Download Final Exam, Chem 111 2012 Study Guide

Study Detail #1 – Kinetics
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Background reading: Chemistry of smog. {14.1} Reaction rates. {14.2} Effect of concentration on rates. {14.3, to page 690} Reaction half-­‐lives. {covered in section 14.3, but in more detail than we need} Arrhenius equation. {14.4} Reaction mechanisms. {14.5, up to page 710} Catalysis. {14.6} Lab preview: The Clock Reaction. Problems in GKFD.
Problem Description 14.1 Recognizing products and reactants in concentration vs time graphs 14.3 Rate dependence on concentration for 2nd order kinetics 14.4 Rate dependence on concentration for 1st order kinetics 14.21 Stoichiometry of nitrogen-­‐oxygen reactions 14.23 Average vs instantaneous rates 14.25 Qualitative: why rates change over time 14.27 Predicting relative rates from stoichiometry 14.33 Predicting relative rates from stoichiometry 14.35 Calculating average rates from time-­‐concentration tables. 14.45 Inferring orders from a rate law 14.47 Inferring a rate law from orders 14.49 Inferring a rate law from kinetic data 14.61 Inferring a rate law & rate constant from kinetic data Getting t1/2 & k from time-­‐concentration tables. Method: Use Excel to graph [ClO] over time, inspect this graph to figure out t1/2, then use Eq. 14.25 to 14.71 get a value for k. Turn in a hardcopy of your graph with t1/2 and k hand-­‐
written on it. 14.76 Activation energies in the forward and reverse directions 14.82 Activation energy and temperature Arrhenius analysis. Method: in Excel, plot ln(k) as a function of 1/T, and use E
14.85 trendline to get the slope and intercept. Using slope = − a and R
(a,b) intercept=ln(A), get Ea and A. Turn in a hardcopy of your graph and Ea (kJ/mol) and A. 14.93 Drawing reaction profiles €
14.95 Writing rate laws & and identifying molecularity 14.97 Deducing overall reaction stoichiometry Additional Problem: Evaluating the plausibility of a given multistep reaction mechanism
Suppose the observed stoichiometry for a reaction is 2NO + Cl2 → 2NOCl , and the
observed rate law is Rate = k[NO][Cl2 ]. A proposed mechanism is as follows:
Step 1: NO + Cl2 → NOCl2
€
€ + NO → 2NOCl
Step 2: NOCl
2
€ Show whether this is/is not a plausible mechanism, given that the reaction progress
diagram looks like the one below. Hint: have to consider overall stoichiometry and
€ predicted rate law, which in turn depends on which step is rate-limiting.
2 Learning goals and assessments:
1. Be familiar with reaction rates as they relate to concentrations.
a) State the principal factors that control reaction rates.
b) Given the rate of change of one reactant or product, calculate rates of change of
other reactants or products, and the overall reaction rate.
c) From graphical depictions of the decay of A over time (i.e., A→products), infer
whether the decay is 1st- or 2nd-order, and calculate half-lives and rate constants.
d) Calculate reaction rates from time-concentration tables.
e) Explain the difference between an average rate and an instantaneous rate.
f) Write rate laws appropriate for a given set of orders (e.g., when a reaction is first
order in A and second order in B, then the rate law is Rate=k[A][B]2).
g) Infer reaction orders from rate laws (the reverse of item (e)).
h) Deduce rate laws from initial reaction rates.
i) Infer rate constants (including units) from rates.
2. Be familiar with reaction rates as they relate to temperature.
a) Write the Arrhenius equation, and define all terms in it.
b) Construct an Arrhenius plot from rate and temperature data.
c) Given an Arrhenius plot, infer the energy of activation and frequency factor.
3. Understand the relationship between rate laws and reaction mechanisms.
a) Sketch a 1-step reaction profile and label reactants, products, transition states, and
the energy of activation.
b) Sketch a 2-step reaction profile and label reactants, products, transition states,
reaction intermediates, and the energies of activation.
c) Explain what is meant by the rate-limiting step in a multistep reaction.
d) State the molecularity and rate law for each step in a given multistep reaction.
e) Evaluate the plausibility of a given multistep reaction mechanism.
f) Suggest ways in which a catalyst might speed a reaction (e.g., via changes in
energy of activation and the frequency factor).
Study Detail #2 – Equilibria
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Introduction to equilibrium. {15.1} Equilibrium constant, K. {15.2} Manipulating K. {15.3} K & Q. {15.4} Heterogeneous equilibria. {15.6} Le Chatelier’s Principle. {15.7} Using initial-­‐change-­‐equilibrium (ICE) tables. {15.8} Problems in GKFD.
Problem Description 15.1 Equilibria and forward/backward reaction rates 3 15.7 Equlibrium concepts 15.15 Equilibria and forward/backward rate constants 15.23 Calculating K from equilibrium concentrations 15.31 KC and KP 15.37 K and stoichiometry 15.39 K and reactions multiplied by a factor 15.47 K for combined reactions 15.54 Using Q to predict equlibrium shifts 15.83,86,89 Le Chatelier 15.95 Calculating equilibrium concentrations. Note, the “drop the x” approximation (.00432-­‐x) is not accurate enough here – but you can use the “square root” trick described on p. 769. 15.96 Calculating equilibrium concentrations. Note, the “drop the x” approximation is appropriate here, but you need to justify it by showing the error is < 5%. Learning goals and assessments
1. Be familiar with descriptions of equilibrium.
a. Define the Law of Mass Action.
b. Explain what is meant in saying equilibrium lies “to the right” or “to the left”.
c. Calculate equilibrium constants from forward and reverse rates.
d. Write equations describing the steam-reforming and water-gas shift reactions.
e. Given a chemical equation, write the expression for KP or KC.
f. Convert values of KP to KC, and vice versa.
g. Find K for reversed or combined reactions, and reactions multiplied by a number.
h. Calculate Q for given conditions.
i. Use the value of Q to predict whether equilibrium will shift right or left.
j. Given initial concentrations, use ICE to find equilibrium concentrations,
justifying any approximations using the error<5% criterion.
k. Describe systems for which the term “heterogeneous equilibrium” applies, and
write mass balance action equations for such systems.
2. Be familiar with how a system responds to perturbations from equilibrium.
a. State Le Chatelier’s Principle.
b. Use Le Chatelier’s Principle to predict how equilibrium shifts because of changes
in temperature, volume, or additions of reactant or product.
4 Study Detail #3 – Bronsted-Lowry acids and bases
1. Introduction to acids and bases. {16.1} 2. The Bronsted-­‐Lowry model. {16.2} 3. pH and autoionization. {16.3} Problems in GKFD
Problem Description 16.13 Identifying acids and bases based on chemical equations 16.15 Conjugate acid-­‐base pairs 16.17 Calculation of [H+] in solutions with a strong acid 16.19 Calculation of [OH-­‐] in solutions with a strong base; for this problem, assume that Sr(OH)2 dissociates to Sr2+ + 2 OH-­‐. 16.27 Calculating pH and pOH from [H+] and [OH-­‐] 16.29,31 Calculating pH and pOH from solutions with strong acid or base 16.43 Chemistry of B-­‐L bases Additional Problem
NaOH and NH3 both produce OH- when dissolved in water, but only NaOH is considered
a strong base. Explain why.
Learning goals and assessments
1. Be familiar with the Bronsted-Lowry model of acids and bases.
a. Give examples of a few common household acids and bases.
b. State the Bronsted-Lowry acid/base definition.
c. For a given reaction, identify the acid and the base.
d. Define and give examples of amphoterism.
e. Identify the conjugate base or acid of a given acid or base.
f. Given stock concentrations, calculate the concentration of H+ ions in solutions of
strong acids (e.g., 0.1 M HCl).
g. Given stock concentrations, calculate the concentration of OH- ions in solutions of
strong bases (e.g., 0.1 M NaOH or Sr(OH)2).
h. Explain the leveling effect.
i. Explain differences between strong vs weak acids and bases.
5 j. Identify given acids as organic vs inorganic, identify the labile proton in it, and
make reasonable predictions about its strength (strong or weak).
k. Write acid dissociation reactions in the “standard” form (i.e., HA→A-+H+) and in
the water-explicit form (i.e., HA+H2O→A-+H3O+).
2. Be familiar with the pH and pOH scales.
a. Define, with a chemical equation, what is meant by autoionization of water.
b. Write equations relating pH to [H+], pOH to [OH-], [H+] to [OH-], and pH to pOH.
c. Calculate pH and pOH for solutions with given [H+] or [OH-] concentrations.
Study Detail #4 –Acid-base titrations & buffers
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pH calculations for weak acids and bases. {16.4, to middle of page 808} Polyprotic acids. {16.5, through page 811} Henderson-­‐Hasselbalch equation. {16.8} Buffers and buffer capacity. {16.9} Acid-­‐base titrations {16.10} Problems in GKFD.
Problem Description 16.35 Ranking [H+] using Ka 16.36 Ranking Ka using degree of dissociation 16.39 Writing the equilibrium constant expression for weak acid 16.47 Getting Ka and % dissociation using an ICE table 16.49 Getting pH using an ICE table. Note: the “drop the x” approximation is OK here, barely (gives 5.4% error). 16.83 Getting pH using Henderson-­‐Hasselbalch 16.87 Getting base/acid ratios using Henderson-­‐Hasselbalch 16.98 Properties of acid-­‐base indicators Additional problems.
1. Suppose you have 100 mL of an acetic acid/sodium acetate buffer (a buffer
containing 0.5 M acetic acid and 0.5 M sodium acetate). Will the addition of 45
mmol of NaOH exceed the capacity of the buffer? Use the 10:1 rule for buffer
capacity.
6 2. Find the pH when 20 mL of 0.1 M HA(aq) (the analyte, pKa=4.5) is titrated with 5,
10, and 15 mL of titrant, 0.1M NaOH(aq) .
Learning goals and assessments
€
€pH properties of various kinds of acid-base mixtures.
1. Be familiar with
a. Define degree of dissociation of an acid, with examples.
b. Write the equilibrium constant expression for a given acid-base ionization.
c. Calculate pKa and pKb from Ka, and Kb, and vice versa.
d. Carry out calculations for a weak acid relating degree of dissociation
concentration, pH, and pKa.
e. Deduce the predominant form of an acid (i.e., HA or A-), at a given pH, given the
pKa of the protonated form.
f. Write the chemical formulas of carbonic acid and its conjugate base (bicarbonate
ion), nitrous acid and its conjugate base (nitrite ion), and acetic acid and its
conjugate base (acetate ion).
g. Use the Henderson-Hasselbalch equation in calculations.
h. Describe essential characteristics of acid-base buffers (e.g., preparation, capacities,
uses).
i. Carry out buffer capacity calculations.
j. Explain what polyprotic acids are, and write ionization equilibria for them.
2. Be familiar with how to plan and interpret acid-base titrations.
a. Define titrant and analyte, and write a balanced neutralization reaction.
b. Sketch a titration curve for the titration of a given weak acid with a strong base,
identify the equivalence point, ½-equivalence point, and buffer region (possibly
several, if the acid is polyprotic).
c. Predict pH at given points along a titration curve.
d. Interpret a given titration curve for the pKa.
e. Explain how a pH color indicator works.
Study Detail #5 – Electrochemistry
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Architecture of electrochemical cells. {18.2} Standard potentials E  . {18.4} Effect of concentration on electrochemical cells. {18.6, to page 906 only} Electrolytic cells. {18.8, to page 914 only} €
Problems in GKFD.
Problem Description 18.7 Using standard potentials to identify cathodes and anodes 18.8 Using standard potentials to identify cathodes and anodes 7 18.9 Porous separator 18.11 Porous separator 18.13 Deducing ½-­‐reactions and diagramming cells 18.14 Silver and Nickel voltaic cell 18.17 Assigning oxidation numbers and diagramming cells 18.33 Predicting spontaneity 18.34 Predicting spontaneity 18.41 Writing balanced cell reactions 18.46 Using the Nernst equation 18.47 Using the Nernst equation 18.48 Using the Nernst equation Learning goals and assessments
1. Be familiar with the design and chemistry electrochemical cells {18.1-­‐4}. a. Assign oxidation numbers to atoms in given molecular formulas. b. Define anode, cathode, electrode, porous separator, electrolyte, and voltaic cell. c. Given a spontaneous redox equation, identify the oxidizing and reducing agents, and write the corresponding redox ½-­‐reactions. d. Given a cell diagram, sketch the electrochemical cell (and vice versa). e. Given a cell reaction, write ½-­‐reactions occurring at the cell’s anode and cathode. 
f. Use E  tables to calculate E cell
for a given cell. 
g. Use E tables to identify the anode and which is the cathode in a given cell. h. Use E  tables to determine whether a given reaction will proceed spontaneously at standard conditions. €
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€i. Use E tables to predict the direction of electron flow in a given cell. 2.€Be familiar with the use of the Nernst equation {18.6}. a. Write the Nernst equation. €b. State conditions under which it is necessary to use the Nernst equation. c. Use the Nernst equation in calculations 3. Be familiar with electrolytic cells {18.8}. 8 a. Explain differences between voltaic cells and electrolytic cells. b. Describe how to convert a voltaic cell into an electrolytic cell. c. Identify the cathode and anode in a given electrolytic cell. Rules for assigning oxidation numbers (see p. 158 of text) 1) The oxidation number of elements in a compound sum to zero for a neutral molecule, or to the net charge for an ion. 2) Oxidation number of an atom in an element = 0. 3) Oxidation number of monatomic ion = charge. 4) Oxidation number of fluorine = -­1 in all compounds. 5) Oxidation number of hydrogen=+1, oxygen = -­2 in most compounds. Exceptions: Oxygen in peroxides ( O⎯O ) have -­1. Hydrogen bonded to metals have -­1. 6) Halogens are -­1. Exception: If combined with O or F, it’s positive. …
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Study Detail #6 – Alkanes
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Molecular and structural formulas. {3.1 & 3.4} Families of organic compounds. {3.2} Bonding in normal alkanes. {3.3} Constitutional isomers and nomenclature of alkanes. {3.5 & 3.6} Cycloalkanes. {3.7} Cis-­‐trans isomerism in cycloalkanes. {3.8} Review of secondary forces. {1.9 (pages 27-­‐32 only)} Physical properties. {3.9} Chemical reactions (combustion & halogenation). {3.10} Problems in Blei & Odian. Answers to odd-numbered problems are in back of book.
Problem Description Section 3.2 Combining power 3.1 3.5 Deducing correct molecular formulas 3.1 9 3.9 Identifying the functional group 3.2 3.13 Drawing expanded structural formulas 3.4 3.15 Drawing line structural formulas 3.4 3.17 Constitutional isomers 3.5 3.21(a,b,c) IUPAC naming 3.6 3.23(a,b,c) IUPAC naming 3.6 Identifying primary, secondary, etc. carbons 3.6 3.28 Cycloalkanes 3.7 3.29 Cycloalkanes 3.7 3.33 Physical properties of alkanes 3.3 3.37 Complete combustion of alkanes 3.10 Halogenation of alkanes 3.10 3.24(a,b) 3.38(a,b,c) Learning goals and assessments 1. Be familiar with structure and nomenclature of alkanes. a) State the combining power for elements C, N, O, H, halogens, P, and S. b) Explain the difference between various types of structural formulas (normal, condensed, and expanded). c) Translate given structural formulas into line drawings, and vice versa. d) Use combining power to predict whether a given molecular formula is correct. e) Identify the family of a given hydrocarbon and make a line drawing of an example of each. f) Name the first ten normal alkanes based on a given molecular formula. g) Describe essential properties (hybridization, geometry, location of electron density buildup, ease of rotation) of sigma bonds made by a carbon atom. h) Describe key differences between conformational isomers and constitutional isomers. i) Make line drawings of all distinct constitutional isomers of a given alkane. j) Explain the 1o-­‐4o classification of alkanes and identify given carbon atoms using this classification. k) Sketch methyl through butyl alkyl groups, including their n-­‐, iso, s-­‐, and t-­‐ variants. l) Use IUPAC nomenclature to name given alkanes. m) Recognize and sketch cis-­‐trans stereoisomers of cycloalkanes. n) Explain differences between stereoisomerism, and constitutional isomerisim, and conformational isomerism. 2. Be familiar with physical properties of alkanes. a) Explain the difference between “primary forces” and “secondary forces”. b) Name the main type of secondary force that pertains to alkanes, and how the strength of that force depends on molecular mass and molecular shape. 10 c) Describe the relationship between secondary force strength and boiling point. 3. Be familiar with chemical properties of alkanes. a) Write balanced equations for combustion of hydrocarbons. b) Write chemical equations for halogenation of a given alkane, and describe the conditions needed for the reaction to take place. Detail #7 – Alkenes, alcohols, acids, and esters
1. Structure of alcohols, phenols, and ethers. {5.1}
2. Constitutional isomerism in alcohols. {5.2}
3. Alcohol nomenclature. {5.3}
4. Physical properties of alcohols. {5.4}
5. Acid-base properties of alcohols. {5.5}
6. Physical properties of carboxylic acids. {7.4}
7. Acidity of carboxylic acids. {7.5}
8. Soaps. {7.7}
9. Synthesis of esters. {7.8}
10. Ester nomenclature. {7.9}
Problems in Blei & Odian
Problem Description Section 5.2 Distinguishing alcohols, phenols, and ethers 5.1 5.10 Hydrogen bonding in methanol 5.4 5.13 Reaction of alcohols with strong acid 5.5 5.29 Physical properties of alcohols 5.4 5.42 Solubility of alcohols, ethers, and alkanes 5.4 7.10 Hydrogen bonding in carboxylic acids 7.4 7.21 Drawing carboxylate salts 7.6 7.27 Identifying soaps 7.7 7.29(a) Synthesis of esters 7.8 7.34(a) Naming esters 7.9 7.39 Acid-­‐ and base-­‐cataylyzed hydrolysis of esters 7.11 7.55 Geometry of an ester. 7.9 11 Learning goals and assessments
1. Be familiar with the structure of alkenes and alkynes. a) Explain structural differences between alkenes, alkenes and alkynes. b) Explain structural differences between aliphatic and aromatic hydrocarbons. 2. Be familiar with structure, nomenclature, and reactivity of alcohols. a) Distinguish aliphatic alcohols from aromatic alcohols. b) Draw structural formulas for specified alcohols. c) Sketch hydrogen-­‐bonding configurations between alcohols and other alcohols, and between alcohol and water. d) Explain what is meant by “hydrophilic” and “hydrophobic” parts of a molecule. e) Given two alcohols, identify the more water-­‐soluble one. 3. Be familiar with structure and reactivity of carboxylic acids and their salts. a) Draw structural formulas for specified carboxylic acids. b) Sketch hydrogen-­‐bonding configurations between carboxylic acids and other carboxylic acids, and between carboxylic acids and water. c) Predict results of acid-­‐base reactions involving carboxylic acids and their salts with water, bases, and acids. d) Sketch and name carboxylate anions conjugate to given carboxylic acids. e) Explain how the solubility of carboxylic acids is different from solubility of their conjugate anions. f) Explain how soaps work. 4. Be familiar with structure, nomenclature, and reactivity of esters. a) Draw and name the ester resulting from the reaction of a given carboxylic acid with an alcohol. b) Draw and name the results of acid-­‐ and base-­‐catalyzed hydrolysis of an ester. c) Describe the hydrogen bonding capabilities of esters (i.e., as H-­‐bond acceptors or H-­‐bond donor). 5. Be familiar with general oxidation and structural properties of organic molecules. a) Given two organic molecules, identify the more oxidized or reduced one. b) Explain the difference between selective oxidation and combustion. c) Name a few oxidizing agents useful in selective oxidation of organic compounds. d) Identify sp1, sp2 and sp3 carbons in a given organic compound, and describe the bond angles associated with each. 12 Detail #8 – Amino acids and proteins
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Introduction to α-amino acids. {12.1}
Zwitterions. {12.2}
Peptides. {12.3}
The disulfide bridge. {12.4}
Protein structure. {12.5-7}
Problems in Blei & Odian
Problem Description 12.6 Section Definition of pI 12.2 Predominant forms of amino acids. For Glutamic acid, 12.7(a-­‐ Carboxylic
Ammonium
Substituent
b) acid pKa
pKa
pKa
2.2
9.9
4.4
12.2 12.8(a-­‐
Predominant forms of amino acids. b) 12.15 Constitutional isomers of a tripeptide 12.3 12.18 Cis-­‐trans isomers of peptides 12.3 12.19 Electrophoresis 12.2 12.31 Forces responsible for secondary structure 12.5 12.32 Forces responsible for tertiary structure 12.5 12.33 Types of attractions between side groups 12.5 12.35 Factors affecting protein structure 12.5 12.43 α-­‐keratins 12.6 12.44 Collagens 12.6 12.53 Hydrophilicity of the protein exterior 12.7 12.89 Structure and function of proteins 12.6-­‐7 12.92 Structure and function of proteins 12.6-­‐7 Learning goals and assessments
1. Be familiar with the structure and properties of amino acids. a) Sketch selected amino acids (glycine, alanine, serine, cysteine, lysine, histidine, and aspartic acid) in a pH-­‐appropriate form. b) Define “zwitterion”. 13 c) Explain how chirality (a.k.a. enantiomeric stereioisomerism) is different from cis-­‐trans stereoisomerism. d) Identify whether a given amino acid is the D-­‐ or L-­‐ enantiomer. e) Identify whether a given amino acid is chiral. 2. Be familiar with the structure and properties of peptides. a) Predict the pI of a given peptide. b) Predict the solubility of a given peptide. c) Given a horizontal expanded structure of a peptide, give its name according to the N-­‐C convention (e.g., Ala-­‐ser-­‐gly). d) Given the N-­‐C name of a peptide, draw the corresponding horizontal expanded structure and identify the peptide bonds in it. e) Identify the amide planes, including their four corners, in a given peptide. f) Identify the angles φ and ψ between amide planes. 3. Be familiar with the structure and properties of proteins. a) Explain the difference between a peptide and a protein. b) Infer the N-­‐ and C-­‐terminus ends of a given protein fragment. c) State essential characteristics of α-­‐helices and β-­‐sheets. d) Sketch examples of various types of secondary interactions of side groups (see Fig. 12.4 in the text). e) Sketch an example of a disulfide bridge (see Fig. 12.5 in the text). f) Describe distinctions between primary through quaternary protein structure. g) Describe distinctions between globular vs fibrous proteins. Detail #9 – Nucleic acids
1. Introduction to nucleotides. {13.1}
2. Nucleic acid formation from nucleotides. {13.2}
3. Nucleic acid structure. {13.3}
Problems in Blei & Odian
Problem Description Section 13.3 Naming and drawing a nucleotide 13.1 13.9 Drawing a nucleic acid 13.2 13.13 Naming and drawing a nucleic acid 13.2 13.17 What’s a histone 13.3 13.20 Complementary base pairs 13.3 13.25 Complementary base pairs 13.3 13.28 The central dogma 13.4 14 Learning goals and assessments
1. Be familiar nucleic acid fundamentals. a) Identify the three major parts of a nucleotide. b) Describe characteristic differences between the two major classes of heterocyclic bases (pyrimidines and purines). c) Describe characteristic differences between the two forms of pentose sugars (ribose and deoxyribose). d) Supply names and/or line drawings of a given nucleotide (e.g., dCMP). e) Supply names and/or line drawings for a given nucleic acid in the 5’ à 3’ ordering convention (e.g., dCMP-­‐dAMP-­‐dTMP). f) Explain the purpose of histones. g) Describe characteristic 3-­‐d aspects of DNA, including its double-­‐helix structure, the reverse ordering of the two strands, and complementary base-­‐
pairing. 15