CHEM 110 Final Review: Atomic Structure, Nuclear Chemistry, Bonding and Metal Complexes

Terms in this set (32)

1. Energy of a photonE = h nu = hc / lambda, where h = 6.626 x 10^-34 J·s is Planck's constant, nu is frequency, c = 3.00 x 10^8 m/s and lambda is wavelength.
2. Mini-problem: energy of one photon of 500 nm lightE = hc/lambda = (6.626 x 10^-34)(3.00 x 10^8) / (500 x 10^-9 m), about 3.98 x 10^-19 J.
3. Photoelectric effect equation and its significanceKE of ejected electron = h nu - work function. No electrons are ejected below a threshold frequency regardless of intensity, showing light comes in quanta (photons).
4. de Broglie wavelengthlambda = h / (m v). All moving matter has wave properties; the effect is noticeable only for very small masses such as electrons.
5. Bohr model energy levels for hydrogenE(n) = -2.18 x 10^-18 J (1/n^2). When an electron drops from a higher to a lower n, a photon is emitted with energy equal to the difference, giving line spectra.
6. Heisenberg uncertainty principleIt is impossible to know both the exact position and exact momentum of a particle at once: (delta x)(delta p) >= h/(4 pi). This is why electrons are described by orbitals (probability distributions).
7. The four quantum numbersn: principal (energy level, size). l: angular momentum (shape, 0 to n-1; s, p, d, f). ml: magnetic (orientation, -l to +l). ms: spin, +1/2 or -1/2.
8. How many nodes does an orbital have?Total nodes = n - 1. Angular nodes = l. Radial nodes = n - l - 1. For example, a 3p orbital has 1 angular and 1 radial node.
9. Aufbau principle, Pauli exclusion principle and Hund's ruleAufbau: fill lowest-energy orbitals first. Pauli: no two electrons in an atom share all four quantum numbers, so at most two per orbital with opposite spins. Hund: fill degenerate orbitals singly with parallel spins before pairing.
10. Electron configurations of Cr and CuCr is [Ar] 3d5 4s1 and Cu is [Ar] 3d10 4s1, exceptions caused by the extra stability of half-filled and filled d subshells.
11. Trend in atomic radius and whyRadius decreases across a period because effective nuclear charge increases while electrons enter the same shell. It increases down a group as electrons occupy higher shells.
12. First ionization energy trend and two common exceptionsGenerally increases across a period and decreases down a group. Exceptions: B is lower than Be (removing a higher-energy 2p electron) and O is lower than N (electron pair repulsion in a 2p orbital).
13. Electronegativity trendIncreases across a period and decreases down a group. Fluorine is the most electronegative element.
14. Alpha, beta and gamma decayAlpha: emits a helium-4 nucleus, mass number drops by 4 and atomic number by 2. Beta (minus): a neutron becomes a proton and emits an electron, atomic number up by 1. Gamma: emits a high-energy photon with no change in mass or atomic number.
15. Positron emission vs electron captureBoth convert a proton into a neutron, lowering atomic number by 1 with no change in mass number. Positron emission releases a positron; electron capture absorbs an inner electron.
16. Mini-problem: what forms when U-238 undergoes alpha decay?Th-234: mass number 238 - 4 = 234 and atomic number 92 - 2 = 90 (thorium), plus a helium-4 nucleus.
17. Mini-problem: 80 g of an isotope decays for three half-lives. How much remains?80 -> 40 -> 20 -> 10 g, so 10 g remain.
18. Mass defect and nuclear binding energyA nucleus has less mass than its separate protons and neutrons. The missing mass is the binding energy, E = delta m c^2. Iron-56 is near the peak of binding energy per nucleon, so it is among the most stable nuclei.
19. Nuclear fission vs fusionFission splits a heavy nucleus into smaller nuclei, releasing energy (nuclear reactors). Fusion combines light nuclei into a heavier one, releasing energy (stars). Both move toward higher binding energy per nucleon.
20. Predicting decay from the neutron-to-proton ratioNuclei with too many neutrons (above the band of stability) tend to undergo beta minus decay. Those with too few tend toward positron emission or electron capture. Very heavy nuclei often alpha decay.
21. What is resonance in Lewis structures?When two or more valid Lewis structures differ only in electron placement, the real molecule is an average (resonance hybrid). For example, both O-O bonds in ozone are equal in length.
22. VSEPR shapes for five and six electron groups with lone pairsFive groups: SF4 (1 lone pair) is seesaw, ClF3 (2) is T-shaped, XeF2 (3) is linear. Six groups: BrF5 (1 lone pair) is square pyramidal, XeF4 (2) is square planar.
23. Hybridization from electron groups2 electron groups: sp (linear). 3: sp2 (trigonal planar). 4: sp3 (tetrahedral). 5: sp3d. 6: sp3d2.
24. Sigma vs pi bondsA sigma bond forms from head-on orbital overlap and allows rotation. A pi bond forms from side-by-side p orbital overlap and restricts rotation. A double bond is 1 sigma + 1 pi; a triple bond is 1 sigma + 2 pi.
25. Bond order formula in molecular orbital theoryBond order = (bonding electrons - antibonding electrons) / 2. Higher bond order means a stronger, shorter bond; bond order 0 means the molecule is not expected to exist (He2).
26. Why does MO theory explain that O2 is paramagnetic?O2 has two unpaired electrons in degenerate pi* antibonding orbitals, giving bond order (10 - 6)/2 = 2 and paramagnetism. A simple Lewis structure shows all electrons paired.
27. Coordination complex vocabulary: ligand, coordination number, chelateA ligand is a Lewis base donating an electron pair to the metal. Coordination number is the number of donor atoms bonded to the metal. A chelating ligand binds through two or more donor atoms (ethylenediamine is bidentate).
28. Mini-problem: oxidation state of Co in [Co(NH3)5Cl]Cl2Two Cl- counter-ions mean the complex ion is 2+. NH3 is neutral and the bound Cl is -1, so Co + (-1) = +2 and Co is +3.
29. Crystal field splitting in an octahedral complexThe five d orbitals split into a lower-energy t2g set (dxy, dxz, dyz) and a higher-energy eg set (dz2, dx2-y2), which point at the ligands. The gap is delta o.
30. High-spin vs low-spin complexesWeak-field ligands give a small delta o, so electrons occupy eg before pairing (high spin, more unpaired electrons). Strong-field ligands give a large delta o, so electrons pair in t2g first (low spin).
31. Spectrochemical series (selected ligands, weak to strong field)I- < Br- < Cl- < F- < OH- < H2O < NH3 < en < CN- < CO.
32. Why are many transition metal complexes coloured?Electrons absorb visible light to jump from t2g to eg orbitals. The colour we see is complementary to the wavelength absorbed; larger delta o means absorption of shorter wavelengths.