CHEM 14A Exam Review: Quantum Atom, Periodic Trends, Chemical Bonding and Acids and Bases

Terms in this set (32)

1. Energy of a photon: formula and constantsE = hν = hc/λ, where h = 6.626×10^−34 J·s (Planck's constant), ν is frequency, c = 3.00×10^8 m/s, and λ is wavelength. Shorter wavelength means higher energy.
2. What did the photoelectric effect show, and what is its equation?Electrons are ejected only if the light's frequency exceeds a threshold, regardless of intensity, showing light comes in photons. KE of ejected electron = hν − Φ, where Φ is the metal's work function.
3. In the photoelectric effect, what happens if you increase light intensity at a fixed frequency above threshold?More electrons are ejected per second, but each electron's maximum kinetic energy stays the same, because KE depends only on photon frequency.
4. de Broglie wavelength formulaλ = h/(mv), where m is mass and v is speed. Matter has wave properties, noticeable only for very small masses like electrons.
5. State the Heisenberg uncertainty principle.Δx·Δp ≥ h/(4π). The more precisely a particle's position is known, the less precisely its momentum can be known, so electrons are described by orbitals (probability distributions), not exact paths.
6. Energy levels of the hydrogen atomE_n = −hR/n^2, with R = 3.29×10^15 Hz (equivalently −2.18×10^−18 J/n^2). A photon is emitted when the electron drops to lower n and absorbed when it rises to higher n.
7. List the four quantum numbers and their allowed values.n = 1, 2, 3, ... (shell, size). l = 0 to n−1 (subshell shape: s, p, d, f). ml = −l to +l (orientation). ms = +1/2 or −1/2 (spin).
8. How many orbitals are in a subshell and in a shell?A subshell with quantum number l has 2l + 1 orbitals (s 1, p 3, d 5, f 7). Shell n has n^2 orbitals and can hold 2n^2 electrons.
9. State the Pauli exclusion principle and Hund's rule.Pauli: no two electrons in an atom share all four quantum numbers, so an orbital holds at most 2 electrons of opposite spin. Hund: degenerate orbitals fill singly with parallel spins before any pairing.
10. Ground-state electron configurations of Cr and CuCr is [Ar]3d5 4s1 and Cu is [Ar]3d10 4s1. A half-filled or filled d subshell is extra stable, so one 4s electron moves into 3d.
11. What is effective nuclear charge and how does it explain atomic radius trends?Zeff is the nuclear charge felt by an electron after shielding by inner electrons. Zeff rises across a period, pulling electrons closer, so radius decreases left to right; radius increases down a group as electrons occupy higher n shells.
12. General ionization energy trend and its two common exceptionsIE increases up a group and left to right across a period. Exceptions: B is lower than Be (removing a higher-energy 2p electron) and O is lower than N (removing a paired 2p electron with extra repulsion).
13. Electronegativity trendElectronegativity increases toward the upper right of the periodic table (excluding noble gases). Fluorine is the most electronegative element, about 4.0 on the Pauling scale.
14. Formal charge formula and how to use itFC = V − (L + B/2), where V = valence electrons of the free atom, L = lone-pair electrons, B = bonding electrons. Prefer structures with formal charges closest to zero and any negative charge on the more electronegative atom.
15. What is resonance? Use carbonate as an example.When several valid Lewis structures differ only in electron placement, the real molecule is a hybrid with delocalized electrons. In CO3^2−, all three C–O bonds are identical with bond order 4/3.
16. Which atoms can have an expanded octet, and why?Atoms in period 3 and below (such as P in PCl5, S in SF6, Xe in XeF4). They are larger and have more room around them for more than 8 valence electrons; period 2 atoms like C, N, O, F cannot.
17. VSEPR electron geometries for 2, 3, 4, 5 and 6 electron domains2 linear (180°), 3 trigonal planar (120°), 4 tetrahedral (109.5°), 5 trigonal bipyramidal (90° and 120°), 6 octahedral (90°).
18. Molecular shapes of NH3, H2O, SF4, ClF3 and XeF4NH3 trigonal pyramidal (~107°), H2O bent (~104.5°), SF4 seesaw, ClF3 T-shaped, XeF4 square planar. Lone pairs repel more than bonding pairs and compress bond angles.
19. When is a molecule polar? Compare CO2 and H2O.It needs polar bonds arranged so their dipoles do not cancel. Linear CO2 has polar bonds that cancel, so it is nonpolar. Bent H2O has dipoles that add, so it is polar.
20. Match steric number to hybridization.Steric number 2: sp. 3: sp2. 4: sp3. 5: sp3d. 6: sp3d2. Steric number = number of atoms bonded to the central atom plus its lone pairs.
21. Sigma vs pi bonds in single, double and triple bondsA single bond is 1 σ. A double bond is 1 σ + 1 π. A triple bond is 1 σ + 2 π. Pi bonds come from side-on p-orbital overlap and prevent rotation around the bond.
22. Bond order from molecular orbital theory, and why O2 is paramagneticBond order = (bonding electrons − antibonding electrons)/2. O2 has bond order 2 and two unpaired electrons in degenerate π* orbitals, so it is attracted to a magnetic field, which Lewis structures fail to predict.
23. Why does He2 not exist as a stable molecule?Its 4 electrons fill both σ1s (bonding) and σ*1s (antibonding), giving bond order (2 − 2)/2 = 0, so there is no net bond.
24. What is a coordination compound? Define ligand, coordination number and chelate.A central metal ion bonded to ligands, which are Lewis bases donating lone pairs. Coordination number is the number of donor atoms bonded to the metal. A chelating ligand, like ethylenediamine, binds through two or more donor atoms.
25. Brønsted-Lowry vs Lewis acids and basesBrønsted-Lowry: acid donates H+, base accepts H+. Lewis: acid accepts an electron pair, base donates an electron pair. Every Brønsted base is a Lewis base, but Lewis acids include species without H, like BF3 or metal ions.
26. How are Ka and Kb of a conjugate pair related?Ka × Kb = Kw = 1.0×10^−14 at 25 °C (pKa + pKb = 14). The stronger the acid, the weaker its conjugate base.
27. List the common strong acids.HCl, HBr, HI, HNO3, H2SO4 (first proton), HClO4 and HClO3. They ionize essentially completely in water, so [H3O+] equals the acid concentration.
28. What makes one oxoacid stronger than another? Rank HClO, HClO2, HClO3, HClO4.More oxygens (more electronegative atoms) pull electron density away and stabilize the conjugate base by delocalizing charge. Strength: HClO4 > HClO3 > HClO2 > HClO.
29. Why is HI a stronger acid than HF?For binary acids in the same group, bond strength dominates. The H–I bond is much longer and weaker than H–F, so HI ionizes more easily, even though F is more electronegative.
30. Mini-problem: pH of 0.10 M acetic acid (Ka = 1.8×10^−5)[H3O+] ≈ sqrt(Ka × C) = sqrt(1.8×10^−6) ≈ 1.3×10^−3 M, so pH ≈ 2.87. The approximation is valid because x is under 5% of 0.10 M.
31. Henderson-Hasselbalch equation and the effective range of a bufferpH = pKa + log([A−]/[HA]). A buffer works best within about pKa ± 1, and has pH = pKa when [A−] = [HA].
32. Weak acid titrated with strong base: pH at the half-equivalence and equivalence pointsAt half-equivalence, [HA] = [A−], so pH = pKa. At equivalence, only the conjugate base remains, so the pH is above 7.