AP Chemistry Review Cards Flashcards

I took the flash cards from this weeks materials and put them into this set.

13857453058	Diatomic Elements and bonding	H2 - single covalent bond F2, Cl2, Br2, I2 - single covalent bond O2 - double covalent bond N2 - triple covalent bond	0
13857453059	The 6 strong acids that 100% ionize	HCl -> H(+) + Cl(-) HBr -> H(+) + Br(-) HI -> H(+) + I(-) HClO4 -> H(+) + ClO4(-) HNO3 -> H(+) + NO3(-) HSO4 -> H(+) + SO4(-)	1
13857453060	The strong bases	All soluble metal hydroxides	2
13857453061	Celsius to Kelvin conversion	273 + C = K	3
13857453062	Percent Error	[(Experimental - Correct)/Correct]*100%	4
13857453063	# of SD that will be considered correct for 99% of all AP Exam Questiosn	3/three	5
13857453064	Only situation where you can use exact number of SDs	Lab measurements involving subtraction and addition	6
13857453065	Latitude in Sig Figs	Greater # of SDs, the greater the accuracy Most accurate lab devices: graduate cylinder, blance, volumetric flask, burette, pipette Least accurate lab device: beaker, Erlenmeyer flask	7
13857453066	Molecular Mass of molecules	H2 = 2.0 g/mol N2 = 28 g/mol O2 = 32 g/mol H2O = 18 g/mol CO2 = 44 g/mol NaOH = 40.0 g/mol	8
13857453067	Atomic Number	the number of protons in the nucleus of an atom also the number of electrons in a neutral atom	9
13857453068	Mass Number	the sum of the number of neutrons and protons in an atomic nucleus	10
13857453069	Average Atomic Mass	average mass in amu or g/mol of the atoms in an element	11
13857453070	Mass spectrometer	sorts isotopes of elements by mass and shows the relative abundance of each isotope	12
13857453071	Similarities between isotopes of an element	All isotopes of an element have same number of protons	13
13857453072	Differences between isotopes of an element	Density Atomic Mass Number of Neutrons	14
13857453073	How are specific isotopes of an element written	name-mass number uranium-235	15
13857453074	Coulombic forces	Attractions between opposite charges and repulsions of similar charges	16
13857453075	Avogadro's number	6.02 x 10^23	17
13857453076	Units for molar mass and mol-mass conversion	molar mass: g/mol g of substance * 1mol/g molar mass = mol substance mol of substance * g molar mass/1mol = g mass of substance	18
13857453077	How to find a limiting reactant	Set up an ICE chart Divide the mol amounts of reactants by the coefficients of the reactants the smaller molar amount will be the limiting reacting.	19
13857453078	Empirical Formula	mol ratio of the elements in a compound reduced to the lowest whole numbers Find the mol of each element Divide each mol amount by the lowest mol amount	20
13857453079	Mole Fraction	molA/total mol mixture = mole fraction Add up all the moles to get the denominator	21
13857453080	Cations which are soluble	Na+, K+, NH4+ and other group 1 ions	22
13857453081	anions which are soluble	NO3-, ClO4-, and SO4 (2-)	23
13857453082	molarity units	mole/L	24
13857453083	Molarity x Volume (L)	Moles of solute	25
13857453084	Molarity x Volume (mL)	Millimoles of solute	26
13857453085	Millimole of solute/volume (mL)	Molarity in M	27
13857453086	Molecular, ionic, and net ionic reaction for the formation of a precipitate of sodium chloride reacting with silver nitrate to form a precipitate	Molecular: NaCl + AgNO3 -> NaNO3 + AgCl Ionic: Na+ + Cl- + Ag+ +NO3- -> Na+ +NO3- +AgCl Net Ionic: Ag+ + Cl- -> AgCl Note: All sodium, potassium, ammonium, and nitrate compounds are soluble and will not be found in NIE's. Weak acids and bases are shows as molecules in net ionic equations.	28
13857453087	Solubility rule:	All sodium, potassium, ammonium, and nitrate compounds are soluble and will not be found in NIE's.	29
13857453088	torr or mmHg to atm	760 mm Hg = 1 atm	30
13857453089	STP for gases	1 atm, 760 mmHg, 0 C, 273 K	31
13857453090	Density of gas	molar mass = (Density x R x T)/Pressure	32
13857453091	Density of a gas @STP	Density g/L = molar mass/22.4 L/mol	33
13857453092	Molecular speed of gases and molecular mass and Maxwell Boltzmann curve	At a given temperature, lighter means faster Hydrogen is always the fastest. On Maxwell-Boltzmann curves, the average speed is slightly to the right of the peak of its curve. The faster the speed, the more spread out it will be.	34
13857453093	Non-ideal T and P conditions	Higher pressure, lower temperature	35
13857453094	Causes of deviations from Ideal-gas	Higher than expected pressures (large molecular volume), lower than expected pressure (condense).	36
13857453095	Partial pressures of gases in a mixture	Pa = Pt x mole fraction	37
13857453096	Specific heat	Water = 4.18 J/gC metals = low specific heats (less than 1).	38
13857453097	Energy from experimental reaction in calorimeter	Calorimetry is used to find the enthalpy of a reaction: qcalorimeter = mcdeltaT	39
13857453098	What is deltaH?	Change in enthalpy of a reaction	40
13857453099	deltaH units	kJ/mol	41
13857453100	deltaH = +	Endothermic - thermodynamically unfavorable and will only happen if there is an increase in S	42
13857453101	deltaH = -	Exothermic - thermodynamically favorable.	43
13857453102	deltaH where heat (KE) is a product	Exothermic and deltaH is negative	44
13857453103	deltaH where heat (KE) is a reactant	Endothermic and deltaH is positive	45
13857453104	enthalpy of formation	Reaction enthalpy to make 1 mole of substance is made from its elements in the most common form.	46
13857453105	enthalpy of formation for elements	0 kJ/mol by definition	47
13857453106	enthalpy of formation of elements exceptions	May not be zero if it is not in its most common form at 25 C	48
13857453107	Enthalpy of reaction from enthalpy of formation	enthalpy of reaction = deltaHf - deltaHf rxn	49
13857453108	Photons and wavelengths and energy	E=hv c=lambda(v) Shorter wavelengths have higher frequencies and greater photon energies	50
13857453109	Unit for frequency	Hertz (Hz) 1/s	51
13857453110	nm and m relationship caution in calculations	nm = 1e-9 meter speed of light is given in meters Speed of light is given in m, must convert from nm to m.	52
13857453111	Nanometer wavelengths UV X-ray Visible light IR	Xrays < 10 nm UV 10 nm - 400 nm Visible 400nm (violet) - 700 nm (red) IR 700 nm - 1 mm	53
13857453112	Generalization of how the different photon energies affect substances	XRays ionize atoms Ultraviolet and visible excite electrons to different energy levels Infrared and microwave cause molecular vibrations and rotations.	54
13857453113	Electron order of orbital filling	1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s2 You won't have to write configurations beyond 5s2 This works with the periodic table.	55
13857453114	Abbreviated electron configuration	Symbol of noble gas preceding element is placed in brackets.	56
13857453115	Hund's rule	Electrons will half fill orbitals in a subshell before doubling up	57
13857453116	Pauli Exclusion Principle	Paired electrons must have opposite spin within an orbital.	58
13857453117	Quantum numbers	Quantum numbers are used to provide a more detailed description of electrons in an atom.	59
13857453118	Atomic radius decreases as atomic number increases from left to right because	Nuclear charge attraction to electrons, Z, is produced by protons in the nucleus attracting outer electrons. Coulombic attractions	60
13857453119	Reason why atomic radius increases as period increases down families	Electrons in higher numbered periods are in successively higher energy levels placing the valence electrons farther from the attractions of the protons in the nucleus.	61
13857453120	Summary of periodic table trends for atomic radii	Smallest in the upper right-hand corner Largest in the lower left hand corner	62
13857453121	Knowing the trends vs explaining the trends	Explaining requires knowing the cause of the trends Referencing a periodic trend does not constitute an explanation of atomic property differences Zeff, number of shells, distance of principle energy levels from the nucleus can be valid explanations.	63
13857453122	PES graphs	PES is an experimental method for determining the binding energies and electronic structure of an atom The data is produced by kicking electrons out of atoms using high energy photons. AP Question PES graphs will always have the highest binding energies of interior electrons to the left with lower binding energies to the right in a logarithmic scale. PES graphs will group the subshells in their respective energy levels. The number of electrons in the subshell determines the height of each subshell peak. Electron configurations mirror PES graphs	64
13857453123	PES graph shifts with increasing atomic size	As the number of protons increases the peaks shift left indicating the greater binding energy of the protons to the electrons.	65
13857453124	# of Valence electrons in atoms	Electrons in the highest energy level These are determined using the columns in the periodic table.	66
13857453125	How non-transition metals ionize	Oxidation, losing electrons to form a cation Electron loss is down to the p6 of the lower energy level	67
13857453126	How transition metals ionize	Oxidation, losing electrons to form a cation Electron loss is from the higher energy s orbital	68
13857453127	How nonmetals ionize	Reduction, gain electrons to p6	69
13857453128	Three elements when bonded with hydrogen can form hydrogen bonding?	Fluorine, Oxygen, and Nitrogen	70
13857453129	Ionization energy	Energy needed to remove an electron from a single atom of an element in its gaseous form Always endothermic, deltaH = +	71
13857453130	Electron affinity energy	Attraction of electron to neutral atom Usually exothermic, deltaH = -	72
13857453131	Factors used to explain increasing ionization energy	Smaller atomic size Greater number of protons in nucleus	73
13857453132	Factors used to explain decreasing ionization energy	smaller atoms nonmetals upper right of the periodic table	74
13857453133	Dramatic increases in ionization energy indicate the limit of the + charge of the ion	Atom with ionization energy sequence 1st 1801 kJ to become X+ 2nd 2430 kJ to become X2+ 3rd 3660 kJ to become X3+ 4th 25000 kJ (dramatic increase in IE prevents further ionization)	75
13857453134	How is an ionic compound formed?	An ionic bond is formed between a metal that loses electrons and a nonmetal which gains those electrons.	76
13857453135	What is the lattice energy?	Kinetic energy released when the ions come together to form a crystal lattice	77
13857453136	What is the difference between ionic bonds and covalent bonds?	In an ionic compound the difference in electronegativity is so great that the particles can be considered ions In a covalent compound the atoms 'share' the electrons rather than taking it away from each other.	78
13857453137	Covalent Bond formation PE graph	Two atoms combine to form a bond Bond formation is always exothermic Bond energy is at the bottom of the PE well	79
13857453138	Bond energy value	Bond energy is always a positive number because it is defined as the kinetic energy needed to break a bond Breaking a bond is always endothermic since energy is required to separate attracted atoms.	80
13857453139	deltaBE and deltaH	deltaBE = BEproducts - BEreactants Endothermic reactions have a decrease in BErxn Exothermic reactions have an increase in BErxn	81
13857453140	VSEPR & Molecular Orbital Model	VSEPR or Valence Shell Electron Repulsion Model uses valence electrons and the lewis dots to predict the structure of covalently bonded molecules. This is in AP Chem Molecular Orbital Model is more complex and more accurate, but will not be in AP Chem	82
13857453141	Formal Charge	the difference between the normal number of lewis dots and the number of electrons controlled by an atom in a molecule (one per bond, two per lone pair)	83
13857453142	Molecular Shapes and Domains	The number of unshared electrons plus the number of sigma bonds is the # of electron domains. The number of electron domains determine the Steric number The SN determines the shape of the molecule	84
13857453143	Which atoms are stable with three domains?	Boron and Aluminum need only 6 pairs	85
13857453144	Which atoms can have expanded octets with more than 4 domains?	Atoms that have electrons in d orbitals, in the 4th, 5th, 6th, and 7th period	86
13857453145	Molecular shapes The symmetrical shape	- linear - trigonal planar - tetrahedral - trigonal bipyramidal - octahedral - square planar	87
13857453146	Molecular shapes The unsymmetrical shape	- nonlinear - trigonal pyramidal - seesaw - square pyramidal	88
13857453147	What is electronegativity?	Electronegativity is an atom's attraction to the shared pair of electrons in a covalent bond Differences in electronegativity can lead to polar bonds and result in a dipole	89
13857453148	Hybridization	s and p orbitals combine in sigma bonding 2 domains = sp hybrid 3 domains = sp2 hybrid 4 domains = sp3 hybrid Each double bond only has one sigma bond, so double bonds only count for one domain.	90
13857453149	Single vs Double bonds	Single bonds (sigma) are longer and weaker than double bonds Double bonds (sigma with pi) are shorter and stronger than single bonds and have a greater bond energy In a double bond, the pi bond is weaker	91
13857453150	Resonance structures	If a double bond can be placed in alternative locations on the molecule, resonance structures are used to explain bonding. The double bond is averaged over all the resonance structures SO2 would have two resonance structures Sulfur oxygen's two bond lengths would be the same less, than a single but longer than a double.	92
13857453151	Lewis Dot structure of H2O	4 electron domains, 2 unbonded Nonlinear or bent (unsymmetrical) predicted angle 109 degrees Hybrid sp3	93
13857453152	Lewis dot structure of NH3	4 electron domains, 3 bonded Trigonal pyramidal, predicted angle 109 degrees Hybrid explanation sp3	94
13857453153	Lewis dot structure of CO2	2 electron domains, 2 bonded Linear, predicted angle 180 degrees 2 sigma bonds, 2 pi bonds Hybrid explanation, sp	95
13857453154	Unshared pairs of electrons and molecular shape	Unshared pairs have greater coulombic repulsive forces than electrons locked in covalent bonds Unshared pairs of electrons will decrease the predicted angles between atoms in a molecule Two unshared pairs of electrons warp the predicted angle of 109 degrees down to 105 degrees.	96
13857453155	Types of intermolecular attractions	London Dispersion Forces Dipole-dipole forces Dipole-ion Hydrogen bonding	97
13857453156	Hydrogen bonding	Intermolecular attraction Special dipole-dipole where molecule has H-N, H-O, or H-F The hydrogen bond is not a covalent bond The hydrogen atom develops a concentrated + charge due to the loss of its only electron. The attraction is the hydrogen dipole-intermolecular bond. This intermolecular attraction forms molecular solids and liquids.	98
13857453157	How are physical properties affected by intermolecular attractions	High attractions: - Increase freezing points - Decrease vapor pressures - Increase boiling temperature - Increase deltaHfusion - Increase deltaHvaporization - Increase viscosities	99
13857453158	Type of crystalline solid with low melting temperatures and no electrical conductivity as a solid or molten	Molecular solid: a solid made of individual molecules held together in a crystalline lattice by any of the following: LDFs Dipole Dipole forces Hydrogen bonding	100
13857453159	Type of crystalline solid with high melting temperatures and no electrical conductivity when solid but conducts when molten	Ionic solid: A solide made of individual ions held together in a crystalline lattice by opposite charges of ions.	101
13857453160	Type of crystalline solid which is malleable and electrically conductive as a solid and when molten	Metallic solid: Atoms held together by attraction to mobile valence electrons - sea of electrons Molecular motion interferes with electron movement	102
13857453161	Impurities put in between atom lattice in metals	interstitial alloys, increase hardness and strength by stressing crystal lattice, much smaller atoms	103
13857453162	Impurities replacing atoms in metallic lattice	substitutional alloy Different size atoms used, increases strength by preventing movement in lattice	104
13857453163	Alloys in chemistry	Sometimes an alloy will alter the chemistry of pure metal Adding Cr to Fe will prevent Fe from rusting.	105
13857453164	Type of crystalline solid with a very high melting temperature composed of nonmetal atoms	Network solids - covalently bonded macromolecules 3d: Diamond, silicon carbide, quartz 2d: graphite, mica, asbestos	106
13857453165	Type of crystalline solid with a very high melting temperature composed of metalloid atoms What are they and what special electrical properties do they have?	The metalloids silicon and germanium are network molecular solids with 4 valence electrons Metalloids conduct electricity poorly as a solid but increase in conductivity with increased temperature	107
13857453166	p-doping	p-semiconductors Si or Ge 4-valence electron solid is doped with elements with 3-valence electrons The missing electron produces positive charged holes in crystal lattice to allow for current flow	108
13857453167	n-doping	n-semiconductors Si or Ge 4-valence electron solid is doped with elements with 5 valence electrons. The extra electron produces conductive negative charges in crystal lattice to allow for current flow. Electronic devices such as transistors and diodes are formed using n and p semiconductors.	109
13857453168	Rf	ratio of the distance moved by the solute to the distance moved by the solvent	110
13857453169	How to calculate K3 given K1 and K2	Take each K to the power of the coefficient, then multiply.	111
13857453170	K of the reverse reaction	1/K	112
13857453171	Factor for increased melting temperatures of ionic compounds	The smaller the ion and greater the charge, the higher the MT	113
13857453172	Signs for deltaH and deltaS for fusion, vaporization, and sublimation	Vaporization: deltaH + and deltaS + Fusion: deltaH + and deltaS + Sublimation: deltaH + and deltaS + These processes break bonds and are endothermic	114
13857453173	Signs for deltaH and deltaS for freezing and condensation	Freezing: deltaH - and deltaS - Condensation: deltaH - and deltaS -	115
13857453174	Alkane, alkene, and alkyne's composition, intermolecular attractions, and solubility in water	Only C and H are in the formula, no dipoles. IMF's: only LDFs Not soluble in water	116
13857453175	Alkane, alkene, and alkyne's bonding and hybridization	alkanes: C-C sp3 hybridization alkenes: C=C sp2 hybridization alkynes: C=-C sp hybridization	117
13857453176	Alcohol functional group, imf, and solubility	C-OH IMFs: Hydrogen bonding and London Forces smaller changes are soluble in water (neither acids nor bases)	118
13857453177	Carboxylic acids' functional group and intermolecular attractions	Carboxylic COOH IMFs: hydrogen bonding and london forces	119
13857453178	Carboxylic acids' acid reactions and name change	Weakly ionize with water, turns from -cooh to -coo-	120
13857453179	Amines' intermolecular attractions	-NH2 IMFs: Hydrogen bonding and London forces	121
13857453180	Amines base reactions and name change	C-NH2 + H2O --> C=NH3 + OH- The weak base, amine, turns into an amide, a weak conjugate acid	122
13857453181	Functional groups	aldehyde - OCH Ester - (embedded oxygen with double bond o) ether - embedded oxygen ketone - c double bond o	123
13857453182	What are mers and polymers?	Polymers (aka plastics) are long-chain carbon chains with repeating units (mers)	124
13857453183	What type of molecular substances will dissolve in water? What type of ionic substances will dissolve in water?	Soluble molecular substances: Molecular substances with strong dipoles and low LDFs will dissolve in water Molecular substances with -OH for hydrogen bonding and low LDFs will dissolve in water Soluble ionic substances: All substances with sodium, potassium and ammonium cations, and nitrate anions will dissolve in water to form dipole-ion attractions with water Other ionic substances may dissolve in water if their ion-dipole attractions to water are greater than their cation-anion attractions	125
13857453184	What is distillation and the distillate?	Distillation is used to separate solution mixtures by evaporating the most volatile components. The condensed vapor is called the distillate.	126
13857453185	What is chromatography?	Chromatography is a method of separating small quantities of components of a mixture using differences in intermolecular attractions. Typically, there is a solvent and a fixed media. The solvent will carry the substances in the mixture with similar IMFs to the solvent, leaving behind the substances that are more attracted to the fixed media.	127
13857453186	What does Rf indicate?	Rf is an indicator of the mixture component's attraction to the solvent. If the Rf is close to 1, then the solute's IMF's are the same as the solvent. If the Rf is small, then the solute's IMF's will be similar to the fixed media.	128
13857453187	What is a spectrophotometer and how does it work?	A spectrophotometer uses the absorption of light to determine the concentration of solution.	129
13857453188	What wavelength of light is appropriate for use in spectrophotometry?	The appropriate wavelength of light is the set of wavelengths that are absorbed most strongly by the solution.	130
13857453189	[concentration] with absorbance	The concentration is directly proportional to the absorbance	131
13857453190	Units for rate	Loss of reactant concentration/time rate = M/s or atm/s or torr/s rate law units must multiply out to make M/s	132
13857453191	What is the rate law expression for a reaction?	rate = k[A]x[B]y x and y can be determined experimentally or from a reaction mechanism	133
13857453192	What is the instantaneous rate law expression, rate constant unit, and concentration-time graph for a zero-order reaction?	rate = k[A]0 unit = M/s line is straight	134
13857453193	What is the instantaneous rate law expression, rate constant unit, and concentration-time graph for a first-order reaction?	rate = k[A]1 unit = 1/s logarithmic - ln(x) gets you a straight line.	135
13857453194	What is the ln[concentration]-time graph for a first order reaction? What does the slope of this line represent?	slope = rate constant, k, for rate = k[A]1	136
13857453195	When dealing with time and concentration, how does the half-life of a 1st order reaction relate to the rate constant?	half-life time = 0.693/k or k=0.693/half-life time Always look for half-life in 1st order rate law problems to quickly determine rate constant.	137
13857453196	2nd order reaction curve of 1/x vs time What does the slope of this line represent?	Slope of 1/[A]=k: rate = k[A]2	138
13857453197	Units for reaction constant k	Zero order: k unit = M/s First order: k unit = 1/s Second order: k unit = 1/Ms The molarities of the reactants must multiply out to produce M/s	139
13857453198	Energy barrier to the formation of products that determines the reaction rate	Ea, the activation energy, is always endothermic Activation energy does not change the deltaH of the reaction	140
13857453199	What two factors are required for a successful activated complex collision?	1. The collision must have sufficient energy to create the activated complex. 2. The collision must have the proper orientation for the collision to make the activated complex. On FRQs regarding successful collisions, both factors must be mentioned for credit.	141
13857453200	What is kinetic control of a reaction?	High activation energies may slow the rate of a thermodynamically favored reaction so much that it may not occur because none of the collisions can be successful to make the product	142
13857453201	What does a catalyst do?	Catalyst lowers the Ea for both forward and reverse reactions. A catalyst allows the reaction to reach equilibrium faster.	143
13857453202	What doesn't a catalyst do?	A catalyst doesn't change the equilibrium constant K. Catalyst does not change deltaG or deltaH	144
13857453203	How is a catalyst identified in a multi-step reaction? How is an intermediate identified in a multi-step reaction? Which of the two can be included in a rate law?	Step 1: A+B=C Step 2: C+D = E+B Net: A+D = E B, the catalyst, is present as a reactant and produced as a product later in a multi-step reaction C is an intermediate, and even if it were in the slow step, would never be included in a rate law expression B is a catalyst and may be included in a rate law rate = k[a][b]	145
13857453204	How to calculate deltaH from BE	BErxn = BEreactants = BEproducts	146
13857453205	In a reaction pathway energy diagram, how is the slow step identified?	The slow step has the higher speed bump (bigger activation energy)	147
13857453206	In a multi-step reaction, which step determines the rate law expression?	The slowest step in the reaction determines the reactants in the rate law expression	148
13857453207	How does an equilibrium reaction affect the reaction rate?	An equilibrium reaction preceding a slow step will result in the reactants of the fast equilibrium step to be included in the rate law expression Step 1: A=X fast Step 2: X+A=>B slow Step 1 has an equilibrium that will affect [X]. A can be used to substitute [X] rate = k[A]^2	149
13857453208	What does the Boltzmann Molecular Speed graph look like at different temperatures? What is plotted on the X and Y axis, and how does it explain the increase of a reaction rate?	At a higher temperature, the total number of molecules is unchanged, but a greater percentage have the energy for successful collisions. y axis is the number of molecules at a given speed x.	150
13857453209	Keq	The equilibrium constant is the ratio of products/reactants Keq > 1 means a greater concentration of products than reactants when the reaction has reached equilibrium Keq < 1 means a greater concentration of reactants than products when reaction reaches equilibrium An equilibrium constant in the area of millions means the reaction will go to completion A really tiny Keq means very little product will be made	151
13857453210	Kc and Kp	Kc = Kp only when mole gas reactant = mol gas product	152
13857453211	Solids and liquids in the Keq expression	(s) and (l) are never included in the equilibrium expression because the concentrations of pure substances in a reaction rarely change	153
13857453212	Temperature and the equilibrium constant	K is the value of the equilibrium expression at equilibrium. It is only changed by temperature. In exothermic rxn, K decreases with increases in pressure In endothermic rxn, K increases with increases in pressure	154
13857453213	Eq quotient and Eq constant	Q < K Reaction will approach the products side Q > K Reaction will approach the reactants side Q=K equilibrium has been reached	155
13857453214	K for a multi-step reaction where equations add up to make an overall reaction	Koverall = Kstep1 x Kstep2	156
13857453215	Equilibrium constant relationship between forward and reverse reactions	Kforward = 1/Kreverse	157
13857453216	Common ions	The common ion is either the cation or anion of the dissolving substance that is added separately to the equilibrium system.	158
13857453217	What is a polyprotic substance?	A polyprotic acid can donate more than one proton A polyprotic acid will have more than one vertical line in its titration curve The Ka of the removal of all the protons is the product of each proton's Ka	159
13857453218	What is an amphiprotic substance?	A substance that can donate or accept a proton and thus act as an acid or base. H2O and HCO3- are amphiprotic	160
13857453219	Neutral water Relationship between [H] and [OH] pH and pOH Kw @ 25 C pKw @ 25 C Temperature changes for Kw and pKw	Neutral water is when [H+] = [OH-] and pH = pOH At all temperatures Kw = [H] x [OH] and pKw = pH + pOH neutral water: pH = 7, pKw = 14, Kw = 1e14 As temp increases, pKw decreases and Kw increases	161
13857453220	Bronsted acids	proton donors	162
13857453221	Bronsted bases	proton acceptors	163
13857453222	How are conjugates related to the original acid-base and which will be favored at equilibrium	An acid reactant will become the conjugate base. A base reactant will become the conjugate acid. The stronger of the two acids (acid or conjugate acid) will be present in lower concentrations at equilibrium)	164
13857453223	Ka and Kb relationship in an acid and its conjugate base	1e-14 = Ka x Kb 14 = pKa + pKb	165
13857453224	Molecular, Ionic and Net Ionic reaction	Molecular: HA + BOH -> H2O + AB Ionic: H+ + A- + B+ + OH- -> H2O + A- + B+ Net Ionic: H+ + OH- -> H2O	166
13857453225	Net ionic reaction for a weak acid and a strong base	Weak acid, strong base Net ionic: HA+ OH- -> A- + H2O	167
13857453226	NIE for a weak base and a strong acid	Weak base, strong acid H+ + B -> HB+	168
13857453227	Equivalence point in titration and indicators	Equivalence point means that the solutions have been mixed and all the acid HA has reacted and been changed into A- An indicator changes color at a specific pH range which is the pKa of the indicator The idea indicator has a pKa = to the pH at the equivalence point.	169
13857453228	Overtitration	Once past the equivalence point, the excess strong acid or base moles and volume of solution are used to determine the pH or pOH directly without the need for a Keq expression	170
13857453229	Half-equivalence point	When o.5 of the [HA] has turned into [A-] remaining [HA] = [A-] Ka = [H+] pKa = pH	171
13857453230	Buffers calculations	Henderson Hasselbalch equation pH = pKa + log [A+]/[HA]	172
13857453231	strong acid titrated with strong base	pH equivalence = 7	173
13857453232	weak acid titrated with strong base	pH>7 large amounts of conj. base	174
13857453233	weak base titrated with strong acid	pH<7 large amounts of conj. acid	175
13857453234	Diprotic Acid titration curve	Diprotic acid curve H2A titrated with strong base Two equivalence points Two half-titration points	176
13857453235	What is S?	S is entropy S can be measured in absolute value S of elements is not 0	177
13857453236	How does the S of phases of a substance compare?	Comparative entropy of the phases s	178
13857453237	What reactions increase S?	Decomposition reactions Reactions that produce gases	179
13857453238	What are the units of S?	J/K	180
13857453239	deltaSrxn	deltaSrs = Sproducts - Sreactants	181
13857453240	What is deltaG	deltaG is standard Gibbs free energy deltaG = 0 equilibrium deltaG is negative, thermofavored deltaG is positive, nonfavored	182
13857453241	What are its units of deltaG	kJ/mol	183
13857453242	deltaG equation	deltaG = deltaH - tdeltaS	184
13857453243	Oxidation number of any pure element	Zero	185
13857453244	What happens in oxidation	Loss of electrons, at the anode	186
13857453245	What happens in reduction	Gain of electrons, at the cathode	187
13857453246	Standard voltage calculation	Ecell = Ecathode - Eanode	188
13857453247	Galvanic cell equation and standard voltage for hydrogen reduction	Half reaction 2H+ + 2e- = H2 Voltage = 0	189
13857453248	Nernst equation	Ecell = E˚cell - (RT/nF) * ln Q	190
13857453249	Galvanic vs Electrolytic Cells	Galvanic cells do work and run spontaneously deltaG = - Ecell = + Keq is large Galvanic cells will usually be shown with the solutions of the two half-cells separated with a salt bridge Electrolytic cells must have a power source that forces the electrons to the cathode Reduction occurs at the cathode Oxidation occurs at the anode deltaG is positive Ecell is negative Keq is close to 0 Salt bridge is not necessary	191
13857453250	Where to find the number of coulombs of electrons in a mole	Faraday's constant: 96485 coulombs/mol e-	192
13857453251	Electrolytic cell current-time calculations for mol of substance reduced	1) amps x seconds = coulombs of electrons 2) coulombs of electrons are divided by Faraday's constant 3) mole e- from electrolysis/ electrons used in reduction reaction	193
13857453252	deltaG and E	deltaG = -nFE Equation is on equation sheet deltaG will be in J/mol	194
13857453253	IMF's	There are always IMFs between molecules.	195
13857453254	PV=PV	Boyles Law	196
13857453255	V/T = V/T	Charles law	197
13857453256	As Kw increases	pH decreases	198
13857453257	In combustion reactions, remember	that H2O has two hydrogen atoms	199
13857453258	Stronger acids	weaker conjugate base	200
13857453259	Stronger base	weaker conjugate acid	201
13857453260	Peroxide oxidation	Oxygen's oxidation number is always -1, rather than -2.	202
13857453261	Intermolecular forces present	Only in covalent molecules, because ionic ones are bonded by ionic forces only.	203
13857453262	HI BrONClF	Mnemonic for the diatomics: Hydrogen Iodine Bromine Oxygen Nitrogen Chlorine Fluorine	204