Download Ap Chem Seminar Notes

ap chem seminar notes
odds and ends:
fractional reaction orders exist when:
For example: find the order of the reaction for
A] CH4
b] Cl2
In the mechanism:
St 1
St 2
St 3
St 4
St 5
Cl2 ↔ 2 Cl
CH4 + Cl  CH3 + HCl
CH3 + Cl2  CH3Cl + Cl
CH3Cl + Cl  CH2Cl2 + H
H + Cl  HCl
fast equilibrium
slow
fast
fast
fast
What this shows is that we can solve for the concentration of an ________________ by
assuming that an _________________ is established in the __________ step.
There is another way to classify chemical reactions.
Reaction type 1: Synthesis
Description
General formula:
Example
Reaction type 2: Decomposition
Description
General formula:
Example
Note: hydroxides from a metal oxide and water when decomposed
Carbonates form a metal oxide and ________ when decomposed.
Reaction type 3: Displacement
Description
General formula:
3 types of displacement:
Examples
Reaction type 4: double displacement
Description
General formula:
Example
Note: When H2CO3 forms from an acid and a carbonate compound, the H2CO3
decomposes into: ___________ and ______________.
Reaction type 5: combustion
Description
General formula:
Example
Reaction type 6: Complex ion equilibria: (see chart, p. 649 Brown LeMay)
TTP 1: calculate the concentration of the silver ion present in solution at equilibrium
when concentrated ammonia is added to a 0.010 M solution of silver nitrate to give an
equilibrium concentration of ammonia = 0.20 M. Neglect the small volume change that
occurs on addition of ammonia.
TTP 2: (#108 p. 786 Zumdahl) A solution is formed by mixing 50.0 mL of 10.0 M NaX
with 50.0 mL of 2.0 x 10-3 M CuNO3. Assume that Cu(I) forms complex ions with X- as
follows:
Cu+ + X- ↔ CuX
K1 = 1.0 x 102
CuX + X- ↔ CuX2-
K2 = 1.0 x 104
CuX2- + X- ↔ CuX3-2
K3 = 1.0 x 103
Calculate the following concentrations at equilibrium: a] CuX3-2 b] CuX2- c] Cu+
Predict the products for:
(syn)
1] Al + Cl2 
(comb)
2] C3H8 + O2 
(dec)
3] CuBr 
(dis)
4] AlCl3 + F2 
(dec)
5] Mg(OH)2 
(dec)
6] MgCO3 
(dis)
7] AlCl3 + K 
(double dis)
8] AlCl3 + AgNO3 
(dis)
9] Al + HF 
(dec)
10] NaCl 
Buffered solution: chapter 15 Zumdahl
A buffered solution is one that:
Buffers contain:
Derivation of Henderson-Hasselbach equation:
TTP 1: A buffered solution contains 0.10 M acetic acid and 0.25 M sodium acetate. A]
Find it’s pH b] find the pH after 0.010 moles of gaseous HCl is added to 1.0 L of the
buffer solution c] find the pH after 0.010 moles of solid NaOH is added to 1.0 L of the
buffer solution
Buffer capacity:
Selecting an ideal buffer:
Titrations of strong acids and strong bases:
TTP 2: calculate the pH when the following quantities of 0.100 M NaOH have been
added to 50.0 mL of 0.100 M HCl solution (p. 633 Brown-LeMay): a] 49.00 mL
B] 49.90 mL c] 50.10 mL d] 51.00 mL
Titration of strong base with a weak acid:
Since the ______________ _______ of a weak acid affect the pH, these problems are a
bit more complicated.
TTP 3: Calculate the pH in the titration of acetic acid by NaOH after 30.0 mL of 0.100 M
NaOH has been added to 50.0 mL of 0.100 M acetic acid (p. 636 Brown-LeMay)
Titration curves for weak acid with strong base
Titration curve for weak base with a strong acid:
Note: at the half-way point of the titration:
Note: the pH is NOT ____ at the equivalence point for any strong acid/base with weak
acid/base titration. For weak acid, strong base titration, the pH at the equivalence point is
_________ 7. For a weak base, strong acid titration, the pH at the equivalence point is
___________ 7.
Titrations of polyprotic acids:
See curve of 25.0 mL of 0.10 M Na2CO3 with 0.10 M HCl (p. 641 Brown-LeMay)
Acid-base indicators (p. 756 Zumdahl)
The color change of an indicator takes place when:
[In-] / [HIn] =
TTP #1 Bromthymol blue, an indicator with a Ka value of 1.0 x 10-7, is yellow in its HIn
form and blue in its In- form. Suppose we put a few drops of this indicator in a strongly
acidic solution. If the solution is then titrated with NaOH, at what pH will the indicator
color change first be visible? (p. 753 Zum)
Chapter 6 – Thermochemistry
Nature of Energy:
State functions:
_________ is a state function, but ________ and ________ are not state functions.
First Law:
Internal energy:
Enthalpy and calorimetry
Specific heat capacity
Molar heat capacity
Hess’s Law:
Standard Enthalpies of formation
TTP 1: For the reaction: CH4(g) + 2O2(g)  CO2(g) +2H2O(g) H = -891 kJ
Calculate the enthalpy change for each of the following:
A] 1.00 g of methane is burned in excess oxygen
B] 1.00 x 103 L of methane gas at 740. torr and 25oC is burned in excess oxygen.
(p. 284 #36 Z)
TTP 2: It takes 78.2 J to raise the temperature of 45.6 g lead by 13.3oC. Calculate the
specific heat capacity and molar heat capacity of lead. (p. 284 #39 Z)
TTP 3: In a coffe-cup calorimeter, 50.0 mL of 0.100 M AgNO3 and 50.0 mL of 0.100 M
HCl are mixed to yield the reaction: Ag+ + Cl-  AgCl
The two solutions were initially at 22.60oC and the final temperature is 23.40oC.
Calculate the heat that accompanies this reaction in kJ/mol of AgCl formed. Assume that
the combined solution has a mass of 100.0 g and has a specific heat capacity of 4.18
J/oCg. (p. 284 #45)
TTP 4: Given the following data:
S + 3/2 O2  SO3 H = -395.2 kJ
2SO2 + O2  2SO3
H = -198.2 kJ
calculate H for the reaction: S + O2  SO2
(p. 285 #53 Z)
TTP 5: Using appendix 4, find Ho for C2H5OH(l) + 3O2(g)  2CO2(g) + 3H2O(g)
Chapter 10 – Liquids and Solids
Intermolecular Forces:
Dipole-Dipole forces
Typically these are about _____ percent as strong as __________ or ________ bonds.
Hydrogen bonding:
(see p. 454 Z)
London Dispersion forces:
Solids: two big categories: _______________ and ____________ .
Smallest repeating unit of a lattice of a crystalline solid is called the ______ _____.
p. 459.
Bragg equation:
p. 463 table 10.3
Metallic crystals pack together as either:
The ccp structure has a unit cell that is a:
Note: each corner of a unit cell is _____ of an atom. There are ____ corners of a unit cell.
The center of each face contains _____ of an atom. There are ____ faces of a unit cell
If there is a body-centered cubic unit cell, there is ___ atom in the center. See p. 465 Z.
Bonding model for metals:
Network atomic solids:
Molecular solids: These have _______ forces between atoms in the molecule, but
_______ forces between molecules. Examples:
Ionic solids: Typically, the larger ions are the ________ and they are packed in either an
hcp or ccp arrangement. The smaller __________ fit into the holes among the closest
packed anions.
There are three holes:
In order of size:
Table 10.7 p.483 Z.
Vapor pressure and changes in state
Eq. 10.4 p. 486 Z.
Changes of state:
Phase Diagrams:
TTP 1: Identify the most important type of interparticle forces present in the solids of
each of the following: (p. 504 # 35 Z.)
A] Ar
B] HCl
c] HF
d] CaCl2
e] CH4
f[ CO g] NaNO3
TTP 2] Predict which substance in each of the following pairs would have the greater
intermolecular attraction: (p. 504 #37 Z.)
A] CO2 or OCS
b] PF3 or PF5
c] SF2 or SF6
d. SO3 or SO2
TTP 3] Iridium (Ir) has a face-centered cubic unit cell with an edge length of 383.3 pm.
Calculate the density of solid Ir. (p. 505 #51 Z.)
TTP #4] What type of solid will each of the following substances form? (p. 506 #71 Z.)
A] CO2
b] SiO2
c] Si
d] CH4
e] Ru
f] I2
G] KBr
h] H2O
i] NaOH
j] U
k] CaCO3
l] PH3
TTP #5] How much energy does it take to convert 0.500 kg ice at -20.oC to steam at
250oC? The specific heat of ice is 2.1 J/oCg ; water is 4.2 J/oCg; steam is 2.0 J/oCg;
Hvap = 40.7 kj/mol; Hfus = 6.02 kj/mol
Chapter 16: Spontaneity, Entropy, and Free Energy
Spontaneous process:
Entropy:
Entropy and the second law of thermo:
The effect of temperature on spontaneity:
Free Energy:
The third law of thermo:
The dependence of Free Energy on Pressure:
Free energy and equilibrium:
The temperature dependence of K
TTP 1] For mercury, the enthalpy of vaporization is 58.51 kj/mol and the entropy of
vaporization is 92.92 j/Kmol. What is the normal boiling point of Hg?
TTP 2] For ammonia, the enthalpy of fusion is 5.65 kj/mol and the entropy of fusion is
28.9 j/Kmol.
A] Will NH3(s) spontaneously melt at 200. K? b] What is the approximate melting point
of ammonia?
TTP 3] Predict the sign of So for each of the following changes:
A] AgCl(s)  Ag+ + Clb] 2H2(g) + O2(g)  2H2O(l)
C] H2O(l)  H2O(g)
TTP 4] Predict the sign of So and then calculate So for each of the following reactions:
A] 2H2S(g) + SO2(g)  3Srhombic(s) + 2H2O(g)
B] 2SO3(g) 2so2(g) + O2(g)
C] Fe2O3(s) + 3H2(g)  2Fe(s) + 3H2O(g)
TTP #5] For the reaction at 298 K, 2NO2(g) ↔ N2O4(g)
The values of Ho and So are -58.03 kj and -176.6 j/K respectively. What is the value
of Go at 298 K? Assuming that Ho and So do not depend on temperature, at what
temperature is Go = 0? Is Go negative above or below this temperature?
TTP #6] For the reaction: SF4(g) + F2(g)  SF6(g)
The value of Go is -374 kj. Use this value and data from appendix 4 to calculate the
value of Go for SF4(g)
TTP 7] Consider the autoionization of water at 25oC,
H2O(l)  H+ + OH- ; Kw = 1.00 x 10-14
A] calculate Go for this reaction at 25oC b] at 40.oC, Kw = 2.92 x 10-14 . Calculate Go
at 40.oC.
p. 829 #’s 30, 31, 33, 33, 37, 46, 51, 61.
Chapter 17 – electrochemistry
Galvanic Cells:
In a galvanic cell, _________________ energy is converted into ______________
Energy.
In a galvanic cell, the emf (cell potential - €cell) is________ than zero.
For example, if copper and zinc are used, the diagram for a galvanic cell is:
Line notation
Cell potential, electrical work, and free energy
Faraday(F)
G = -nF€o
Concentration cells
Nernst equation
Ion-selective electrodes:
Calculations of Equilibrium constants for redox reactions
Batteries
Fuel cells:
Corrosion
Electrolysis:
TTP 1] Sketch the galvanic cells based on the following overall reactions. Show the
direction of electron flow and identify the anode and cathode. Assume all concentrations
are 1.0 M and all pressures are 1.0 atm. Determine the value of €o. Write the line
notation. Find the value of K. (Z #’s 25, 27, 31, 69)
A] Cu+2 + Mg(s) ↔Mg+2 + Cu(s)
B] Cr+3 + Cl2 ↔ Cr2O7-2 + Cl-
TTP 2] Place in order of increasing strength as oxidizing agents (all under standard
conditions)
Cd+2, IO3- ,
K+, H2O, AuCl4- ,
I2
TTP 3] Place in order of increasing strength as reducing agents (all under standard
conditions)
Cr+3 , H2,
Zn,
Li,
F- ,
Fe+2
TTP 4] A galvanic cell is based on the following half-reactions at 25oC:
Ag+ + e-  Ag
H2O2 + 2H+ + 2e-  2H2O
Predict whether €cell is larger or smaller than €ocell for the following cases:
A] [Ag+] = 1.0 M , [H2O2] = 2.0 M , [H+] = 2.0 M
B] [Ag+] = 2.0 M , [H2O2] = 1.0 M , [H+] = 1.0 x 10-7 M
C] for each case, find €cell (Z. #’s 55, 57)
TTP 5] What volume of F2 gas, at 25oC and 1.00 atm, is produced when molten KF is
electrolyzed by a current of 10.0 A for 2.00 hours? What mass of potassium metal is
produced? At which electrode does each reaction occur? (Z #85)
TTP 6] Electrolysis of a molten metal chloride, (MCl3) using a current of 6.50 A for
1397 seconds deposits 1.41 g of the metal at the cathode. What is the metal?
Chapter 18 – The representative elements: Groups 1A – 4A
The oxides of the 2A elements form bases in water, except:
Carbon dioxide is very different from the oxide of Silicon because:
Fluorine, F2, has a smaller electron affinity that Cl2 because:
Lithium is a stronger reducing agent than the elements below it because:
Lithium reacts more slowly with water than the elements below it b/c:
The bonds aluminum forms with nonmetals are significantly___________________,
causing it to have properties of both ____________ and bases. Substances that can act as
both an acid and a base are said to be:
Reaction:
Carbon has 3 allotropes:
Chapter 19 – The Representative Elements: Groups 5A – 8A
All the 5A can form 5 covalent bonds except:
This is because:
Nitrogen fixation: chart p. 925
Reactions of nitrogen:
Chemistry of phosphorus
Chemistry of Sulfur:
Why is HF a weak acid, but HCl is a strong acid?
Disproportionation:
Chapter 20 – Transition metals and Coordination Chemistry
Ions of elements #’s 21-30:
Table 20.5
Table 20.6
Coordination number:
Table 20.12
Table 20.13
Rules for naming coordination compounds: (p. 979)
The crystal field model:
Strong field and weak field cases:
Chapter 21: Nuclear chemistry
Figure 21.1
Table 21.2
Chapter 22: Organic chemistry
Alkanes
Saturated vs. unsaturated
Nomenclature:
Alkenes and alkynes:
Aromatic hydrocarbons:
Common functional groups: table 22.5
Alchohols:
Aldehydes and ketones:
Carboxylic acids and esters
amines
polymers