CHEMIESTUDYSHEET
Chemiestudysheet.><
[fall]
the basics
~ nuclear model
~ avogadro’s number: 6.02 x 1023 atoms = 1 mol
~ % composition by mass: what % of molar mass belongs to each element
~ significant figures
▪ multiply/divide: keep lowest # of sig figs
▪ add/subtract: keep lowest # of decimal places
▪ decimal places in log function = sig figs in answer
~ molecular vs. empirical
▪ molecular: actual amount of atoms
→ molar mass / empirical mass = ratio of molecular to empirical
▪ empirical: simplest ratio of atoms in formula or moles of each element
~ concentration
▪ % by mass (g / g)
▪ molarity (mol / L)
~ dilution
▪ (Lbefore)(Mbefore) = (Lafter)(Mafter)
~ limiting reagent
▪ convert to mols; whichever produces less product = limiting reagent
▪ % yield = actual / theoretical * 100%
atomic theory
~ dalton’s atomic theory
▪ elements consist of atoms
▪ all atoms of an element are alike in mass and other properties; atoms of one element differ from those of another
▪ different elements combine in a simple numerical ratio
~ thomson & cathode ray experiment
▪ “plum pudding model”
~ rutherford’s gold foil experiment
▪ alpha particles (He2+)
~ average atomic mass = (fraction isotope 1)(mass 1) + (frac 2)(mass 2) + …
light
~ 1 nm = 10-9 m
~ 1 Å = 10-10 m
~ C = 3.00 x 10-8 ms-1 = (λ)(υ)
~ frequency (υ) = number of peaks per second
▪ 1 Hertz (Hz) = 1 s-1
~ wavelength (λ)
▪ visible light: 400 nm – 700 nm
~ Ephoton = h υ(s-1) = hc / λ(m)
▪ h (Planck’s constant) = 6.63 x 10-34 J*s
~ bohr’s H atom
▪ for H & 1e- ions:
→ En = -z2RH / n2
o z: atomic number
o RH: 2.18 x 10-18 J
o n: 1, 2, 3, 4…
→ ΔE = z2RH(1/ni2 – 1/nf2)
electron configuration & orbital diagrams
~ e- configuration
~ anions: add to valence e-
~ cations: remove valence first (p before s); then d from inner level
~ orbital diagrams
▪ pair ↑↓ (Pauli exclusion principle)
▪ spread out e- if possible (e- repel)
▪ unpaired e- have parallel spins
→ paramagnetic: unpaired e- vs. diamagnetic: all e- paired
▪ exceptions: d5 (Cr, Mo) / d10 (Cu, Ag, Au)
quantum numbers
~ principle (n) = E levell
~ orbital (l): s = 0, p = 1, d = 2, f = 3
▪ l = 0, 1, 2, … (n-1)
~ magnetic (ml)
▪ ml = -l … 0 … +l
~ spin (ms)
▪ +1/2 or -1/2
periodic trends and molecular structures
~ atomic radius
▪ energy level
▪ greater zeff = smaller radius
▪ cations: less e- repulsion; smaller
▪ anions: more e- repulsion; larger
▪ isoelectronic: same # e-
~ ionization E (I) = E added to remove e-
▪ radius ↑ I1↓
▪ change in energy level = huge jump in E needed
▪ exceptions:
→ I1 (Al) < I1 (Mg); I1(B) < I1 (Be)
~ electron affinity (EA) = ΔE when e- added
▪ radius ↓, EA more (-) = easier to add
▪ s2, p6: EA >>> 0
▪ exceptions:
→ N, P, As, Sb: add e- repulsion to p3 = EA less (-)
→ F vs Cl: orbital of F too small = extra e- repulsion = EA less (-)
~ electronegativity (EN) = ability of atom to attract e- in chemical bond
~ bonding:
▪ ionic (metal & nonmetal; large ΔEN)
→ “formula unit”
▪ molecular (nonmetal & nonmetal; small ΔEN)
→ nonpolar covalent bond: ΔEN = 0
→ polar covalent bond: ΔEN > 0
→ bond length:
o radius ↑ bond length↑
o # bonds ↑; bond length ↓; bond energy ↑
→ bond order = total # bonds / # atoms bonded to center
o bigger = shorter
→ “free radical”: paramagnetic – one electron on center atom; bonds with similar electron on same kind of other atom through dimerization
→ “expanded octets”: only if center atom not in row 2
~ formal charge = (#valence e-) – (#nonbonding e-) – ½(#bonding e-)
▪ sum of formal charges = overall charge
▪ formal charges closest to 0 = most important
~ VSEPR theory: refer to attachment
▪ bond angles
→ dipole moment (μ): >0, polar; =0, nonpolar
→ no lone pairs on center = nonpolar (unless different outer elements)
→ lone pairs on center = polar (except AX2E3/linear; AX4E2/square planar)
pressure
~ barometric P
~ ideal gas law
▪ pV = nRT
▪ d = mmP / RT
~ effusion
▪ μA/ μB (speed) = √mmB/mmA = nA/nB = dA/dB = tB/tA
~ manometer
▪ d1h1 = d2h2
~ dalton’s law
▪ PA = Ptotal(nA / ntotal)
~ collecting gas “over water”
▪ Pbar = Pgas + Pwatervapor
thermodynamics
~ kinetic molecular theory (KMT): ideal gases
▪ V = L of empty space; assume no molecular V
▪ P = F/A; assume no IMFs
~ gas laws
▪ boyle’s law: P↑V↓
▪ charles’ law: T↑V↑
▪ avogadro’s law: n↑V↑
~ 1st law of thermodynamics: conservation of E
~ calorimetry
▪ q(J) = m(g)s(J/gºC)Δt
~ constant P “coffee cup” calorimetry
▪ ionic (s) → ionic (aq) or X(aq) + Y(aq) → products
▪ ΔHrxn = +/- (total gsoln)(Ssoln)(Δt) / #mols dissolved or # mols product
→ + endo; - exo
~ qcalorimeter = (htcap(kJ/ºC or J/ºC))(Δt)
~ constant V “bomb” calorimetry
▪ not in water: heat of combustion = -(htcap)(Δt) / #g or mol burned
▪ in water: heat of combustion = -[(gH2O)(SH2O)(Δt)+(htcap)(Δt)] / #g or mol burned
~ thermochemical equation
▪ based on molar ratio
▪ multiply rxn by x = multiply ΔH by x
▪ flip rxn = -ΔH
▪ hess’s law: add rxns = add ΔH
~ ΔHºrxn = ΣnΔHºf(products) - ΣnΔHºf(reactants)
▪ ΔHºf = 0 for standard state elements
~ breaking bonds: (+); forming bonds: (-)
▪ coefficient(type of bond) LEFT – coefficient(type of bond) RIGHT
~ work(w) = +/- P|ΔV|
▪ +: compressed; -: expands
~ Δu (internal E) = q + w
~ born-haber cycle
▪ standard state elements → 1 mol compound; ΔHºf
▪ X(g) → X+ (g) + 1e-; I1
▪ X(g) → X- (g); EA1
▪ X2 (g) → 2X (g); ΔHdissociation (bond energy)
▪ ions (g) → 1 mol ionic (s); lattice E (LE) [always K = goes left
▪ Q < K = goes right
entropy, spontaneity, & more thermodynamics
~ entropy (S) = disorder
▪ substance
→ T↑ = ΔS > 0
→ phase change
▪ rxn
→ ΔS > 0 if # mols (g) ↑
→ ΔSºrxn = ΣnΔSº(products) - ΣnΔSº(reactants)
→ Sº of standard elements IS NOT 0!
→ more bonds, S ↑
~ 2nd law of thermodynamics
▪ ΔSuniverse = ΔSrxn + ΔSsurr; if >0 = spontaneous
▪ ΔSsurr = -ΔHrxn / T
▪ ΔG(kJ/mol) = ΔHrxn (kJ/mol) – T(K)ΔSrxn(kJ/mol*K)
→ > 0, nonspontaneous; =0, equilibrium; 0 = spontaneous
▪ don’t multiply voltages!
▪ ox. || red
▪ Al (s) | Al3+ (aq) || Cl- (aq) | Cl2 (g) | Pt(s)
→ Pt(s): unreactive, so used as a cathode
~ ΔGº = -nFEºcell
▪ F = 96500 C/mol
~ Eºcell (J/C) = (RET / nF)lnK
▪ RE = 8.31 J
▪ if at 25ºC (298K);
→ Eºcell = (0.0257 / n)lnK
▪ if K is very high, reacts well
~ metal + strong acid (HCl, HBr, HI)
▪ red: 2H+ + 2e- → H2; 0V
~ metal + HNO3
▪ red: NO3- + 4H+ + 3e- → NO + 2H2O; 0.96V
~ nernst equation
▪ Ecell = Eºcell – (0.0257 / n)lnQ
~ electrolysis
▪ choose best ox & red (aq)
▪ spectator ions: Na+, K+, SO42-, NO3-
▪ ox: 2H2O → O2 + 4H+ + 4e-; -1.229V
→ exception: 2Cl- → Cl2 + 2e-; -1.358V
▪ red: 2H2O + 2e- → H2 + 2OH-; -0.828V
▪ time(s) x current(A) = charge(C)
valence bond theory
~ hybridization
▪ atoms bonded to center / hybridization
→ 2 / sp
→ 3 / sp2
→ 4 / sp3
→ 5 / sp3d
→ 6 / sp3d2
~ single bond: σ bond
~ double, triple bond: π bond
▪ planar
~ resonance structures: delocalized molecular orbitals
~ molecular orbital theory
▪ σ1s σ1s* σ2s σ2s* σ2p π2p π2p* σ2p*
▪ switch σ2p π2p if no O, F, or N present
▪ bond order = #b – a(*) / 2 = bonding – antibonding / 2
crystal types
~ network covalent
▪ C(diamond), Si, SiC, SiO2 (quartz)
→ very high melting point (need to break many covalent bonds)
→ very hard (rigid structure of bonds)
→ poor conductor of electricity (localized e-)
▪ C(graphite)
→ trigonal planar structure
→ delocalized e- between layers of graphite = good conductor
→ weak attraction = flakes off easily
~ ionic (contains metal ion or ammonium ion)
▪ (s)
→ high melting point
→ brittle
→ poor conductor of electricity (ions are held in place)
▪ (l), (aq): good conductor; ions move freely
▪ coulomb’s law
→ charge difference ↑, size of ions ↓ = LE more (-), melting point ↑, solubility ↓
~ metallic (metal cation in a sea of e-)
▪ delocalized e- = good conductors
▪ size of metal cation ↓ = melting point ↑
~ molecular
▪ low melting/boiling points
▪ usually soft or brittle
▪ poor conductors
▪ van der waal’s forces
▪ hydrogen bonding (strongest)
→ occurs between two molecules containing hydrogen bonded directly to N, O, or F
→ density of ice less than density of water because H-bonds form in a way that leaves empty space
▪ london (dispersion) forces
→ instantaneous vs induced dipole
→ molar mass ↑, stronger london, mp/bp ↑
▪ dipole-dipole (weakest): for polar molecules only
→ hydrocarbons: nonpolar
→ has central oxygen: polar
▪ IMFs ↑, vapor P ↓
random topics
~ electrolytes
▪ nonelectrolytes: cannot conduct electricity at all
→ molecular compounds, pure acids (not ionic)
→ no ions to conduct electricity
▪ strong electrolytes: conduct very well
→ soluble ionics, strong acids
→ lots of ions to conduct
▪ weak electrolytes: weakly conduct
→ weak acids
~ real gases: low T, high P
▪ Vreal > Videal (empty space)
▪ Preal < Pideal (assume no IMFs)
▪ van der waal’s equation
→ (P + n2a/V2)(V – nb) = nRT
→ IMFs ↑ = a ↑
→ molecular V ↑ = b ↑
→ the higher a & b, the less ideal the gas is
~ “like dissolves like”
▪ polar / polar
▪ nonpolar / nonpolar
→ (polar / nonpolar = immiscible)
▪ H-bond / H-bond
▪ ionic / polar
→ * remember solubility rules first!
→ ion dipole attraction
→ ionic / nonpolar doesn’t work because system won’t receive lost energy, so nonspontaneous
→ benzene (C6H6): nonpolar; resonance: every other is double-bond
→ cyclohexane (C6H12): nonpolar
→ T ↑, solubility of most salts ↑
→ T ↑, solubility of most gases ↓
→ henry’s law: P↓solubility of gases↓
heating curves & phase diagrams
~ heating curve (refer to attachment)
~ phase changes: q = nΔHfusion or nΔHvaporization
▪ ΔHfusion Ea
~ boltzmann distribution & temperature changes (refer to attachment)
~ root-mean-square speed = μrms = √3RET(K) / mm(kg/mol)
▪ RE = 8.31 kg*m2 / s2*mol*K
~ rxn mechanism
▪ overall rate = (k of slow step)[reactants]x
▪ forward rate = reverse rate
~ effects of catalysts
▪ graphically (refer to attachment)
▪ graph (1/T K-1, lnk): (1/T)
→ slope = -Ea / RE
▪ arrhenius equation
→ ln(k2 / k1) = (Ea / RE)(1/T1 – 1/T2)
→ ln(t1 / t2) = (Ea / RE)(1/T1 – 1/T2)
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