CHEM 121 Midterm Review: Atomic Structure, Bonding, Stoichiometry, Gases and Thermochemistry

Terms in this set (32)

1. How does atomic radius change across a period and down a group?It decreases left to right across a period (more protons, similar shielding, higher effective nuclear charge) and increases down a group (electrons occupy higher shells).
2. General trend in first ionization energy, and two common exceptionsIt increases across a period and decreases down a group. Exceptions: B is lower than Be (removing a 2p electron) and O is lower than N (pairing repulsion in a 2p orbital).
3. Electronegativity trendIt increases toward the top right of the periodic table (excluding noble gases). Fluorine is the most electronegative element.
4. What are the four quantum numbers and their allowed values?n = 1, 2, 3...; l = 0 to n − 1 (s, p, d, f); ml = −l to +l; ms = +1/2 or −1/2.
5. Aufbau principle, Pauli exclusion principle and Hund's ruleAufbau: fill lower-energy orbitals first. Pauli: no two electrons in an atom share all four quantum numbers, so at most 2 per orbital with opposite spins. Hund: fill degenerate orbitals singly before pairing.
6. Write the electron configurations of Fe and Fe3+.Fe: [Ar] 4s2 3d6. Fe3+: [Ar] 3d5. When transition metals form cations, the 4s electrons are removed before the 3d electrons.
7. Why are Cr and Cu exceptions to the expected configurations?Cr is [Ar] 4s1 3d5 and Cu is [Ar] 4s1 3d10, because half-filled and completely filled d subshells are especially stable.
8. Relate the energy of a photon to its frequency and wavelength.E = hν = hc/λ, where h = 6.626 × 10^−34 J·s and c = 3.00 × 10^8 m/s. Shorter wavelength means higher energy.
9. Rank the isoelectronic species O2−, F−, Na+, Mg2+ by size.O2− > F− > Na+ > Mg2+. They all have 10 electrons, so the ion with more protons pulls them in more tightly and is smaller.
10. Formula for formal chargeFormal charge = (valence electrons) − (nonbonding electrons) − (bonding electrons / 2). The best Lewis structure keeps formal charges small, with negative charge on the more electronegative atoms.
11. VSEPR shapes of H2O and NH3Both have 4 electron domains (tetrahedral arrangement). H2O has 2 lone pairs, so it is bent (about 104.5°); NH3 has 1 lone pair, so it is trigonal pyramidal (about 107°).
12. Why is CO2 nonpolar while H2O is polar?Both have polar bonds, but CO2 is linear so its bond dipoles cancel. H2O is bent, so its bond dipoles add up to a net dipole.
13. Match hybridization to electron geometry: sp, sp2, sp3.sp: 2 domains, linear, 180°. sp2: 3 domains, trigonal planar, 120°. sp3: 4 domains, tetrahedral, 109.5°.
14. How many sigma and pi bonds are in a double bond and a triple bond?A double bond has 1 sigma and 1 pi bond; a triple bond has 1 sigma and 2 pi bonds. Every single bond is a sigma bond.
15. Rank the main intermolecular forces and their effect on boiling point.London dispersion forces (all molecules, stronger with more electrons), then dipole-dipole, then hydrogen bonding (H bonded to N, O or F). Stronger forces give higher boiling points.
16. What is a mole, and how do you convert grams to moles?One mole is 6.022 × 10^23 particles (Avogadro's number). Moles = mass in grams / molar mass in g/mol.
17. 2 H2 + O2 → 2 H2O. With 4 mol H2 and 3 mol O2, what is the limiting reactant and how much water forms?4 mol H2 needs only 2 mol O2, so H2 is limiting. It makes 4 mol H2O, and 1 mol O2 is left over.
18. How is percent yield calculated?Percent yield = (actual yield / theoretical yield) × 100%. The theoretical yield comes from the limiting reactant.
19. What is the empirical formula of glucose, C6H12O6?CH2O, the simplest whole-number ratio of atoms. The molecular formula is a whole-number multiple of it (here 6).
20. Define molarity and give the dilution equation.Molarity M = moles of solute / litres of solution. For dilution, M1V1 = M2V2, because the moles of solute stay the same.
21. Ideal gas law and the two common values of RPV = nRT, with T in kelvin. R = 0.08206 L·atm/(mol·K) or 8.314 J/(mol·K).
22. What volume does 1 mol of ideal gas occupy at 0 °C and 1 atm?About 22.4 L. This only applies at those conditions; at other conditions use PV = nRT.
23. State Dalton's law of partial pressures.Total pressure is the sum of the partial pressures: P_total = P1 + P2 + ... Each gas's partial pressure equals its mole fraction times P_total.
24. State Graham's law of effusion.Rate of effusion is inversely proportional to the square root of molar mass: rate1/rate2 = √(M2/M1). H2 effuses 4 times faster than O2 because √(32/2) = 4.
25. When do real gases deviate most from ideal behaviour?At high pressure and low temperature, when molecular volume and intermolecular attractions become significant.
26. Sign of ΔH for exothermic and endothermic reactionsExothermic reactions release heat to the surroundings, so ΔH < 0. Endothermic reactions absorb heat, so ΔH > 0.
27. Mini-problem: how much heat raises 100 g of water from 20 °C to 30 °C? (c = 4.184 J/g·°C)q = m·c·ΔT = 100 × 4.184 × 10 = 4184 J, or about 4.18 kJ.
28. What does Hess's law state?The enthalpy change of a reaction is the same whether it happens in one step or many, so ΔH values of steps can be added. Reversing a reaction flips the sign of ΔH; multiplying it scales ΔH.
29. How do you calculate ΔH°rxn from standard enthalpies of formation?ΔH°rxn = Σ n·ΔH°f(products) − Σ n·ΔH°f(reactants). ΔH°f of an element in its standard state (e.g., O2 gas) is zero.
30. Estimating ΔH from bond enthalpiesΔH ≈ Σ(bonds broken) − Σ(bonds formed). Breaking bonds costs energy; forming bonds releases it.
31. What makes lattice energy larger?Higher ion charges and smaller ions. MgO (2+ and 2−) has a much larger lattice energy than NaCl (1+ and 1−).
32. Compare ionic, molecular, covalent network and metallic solids.Ionic: high melting, brittle, conduct when molten. Molecular: low melting. Covalent network (diamond): very high melting and hard. Metallic: conduct as solids, malleable.