Chemistry: Content Knowledge (0245 or 5245)I. Matter and Energy; Heat and Thermodynamics, and Thermochemistry A. Matter and Energy

a. Organization of matteri. 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 composition1. Homogenous – having visibly indistinguishable parts2. Heterogeneous – having visibly distinguishable parts3. Solutions – a homogeneous mixture 4. Suspensions – heterogeneous mixture containing solid particles; will eventually

settleiii. States of matter

1. Solid – rigid; fixed volume, shape2. Liquid – definite volume, no specific shape; assumes shape of container3. 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 matteri. 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 chargeiii. Molecules – unit of atoms held together by equal sharing of electrons (covalent bonds)

c. Differences between chemical and physical properties and chemical and physical changesi. Chemical vs. physical properties

1. Chemical property – a property that can only be observed by changing the identity of the substancea. Changes happen on a molecular level

2. Physical property – can be observed without changing the identity of the substancea. About energy and states of matter

ii. Chemical vs. physical changes1. Chemical change – changes that alter the identity of a substance2. Physical change – changes in matter that do not alter the identity of the substance

iii. Intensive vs. extensive properties1. 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 potential1. Kinetic energy – (½mv2) energy due to the motion of an object2. Potential energy – energy due to position or composition

ii. Chemical, electrical, electromagnetic, nuclear, and thermal energy1. Chemical energy – energy stored in the bonds of chemical compounds2. 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 space4. Nuclear energy – energy of an atomic nucleus, which can be released by fusion or

fission or radioactive decay5. 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 electrical2. Chemical explosion: chemical to kinetic and thermal* Transducer – a device that converts one form of energy to another

B. Thermodynamics in Chemistrya. Temperature, thermal energy, and heat capacity, including temperature scales, units of

energy, and calculations involving those conceptsi. 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 heat1. 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/ΔT2. Specific heat capacity, q = mcΔT3. ΔE = q + w, where w = -PΔV

b. Concepts and calculations involving phase transitions between the various states of matteri. 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 sublimation1. 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 = mciceΔT2. B – ΔHfusion

3. C – q = mcwaterΔT4. D –ΔHvaporization

5. E - q = mcsteamΔT*Heat added = q

c. Kinetic molecular theory and ideal gas lawi. Assumptions of kinetic molecular theory – simple model that attempts to explain the

properties of an ideal gas1. Gas molecules have zero volume2. Gas molecules exert no forces on each other3. 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 = 3/2RT3. 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 reactionsi. Exothermic and endothermic reactions

1. Exothermic – refers to a reaction where energy (as heat) flows out of the system2. Endothermic – refers to a reaction where energy (as heat) flows into the system

ii. Bond energy; Hess’s law1. Bond energy – the energy required to break a given chemical bond2. 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 constant2. In any spontaneous process there is always an increase in the entropy of the

universe3. The entropy of a perfect crystal at 0 K is zero

ii. Spontaneous/reversible processes1. Spontaneous – a process that occurs without outside intervention; may be fast or

slow2. 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 processes1. 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)

SΔ sys SΔ surr SΔ univ Spontaneous?+ + + Yes

- - - No (opposite direction)

+ - ? Yes, if SΔ sys > SΔ surr

- + ? Yes, if SΔ surr > SΔ sys

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 StructureA. 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 atomiv. 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

atomii. Average atomic mass – average total mass of protons, neutrons, and electrons in a

single atomB. Electron configuration of the elements based on the periodic table

a. Aufbau principle, Hund’s rule, Pauli exclusion principlei. 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 bondingC. Radioactivity

a. Characteristics of alpha particles, beta particles, and gamma radiationi. Alpha particle – a helium nucleus

ii. Beta particle – an electron produced in radioactive decayiii. Gamma particle – a high-energy photon

b. Radioactive decay processes; half lifei. Alpha-particle production – involves a change in mass number for the decaying

nucleusii. 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 valuec. 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 levelsa. 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 orbitalb. Energy of electronic absorption/emission spectral lines in various regions of the

electromagnetic spectrumi. 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 wavelengthi. 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 spaceiii. × =c (c=speed of light=3.0 × 10λ ν 8 m/s)iv. E = h (h=Planck’s constant=6.26 × 10Δ ν -34 J×s)

III. Nomenclature; the Mole, Chemical Bonding, and GeometryA. Nomenclature and Chemical Composition

a. Systematic names and chemical formulas of simple inorganic compoundsi. Binary compounds

1. The cation is always named first and the anion second2. A monatomic cation takes its name from the element3. A monatomic acnion takes its root and adds –ide4. 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 –ous5. Prefixes: mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca

ii. Acids, bases, and salts1. 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 salt2. Determine the prefix to be used based on the number of water molecules3. Add the word “hydrate” to the end of the prefix

b. Names of common organic compounds based on functional groupsi. 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 (2,3-Dimethylpentane)

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, amines1. 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 –one4. Carboxylic acid (COOH)– drop –e for –oic acid 5. Esters (COO)– drop –e for –oate6. 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 compositioni. 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 totalb. Convert the mass of each element to moles using molar mass c. Divide each mole value by the smallest number of moles calculatedd. 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 composition1. 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 Structurea. Common properties of bonds

i. Relative bond lengths1. As the number of shared electrons increases, the bond length shortens

ii. Relative bond strengths1. Inversely related to bond length; as bond length increases, bond strength decreases

b. Bond typesi. 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 sharinga. 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 atoms2. Use a pair of electrons to form a bond between each pair of bound atoms3. 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 moleculea. 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 molecule1. Can utilize electronegativity to determine polarity, as a great difference in

electronegativity will result in polar moleculesii. Symmetry of molecular structure

e. Intermolecular interactionsi. 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 propertiesi. 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 ChemistryA. Periodicity

a. Basis of the periodic table and general layouti. Arranged in groups and periods

1. Periods = rows2. Groups = columns

ii. Atomic number – the number of protons found in the nucleus of an atom iii. Symbols of the elementsiv. Metals, nonmetals, metalloids

1. Group 1: alkali metals2. Group 2: alkaline earth metals3. Metalloids: B, Si, Ge, As, Sb, Te, Po4. Group 17: Halogens (nonmetals) 5. Group 18: Noble gases (nonmetals) 6. Other nonmetals: C, N, O, P, S, Se

v. Transitions elements1. 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 Principlesa. Balancing chemical equations

i. Simple chemical equations1. Start with the most complicated molecule

ii. Chemical equations involving oxidation-reduction1. Oxidation – loss of electrons or an increase in oxidation state2. 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 volume1. Balance the equation for the reaction2. Convert the known mass of the reactant or product to moles of that substance3. 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 reaction4. Use the appropriate mole ratios to calculate the number of moles of the desired

reactant or product5. Convert from moles back to grams if required by the problem

ii. Limiting reagent calculations and percent yield1. Limiting reagent – the reactant that is consumed first and which therefore limits

the amounts of products that can be formed a. Convert mass to molesb. 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 typesi. 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 speciesa. 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-producta. YOH + HX XY + H2O

ii. Decomposition, dehydration1. Decomposition – the separation of a chemical compounds into elements or simpler

compoundsa. 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 + H2Oiii. Single and double replacement

1. Single replacement – a type of redox reaction when an element or ion moves out of one compound and into anothera. 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-reductiond. Chemical kinetics

i. Rate laws, rate constants, and reaction order1. Reaction rate – change in concentration of a reactant or product per unit time2. 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 constant4. Order (n) – experimentally determined 5. First-order reaction - Rate = -(ΔA/Δt) = k[A]; ln[A] = -kt + ln[A]o; t½=ln(2)/k6. 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/2kii. 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 itselfa. Lowers the activation energy

iii. Factors affecting reaction rate such as concentration, surface area, and temperature1. 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 statesi. 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 +15. Fluorine is -16. The sum of O.S.’s for an electrically neutral compound must be zero; for an ion, it

must equal the charge on the ionii. 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 seriesv. 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 Chemistrya. Important biochemical compounds

i. Carbohydrates, including simple sugars – empirical formula CH2Oii. Lipids – esters (OC=O)

iii. Proteins and amino acids – nitrogen; amine and carboxylic acidiv. 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. Alcoholsii. Ketones and aldehydes

iii. Alkanes, alkenes, and alkynesiv. Ethersv. Carboxylic acids

vi. Aminesvii. Benzene – 6 Carbons, 6 Hydrogens, 3 double bonds

V. Solutions and Solubility; Acid-Base ChemistryA. Solutions and Solubility

a. Solution terminology and calculationsi. Dilute, concentrated

1. Dilute – small amount of solute and large amount of solvent2. Concentrated – large amount of solute and small amount of solvent

ii. Saturated, unsaturated, supersaturated1. 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 phase2. Unsaturated – the solution is still capable of dissolving more substance3. 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, solute1. Solvent – a substance that dissolves another substance, resulting in a solution2. 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 solution2. Molality = moles of solute / mass of solvent3. Mole fraction = moles of solute / total moles of solution

v. Preparation of solutions of varying concentrationsb. 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 dissolutionb. Temperature directly affects the rate of dissolutionc. 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 solutionsi. Solubility of ionic compounds (e.g., solubility rules, slightly soluble compounds)

1. Solubility decreases with increased ionic strength2. 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 occursb. 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 conductivity1. Electrolyte – a material that dissolves in water to give a solution that conducts an

electrical current2. Nonelectrolyte – a substance which, when dissolved in water, gives a nonconducting

solution3. 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 point1. Neutralization – results from a reaction between an acid and a base2. 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, pH = 1)Δ1. The pH decreases as [H+] increases, and vice versa2. pH=0, very acidic; pH=14, very basic

ii. Calculation of pH and pOH1. 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 titrationsi. Use and selection of indicators (e.g., phenolphthalein, litmus paper)

1. Phenolphthalein – colorless in acidic solutions; pink in basic solutions2. Litmus paper – red in acidic solutions; blue in basic solutions

ii. Endpoint determination – the point in a titration at which the indicator changes coloriii. 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 PerspectivesA. History and Nature of Scientific Inquiry

a. Processes involved in scientific inquiryi. 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

happensiv. Process skills including observing, concluding, comparing, inferring, categorizing and

generalizingb. Experimental design

i. Testing hypothesesii. Significance of controls

iii. Use and identification of variablesiv. Data collection planning

c. Nature of scientific knowledgei. Subject of change

ii. Consistent with experimental evidenceiii. 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

figuresi. 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 Environmenta. Impact of chemistry and technology on society and the environment

i. Pharmaceuticalsii. Acid rain

iii. Medical imagingiv. Air and water pollutionv. Greenhouse gases

vi. Ozone layer depletionvii. Waste disposal and recycling

viii. Nanotechnology b. Applications of chemistry in daily life

i. Plastics, soap, batteries, fuel cells, and other consumer productsii. Water purification

iii. Chemical properties of household propertiesc. Advantages and disadvantages associated with various types of energy production

i. Renewable and nonrenewable energy resourcesii. 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 SafetyA. Collect, evaluate, manipulate, interpret, and report data

a. Significant figures in collected data and calculationsb. Organization and presentation of datac. 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

a. Standard units of measurementb. Unit conversionc. Scientific notationd. Measurement equipment

C. Basic error analysisa. Determining meanb. 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 errord. Percent error = |your result–accepted value|/accepted value x 100%

D. Appropriate preparation, use, storage, and disposal of materials in the laboratorya. Appropriate use and storageb. Safe disposalc. Preparation for classroom used. Safe procedures and safety precautions

E. Appropriate use, maintenance, and calibration of laboratory equipmenta. Appropriate use and storageb. Maintenance and calibrationc. Preparation for classroom used. Safety procedures and precautions when using equipment

F. Safety procedures and precautions for the high school chemistry laboratorya. Location and use of standard safety equipment such as eyewash and showerb. Laboratory safety rules for studentsc. Appropriate apparel and conduct in the laboratory, such as wearing gogglesd. Emergency procedures