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142 Cards in this Set
- Front
- Back
energy change from electron transitions
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E=R(H)(1/ni^2-1/nf^2)
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angular momentum quantum number
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l, takes any values from 0 to n-1, a given n shell contains all the orbitals from l=0 to l=n-1
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magnetic quantum number
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describes orbital orientation, each can hold up to 2 electrons, each orbital has 2l+1 orientations, magnetic quantum number takes on values from -l to +l (and zero)
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spin quantum number
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±1/2, electrons in each orbital with a particular orientation don't repel each other because they have opposite spins, the magnetic fields they create causes a magnetic attraction
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Pauli exclusion principle
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every electron can be described by a unique set of quantum numbers n, l, ml, ms
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maximum electrons in a shell
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max = 2n^2
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Aufbau principle
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electron filling of orbitals based on energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s...
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Hund's rule
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when filling orbitals an electron is always added to an empty orbital before pairing up and has the same spin as other unpaired electrons
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stability of very large orbitals
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especially stable when full or half filled (d or f), sometimes s electrons from a higher n value than that larger orbital can donate an electron to give a more stable large orbital
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representative elements
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elements with either s or p orbitals for valence orbital (IA through VIIIA)
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metals
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lust, malleable, conduct electricity/heat, lose electrons easily, most bonding is ionic
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nonmetals
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lower molar masses tend to be gases in elemental state, groups V, VI, VII can gain electrons to form anions, most bonding is covalent
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metalloids
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silicon, germanium, arsenic, antimony, tellurium, polonium, astatine; down a column (groups III to VII) increase metallic character (used as semiconductors)
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alkali metal group
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hydrogen is an exception to the properties, generally soft, low-melting, lustrous metals, react violently with water (form hydroxide salt and H2 gas), more reactive as atomic number increases
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alkaline earth metals
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meryllium has different properties (nonmetal), generally fairly soft with low density, readily form oxides and hydroxides, oxides and hydroxides are generally insoluble (barium hydroxide is a soluble strong base)
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oxygen group
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group VIA have their number of possible oxidation states increase with atomic number, oxygen reacts well with most metals but reactivty decrease down the column, selenium/tellurium are semiconductors, sulfur is an electrical insulator
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halogens
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elementally they exist as diatomic molecules
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transition metals
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lose electrons easily and form ionic compounds, ionic states can vary from +1 to +8, they lose electrons from a higher n value shell first, cations with half or full d orbital are very stable, many have more than one stable oxide (oxide with higher % oxygen forms most easily)
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ionization energy trends
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decreases down a group, increases from left to right
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electron affinity
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energy change that occurs when electrons are added to the valence orbital (the greater the Ea, the more stable the anion that formed is), increases across period to VII then drops at the noble gases and then increases again; each group has approximately equal Ea's, Ea is higher if an electron is needed for a half or full d orbital
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electronegativity
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atom's ability to pull electron density toward itself
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effective nuclear charge
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Zeff increases going across a row and up a column, the attractive force on the electrons increases and atomic radius decreases
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atomic radius
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size decreases across period (electron number increases but Zeff does too), Zeff decreases down a group so radius increases
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ionic radius
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the more electrons an ion has gained, the larger its radius; isoelectronic (ions with same number of electrons) have decreased radius across a row and increased down a column
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lattice energy
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energy required to break up the lattice into individual ions, energy given off when lattice is created is -U where U=kz1z2/d (d is avg. distance between nuclei, z1 and z2 are ion charges)
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sp3d2 geometry
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octahedral, 90 degree bond angle
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sp3d geometry
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trigonal bipyramidal, 120 and 90 degree bond angles
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lone pairs and bond angles
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as the number of lone pairs increase the bond angle between atoms on the central atom decreases
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trigonal pyramidal geometry
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three bonds and a lone pair
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bent geometry
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two bonds and two lone pairs
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kinetic molecular theory
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1. a gas consists of very small particles that move randomly
2. the volume of each gas particle is negligable compared to the spaces between particles 3. there are no intermolecular attractive forces between the gas particles 4. when gas particles collide with each other or with the walls of the container, there is no net gain or loss of kinetic energy 5. the average kinetic energy of each gas particle is proportional to the temperature |
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kinetic energy of gas particle
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Ek=(1/2)mv^2
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velocity of gas particle
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v=(3RT/M)^(1/2)
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real gas conditions
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high pressure and/or small volume can cause intermolecular attractive forces, when pressure increases greatly then repulsive forces become important (makes actual volume slightly above that predicted by ideal gas law)
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real gas equation
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(P+an^2/V^2)(V-nb)=nRT
(1st part is correction for pressure/intermolecular forces; 2nd part is correction for volume/size of gas particles) |
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typical phase diagram
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solid/liquid line has a positive linear slope; liquid/gas line has a concave up curve; solid/gas line has a slighter concave up
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vapor pressure equation
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ln(P2/P1)=(∆Hvap/R)(1/T1-1/T2)
(vapor pressure is constant at any given temperature, boiling point is T1) |
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molality
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mole solute/kg solvent
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mole fraction
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Xsolute=mole solute/mole total
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vapor pressure depression of solution
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∆P=XsolutePº
(Pº=vapor pressure of pure solvent) actual vapor pressure is found by using P=Pº-∆P |
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boiling and freezing point elevation and depression in solution
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∆Tb=Kb x molality
∆Tf=Kf x molality (can be multiplied by the number of particles per molecule when dissolved) |
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osmotic pressure
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πV=nRT (π=osmotic pressure)
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Henry's law for dissolved gases
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amount of gas dissolved is directly proportional to the pressure of the gas above the solution (P=kX, X is mole fraction of gas and P is pressure above solution), law works best when gas molecules remain intact in solvent and don't react
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colloids
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suspension of small particles in a medium, distinguished from solids by Tyndall effect where they scatter light
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strong acid with weak base reaction
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ionic compound that forms is acidic in solution
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weak acid with strong base reaction
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ionic compound that forms is basic in solution
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enthalpy of reaction
(aA+bB-->cC+dD) |
∆Hºrxn=[∑c∆Hf(C)+d(∆Hf(D))]
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enthalpy of formation
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∆Hºf=∑BE bonds broken-∑BE bonds formed
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Hess's law
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the sum of the ∆H's for each individual equation sum to the ∆Hrxn for the net equation (reverse the sign of ∆H if reverse reaction, multiply entire reaction and ∆H by scalar)
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half life of 1st order reaction
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ln(2)=kt(1/2)
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Arrhenius equation
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ln(k2/k1)=(Ea/R)(1/T1-1/T2)
(k1 is rate constant at T1, k2 is rate constant at T2, greater at higher T) |
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equilibrium constant factors
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temperature affects K, amount of starting material and products affects K
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problems involving small K
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starting material change is negligable, product amounts are small
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Le Chatelier and volume
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increase V (decrease P), make more gas molecules; decrease V (increase P), remove gas molecules
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effect of pH on basic compound solubiltiy
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lower pH increases solubility of basic ionic compounds
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six strong acids
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HCl, HBr, HI, HNO3, H2SO4, and HClO4
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common strong bases
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LiOH, NaOH, KOH, Ba(OH)2 (for barium hydroxide the base concentration is half the hydroxide concentration, pH greater)
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buffer
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solution that resists change in pH and has approximately equal concentrations of an acid and its conjugate base
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Henderson-Hasselbach equation
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pH=pKa+log([A-]/[HA])
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enthalpy and internal energy and work
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H=U+PV, at atmospheric pressure ∆H=q
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second law of thermodynamics
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for a sponataneous process ∆S must be > q/T, at equilibrium ∆S=q/T=∆H/T
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third law of thermodynamics
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a substance that is perfectly crystalline at 0K has an entropy of zero
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processes with increasing entropy
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more gas produced than used up, solid converted to liquid, solid dissolved in a solvent
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processes with decreasing entropy
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less gas produced than used up, liquid converted to solid, solid precipitation from a solution
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standard free energy
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∆Gº=∆Hº-T∆Sº
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work and ∆G
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maximum work for a spontaneous reaction is equivalent to ∆G, the free energy change is the maximum energy available to do work
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K and ∆G
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not at equilibrium: ∆G=∆Gº+RTlnQ
at equilibrium: ∆Gº=-RTlnK (∆G=0) |
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relationship between equilibrium constant and temperature
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ln(K2/K1)=(∆H/R)(1/T1-1/T2)
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balancing redox
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start with main atoms, add H2O to balance oxygen, add protons to balance hydrogen
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voltaic/galvanic cells
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the redox reactions in two compartments produce an electric current
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anode half-cell
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oxidation, electrode wears away becoming ion in solution, gives up electrons
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cathode half-cell
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reduction, electrode builds up with metal ions in solution plating out as pure metal, receives electrons and uses them
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cell notation for galvanic cells
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anode|anode's ion||cathode's ion|cathode
(single line means phase change, double line means salt bridge) |
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work produced by a cell
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w(max)=-nFEcell
work is negative when work is produced |
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Eºcell
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cell potential uner standard conditions of 25 C, 1M concentrations of solutions, 1 atm pressure of gases
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spontaneous reaction listing order
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anode reaction is placed above cathode reaction
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sign of standard potential and reaction manipulation
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only sign changes, the value never changes with differening numbers of electrons
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free energy and reduction potential
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∆Gº=-nFEºcell
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reduction potential and equilibrium
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Eºcell=(RT/Fn)lnK
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Nernst equation
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calculate nonstandard cell potential when concentrations are other than 1M: Ecell=Eºcell-(RT/nF)lnQ
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electrolytic cells
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constructed for nonspontaneous redox reactions that require electricity to make them run
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calculating current needed to deposit a given mass of metal by electroplating
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need Faraday constant, electrons (n), time in seconds, molar mass of metal
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radioactive decay equation
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[A(t)]=[Ao]e^(-kt)
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collision theory
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rxn rate = fZ
(f=fraction of effective collisions with enough KE, z=total collisions) |
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factors that affect rxn rate
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1. reactant concentrations (higher z in r=fZ)
2. higher temp (greater KE, higher f) 3. medium polarity/phase 4. catalyst addition |
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equilibrium and reaction rates
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forward rate equals reverse rate
Keq is equal to the ratio of the rate constants (multiply if mroe than one step) |
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systems
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isolated, closed, open
isothermal, adiabatic, isobaric |
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state functions
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pressure, volume, temperature, enthalpy, entropy, free energy, internal energy
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heat of formation
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enthalpy change if one mole of a compound is formed directly from its elements in their standard states
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STP conditions
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0 C (273K), 1 atm
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density of ideal gas
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d=(MM x P)/RT
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molecular speeds
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avg. KE=(3/2)kT
avg speed=(3RT/MM)^(1/2) at higher temp, peak shifts toward higher KE/speed and down at peak |
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diffusion/effusion
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(r1/r2)=(MM2/MM1)^(1/2)
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liquid attractive forces
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the forces overcome the KE that can keep them apart as in the gas phase
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solids
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strong attractive forces that only have vibrational motion of atoms
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ionic solids
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high MP and BP, poor solid state electrical conductivity
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metallic solids
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metal atoms are packed together with high MP, BP (covalent attractions), layer of spheres shape
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crystal unit cells
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simple cubic (cube), body-centered cubic (simple plus 1 in middle), face-centered cubic (simple plus one on each face)
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gas-liquid equilibrium
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gas forms when high KE particles leave liquid and pressure over gas can force some back into liquid, vapor pressure is pressure gas exerts over liquid when condensation-evaporation equilibrium occurs
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vapor pressure multiple components
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gas pressure exert by A and B over a combined solution are the same as each exerts over its pure solution
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Raoult's law
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solute B added to solvent A causes vapor pressure of A to decrease where the change is ∆P=Paº-Pa=XbPaº (pure solvent is when Pa=XaPaº), attraction of A and B molecules are the same as with themselves
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solubility: always soluble
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salts of alkali metals, ammonium salts
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halide solubility
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soluble except with Ag+, Pb+, Hg2(2+)
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sulfate solubility
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soluble except Ca2+, Sr2+, Ba2+, Pb2+
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metal oxide solubility
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insoluble except with alkali metals, ammonium, CaO, SrO, BaO (hydrolyze to soluble hydroxides)
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hydroxide solubility
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insoluble except with alkali metals, ammonium, Ca2+, Sr2+, Ba2+
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carbonates, phosphatess, sulfides, sulfites
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insoluble except with alkali metals and ammonium
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Arrhenius acid/base
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acid produces H+, bases produce OH- (only true in aqueous media)
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Bronsted-Lowry acid/base
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acid is a proton donor, base in a proton acceptor (works for nonaqueous)
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Lewis acid/base
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acid is electron pair acceptor, base is electron pair donor
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angular momentum of electron
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ang. mom.=nh/2π
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Lyman, Balmer series
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n>1 to n=1, n>2 to n=2
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paramagnetic atoms
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have unpaired electrons, a magnet will align the spins and weakly attract the atom
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dimagnetic
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have no unpaired electrons, a magnet will slightly repel it (antimagnetic)
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complex ions
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transition metals form them with water or nonmetals, d orbitals split and d electrons can absorb light
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octet exceptions
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hydrogen (2), lithium (2), beryllium (4), boron (6), 3rd row (+8)
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formal charge
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valence-.5bonding-nonbonding
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van der waals forces
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1. dipole-dipole (solid, liquid, not gas), higher BPs
2. H-bonds, high BPs 3. London dispersion (electron density shifts around randomly, close molecules can interact) |
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reaction rate (aA+bB->cC+dD)
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rate=-(1/a)(∆[A]/∆t)=-(1/d)(∆[D]/∆t)
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molar volume of gas at STP
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1 atm=22.4mol/L
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Dalton's laws of partial pressure
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partial pressure of gas equals mole fraction times total pressure, total pressure equals all gas components partial pressure added up
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disproportionation reactions
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reaction where an element is both oxidized and reduced at the same time
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zeroth law of thermodynamics
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if two thermodynamic systems are in thermal equilibrium with a third, then they are also in thermal equilibrium with each other
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energy equivalence
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same units for mechanical (KE/PE), chemical, electrical, and thermal energy
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temperature scales
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Tk=Tc+273
Tf=32+(9/5)Tc |
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PV work
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pressure versus volume curve where work done is area under or enclosed by curve
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thermodynamic control
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depends on stability of the product, occurs either in vigorous reaction conditions (heat) or over long time when equilibrium is allowed to settle out
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kinetic control
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occurs best in non-equilibrium situations when present with a thermodynamically stable reaction because its low Ea will be overcome more easily even if its product is less stable
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Law of mass action
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can explain thermodynamic or kinetic situations (the Keq for thermodynamics and the rxn rate for kinetics)
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common ion effect and separation
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addition of soluble salt with common ion of slightly soluble salt will push the slightly soluble salt further towards precipitation
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complex ion formation
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electrophilic metal can bind to electron-donating groups like water and nonmetals, metals become more soluble in aqueous phase as complexes
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electrolysis
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uses current to separate bonded elements and compounds
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electrolytic electrolyte function
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because water is not a good enough conductor, an electrolyte must be added to solution
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electrolytic cell electron pathway
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inert electrodes are used, electrons flow from positive anode to negative cathode, negative cathode precipitates a positive ion when it collides with negative electrode and picks up an extra electron, positive anode takes an electron from an anion to precipitate it and complete the circuit
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Faraday's laws of electrolysis
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the mass of a substance produced at an electrode during electrolysis is proportional to the number of moles of electrons transferred at that electrode; the number of faradays of electric charge required to discharge one mole of substance at an electrode is equal to the number of excess elementary charges on thta ion
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mass of substance deposited or gas liberated at electrolytic electrode
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m=(1/F)(QM/n)
(m=mass, F=constant, Q=current, M=molar mass, n=number of electrons on ion) |
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electron flow in galvanic cells
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anode is where oxidation takes place (solid to ion) and electrons are produced, moved in wire to positive cathode where reduction takes place (ion to solid), salt bridge prevents charge build-up
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Henry's law
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C=k x Pgas
(C=gas solubility in solvent, k=constant, Pgas=partial gas pressure of that gas of itnerest Higher pressure of that gas over the liquid phase causes greater solubiltiy |
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Raoult's law
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dissolved solute lowers vapor pressure of a solvent, the partial pressure of each component equals the mole fraction times the pressure of that liquid only
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osmotic pressure qualitative
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the pressure necessary to stop osmotic flow, depends only on the concentration (not nature) of the solute (the concentration in the osmotic equation is based on the solution concentration)
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∆H and bond energy
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∆Hrxn=∑BE(reactants)-∑BE(products)
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redox titrations
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determination of the concentration of an oxidant or reductant based on titration with opposite redox partner
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