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223 Cards in this Set
- Front
- Back
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mass of a proton
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1 amu
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mass of a neutron
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slightly larger than a proton
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valence electrons
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furthest electrons from nucleus; determine reactivity of an atom
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1 amu
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1/12 the mass of C-12 = 1.66 E-24 g, or 1 gram = 6.022 E 23 amu
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Avogadro's number
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6.022 E23
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atomic weight
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weight (in grams) of one mole of an element (g/mol)
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isotopes
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same atomic number, different mass # (different # of neutrons)
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Planck proposed...
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energy emitted as electromagnetic radiation from matter comes in bundles called quanta, where 1 quanta has E = hf
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Planck's constant
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a proportionality constant, = 6.626 E-34 J*s
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Rutherford
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discovered that atoms have a dense, low volume, positive center
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Bohr proposed...
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Electrical force holds electrons in orbit and angular momentum (L=mvr) is quantized for electrons, so L=nh/(2*pi). Allowed values of L were equated with allowed electron energies to give E=-R(H)/n^2
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Rydberg constant
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R(H) = 2.18 E-18
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Energy of an electron according to Bohr
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E = -R(H) / n^2 (energy of an electron is quantized with respect to its principle quantum number)
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energy of a photon
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E = hc/(lambda) or E = hf
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atomic emission spectrum
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a line spectrum; each line corresponds to a different electronic transition
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Balmer series
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group of hydrogen emission lines corresponding to transitions form n>2 to n=2
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Lyman series
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transitions from n>1 to n=1 in the hydrogen emission spectrum
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energy of an electronic transition
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E = hc/lambda = E(i) - E(f) = R(H) * (1/n(f)^2 - 1/n(i)^2) ... positive if energy is released
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what does the Bohr model not explain?
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atomic structures with more than one electron b/c it does not consider electron-electron repulsion
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Heisenberg uncertainty principle
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impossible to determine with perfect accuracy the momentum and position of an electron simultaneously
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Pauli exclusion principle
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no two electrons in a given atom have the same set of 4 quantume numbers
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energy state
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the position and energy of an electron as described by its quantum numbers
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the quantum numbers
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principal, azimuthal (or angular momentum), magnetic, and spin
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n
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principal quantum #; larger means increased radius and energy; max number of e- in energy level n is 2n^2
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energy difference between adjacent electron shells
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decreases as distance from nucleus increases
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azimuthal quantum number
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the subshells within each principal energy level; has a range of 0 to n-1 (s, p, d, and f); max number of e- within a subshell is 4l+2
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magnetic quantum number
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specific orbital within a subshell; has range l to -l
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spin quantum number
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the intrinsic angular momentum, + or -1/2
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parallel spins
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electrons in different orbitals with same spin quantum number
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ranking energy level of a subshell?
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use (n+l); if 2 subshells have same value, the subshell of lower n has lower energy
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Hund's rule
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within a given subshell, orbitals are filled such that there are a max number of 1/2 filled orbitals with parallel spins (electrons prefer empty orbitals to avoid a pairing energy barrier)
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paramagnetic
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unpaired electrons = a magnetic field can align the spins and weakly attract the atom)
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diagmagnetic
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no unpaired electrons = slightly repelled by a magnetic field
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periodic law
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chemical properties of the elements are dependent in a systematic way upon their atomic numbers
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A elements
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representative = have either s or p sublevels as valence orbitals
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B elements
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nonrepresentative = partially filled d or f sublevels are valence orbitals
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atomic radii
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half the distance between centeres of 2 atoms that are just touching each other; decreases from L to R and up a group
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why does atomic radius decrease across a period?
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electrons within a shell cannot shield one another form attractive pull of the nucleus; and as the number of protons increases, the effective nuclear charge also increases
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ionization energy
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energy required to remove e- from gaseous atom or ion; increases from L to R and up a group
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electron affinity
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change in energy that occurs when e- is added to a gaseous atom (a greater Z(eff), a greater EA); if positive, then energy is released; Group VIIA have highest
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Z(eff)
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effective nuclear charge
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electronegativity
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measure of attraction for e- in a chemical bond (a result of the perceived Z(eff) by the e- in a bond); most common system is the Pualing scale where 0.7 is most electropositive (Cs) to 4.0 for most electronegative (F); increases from L to R and down a group
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Cesium
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most metallic, least electronegative of all naturally occuring elements; also has smallest ionization energy and least exothermic electron affinity
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Florine
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most electronegative, largest/most exothermic electron affinity
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malleability vs. ductility
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ability to be hammered into shapes vs. drawn into wires
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the metalloids
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B, Si, Ge, As, Sb, Te
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Alkali metals
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group IA, lower densities than other metals and largest atomic radii in each of their respective periods
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Group IIA
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Alkaline earths
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Group VIIA
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Halogens
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Transition elements
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groups IB to VIIIB, have high bp and mp, low ionization energy (various oxidation states poss)
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exceptions to the octet rule
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H, Li, Be, B, and elements beyond the second row
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ionic bonding
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held together by electrostatic forces; occurs between atoms with a difference in electronegativity >1.7; less electronegative atom forms the cation
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ionic compounds
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high mp and bp due to strong electrostatic forces; can conduct electricity in liquid and aqueous states; forms crystal lattices of infinite arrays (max attractive forces, min repulsive)
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covalent compounds
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discrete molecular units with weak intermolecular forces, therefore have low melting solids and do not conduct electricity
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bond order
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# of shared electron pairs between 2 atoms
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bond length
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average distance between the nuclei of 2 atoms in a bond (decreases with increased bond order)
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bond energy
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energy required to separate two bonded atoms; increases with increased bond order
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formal charges
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#valence e- in free atom - 1/2 bonding electrons - nonbonding e-
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stabiltiy of a resonance structure
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small or no formal charges; place negative formal charges on more electronegative atoms
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dipole moment
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charge * distance between partial charges = q*r, measured in Debye units (C*m); a vector quantity
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coordinate covalen bond
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shared electron pair comes from the lone pair of a molecule (e.g. between Lewis acids and bases)
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VSEPR
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valence shell electron pair repulsion theory = uses Lewis structures to predict molecular geometry; arrange electron pairs as far apart as possible)
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linear
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two regions of e- density, 180 angle, BeCl2
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trigonal planar
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3 regions or e- density, 120 angles, BH3
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angular
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3 regions of e-density, water
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tetrahedral
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4 regions of e-density, 109.5 angles, CH4
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trigonal pyramidal
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4 regions of e-density, NH3
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trigonal bypyramidal
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5 regions of e- density, 90/120/180 angles, PCl5
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octahedral
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six regions of e- density, 90 and 180 angles, SF6
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molecular orbitals
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obtained by adding the wave functions of atomic orbitals
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bonding vs antibonding
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signs of 2 atomic orbitals match vs. don't match
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sigma vs. pi bonds
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head to head overlap vs. parallel
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van der Waals forces
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dipole-dipole interactions and dispersion forces
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intermolecular forces
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van der Waals and hydrogen bonding
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dipole-dipole interactions
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between polar molecules in solid and liquid phases (gas phase puts atoms too far apart); leads to higher bp than np species of comparable weight
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hydrogen bonds
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specific form of dipole-dipole and can be intra- or intermolecular; H bound to F, O, or N carries little e- density (leads to much higher bp)
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London forces
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also dispersion forces - random location of e- allows unequal sharing of e- and leads to rapid polarization or e- cloud; Strength depends on how easily e- can move... allows noble gases to liquify
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1 atm is equivalent to how many mm Hg or torr?
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1 atm = 760 mm Hg = 760 torr
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standard temperature and pressure (STP) vs. standard conditions
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273.15 K (0*C) and 1 atm vs. 25*C (298 K)
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an ideal gas
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a hypothetical gas whose molecules have no intermolecular forces and occupy no volume; many gases behave in a nearly-ideal fashion at low pressures (1 atm) and high temperatures.
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Boyle's law
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PV = constant; or P1V1 = P2V2
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Charles' law
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or the Law of Charles and Gay-Lussac; V/T = constant or V1/T1 = V2/T2
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Avogadro's Principle
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n/V = constant; or n1/V1 = n2/V2
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Ideal Gas Law
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PV = nRT
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gas constant
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R = 8.21 E-2 L*atm / mol*K = 8.314 J/K*mol
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volume of 1 mol of gas under STP conditions
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22.4 L
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moles from mass and molar mass
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n = m (in grams) / MM (molar mass)
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quick calculation of density for a simple gas at STP
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density = MM (g/mol) / 22.4 L/mol
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find molar mass of sample, given measurements at nonstandard T and P
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find volume at STP, using V1 = V2*(P2/P1)*(T1/T2) -- derived from PV/T = constant -- then mass/V1 = density at STP; multiply density at STP by 22.4 L/mol to find molar mass
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deviations from ideal gas behavior due to pressure
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as pressure increases, the gas approaches the condensation pressure for a given T, and intermolecular attraction forces become more significant until the gas condenses into the liquid state. At moderately high pressure (a few 100 atm), a gas' volume is less than would be predicted due to intermolecular attraction. BUT, at extremely high P, the size of the particles becomes relatively large compared to the distance between them, and the gas takes up a larger volume than predicted.
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deviations from ideal gas behavior due to temperature
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decreased T results in a decreased average velocity, and attractive intermolecular forces become increasingly significant. As the T approaches the condensation point, the intermolecular forces cause the gas to have a smaller volume than predicted.
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van der Waals equation of state
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(P + n^2a / V^2)*(V - nb) = nRT, where a and b are physical constants, experimentally determined for each gas. 'a' corrects for attractive forces (so will be smaller for a gas such as He), and 'b' corrects for the volume of the molecules themselves (so larger b for larger molecules). a is generally larger than b.
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Dalton's law of partial pressures
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total pressure of a gaseous mixture is equal to the sum of their partial pressures
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determining partial pressure of a gas
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P(a) = P(total)*X(a) where X(a) = n(a) / n(total) = moles of a / total moles
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assumptions of the kinetic molecular theory of gases
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gas particles have negligible volume compared to their container; no intermolecular attractions or repulsions; continuous, random motion; collisions between any 2 gas particles are elastic, i.e. conservation of energy; average KE of gas particles is proportional to the absolute T of the gas, and is the same for all gases at a given T.
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average kinetic energy (of a gas)
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KE = 1/2 mv^2 = (3/2) *kT, where k is the Boltzmann constant
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average molecular speed (of a gas)
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root-mean-square speed, u(rms) = sqrt (3RT / MM), where R is the gas constant and MM is molecular mass
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Maxwell-Boltzmann distribution curve
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shows the distribution of speeds of gas particles at a given T (# molecules vs. molecular speed); the curve flattens and shifts to the right as T increases, indicating that more molecules are moving at high speeds at higher T.
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Graham's law of diffusion
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under isothermal and isobaric conditions, the rates at which two gases diffuse are inversely proportional to the square root of the molar masses; r1/r2 = sqrt (MM2 / MM1)
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Graham's law of effusion
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for two gases at the same T, the rates of effusion are proportional to the average speeds; in terms of molar mass, the relationship is the same as that for diffusion: r1/r2 = sqrt (MM2/MM1)
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effusion
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the flow of gas particles under pressure from one compartment to another through a small opening
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condensed phases
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liquids and solids, due to their smaller volumes relative to gases
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miscibility
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degree to which two liquids can mix
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emulsion
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a fairly homogenous mixture of two immiscible liquids due to extremem conditions (e.g., violent shaking of oil and water); emulsions are actually mixtures of particles too small to be seen distinctly
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two most common forms of crystals
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metallic and ionic
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ionic solids
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high mp, high bp, and poor electrical conductivity due to the compound's strong electrostatic interactions, which also cause ions to be relatively immoble.
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metallic solids
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high mp and bp as a result of their strong covalent attractions; pure metallic structures are usually described as layers of spheres of roughly similar radii.
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unit cells
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the repeating units of crystals (both ionic and metallic); many types, including simple cubic, body-centered cubic, and face-centered cubic.
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phases changes in an isolated system
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are reversible; and an equilibrium exists between phases
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evaporation
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molecules near the surface of a liquid may have enough energy to leave the liquid phase and escpae into the gaseous phase; this is a cooling process, since the temperature of the remaining liquid decreases every time it loses a high-energy particles.
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condensation
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escaping molecules from a liquid are trapped by a cover above the solution. These molecules exert a countering pressure, which forces some of teh gas back into the liquid phase.
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vapor pressure
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the pressure that gas exerts over the liquid, once an equilibrium between condensation and evaporation has been reached. Vapor pressure increases as temperature increases, since more molecules have sufficient kinetic energy to escape into the gas phase.
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boiling point
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the point at which the vapor pressure equals the external pressure
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melting point of amorphous solids
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tend to melt over a large range of temperatures due to the less-ordered molecular distribution of an amorphous solid (e.g., glass)
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sublimation
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a phase change from solid directly to gas (e.g. dry ice, solid CO2)
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deposition
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the reverse of sublimation, from gas to solid
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phase equilibria and Gibbs function
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the change in Gibbs free energy must equal zero for an equilibrium between two phases to exist
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phases and their relative T and P
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in general, gas = high T, low P; solid = low T, high P; liquid = high T, high P
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lines of a phase diagram
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represent the T and P at which two phases are in equilibrium
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triple point
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the intersection of the three lines indicates the T and P at which all three phases are in equilibrium
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critical point
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the point at the end of the evaporation/condensation line that indicates the T and P above which no distinction between liquid and gas is possible
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phase diagram of water is different in that...
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the freezing/melting line slopes to the left (instead of right) b/c the solid state is less dense than the liquid
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Raoult's Law
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enables one to determine the relationship between the vapor pressure of vapor A and the concentration of liquid A in solution;
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colligative properties
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physical properties dervied solely from the number of particles present, not the nature of those particles (these properties are usually associated with dilute solutions)
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freezing-point depression
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solutes lower the mp of the solvent b/c the solute particles interfere with the process of crystal formation; specifically, delta T(f) = K(f)*m where delta T(f) is the freezing-point depression, K(f) is a proportionality constant characteristic of a particular solven, and m is the molality of the solution.
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molality
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mol solute / kg of solvent; for dilute aqueous solutions at 25*C, the molality is approximately equal to the molarity b/c the density of water at this T is 1 kg / L
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boiling-point elevation
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solutes increase the bp of a solvent because a solution has a lower vapor pressure than a pure liquid; delta T(b) = K(b)*m where delta T(b) is the boiling-point elevation, K(b) is the proportionaliy constant characteristic of a particular solvent, and m is the molality of the solution.
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osmotic pressure
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the pressure of a solution against a concentration gradient; pi = MRT, where M is the molarity of the solution, R is the ideal gas constant, and T is temperature in K.
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vapor-pressure lowering
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(also Raoult's Law) delta P = P(o, a) - P(a), where P(o, a) is the vapor pressure of A above pure solvent A, and P(a) is the vapor pressure of A above a solution containing B; furthermore, delta P = X(b)P(o,a) where X(b) is the mole fraction of the solute B in solvent, and delta P = X(a)P(o,a). This holds only when the attraction between molecules of the different components is equal to the attraction between the molecules of any one component in its pure state.
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solvent
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the component of the solution whose phase remains the same after mixing; if the two substances are already in the same phase, the solvent is the component present in greater quantity.
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solvation and attractive forces
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solvation is possible when the attractive forces between solute and solvent are stronger than those between solute particles; the attractive forces are either ion-dipole interactions or van der Waals forces
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solubility
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the maximum amount of that substance that can be dissolved in a particular solvent at a particular T
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dilute vs. concentrated
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ratio of solute to solvent is small vs. large
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solubility of alkali metals
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all salts of alkali metals are water soluble
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solubility of ammonium
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all salts of ammonium (NH4+) are water soluble
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solubility of chlorides, bromides, and iodides
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all are water solubles, except with Ag+, Pb2+, and Hg(2)2+
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solubility of sulfates
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all salts of the sulfate ion (SO4, -2) are water soluble, except Ca+2, Sr+2, Ba+2, and Pb+2
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solubility of metal oxides
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all are insoluble, with the exception of the alkali metals and CaO, SrO, and BaO (all of which hydrolyze to form solutions of the corresponding metal hydroxides)
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solubility of hydroxides
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insoluble except alkali metals, Ca2+, Sr2+, and Ba2+
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solubility of carbonates, phosphates, sulfides, and sulfites?
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insoluble -- CO3 (-2), PO4 (-3), S (-2), and SO3 (-2) -- except with alkali metals and ammonium
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suffixes -ous and -ic?
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lesser and greater charge, respectively; e.g., ferrous = Fe2+ and ferric = Fe3+
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suffix -ide
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indicates a monatomic anion, e.g. hydride (H-), sulfide (S 2-), phosphide (P 3-)
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suffixes of oxyanions
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suffix -ite has less oxygen than -ate; e.g., nitrite (NO2 -) vs. nitrate (NO3 -) and sulfite (SO3 -2) vs. sulfate (SO4 -2)
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prefixes of oxyanions
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used when a series of oxyanions contains 4 levels; hypo- indicates less oxygen than per-
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the oxyanions with Cl
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hypochlorite (ClO-), chlorite (ClO2 -), chlorate (ClO3 -), perchlorate (ClO4 -)
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the prefix bi-
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indicates the addition of a single hydrogen ion to an anion; e.g., hydrogen carbonate = bicarbonate (HCO3-)
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electrolytes
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solutes whose solutions are conductive; a solute is considered a strong electrolyte if it dissociates completely into its constituent ions
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electrolytes and colligative properties
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since electrolytes ionize in solution, they will produce a larger effect on colligative properties than one would expect from the given concentration
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percent composition by mass
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(mass of solute / mass of solution) * 100
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mole fraction
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equal to the number of moles of the compound divided by the total number of moles of all species within the system
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molarity
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number of moles of solute per liter of solution; note that for dilute solutions, the volume of the solution is approximately equal to the volume of solvent used
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normality
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equal to the number of gram equivalent weights of solute per liter of solution. A gram equivalent weight is a measure of the reactive capacity of a molecule.
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ion product
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I.P. of a compound in solution = [A+]^m *[B-]^n where A(m)B(n) dissociates to mA+ +nB-
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solubility product constant
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again, for a saturated solution at equilibrium, Ksp = [A+]^m[B-]^n
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difference between the ion product and Ksp
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IP is defined with respect to inital concentrations and does not necessarily represent either an equilibrium or saturated solution, while Ksp does; at any point other than at equilibrium, the ion product is often referred to as Qsp.
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saturation occurs when...
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when a salt's IP = Ksp; the rate at which the salt dissolves equals the rate at which it precipitates out of solution. If IP > Ksp, then the solution is supersaturated.
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Ksp for a slightly soluble salt of general formula MX3
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Ksp = 27x^4 where x is the molar solubility
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Ksp for a slightly soluble salt of general formula MX2
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Ksp = 4x^3, where x is the molar solubility
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factors affecting solubility
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temperature of solution, the solvent, and (for a gas-phase solute), the pressure; also affected by the addion of other substances to the solution = common ion effect
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Ksp for a slightly soluble salt of general formula MX
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Ksp = x^2, where x is the molar solubility.
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litmus paper
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turns red in acidic solution; blue in basic
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Bronsted-Lowry definition of acid/base
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an acid donates protons, while a base accepts protons; they always occur in pairs (conjugate acid-base pairs), which are related by the transfer of a proton
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Arrhenius definition of acid/base
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acid produces a proton, and a base produces a hydroxide ion - only works in aqueous solution
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Lewis definition of acid/base
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acid is an electron-pair acceptor, and base is an electron-pair donor
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examples of Lewis acids that are not Bronsted-Lowry acids
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BCl3 and AlCl3
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naming acids formed from oxyanions
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oxyacids: if the anion ends in -ite, then the acid will end with -ous; if the anion ends in -ate, then acid ends in -ic acid
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permanganate
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MnO4- (an exception to the naming rules, since there are no 'manganate' or manganite')
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nitrous vs. nitric acid
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nitrous acid is formed from nitrite (HNO2), and nitric acid is formed from nitrate (HNO3)
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pH
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a measure of hydrogen ion concentration, pH = -log[H+] = log(1/[H+])
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water dissociation constant
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Kw = [H+][OH-] = 10^-14, so pH +pOH = 14
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log(xy) is
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log x + log y
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log (x^n)
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n*log x
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log (10^x)
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x
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what is - log (10^-x) equal to?
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x
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approximating log (n * 10^-m)
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log (n * 10^-m) = -m*log n (if it is a negative log, then = m - log n); note that since n is a number between 1 and 10, its logarithm wil be a fraction between 0 and 1, thus (m - log n) will be between m-1 and m. Further, the larger n is, the larger the fraction log n will be, and the answer will be closer to m-1.
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when can you neglect the dissociation of water?
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when the concentration of acid or base is > 10^-7 M
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strong acids
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perchloric acid (HClO4), nitric acid (HNO3), sulfuric acid (H2SO4), and hydrochloric acid
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strong bases
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sodium hydroxide, potassium hydroxide, other soluble hydroxides of Group IA and IIA metals
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acid dissociation constant
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Ka = [H3O+][A-] / [HA]; the measure of the degree to which an acid dissociates; the weaker the acid, the smaller the Ka.
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base dissociation constant
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Kb = [B+][OH-] / [BOH]
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conjugate acid/base pairs, and equilibrium constants
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Ka * Kb = Kw = 1E10^-14; Ka and Kb are inversely related.
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calculating [H+] of an aqueous solution of weak acid
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Ka = x^2 / ([HA]init - x)... x in the denominator can be ignored as long as x is less than 5% of the initial concentration of HA
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neutralization reaction
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HA + BOH = BA + H20; the salt (BA) may precipitate out or remain ionized in solution, depending on its solubility and the amount produced. Neutralization reactions generally go to completion.
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reverse of neutralization
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hydrolysis: salt + water returns the acid and base
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strong acid + strong base
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neutralize eachother, and the ions formed in the reaction do not react with water. If in equal concentrations, the pH will be 7.
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strong acid + weak base
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usually produces a salt, but no water (since weak bases are usually not hydroxides). In addition, the cation of the first reaction will react with the water, to reform a weak base b/c the cation of the first rxn is stronger than the conjugate base of the strong acid. In sum, the cation of the first reaction will lower the concentration of hydroxide, leading to an excess of H+ = lower pH
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weak acid + strong base
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results in a basic solution, due to the hydrolysis of the salt to reform the acid, with the concurrent formation of hydroxide ion from the hydrolyzed water molecules.
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weak acid + weak base
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depends on the relative strengths of the reactants; if Kb > Ka, then the solution is basic
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If the gram equivalent weight of sulfuric acid is 98 g/mol, then how many acid equivalents are released by 49 g?
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1, because 98/2 = 49.
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amphoteric
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or amphiprotic; a species that can act either as an acid or base, depending on its chemical environment. Examples include water, the partialy dissociated conjugate bases of a polyprotic acid (HSO4 - ), the hydroxides of certain metals (Al, Zn, Pb, and Cr), and species that can act as either oxidzing or reducing agents
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equivalence point (titration)
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the number of acid equivalents equals the number of base equivalents; a polyprotic acid or base can have several equivalence points
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indicators
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weak organic acids or bases that have different colors in their undissociated and dissociated states. The point in a titration at which the solution changes colored is called the end point (and is not the same as the equivalence point)
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buffer solution
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a mixture of a weak acid and its salt (conjugate base + a cation) or a mixture of a weak base and its salt (conjugate acid + an anion).
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Henderson-Hasselbach equation
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used to determine the pH of a solution in the buffer region where the concentrations of the species and its conjugate are present in approximately equal concentrations; pH = pKa + log([A-] / [HA]) or pOH = pKb + log([HB] / [B-])
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when does pH = pKa in a titration?
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when [A-] = [HA]; from the H-H equation, log 1 = 0
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OIL RIG
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oxidation is loss, reduction is gain
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oxidation # of a Group VIIA element in a compound
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is -1, except when combined with an element of higher electronegativity. For example, in HOCl, chlorine has an oxidation number of +1.
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oxidation # of hydrogen
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is +1, but is -1 in compounds with less electronegative elements than hydrogen (GroupsIA and IIA) - for example, NaH and CaH2
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oxidation # of oxygen
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is -2, except in cases such as OF2 (F is more electronegative, so O has +2), and in peroxides: BaO2, oxygen is -1 due to its structure ([O-O]2-)
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placement of cation and anion in formula writing?
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convention puts the cation first and anion second, so NaH implies H-
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ion-electron method
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or half-reaction method; used to balance redox reactions
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types of electrochemical cells
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galvanic (or voltaic) and electrolytic cells
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anode vs. cathode
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the electrode at which oxidation vs. reduction occurs
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AN OX and RED CAT
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anode = oxidation
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galvanic cells
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spontaneous reaction, so deltaG is negative; galvanic cells provide energy and are used to do work.
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Daniel cell
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zinc bar is placed in aqueous ZnSO4 solution, and a copper bar is placed in an aqueous CuSO4 solution. Zn is oxidized to Zn2+ at the anode, while copper is reduced from Cu2+ to Cu at the cathode.
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salt bridge
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permits the exchange of cations and anions between two half-cells; prevents a charge gradient from building up (anions around the cathode, cations around the anode). The salt bridge contains an inert electrolyte, usually KCl or NH4NO3, whose ions will not react with the electrodes or the ions in solution.
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flow of electrons in a Daniel cell
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from the zinc bar (anode) through the wire and voltmeter, to the copper bar (cathode).
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cell diagram
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shorthand notation representing the reactions in an electrochemical cell; e.g., anode | anode solution || cathode solution | cathode, a single vertical line = phase boundary, and the double line indicates the presence of a salt bridge
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electrolytic cells
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nonspontaneous reaction, so deltaG is positive; electrical energy is required to induce the reaction, and the half-reactions are often placed in the same container.
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charge carried by one mole of electrons?
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(1.6 E-19 C per electron)(6.022 E23 electrons per mole) = 96,487 C/mol of electron = Faraday's constant
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one Faraday
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F = the amount of charge contained in one mole of electrons = 96,487 coulombs or J/V
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flow of electrons in an electrochemical cell?
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always from the anode to cathode (even though the anode is positive in an electrolytic cell)
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why is the cathode considered negative in an electrolytic cell?
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because the cathode provides the electrons necessary to reduce the species there; in addition, the cathode is connected to the negative pole of the batter used for the electrolysis.
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in electrophoresis, which way do negatively charged amino acids travel?
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toward the anode (positively charged amino acids travel toward the cathode)
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reduction potential
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the tendency of a species to acquire electrons and be reduced; the more positive a potential, the greater the specie's tendency to be reduced.
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standard hydrgoen electrode
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SHE = arbitrarily given a reduction potential of 0.00; so all reduction potentials are defined relative to SHE
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standard electromotive force of a reaction
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EMF = Eo(red) + Eo(ox); this should be positive for a galvanic cell, and negative for an electrolytic cell
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finding the free energy of an electrochemical reaction
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deltaG = -nFE(cell), where n is the number of moles of electrons exchanged, F is Faraday's constant, and E(cell) is the EMF of the cell. If the rxn takes place under standard conditions (25*C, 1atm, and all solutions at 1M), then the deltaG is the standard Gibbs free energy and E(cell) is the standard cell potential.
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Nernst equation
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E(cell) = Eo(cell) - (RT/nF)(lnQ), used to determine the EMF of a cell at nonstandard conditions (concentrations)
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potentiometer
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a kind of voltmeter that draws no current, and gives a more accurate reading of the difference in potential between two electrodes
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relation of EMF to Keq
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nFEo(cell) = RT ln(Keq), because deltaGo = -nFEo(cell) = -RT ln(Keq)
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if Eo(cell) is positive, then Keq?
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must be greater than 1 (products are favored), because lnK is positive
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