CHEM123 Midterm Review February 27, 2011

1. CHEM123 Midterm Review February 27, 2011

2. About me No clue? NO PROBLEM! irenexliu@gmail.com

3. CHEM123 Midterm • Chp 12: Liquids, Solids, Intermolecular Forces • Chp 14: Chemical Kinetics • Chp 15: Chemical Equilibrium

4. Tips for CHEM123 • Heavily calculation focused - do lots of HARD problems • Watch out for units!!! (i.e. Ideal gas constant, radius of atoms)

5. Chapter 12 Topics include: Phase diagrams and transitions Clausius-Clapeyron equation Intermolecular forces Solids & Cubic packing Ionic solids & Interstitial sites Born-Haber Cycle

6. Phase Transitions • Enthalpy all should be measured at fixed T and P.

7. Phase Diagrams • Phase diagrams show regions of uniform phases for different P and T • And regions where phase equilibria exist. • As P increases for constant T • Pure gas • Solid gas interface (sublimation) (normal sublimation T because P = 1atm). • Pure solid • As T increases for constant P D solid liquid interface (melting) • Pure liquid • As P decreases with constant T • Liquid gas interface (boiling) • Gas Carbon dioxide E D C F Triple point B G A Petrucci Fig 13-19

8. Phase Diagram - water • Many different solid structures • Polymorphism • Slope for S/L is (-) • MP ↓as P↑ • Liquid denser than Solid Water

9. Liquid Properties ↑Intermolecular attraction = ↑ BP, ↑ γ, ↑ μ, ↓VP • Vapour pressure (VP): pressure given off material when system reaches equilibrium • Rate of evaporation = Rate of Condensation • Normal Boiling Point :Temp at which VP = 1 atm • Standard Boiling Point :Temp at which VP = 1 bar • Surface Tension: Energy to increase surface area of liquid • Viscosity (μ): liquid’s resistance to flow

10. Clausius-Clapeyron Eq. * Note The C- C equation can also be applied to other phase transitions, provided that the proper heat of phase transition is used. • Clausius-Clapeyron equation: • Simple ratio for P means P2 and P1 units just have to match • T must be in Kelvin • Hvapassumed to be constant, R is 8.314 J/ mol ∙ K

11. Problem1: Ans: 0.0445 atm

12. Intermolecular forces Intermolecular forces: forces of interaction among molecules Intermolecular forces strength example van der Waals or London dispersion ~ 0.1 kJ mol-1Ar, S8, CH4 Dipole-dipole ~ 10 kJ mol-1 CO, NO2 hydrogen bonding ~20-40 kJ mol-1H2O, HF, NH3 ionic attraction ~100-1000 kJ mol-1NaCl, KBr Intramolecular: covalent bonding ~100-2000 kJ mol-1 diamond Explain each term and correlate properties such as hardness, m.p., b.p., H, of materials with intermolecular forces in them.

13. Intermolecular forces • Dipole-Dipole: • polar molecules (dipole moments μ) • London Dispersion Forces: • All molecules (polarizabilityα) • ↑ with ↑ molecular weight • Hydrogen Bonding: • H and (O, N or F) • Strongest of the 3 4. Ionic

14. Intermolecular forces Types of intermolecular forces

15. Problem2: • Rank from lowest boiling point to highest • NaF, HF, H2 • Acetaldehyde CH3CHO, Acetic acid anhydride (CH3COO)2O, Acetone CH3COCH3, Ethanol C2H5OH, Alkane (CH3CH2CH3) ans: H2<HF<NaF CH3CH2CH3< CH3CHO < CH2COCH3< C2H5OH < (CH3COO)2O

16. Heating Curves Heat Capacity: Csolid, liquid, gas Enthalpy of fusion: ΔHfus Enthalpy of vaporization: ΔHvap

17. Solid Structures • Amorphous • Crystalline • Nature of bonding: Ionic, Network covalent, Molecular, Metalic • Geometry/packing: CUBIC, trigonal, tetragonal, hexagonal, monoclinic, triclinic, orthorhombic

18. Cubic packing (1 type of atom)

19. Closest-packed Structures Closest-packed sphere structure: 2 types Sequence Type ABAB… hcp(hexagonal closest packing) ABCABC... ccp (cubic closest packing) orfcc (face centred cubic) Density of a Crystalline Solid =(ncell* (molar mass/avogadro))/a3

20. Binary Ionic solids • Vapour pressure (VP): pressure given off material when

21. Problem 3: A metal has a face-centered cubic lattice, a density of 11.4 g/mL, and an atomic radius of 175 pm. What is its relative atomic mass? [Avogadro number = 6.022e23] 207.2 g / mol b. 197.0 g / mol c. 107.9 g / mol d. 106.4 g / mol e. 72.6 g / mol Ans: a

22. Problem 4: A Ans: a

23. Chapter 14 Chemical Kinetics

24. Chemical Kinetics Topics include: Rates of chemical reactions Differential & integrative rate laws: 0th, 1st, 2nd order Effect of temperature Measuring reaction rates Reaction mechanisms, steady state approx. Catalysis

25. Rates of reactions aA + bBcC + dD • Measuring rates of reactions by measuring change in property that depends on concentration with time • Colour, density, acidity.

26. Rate laws n : order of reaction with respect to A m: order of reaction with respect to B n + m: overall order of reaction k: reaction rate constant (function of temperature) • Differential rate law: relate rate of reaction to concentration of reactants • Rate law are generally EMPIRICAL, and not derived from stoichiometry • note: we use small “k” to denote rate constant, and big “K” to denote equilibrium constant • Integrative rate laws: relate concentration with time

27. Rate Laws ** Summary of rate laws**

28. Chemical Kinetics Ea - activation energy T - Temperature in K R = 8.314 J/mol ∙ K A and Ea depends on specific rxn • Arrhenius equation: Temp vs K: • Determining rates of reaction: method of initial rates • Set up a series of reactions with varying concentrations of reactants. • keep [A] fixed and vary [B] => Measure initial rates. • Then keep [B] fixed and vary [A] => Measure initial rates. • Take ratios of initial rates with the reagent concentrations to determine n and m. • Once n and m are determined, rate law can be solved for k

29. Chemical Kinetics PROBLEM 5: Experiments were performed with different initial concentrations of A and B (no C) for the reactionA+ B  C The initial reaction rates were determined and are given in the following table. Determine the order of the reaction with respect to A and B and the value of the rate constant. • solution: k=0.844L2/(mol2 ∙s) (see appendix)

30. Chemical Kinetics PROBLEM 6 :

31. Reaction Mechanisms • Reaction Mechanism must be consistent with overall reaction stoichiometry and the overall rate law • Each step is termed “elementary process” • Unimolecular: A → products (dissociation) • Bimolecular: A + B → products (bimolecular collision) • Termolecular: A + B + C → products (very rare) • Exponents for concentration are same as stoichiometric factors • Intermediates are produced in one elementary step and consumed in another • Rate determining step = slowest elementary step

32. Reaction Mechanisms Transition state ≠ intermediates!

33. Reaction Mechanisms H2(g) + ICl(g) → HI(g) + HCl(g) (slow) HI(g) + ICl(g) → I2(g) + HCl(g) (fast) H2(g) + 2 ICl(g) → I2(g) + 2 HCl(g) Rate of reaction: r = k[H2][ICl] HI= Reaction intermediate

34. Chemical Kinetics • Steady State Approximation • Assume concentration of any intermediate is constant • Rate of formation = rate of consumption • Intermediate is classified as neither product or reactant

35. Chemical Kinetics Problem 7: For the reaction X2 + Y + Z --> XY + XZ, the mechanism is believed to be X2 + Y ---> XY + X k1, very slow X + Z ---> XZ k2, very fast What is the reactive intermediate? a. X2 b. Y c. X d. Z e. XY Ans: c

36. Chemical Kinetics Problem 7 con’t: For the reaction X2 + Y + Z --> XY + XZ, the mechanism is believed to be X2 + Y ---> XY + X k1, very slow X + Z ---> XZ k2, very fast What is the concentration of the intermediate, according to the steady state approximation? [XY][XZ] / [X2][Y][Z] k1[X2][Y] - k2[X][Z] k1[X2][Y] / k2[X][Z] k2[XZ] / k1[XY] k1[X2][Y] / k2[Z] Ans: e

37. Chemical Kinetics Rate of production of P If [S] high then rate is zero order If [S] is low then rate is first order • Enzyme kinetics: steady state assumption • General reaction:

38. Chemical Kinetics HCOOH → H2O + CO Catalysis: increase rate of reaction via alternate reaction path with lower Ea. Homogenous catalyst: dissolved in reaction medium through out the reaction (H+ is the catalyst)

39. Chemical Kinetics • Heterogeneous catalysts: solid that speed up the ration rates of gas/liquid • Require interaction with an active site on catalyst surface • 2 CO + 2NO  2CO2 + N2

40. Chapter 15 Chemical Equilibrium

41. Chemical Equilibrium Topics include: The Equilibrium Process Equilibrium Constants Kc, Kp Manipulating K for Complex Reactions. External Effects on Equilibriums

42. Equilibrium • Vapour pressure, Chemical equilibrium, solubility equilibrium • Chemical equilibrium: forward rate = backward rate • Equilibrium constant (Kc) is constant for a given temperature • Kc>1: products in higher amount • Kc<1: reactants in higher amount • Magnitude of Kc does not indicate if a reaction happens OR how fast it will react

43. Equilibrium A B nA n B B A Manipulating Kc:

44. Equilibrium A B (1) B C (2) Overall Reaction A C (3) Kc for sequential reactions:

45. Equilibrium – Gases & Solids • Kp = Kc (RT) ∆ngas • For gas: can also be defined in terms of partial pressure • in general: R = 0.082 atm L/mol K • For solids: only concentration of gas/dissolved reactants is considered

46. Equilibrium – Gases & Solids PROBLEM 8: (1) (2)

47. Equilibrium - Q Reaction Quotient (Q): predict direction of change in Equilib. Q is defined for ANY concentration of reactants/products K is defined ONLY when chemical equilibrium is reached

48. Equilibrium - Q • Therefore, if • Q = Kc  chemical equilibrium Q < Kc  reactants are in excess. • System must change to reduce reactants and increase products Q > Kc  products are in excess. • System must change to reduce products and increase products.

49. Equilibrium • Effects on equilibrium: Le Châtelier Principle: • Change in concentration of reactants/products • Change in volume • Change in pressure • Adding inert gas (no change if constant V, change if constant P) ** the above will only change Q and disturb equilibrium ** ** Temperature will change K **

50. Equilibrium • Kc1 is the equilibrium constant at the temperature T1. • Kc2 is the equilibrium constant at the temperature T2. • T1 and T2 in Kelvin. • ∆ H is the enthalpy of reaction (assumed constant). Van’t Hoff Equation: describes relationship between temperature and equilibrium constant, K