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