CHM 1311 Midterm Review: Stoichiometry, Solutions, Gases and Thermochemistry

Terms in this set (29)

1. What is a mole?The amount of substance containing Avogadro's number of particles, 6.022 x 10^23. The mass of one mole in grams equals the formula mass in atomic mass units.
2. Mini-problem: how many moles are in 36.04 g of water?Molar mass of H2O = 2(1.008) + 16.00 = 18.02 g/mol, so 36.04 g / 18.02 g/mol = 2.000 mol.
3. Empirical formula vs molecular formulaThe empirical formula gives the simplest whole-number ratio of atoms. The molecular formula gives the actual number of each atom, a whole-number multiple of the empirical formula (glucose C6H12O6 has empirical formula CH2O).
4. Mini-problem: a compound is 40.0% C, 6.7% H and 53.3% O by mass. Empirical formula?In 100 g: 40.0/12.01 = 3.33 mol C, 6.7/1.008 = 6.6 mol H, 53.3/16.00 = 3.33 mol O. Dividing by 3.33 gives a 1:2:1 ratio, so CH2O.
5. Limiting reactantThe reactant that is completely used up first, which determines the maximum amount of product. Compare the mole ratio available to the ratio in the balanced equation.
6. Mini-problem: 2 H2 + O2 -> 2 H2O with 4.0 mol H2 and 3.0 mol O2. Limiting reactant and product?4.0 mol H2 needs only 2.0 mol O2, so H2 is limiting. It forms 4.0 mol H2O, with 1.0 mol O2 left over.
7. Percent yield formulaPercent yield = (actual yield / theoretical yield) x 100. The theoretical yield is calculated from the limiting reactant.
8. Molarity and the dilution equationMolarity M = moles of solute / litres of solution. When diluting, moles of solute stay constant, so M1V1 = M2V2.
9. Mini-problem: 50.0 mL of 2.00 M NaCl is diluted to 250.0 mL. New concentration?M2 = (2.00 M x 50.0 mL) / 250.0 mL = 0.400 M.
10. Strong vs weak electrolytesStrong electrolytes dissociate completely in water (soluble ionic compounds, strong acids such as HCl). Weak electrolytes ionize only partly (weak acids such as acetic acid). Nonelectrolytes do not ionize (sugar).
11. Net ionic equation for mixing AgNO3(aq) and NaCl(aq)Ag+(aq) + Cl-(aq) -> AgCl(s). Na+ and NO3- are spectator ions because they stay dissolved and unchanged.
12. Key oxidation number rulesElements alone are 0. Monatomic ions equal their charge. O is usually -2 (-1 in peroxides). H is +1 with nonmetals. The sum equals the overall charge of the molecule or ion.
13. Mini-problem: oxidation state of Mn in MnO4-x + 4(-2) = -1, so x = +7.
14. Ideal gas law and values of RPV = nRT, with T in kelvin. R = 0.08206 L·atm/(mol·K) or 8.314 J/(mol·K).
15. Mini-problem: volume of 1.00 mol of ideal gas at 0 degrees C and 1.00 atmV = nRT/P = (1.00)(0.08206)(273.15)/1.00, about 22.4 L.
16. Dalton's law of partial pressuresThe total pressure of a gas mixture equals the sum of the partial pressures of its components. Each partial pressure is P(i) = X(i) x P(total), where X(i) is the mole fraction.
17. Main assumptions of kinetic molecular theoryGas particles have negligible volume, move randomly in straight lines, have no attractions or repulsions, and collide elastically. Average kinetic energy is proportional to kelvin temperature.
18. Graham's law of effusionrate1 / rate2 = sqrt(M2 / M1). Lighter gases effuse faster; for example, H2 effuses 4 times faster than O2 because sqrt(32/2) = 4.
19. When do real gases deviate most from ideal behaviour?At high pressure (particle volume matters) and low temperature (intermolecular attractions matter). The van der Waals equation corrects for both.
20. Heat equation for a temperature changeq = m c delta T, where m is mass, c is specific heat capacity and delta T is the temperature change. For water, c = 4.184 J/(g·degree C).
21. Mini-problem: heat needed to warm 100.0 g of water by 10.0 degrees Cq = (100.0 g)(4.184 J/g·degree C)(10.0 degrees C) = 4184 J, or about 4.18 kJ.
22. Exothermic vs endothermic: sign of delta HExothermic reactions release heat to the surroundings and delta H is negative. Endothermic reactions absorb heat and delta H is positive.
23. Hess's lawThe enthalpy change of a reaction is the same whether it occurs in one step or several, so delta H values for steps can be added. Reversing a reaction flips the sign; multiplying it scales delta H.
24. Calculating delta H of reaction from enthalpies of formationdelta H rxn = sum of n x delta Hf(products) - sum of n x delta Hf(reactants). The standard enthalpy of formation of an element in its standard state is zero.
25. Coffee-cup vs bomb calorimetryA coffee-cup calorimeter works at constant pressure and measures delta H. A bomb calorimeter works at constant volume and measures delta U (internal energy change).
26. First law of thermodynamics and sign conventionsdelta U = q + w. q is positive when the system absorbs heat; w is positive when work is done on the system. Expansion work against constant pressure is w = -P delta V.
27. VSEPR shapes of CH4, NH3 and H2OAll have four electron groups (tetrahedral electron geometry). CH4 is tetrahedral (109.5 degrees), NH3 is trigonal pyramidal (about 107 degrees), and H2O is bent (about 104.5 degrees) due to lone pairs.
28. Why is CO2 nonpolar while H2O is polar?Both have polar bonds, but CO2 is linear so the bond dipoles cancel. H2O is bent, so its bond dipoles add to give a net dipole.
29. Formal charge formulaFormal charge = valence electrons - nonbonding electrons - (1/2) bonding electrons. The best Lewis structure keeps formal charges near zero, with negative charges on more electronegative atoms.