AP Chemistry Topic Review 2014

1. AP Chemistry Topic Review2014

2. Periodicity • Effective nuclear charge (Z) • Generally speaking, effective nuclear charge is the charge felt by the valence electrons after you have taken into account the number of shielding electrons that surround the nucleus. • 1st level of a neon atom 10 • 2nd level of a neon atom: 8 • 1st level of a sodium level: 11 • 2nd level of a sodium atom: 9 • 3rs level of a sodium atom: 1 • What about a sodium ion?

3. Coulombs Law Coulomb's Law relates the magnitude and sign of the electrostatic force  acting simultaneously on two point charges  and  as follows: F = (ke)q1q2 r2 where r is the separation distance and ke  is Coulomb's constant. If the product q1q2 is positive, the force between the two charges is repulsive; if the product is negative, the force between them is attractive. Compare NaCl to MgCl2

4. VSEPR

5. hybridization sp3: 4 EQUAL sigma bonds sp2: 3 sigma, 1 pi n sp2hybridization the 2s orbital is mixed with only two of the three available 2p orbitals sp: 1 sigma 2 pi

6. Spectroscopy Infrared spectroscopy deals with the interaction of infrared light with matter E = hn where h = 6.6 x 10-34 joule second and n = frequency of the photon. This shows that high energy photons have high frequency The frequency, n, and speed of light, c, are related through the relation c =  where c = 3.0 x 108 meter/second and l = wavelength for the light Molecules are flexible, moving collections of atoms. The atoms in a molecule are constantly oscillating around average positions. Bond lengths and bond angles are continuously changing due to this vibration. A molecule absorbs infrared radiation when the vibration of the atoms in the molecule produces an oscillating electric field with the same frequency as the frequency of incident IR "light"

7. The infrared spectrum for a molecule is a graphical display. It shows the frequencies of IR radiation absorbed and the % of the incident light that passes through the molecule without being absorbed. The spectrum has two regions. The fingerprint region is unique for a molecule and the functional group region is similar for molecules with the same functional groups.

8. Beer’s Law (Colorimetry, Spectroscopy) The graphing method is called for when several sets of data involving STANDARD SOLUTIONS are available for concentration and absorbance. This is probably the most common way of Beer's law analysis based on experimental data collected in the laboratory. Graphing the data allows you to check the assumption that Beer's Law is valid by looking for a straight-line relationship for the data. Question: What is the concentration of a 1.00 cm (path length) sample that has an absorbance of 0.60?

9. Beer’s Law A = εcl, where A is absorbance, c is concentration, l is cuvette width, and ε is the molar absorptivity of the molecule and I is intensity. If you have a set of solutions of known concentration, and you measured their absorbance, then plot A vs c, and the slope of this plot will be the molar absorbtivity ε (since l = 1). Then with ε, the concentration of a solution of this substance can be measured by measuring the absorbance.

10. Spectral Ranges

11. Classification

12. Molecular solids Vs. Ionic solids Intermolecular forces • Ionic bonding Attraction between – and + ions • Covalent bonding and Polarity • Hydrogen “bonding” H- F,O,N • Dipole dipole forces: high en difference • London dispersion forces: get higher with larger molecules Unique properties of water

13. Metals Metallic bonding constitutes the electrostatic attractive forces between the delocalized electrons, called conduction electrons, gathered in an electron cloud and the positively charged metal ions. Understood as the sharing of "free" electrons among a lattice of positively charged ions (cations), metallic bonding is sometimes compared with that of molten salts; however, this simplistic viewholds true for very few metals. In a more quantum-mechanical view, the conduction electrons divide their density equally over all atoms that function as neutral (non-charged) entities. Metallic bonding accounts for many physical properties of metals, such as strength, ductility, thermal and electrical conductivity, opacity, and luster

14. Network Solids In a network solid there are no individual molecules and the entire crystal may be considered a macromolecule Examples of network solids include diamond with a continuous network of carbon atoms and silicon dioxide or quartz with a continuous three-dimensional network of SiO2 units. Graphite and the mica group of silicate minerals structurally consist of continuous two-dimensional layers covalently bonded within the layer with other bond types holding the layers together

15. Ionic Solids Stable, high-melting substances held together by STRONG electrostatic forces that exist between oppositely charged ions.

16. Molecules • Lewis Dot Structures and Polarity of a molecule

17. Solutions Solution processes Polarity Particulate level representation

18. Solution stoichiometry • Acid base • Precipitation reactions • Redox reactions

19. Phases of Matter • KMT • Ideal gases • Ideal gas law

20. Gas laws • Boyles • Charles • Combined • Daltons

21. Graphs Heating curve Phase diagram

22. Critical Temperature: Temperature above which the vapor cannot be liquefied.Critical Pressure: Pressure required to liquefy AT the critical temperature.Critical Point: Critical temperatue and pressure (for water, Tc = 374° C and 218 atm).

23. Physical and Chemical Processes • Physical vs chemical changes • Reaction review • Enthalpy • Activation energy

24. Electrochemistry Redox reactions and cell potential Connections between cell potential and free energy Redox titrations

25. THE COLLISION THEORY OF REACTION RATES Particles must collide. Only two particles may collide at one time. Proper orientation of colliding molecules so that atoms can come in contact with each other to become products.

26. Rate laws • The basics • Example

27. Rate laws

28. Rate laws: A SUMMARY

29. the slowest step is the rate determining step reaction intermediate--produced in one step but consumed in another. catalyst--goes in, comes out unharmed and DOES NOT show up in the final rxn. Rate Mechanisms

30. Molecularityof the Rate Determining Step in the Mechanism

31. ENERGY AND WORK • E = q(heat) + w(work) • Signs of q: • +q if heat absorbed • –q if heat released w = -PV NOTE: Energy is a state function. (Work and heat are not.)

32. Signs of w (commonly related to work done by or to gases) • + w if work done on the system (i.e., compression) • -w if work done by the system (i.e., expansion) When related to gases, work is a function of pressure (pressure is force per unit of area) and volume

33. Exercise 1 Internal Energy Calculate ∆E for a system undergoing an endothermic process in which 15.6 kJ of heat flows and where 1.4 kJ of work is done on the system. ∆E = 17.0 kJ

34. ENTHALPY Measure only the change in enthalpy, H (the difference between the potential energies of the products and the reactants) • H is a state function • H = q at constant pressure (i.e. atmospheric pressure) • (true most of the time for us and a very handy fact!) • Enthalpy can be calculated from several sources including: • Stoichiometry • Calorimetry • From tables of standard values • Hess’s Law • Bond energies

35. Calorimetry The process of measuring heat based on observing the temperature change when a body absorbs or discharges energy as heat. After all data is collected (mass or volume; initial and final temperatures) we can use the specific heat formula to find the energy released or absorbed. We refer to this process as constant pressure calorimetry. ** q = H @ these conditions.**

36. Specific heat capacity (Cp) Confusing terms: Specific heat, specific heat capacity, molar heat capacity, latent heat Heat capacity (C): the amount of energy per rise (or fall) in temperature. This is an intensive property and unique to each substance. The specific heat capacity of water is 1cal/C or 4.18J/ C Energy released or gained: q = CmT Specific heat (S): Same as specific heat but specific to 1 gram of substance It is an extensive property. It’s units are J/g C S= heat transferred (grams of material)(temp change) Latent Heat: energy required for a phase change. Units are usually J/g

37. Equilibrium K = [products] [reactants] Leave out pure liquids and solids Kp = Kc(RT)n K > 1 products favored K < 1 reactants favored Use RICE tables!

38. LeChatelier’s Principle effect of changes in concentration, pressure, & temperature. Equilibrium always “shifts” away from what you add. “Stress” means too much or too little: chemical, heat, or room Ex:

39. Buffers Buffers are a mixture of a weak acid & its conjugate base or a weak base & its conjugate acid. Examples: HC2H3O2 & C2H3O2- or NH3 & NH4+ HA H+ + A-Ka = [H+][A-] [HA] The best buffer, Ka = [H+]; pH = pKa. The pH of a buffer can be adjusted by changing the ratio of acid and base.

40. Titrations • Steps: • Change volumes to moles using molarity • Subtract the moles used from the reactants and add moles gained by the products • Change moles back to molarity using the new volumes • Write and use your RICE table • Solve for x Example:

41. Solubility (Ksp) Solubility Rules Review/memorize these rules. They can be split into four groups: ALWAYS SOLUBLE: alkali metal ions (Na+, K+, Li+, Rb+, Cs+), NH4+, NO3-, C2H3O2-, ClO3-, ClO4- USUALLY SOLUBLE: chlorides, bromides, iodides (Cl-, Br-, I-) except “AP/H” (Ag+, Pb2+, Hg22+) sulfates (SO42-) except “CBS/PBS” (Ca2+, Ba2+, Sr2+, Pb2+) fluorides (F-) except “CBS/PM” (Ca2+, Ba2+, Sr2+, Pb2+, Mg2+) USUALLY INSOLUBLE: oxides/hydroxides (O2-, OH-) except “CBS” ((Ca2+, Ba2+, Sr2+) NEVER SOLUBLE: CO32-, PO43-, S2-, SO32-, CrO42-, C2O42- except alkali metals & NH4+

42. Ksp Example: Co(OH)2(s) Co2+ + 2OH- Ksp = [Co2+][OH-]2= 2.5 x 10-16 What is the pH of a saturated solution? Let x = the amount (moles) of solid that will just saturate 1 L of solution. R Co(OH)2(s) Co2+ + 2OH- • I --- 0 0 • C +x +2x • E x 2x (x) (2x)2= 4x3 = 2.5 x 10-16 x = 3.97 x 10-6 [OH-] = 2x = 7.94 x 10-6 pOH = 5.1 pH = 14- pOH = 8.9

43. Will a Precipitate Form? Ion Product (Qsp) = “reaction quotient”. Qsp < Ksp more solid will dissolve Qsp = Ksp solution is saturated Qsp > Kspppt will form until Qsp = Ksp Note: Be sure to calculate concentration of DILUTED ions. Example: 50. mL of 2.0 x 10-4M Co(NO3)2 is mixed with 200 mL of 1.0 x 10-3MNaOH. Will a precipitate form? [Note:Ksp given in other example problem.] [Co2+] = 2.0 x 10-4 M x = 4.0 x 10-5M [OH-] = 1.0 x 10-3M x = 8.0 x 10-4M Qsp = (4 x 10-5) (8 x 10-4)2 = 2.56 x 10-11 Qsp > Ksp; a precipitate will form!

44. Thermodynamics • Product or Reactant favored reactions depend on H, S, and absolute Temp

45. Free Energy Gibbs Free Energy, G, puts the effects of H, S, and Temperature together. G = H - TS G<0, G -, product-favored reaction G>0, G +, reactant-favored reaction G=0, reaction is at equilibrium

46. Electrochemistry Balance half reactions: • write two separate half-reactions • balance all atoms except H & O • balance O’s (add H2O’s) • balance H’s (add H+’s) • add e- ‘s to more positive side • balance e-‘s between half-reactions • combine half-reactions • adjust for basic solution if needed

47. Galvanic Cells Every atom has a different “potential” to accept electrons… “reduction potential” The reduction with the more positive E value will occur as written; the other reaction will reverse (oxidation). Ex: 2Ag+ + Cd 2Ag + Cd2+ +0.80 v – (-0.40 v) = 1.20 volts The difference in the E values is the voltage of a cell made using these two reactions. Oxidationoccurs at the Anode Reductionoccurs at the Cathode

48. Electrolytic Cells coulomb (C) = an amount of charge amp = current = charge per second 1 amp · 1 second = 1 Coulomb 1 C / amp·s Faraday constant, F: 1 mole e- = 96,500 C Electrolysis calculations begin with amp·s Example: How many moles of copper metal can be plated using a 10 amp circuit for 30 s? 10amp x 30s x 1C x 1 mol e- x 1 mol Ag = 1 amp·s 96500C 1 mol e- = 3.1 x 10-3 mole Ag

49. Organic Chemistry

50. Alkyl groups Cis and trans