by the smallest number of moles calculated
d. Round to the nearest whole number (now have mole ratio)
e. Multiply by integers if you have x.5
2. Molecular formula – the exact formula of the molecule (must know molar mass)
a. Molecular formula = empirical formula x “n”
iii. Percent composition
1. Calculate each element’s (within the molecule) contribution to the mass by
multiplying the molar mass by the number of moles; then add together all of the
contributions; finally, divide a specific element’s contribution by the total mass and
multiply by 100%
B. Bonding and Structure
a. Common properties of bonds
i. Relative bond lengths
1. As the number of shared electrons increases, the bond length shortens
ii. Relative bond strengths
1. Inversely related to bond length; as bond length increases, bond strength decreases
b. Bond types
i. Ionic bonding – an atom that loses electrons relatively easily reacts with an atom that
has a high affinity for electrons
ii. Covalent bonding (polar, nonpolar, hybridization) – electrons are shared by nuclei
1. Polar – unequal sharing
a. The ions form so that the valence electron configuration becomes a noble gas
configuration
2. Nonpolar – equal sharing
a. They share electrons in a way that completes the valence electron configurations
of both atoms
iii. Metallic bonding – mobile valence electrons are shared among atoms (in a usually
crystalline structure)
c. Structural formulas and molecular geometry (shape)
i. Lewis structures including formal charges – shows valence electrons, only
1. Sum the valence electrons from all the atoms
2. Use a pair of electrons to form a bond between each pair of bound atoms
3. Use the remaining electrons to satisfy the duet rule for hydrogen, and the octet rule
for the others
4. Formal charge – the difference between the number of valence electrons on the
free atom (use groups on PT) and the number of valence electrons assigned to the
atom in the molecule
a. Assigned valence electrons = # lone pair + ½ # shared electrons
ii. Resonance structures – occur when more than one valid Lewis structure can be
written for a particular molecule; use brackets and double-headed arrows to indicate
iii. Molecular geometry (shape and approximate bond angles)
d. Identify polar and nonpolar molecules
i. Analysis of bonding in the molecule
1. Can utilize electronegativity to determine polarity, as a great difference in
electronegativity will result in polar molecules
ii. Symmetry of molecular structure
e. Intermolecular interactions
i. Hydrogen bonding – unusually strong dipole-dipole attractions that occur among
molecules in which hydrogen is bonded to a highly electronegative atom (O, N, F)
ii. London dispersion forces – the forces, existing among noble gas atoms and nonpolar
molecules, that involve an accidental dipole that induces a momentary dipole in a
neighbor
iii. Dipole-dipole – the attractive force resulting when polar molecules line up so that the
positive and negative ends are close to each other
iv. Dipole-induced dipole – result when an ion or a dipole induces a dipole in an atom or
molecule with no dipole (weak)
f. How bonding and structure correlate with physical properties
i. The stronger the intermolecular attractions, the higher the temperature at which the
substance will boil/the less soluble the substance will be/the lower the equilibrium
vapor pressure/rate of evaporation/condensation at equilibrium/concentration of
vapor/vapor pressure
IV. Periodicity and Reactivity; Chemical Reactions; Biochemistry and Organic Chemistry
A. Periodicity
a. Basis of the periodic table and general layout
i. Arranged in groups and periods
1. Periods = rows
2. Groups = columns
ii. Atomic number – the number of protons found in the nucleus of an atom
iii. Symbols of the elements
iv. Metals, nonmetals, metalloids
1. Group 1: alkali metals
2. Group 2: alkaline earth metals
3. Metalloids: B, Si, Ge, As, Sb, Te, Po
4. Group 17: Halogens (nonmetals)
5. Group 18: Noble gases (nonmetals)
6. Other nonmetals: C, N, O, P, S, Se
v. Transitions elements
1. Group 3-12: Transition metals
b. Periodic trends in physical and chemical properties of the elements
i. Physical properties (e.g., boiling/melting points, conductivity)
ii. Chemical reactivity – those furthest from an octet will be most reactive
B. Chemical Reactions and Basic Principles
a. Balancing chemical equations
i. Simple chemical equations
1. Start with the most complicated molecule
ii. Chemical equations involving oxidation-reduction
1. Oxidation – loss of electrons or an increase in oxidation state
2. Reduction – gain of electrons or a decrease in oxidation state
3. Compute the oxidation reaction separately from the reduction reaction, and then
combine; make sure that the number of electrons matches
b. Stoichiometric calculations
i. Simple calculations based on balanced chemical equations involving moles, mass, and
volume
1. Balance the equation for the reaction
2. Convert the known mass of the reactant or product to moles of that substance
3. Use the balanced equation to set up the appropriate mole ratios – the ratio of moles
of one substance to moles of another substance in a balanced reaction
4. Use the appropriate mole ratios to calculate the number of moles of the desired
reactant or product
5. Convert from moles back to grams if required by the problem
ii. Limiting reagent calculations and percent yield
1. Limiting reagent – the reactant that is consumed first and which therefore limits
the amounts of products that can be formed
a. Convert mass to moles
b. Multiply each mole amount by the mole ratio; if the result is greater than the
number of moles available for the other substance, that other substance is the
limiting reactant
c. Use the mole amount of the limiting reactant to compute the moles of product
consumed by multiplying it by the mole ratio (product/limiting reactant)
d. Convert back to grams
2. Percent yield = actually yield/theoretical yield x 100%
a. Theoretical yield – the maximum amount of product formed, when the limiting
reactant is completely consumed
c. Identify, write, and predict products of simple reaction types
i. Combustion, neutralization
1. Combustion – the sequence of exothermic chemical reactions between a fuel and an
oxidant accompanied by the production of heat and conversion of chemical species
a. Ex. CH4 + 2 O2  CO2 + 2 H2O + energy
2. Neutralization – an acid and a base react to form a salt, usually with water
produced as a by-product
a. YOH + HX  XY + H2O
ii. Decomposition, dehydration
1. Decomposition – the separation of a chemical compounds into elements or simpler
compounds
a. AB  A + B
2. Dehydration – a chemical reaction that involves the loss of water from the reacting
molecule
a. 2 R-OH  R-O-R + H2O
iii. Single and double replacement
1. Single replacement – a type of redox reaction when an element or ion moves out of
one compound and into another
a. A + BC  AC + B
2. Double replacement – involves the exchange of bonds between the two reacting
chemical species
a. AB + CD  AC + BD
iv. Oxidation-reduction
d. Chemical kinetics
i. Rate laws, rate constants, and reaction order
1. Reaction rate – change in concentration of a reactant or product per unit time
2. Rate law – an expression that shows how the rate of reaction depends on the
concentration of reactants (Rate = -(ΔA/Δt) = k[A]n)
3. Rate constant (k) – proportionality constant
4. Order (n) – experimentally determined
5. First-order reaction - Rate = -(ΔA/Δt) = k[A]; ln[A] = -kt + ln[A]o; t½=ln(2)/k
6. Second-order reaction – Rate = -(ΔA/Δt) = k[A]2; (1/[A] = kt + 1/[A]o; t½=1/k[A]o
7. Zero-order reaction – Rate = k; [A] = -kt + [A]o; t½=[A]o/2k
ii. Activation energy and reaction mechanisms involving catalyst
1. Activation energy – a threshold energy that must be overcome to produce a
chemical reaction
2. Catalyst – a substance that speeds up a reaction without being consumed itself
a. Lowers the activation energy
iii. Factors affecting reaction rate such as concentration, surface area, and temperature
1. Chemical reactions speed up when the temperature, concentration, and surface area
are increased
e. Chemical reaction equilibrium
i. Equilibrium constant (K)
1. jA + kB ⇋lC + mD
2. K = [C]l[D]m/[A]j[B]k
ii. Le Chatelier’s principle – if a change is imposed on a system at equilibrium, the
position of the equilibrium will shift in a direction that tends to deuce that change
1. If a reactant or product is added to a system at equilibrium, the system will shift
away from the added component; if it is removed, the system will shift toward
the removed component
2. If energy is added to the system by heating it, the shift will be in the direction
that consumes energy
f. Oxidation-reduction reactions and how to determine oxidation states
i. Oxidation states
1. The O.S. of an atom in an element is 0 (i.e. Na, O2)
2. The O.S. of a monoatomic ion is the same as its charge (use group designations)
3. Oxygen is usually -2, except in peroxides, where each oxygen is -1, and O2-2
4. Hydrogen is +1
5. Fluorine is -1
6. The sum of O.S.’s for an electrically neutral compound must be zero; for an ion, it
must equal the charge on the ion
ii. Identify oxidation-reduction reactions and half reactions
iii. Standard reduction potential – measure of the tendency of a chemical species to
acquire electrons and thereby be reduced; measured in V or mV; the more positive the
potential, the greater the species’ affinity for electrons
iv. Electrochemical reactivity series
v. Electrochemical cell – device used for generating an electromotive force (voltage) and
current from chemical reactions, or the reverse, inducing a chemical reaction by a flow
of current (i.e. battery, Galvanic cell)
C. Biochemistry and Organic Chemistry
a. Important biochemical compounds
i. Carbohydrates, including simple sugars – empirical formula CH2O
ii. Lipids – esters (OC=O)
iii. Proteins and amino acids – nitrogen; amine and carboxylic acid
iv. DNA and RNA – five carbon sugar, nitrogenous base, phosphoric acid
v. Products of photosynthesis and respiration
1. Photosynthesis – 6 CO2 + 6 H2O + energy  C6H12O6 + 6 O2
2. Respiration - C6H12O6 + 6 O2  6 CO2 + 6 H2O + energy
b. Common organic compounds (i.e., identify functional groups)
i. Alcohols
ii. Ketones and aldehydes
iii. Alkanes, alkenes, and alkynes
iv. Ethers
v. Carboxylic acids
vi. Amines
vii. Benzene – 6 Carbons, 6 Hydrogens, 3 double bonds
V. Solutions and Solubility; Acid-Base Chemistry
A. Solutions and Solubility
a. Solution terminology and calculations
i. Dilute, concentrated
1. Dilute – small amount of solute and large amount of solvent
2. Concentrated – large amount of solute and small amount of solvent
ii. Saturated, unsaturated, supersaturated
1. Saturated – the point at which a solution of a substance can dissolve no more of that
substance and additional amounts of it will appear as a separate phase
2. Unsaturated – the solution is still capable of dissolving more substance
3. Supersaturated – refers to a solution that contains more of the dissolved material
than could be dissolved by the solvent under the solubility amount’ involve a change
in some condition, such as increase in temperature or pressure, or decreasing the
volume of saturated liquid (evaporation)
iii. Solvent, solute
1. Solvent – a substance that dissolves another substance, resulting in a solution
2. Solute – a substance that is dissolved in a solvent to form a solution
iv. Concentration units (e.g., molarity, molality, mole fraction, parts per million (ppm), parts
per billion (ppb), percent by mass or volume)
1. Molarity = moles of solute / liters of solution
2. Molality = moles of solute / mass of solvent
3. Mole fraction = moles of solute / total moles of solution
v. Preparation of solutions of varying concentrations
b. Factors affecting solubility and dissolution rate
i. Dissolution rates (i.e., temperature, pressure, surface area, agitation)
1. Dissolution – process by which a solid, liquid or gas forms a solution in a solvent
a. Agitation directly affects the rate of dissolution
b. Temperature directly affects the rate of dissolution
c. Surface area inversely affects the rate of dissolution
d. Pressure directly affects the rate of dissolution
ii. Solubility and solubility curves (temperature and pressure dependent)
1. Solubility – the amount of a substance that dissolves in a given volume of solvent at
a given temperature
a. Temperature increases, solubility increases or decreases
b. Pressure of gas increases, solubility increases
c. Solution phenomena based on colligative properties (depend only on the number, and
not the identity, of the solute particles in an ideal solution)
i. Freezing point depression – the freezing point/melting point can be depressed by the
adding of a solute such as salt, which is used in the winter to lower the freezing point of
ice on roads; the more solute added to the pure solvent (which has a higher freezing
point), the more the freezing point is depressed
ii. Boiling point elevation – the boiling point of a liquid (solvent) will be higher when
another compound is added, meaning that a solution has a higher boiling point than a
pure solvent
iii. Vapor pressure effects – the presence of a solute reduces the tendency of solvent
molecules to escape, which lowers the vapor pressure of a solvent
d. Common applications of equilibrium in ionic solutions
i. Solubility of ionic compounds (e.g., solubility rules, slightly soluble compounds)
1. Solubility decreases with increased ionic strength
2. Solubility tends to decrease at high temperatures
ii. Ksp calculations including percent dissociation and precipitation
1. Ksp is the solubility product constant; an equilibrium constant
2. Percent dissociation – amount dissociated/initial concentration x 100%
3. Ion product (Q) – use initial concentrations instead of equilibrium concentrations
used in Ksp (ex. CaF2; Q = [Ca2+]o[F-]o2)
a. If Q > Ksp, precipitation occurs
b. If Q < Ksp, no precipitation occurs
iii. Common ion effect – results when two substances, which both ionize to give the same
(common) ion, are involved in a chemical equilibrium; the solubility of a solid is lowered
if the solution already contains ions common to the solid
iv. Electrolytes, nonelectrolytes, and electrical conductivity
1. Electrolyte – a material that dissolves in water to give a solution that conducts an
electrical current
2. Nonelectrolyte – a substance which, when dissolved in water, gives a nonconducting
solution
3. Electrical conductivity – the ability to conduct an electric current
B. Acid-Base Chemistry
a. Define and identify acids and bases and know their properties
i. Arrhenius acids and bases – acids produce hydrogen ions in aqueous solution, while
bases produce hydroxide ions
ii. Bronsted-Lowry acids and bases – an acid is a proton (H+) donor, and a base is a
proton acceptor
iii. Lewis acids and bases – acids are electron-pair acceptors, and bases are electron-pair
donors
iv. Neutralization and equivalence point
1. Neutralization – results from a reaction between an acid and a base
2. Equivalence point – the point in a titration when enough titrant has been added to
react exactly with the substance in solution being titrated
b. The pH scale and calculations involving pH and pOH
i. pH = -log[H+] (Δ[H+] = 10, ΔpH = 1)
1. The pH decreases as [H+] increases, and vice versa
2. pH=0, very acidic; pH=14, very basic
ii. Calculation of pH and pOH
1. pH = -log[H+]
2. pOH = -log[OH-]
3. pH + pOH = 14
iii. Calculation of [H+] and [OH-]
iv. Kw – ion-product constant (dissociation constant); Kw=[H+][OH-]=1.0x10-14mol2/L2
1. Neutral – [H+]=[OH-]
2. Acidic – [H+]>[OH-]
3. Basic – [H+]<[OH-]
c. Concepts and calculations involving acid-base titrations
i. Use and selection of indicators (e.g., phenolphthalein, litmus paper)
1. Phenolphthalein – colorless in acidic solutions; pink in basic solutions
2. Litmus paper – red in acidic solutions; blue in basic solutions
ii. Endpoint determination – the point in a titration at which the indicator changes color
iii. Calculations based on titrations
d. Equilibrium relationships in acid-base chemistry
i. Strong/weak acids and bases, including common examples
1. Strong indicates that the acid/base (nearly) completely dissociates; has weak
conjugate acid/base
a. Acids – HCl, HNO3, H2SO4
b. Bases – NaOH, LiOH, KOH, Ca(OH)2
2. Weak indicates that the acid/base does not dissociate (much); has strong conjugate
acid/base
ii. Monoprotic and polyprotic acids – mono (1 proton), poly (more than 1 proton)
iii. Ka, Kb, and percent dissociation
1. Ka=[H+][A-]/[HA]
2. Kb=[BH+][OH-]/[B]
iv. Hydrolysis – a chemical reaction during which water molecules are split into hydrogen
cations and hydroxide anions; can be used to break down large molecules
v. Buffer solutions – solution that resists a change in its pH when wither hydroxide ions
or protons are added
1. May contain a weak acid and its salt (HF and NaF), or a weak base and its salt (NH3
and NH4Cl)
VI. History and Nature of Science; Science, Technology, and Social Perspectives
A. History and Nature of Scientific Inquiry
a. Processes involved in scientific inquiry
i. Formulating problems
ii. Forming and testing hypotheses
iii. Development of theories, models, and laws (postulates, assumptions)
1. A law summarizes what happens; a theory (model) is an attempt to explain why it
happens
iv. Process skills including observing, concluding, comparing, inferring, categorizing and
generalizing
b. Experimental design
i. Testing hypotheses
ii. Significance of controls
iii. Use and identification of variables
iv. Data collection planning
c. Nature of scientific knowledge
i. Subject of change
ii. Consistent with experimental evidence
iii. Reproductibility – related to the precision of a measurement
iv. Unifying concepts and processes (e.g., systems, models, constancy and change,
equilibrium, form and function)
d. Major historical developments in chemistry and the contributions of major historical
figures
i. How current chemical principles and models developed over time
ii. Major developments in chemistry (e.g., atomic model, ideal gas behavior) including
major historical figures
B. Science, Technology, Society, and the Environment
a. Impact of chemistry and technology on society and the environment
i. Pharmaceuticals
ii. Acid rain
iii. Medical imaging
iv. Air and water pollution
v. Greenhouse gases
vi. Ozone layer depletion
vii. Waste disposal and recycling
viii. Nanotechnology
b. Applications of chemistry in daily life
i. Plastics, soap, batteries, fuel cells, and other consumer products
ii. Water purification
iii. Chemical properties of household properties
c. Advantages and disadvantages associated with various types of energy production
i. Renewable and nonrenewable energy resources
ii. Conservation and recycling
iii. Pros and cons of power generation based on various sources such as fossil and nuclear
fuel, hydropower, wind power, solar power, and geothermal power
VII. Mathematics, Measurement, and Data Management; Laboratory Procedures and Safety
A. Collect, evaluate, manipulate, interpret, and report data
a. Significant figures in collected data and calculations
b. Organization and presentation of data
c. Knows how to interpret and draw conclusions from data presented in tables, graphs, and
charts (e.g., trends in data, relationships between variables, predictions and conclusions
based on data)
B. Units of measurement, notation systems, conversions, and mathematics used in chemistry
C.
D.
E.
F.
a. Standard units of measurement
b. Unit conversion
c. Scientific notation
d. Measurement equipment
Basic error analysis
a. Determining mean
b. Accuracy and precision
i. Accuracy – agreement of a particular value with the true value (cluster at target)
ii. Precision – degree of agreement among several measurements of the same quantity
(cluster, not necessarily at target)
c. Identifying sources and effects of error
d. Percent error = |your result–accepted value|/accepted value x 100%
Appropriate preparation, use, storage, and disposal of materials in the laboratory
a. Appropriate use and storage
b. Safe disposal
c. Preparation for classroom use
d. Safe procedures and safety precautions
Appropriate use, maintenance, and calibration of laboratory equipment
a. Appropriate use and storage
b. Maintenance and calibration
c. Preparation for classroom use
d. Safety procedures and precautions when using equipment
Safety procedures and precautions for the high school chemistry laboratory
a. Location and use of standard safety equipment such as eyewash and shower
b. Laboratory safety rules for students
c. Appropriate apparel and conduct in the laboratory, such as wearing goggles
d. Emergency procedures