CHEM 1A03 Exam Review: Bonding, Molecular Structure, Intermolecular Forces and Phase Changes

Terms in this set (28)

1. Mini-problem: how many moles are in 9.0 g of water (M = 18.0 g/mol)?n = m/M = 9.0 g / 18.0 g/mol = 0.50 mol.
2. How do you find the limiting reagent?Convert each reactant to moles, divide by its stoichiometric coefficient, and the smallest result identifies the limiting reagent, which sets the maximum product.
3. Define molarity and give a dilution formula.Molarity = moles of solute per litre of solution (mol/L). For dilution, C1V1 = C2V2.
4. What is electronegativity and what is its periodic trend?The ability of an atom in a bond to attract shared electrons. It increases across a period and decreases down a group; fluorine is the most electronegative.
5. Rank bond types by electronegativity difference.Roughly: below about 0.4 nonpolar covalent, about 0.4 to 1.8 polar covalent, and large differences (about 2 or more) mostly ionic. It is a continuum, not sharp cutoffs.
6. What is lattice energy and what increases it?The energy released when gaseous ions form an ionic solid (or needed to separate it). It increases with higher ion charges and smaller ion sizes, e.g. MgO > NaCl.
7. Exceptions to the octet ruleOdd-electron species (NO), incomplete octets (BF3, boron with 6 electrons), and expanded octets for period 3 and below (SF6, PCl5).
8. Draw-in-words: the Lewis structure of CO2O=C=O: carbon double-bonded to each oxygen, each oxygen with two lone pairs, no lone pairs on carbon; linear, 180 degrees.
9. VSEPR shape and bond angle of H2O4 electron domains (2 bonds, 2 lone pairs): tetrahedral electron geometry, bent molecular shape, angle about 104.5 degrees.
10. Why does a lone pair compress bond angles?Lone pairs occupy more space near the central atom than bonding pairs, so they repel bonding pairs more strongly. Repulsion order: lone-lone > lone-bond > bond-bond.
11. VSEPR: shape of PCl5 and of ClF3PCl5 (5 domains, no lone pairs) is trigonal bipyramidal. ClF3 (5 domains, 2 lone pairs in equatorial positions) is T-shaped.
12. Hybridization of carbon in ethane, ethene and ethyneEthane: sp3. Ethene: sp2 (one pi bond). Ethyne: sp (two pi bonds).
13. Bond length and bond energy vs bond orderHigher bond order gives shorter and stronger bonds: C-C (longest, weakest) > C=C > C≡C (shortest, strongest).
14. Estimating delta H of reaction from bond enthalpiesdelta H is approximately the sum of bond energies broken minus the sum of bond energies formed. Breaking bonds costs energy; forming bonds releases energy.
15. Name the intermolecular forces from usually weakest to strongest.London dispersion forces, dipole-dipole forces, hydrogen bonding (ion-dipole forces are stronger still, between ions and polar molecules).
16. What is required for hydrogen bonding?An H atom covalently bonded to N, O or F, attracted to a lone pair on another N, O or F atom.
17. What determines the strength of London dispersion forces?Polarizability, which increases with the number of electrons (molar mass) and with more elongated shapes; e.g. n-pentane boils higher than neopentane.
18. Why does water have an unusually high boiling point for its size?Each H2O molecule can form extensive hydrogen bonds with neighbours, which require more energy to overcome than dispersion forces in similar-sized molecules like H2S.
19. Explain why ice floats.Hydrogen bonds lock ice into an open hexagonal lattice that is less dense than liquid water, where molecules pack more closely.
20. Relationship between intermolecular forces and vapour pressureStronger intermolecular forces mean fewer molecules escape to the gas phase, so vapour pressure is lower and the boiling point is higher.
21. Define normal boiling point.The temperature at which a liquid's vapour pressure equals 1 atm (101.325 kPa).
22. Heat for a temperature change vs a phase changeTemperature change within one phase: q = m·c·delta T. Phase change at constant temperature: q = n·delta H (e.g. delta H_fus or delta H_vap).
23. Mini-problem: heat to warm 100 g of water from 20 C to 30 C (c = 4.18 J/g·C)?q = (100 g)(4.18 J/g·C)(10 C) = 4180 J = 4.18 kJ.
24. On a heating curve, why is the temperature flat during melting?Added energy goes into overcoming intermolecular attractions (potential energy), not into raising kinetic energy, so temperature stays constant until the phase change is complete.
25. What do the triple point and critical point on a phase diagram represent?Triple point: the T and P where solid, liquid and gas coexist. Critical point: beyond this T and P, liquid and gas become indistinguishable (supercritical fluid).
26. Why does the solid-liquid line on water's phase diagram have a negative slope?Because ice is less dense than liquid water, increasing pressure favours the liquid, so the melting point decreases slightly as pressure rises.
27. Clausius-Clapeyron equation (two-point form)ln(P2/P1) = -(delta H_vap/R)·(1/T2 - 1/T1), relating vapour pressure to temperature; T in kelvin and R = 8.314 J/(mol·K).
28. Is enthalpy of vaporization usually larger than enthalpy of fusion? Why?Yes. Vaporization must separate molecules almost completely, while melting only loosens the arrangement, so delta H_vap is much larger than delta H_fus.