CHEM 130 Exam Review: Stoichiometry, Gases, Thermochemistry and Chemical Equilibrium

Terms in this set (32)

1. What is a mole?An amount containing 6.022×10^23 particles (Avogadro's number). The mass of one mole of a substance in grams equals its formula mass in atomic mass units.
2. Mini-problem: How many moles are in 36.0 g of water?Molar mass of H2O = 2(1.008) + 16.00 = 18.02 g/mol, so 36.0 g ÷ 18.02 g/mol ≈ 2.00 mol.
3. Mini-problem: 2H2 + O2 → 2H2O with 4 mol H2 and 3 mol O2. Which is limiting, and how much water forms?4 mol H2 needs only 2 mol O2, so H2 is limiting. 4 mol H2O forms and 1 mol O2 is left over.
4. Percent yield formulaPercent yield = (actual yield ÷ theoretical yield) × 100%. Theoretical yield is calculated from the limiting reactant.
5. Empirical vs molecular formulaThe empirical formula is the simplest whole-number ratio of atoms; the molecular formula is the actual number. Glucose is C6H12O6 with empirical formula CH2O.
6. Molarity and the dilution equationMolarity M = moles of solute ÷ liters of solution. When diluting, moles of solute stay constant: M1V1 = M2V2.
7. Write the net ionic equation for mixing AgNO3(aq) and NaCl(aq).Ag+(aq) + Cl−(aq) → AgCl(s). Na+ and NO3− are spectator ions and are left out.
8. How do you identify oxidation and reduction in a reaction using oxidation numbers?The species whose oxidation number increases is oxidized (it is the reducing agent); the species whose oxidation number decreases is reduced (it is the oxidizing agent).
9. Ideal gas law and the two common values of RPV = nRT. R = 0.08206 L·atm/(mol·K) when using liters and atmospheres, or 8.314 J/(mol·K) for energy units. Temperature must be in kelvin.
10. Molar volume of an ideal gas at STPAt STP (0 °C = 273.15 K and 1 atm), one mole of an ideal gas occupies about 22.4 L.
11. Dalton's law of partial pressuresP_total = P1 + P2 + ... and each partial pressure Pi = xi × P_total, where xi is the mole fraction of gas i.
12. What does kinetic molecular theory say about temperature and molecular speed?Average kinetic energy is proportional to absolute temperature. Root-mean-square speed = sqrt(3RT/M), with M in kg/mol, so at the same temperature lighter gases move faster.
13. Graham's law of effusionrate1/rate2 = sqrt(M2/M1). For example, H2 (2 g/mol) effuses about 4 times faster than O2 (32 g/mol).
14. When do real gases deviate most from ideal behavior, and why?At high pressure and low temperature, when molecules are close together: their own volume and intermolecular attractions are no longer negligible.
15. Sign of ΔH for exothermic vs endothermic processesExothermic: heat flows from system to surroundings, ΔH < 0, surroundings warm up. Endothermic: heat flows into the system, ΔH > 0, surroundings cool down.
16. Heat equation for a temperature change, and the specific heat of waterq = m·c·ΔT, where m is mass, c is specific heat capacity and ΔT is the temperature change. For liquid water, c = 4.184 J/(g·°C).
17. Mini-problem: How much heat raises 50.0 g of water from 20.0 °C to 30.0 °C?q = (50.0 g)(4.184 J/g·°C)(10.0 °C) = 2092 J ≈ 2.09 kJ.
18. State Hess's law.Because enthalpy is a state function, ΔH for an overall reaction equals the sum of ΔH values for any series of steps that add up to it. Reversing a step flips the sign; multiplying a step multiplies ΔH.
19. 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 (like O2(g) or C(graphite)) is zero.
20. How do you estimate ΔH from bond enthalpies?ΔH ≈ Σ (bond enthalpies of bonds broken) − Σ (bond enthalpies of bonds formed). Breaking bonds costs energy; forming bonds releases it.
21. Coffee-cup vs bomb calorimeterA coffee-cup calorimeter works at constant pressure and measures ΔH. A bomb calorimeter works at constant volume and measures ΔU (ΔE), often for combustion.
22. Write the equilibrium expression rules. What is left out?K = [products]^coefficients ÷ [reactants]^coefficients. Pure solids and pure liquids are left out because their activity is 1.
23. Reaction quotient Q vs K: which way does the reaction shift?If Q < K, the reaction proceeds forward toward products. If Q > K, it proceeds in reverse toward reactants. If Q = K, the system is at equilibrium.
24. Le Châtelier: effect of decreasing the volume of a gaseous equilibriumPressure increases, and the equilibrium shifts toward the side with fewer moles of gas. Adding an inert gas at constant volume causes no shift.
25. How does raising temperature affect K for an exothermic reaction?K decreases, because heat acts like a product; the equilibrium shifts toward reactants. For an endothermic reaction, K increases with temperature. Temperature is the only factor that changes K.
26. How does K change when you reverse a reaction, multiply it by n, or add reactions?Reverse: K becomes 1/K. Multiply all coefficients by n: K becomes K^n. Add two reactions: multiply their K values.
27. How are Kp and Kc related?Kp = Kc(RT)^Δn, where Δn = moles of gaseous products − moles of gaseous reactants and R = 0.08206 L·atm/(mol·K).
28. Mini-problem: For A ⇌ B with K = 4.0, starting with 1.0 M A and no B, find equilibrium concentrations.ICE: [B] = x, [A] = 1.0 − x. x/(1.0 − x) = 4.0, so x = 0.80. At equilibrium [A] = 0.20 M and [B] = 0.80 M.
29. What does a catalyst do to an equilibrium?Nothing to the position of equilibrium or K. It speeds up forward and reverse reactions equally, so equilibrium is reached faster.
30. Electron configurations of Fe and Fe3+Fe is [Ar]4s2 3d6. Fe3+ is [Ar]3d5: transition metals lose their 4s electrons before 3d electrons when forming cations.
31. Rank O2−, F−, Na+ and Mg2+ by size and explain.O2− > F− > Na+ > Mg2+. They are isoelectronic (10 electrons), so the ion with more protons pulls the same electrons in more tightly and is smaller.
32. Why does MgO have a much larger lattice energy than NaCl?By Coulomb's law, lattice energy grows with the product of ion charges and shrinks with distance. Mg2+ and O2− have double charges and are smaller than Na+ and Cl−.