Download Recaps and Additional Slides Chem 20B: Reference

Recaps and Additional Slides
I will keep updating these. Please check back!
Chem 20B: Reference Slides & Recaps
READ THE UPDATED SYLLABUS BEFORE AND AFTER EACH CLASS
http://www.nano.ucla.edu/_psw/chem20bw17.html or http://bit.ly/chem20Bw17
Syllabus will be adjusted after each lecture. Upcoming assignments are finalized
when date is green and underlined.
Lecture 1, Monday 9 January
Textbook: Principles of Modern Chemistry, 8th ed, Oxtoby, Gillis, & Campion
Discussion: Location & time for your Section on MyUCLA
You will learn both scientific intuition and how to work quantitative problems.
You will also learn some chemistry.
Goal: be an informed citizen, have basic scientific understanding and thinking
skills, and have background to go further in chemistry, science, engineering, etc.
I expect a lot of you, so that we can cover key issues in science and engineering this
quarter
Come prepared by having read material and be ready to discuss it.
Turn in homework in your section folder each lecture.
You will make up a problem (15×) for a set of lectures and answer it.
(The top few of the quarter receive nominal extra points + immortality!)
No late homework submission without prior approval of your TA.
1
Chem 20B: Highlights from the Syllabus, cont.
Grading:
Homework 25% (5% submitted + 10% graded problems + 5% creative problem +
5% reading memos)
Assessments 10% (5 × 2%)
Midterms 40% (2 × 20% each)
Final 25%
Exams:
No notes or calculators or phones or devices
You will receive a periodic table and list of formulas and constants
Recording lectures is not allowed without my explicit permission, and under no
circumstances can be posted online or otherwise transmitted
Office hours:
Tuesday 1030-1130 AM & Thursday 230-330 PM, 3041 Young Hall
(we may move if crowded)
Often on iChat, AIM, WeChat as psweiss
TA office hours are on the web site
READ THE UPDATED SYLLABUS BEFORE AND AFTER EACH CLASS
Key Chemical Knowledge
How big is an atom?
How strong is a chemical bond?
NB- energy is a critical parameter in chemistry and especially this quarter
Choose your favorite energy unit (kcal/mole, kJ/mole, J, eV, cm-1) and learn all the
conversions
How energetic is a visible photon?
How does the rest of the electromagnetic spectrum compare?
Dimensional analysis is very useful
2
Recap of Lecture #1: Intro & Gases
Final exam: Tuesday 21 March 630-930 PM, rooms to be announced
Simplified class web site & syllabus sites (case matters):
bit.ly/chem20Bw17 & bit.ly/Ch20Bw17Syl
Energy and units of energy - kcal/mole, kJ/mole, J, eV, cm-1
Bond strengths, photon energies
Spectroscopies (much more to come here)
Atomic size
Temperature & pressure
(State functions & linearizing relationships)
Average speed of N2 in air ~0.5 km/sec
Density change between solid N2 and N2 in air is ~600×. In general, estimate 1000×
En route today to
Ideal gas equation (of state)
pV = nRT
Dimensional analysis
Recap of Lecture #2: Gases
Final exam: Tuesday 21 March 630-930 PM, rooms to be announced
Simplified class web site & syllabus sites (case matters):
bit.ly/chem20Bw17 & bit.ly/Ch20Bw17Syl
Ideal gas law
pV = nRT
In solving problems:
Make table of initial and final (and sometimes intermediate) conditions
Convert units to match & check that units cancel properly – dimensional analysis
Touchstones:
N2 at room temperature ~0.5 km/sec
Mole of ideal gas (at STP) 22.4 L
Liquids and solids are denser than gases by ~1000×
3
Recap of Lecture #3: Non-Ideal Gases
Mass Spectrometry (start)
Real atoms and molecules interact
Know energy ranges of intramolecular and intermolecular interactions
Molecular, interatomic, and intermolecular potentials
Repulsive wall is steep on close approach
Depth of potential well is (~) bond strength
Position of potential well is bond length
Mass spectrometry
Use fragments and isotope ratios to identify to determine connectedness of
molecules and parts of molecules
Recap of Lecture #4: Phase Transitions
Phases
Solid – Condensed (close together), organized
Has vibrational but not translational nor rotational motion
Liquid – Condensed, not organized
Has translational, vibrational, and rotational motion
Gas – Separated (much more randomness è higher entropy than condensed
phases)
Has translational, vibrational, and rotational motion
Plasma – high energy, ionized gas
Ionization potentials are several eV (UV)
Transitions
Breaking bonds takes energy, SO making bonds gives off energy (heat)
Ionization energies are comparable to bond energies
4
Recap of Lecture #5: Interactions
Interactions and Interaction Strengths
Ionic bonding
Charges and separation
Covalent bonding
Metallic bonding
Delocalized electrons
Comparable cohesive energies (bond strengths)
Weaker Interactions
Ion-dipole
Hydrogen bonding
Dipole-dipole
Ion – induced dipole
Dipole – induced dipole
Dispersion (fluctuating dipole – induced dipole)
Recap of Lecture #6: Thermodynamics I
ΔG = ΔH – TΔS
So, ΔH < 0, making stronger bonds, is favorable
So, ΔS > 0, increased disorder, is favorable
ΔH and ΔS vary little with temperature. ΔG does vary with T è effect of ΔS
ΔG < 0 Keq > 1 Spontaneous
ΔG > 0 Keq < 1 Not spontaneous
ΔG = 0 Keq = 1 System at equilibrium (with equal amts of reactants & products)
Much more entropy in a gas than in a solution, liquid, or solid (in that order)
Polyatomic molecular gases have more entropy than diatomic molecular gases
than atomic gases
5
Recap of Lecture #7: Solutions, Acids, & Bases
Solution composition
Most common – molarity = moles/L
Also – mole fraction, molality, mass fraction (and normality)
Acid/Base Definitions
Arrhenius:
Acid - Proton donor, Base - Hydroxide donor
Bronsted-Lowry:
Acid - Proton donor, Base - Proton acceptor
Lewis:
Acid - Electron pair acceptor, Base - Electron pair donor
Equilibria
Strong acids dissociate essentially completely in water
Strong acids (for our purposes): HCl, HBr, HI, H2SO4, HNO3, HClO4
Strong bases react completely to give OH¯ in water
They are (for our purposes)
Alkali (Group 1A) hydroxides: LiOH, NaOH, KOH, RbOH, CsOH
Heavy alkaline earth (Group 2A) hydroxides: Ca(OH)2, Sr(OH)2, Ba(OH)2
H¯(typically used as NaH), NH2¯ (and there are more)
Ka Kb = [H+][OH¯] = Kw = 10-14
pX = -log10X
pKa + pKb = 14 = pH + pOH
Recap of Lecture #8: Equilibria & Oxidation States
Periodic Trends
Electronegativity ì
Electropositivity í
Ion and atom sizes í
These can be quantified, but be careful of sign conventions – use your intuition
Oxidation-reduction (redox) reactions
Disproportionation reactions – same element is both oxidized and reduced.
With electrochemistry, we can quantitate amount of reaction
Spontaneous reactions produce energy (generally make stronger bonds) Keq>1
Nonspontaneous reactions require energy, e.g., electrolytic reactions, Al reduction
Assign formal oxidation states:
1. Oxygen is almost always -2
2. Halogens (F, Cl, Br, I) always are -1, except when Cl, Br, I are bound to O or F
3. Hydrogen is always +1, except when bound to group I, II, or III metals, then -1
4. Other elements to balance charge on compound (can have mixed valency)
Ka Kb = [H+][OH¯] = Kw = 10-14
pKa + pKb = 14 = pH + pOH
Buffer pH~pKa
6
Recap of Lecture #9:
Oxyacids, Thermodynamics
Oxyacids
For a particular central atom, the more positive the formal oxidation state, the
stronger the acid
For a series of oxyacids where the central atom has the identical formal oxidation
state, the more electronegative the central atom, the stronger the acid
Laws of Thermodynamics
1 Energy is conserved
2 Entropy increases
3 At 0 K, S = 0 for a pure element
ΔG = ΔH – TΔS
So, ΔH < 0, making stronger bonds, is favorable
So, ΔS > 0, increased disorder, is favorable
ΔH and ΔS vary little with temperature. ΔG does vary with T è effect of ΔS
Recap of Lecture #10:
Thermodynamics & Electrochemistry
ΔG = -nFE and for standard states: ΔG° = -nFE°
n = number of electrons transferred in a balanced redox reaction
F = Faraday = 96,500 coulomb/mole e− = 96,500 J/V-mole e−
Standard States:
Solid
Pure solid
Liquid
Pure liquid
Gas
1 atm pressure
Solution
1M
Temperature
Usually 25 °C
ΔG° = -2.303 RT log10Keq
E° = 0.059 log10Keq
n
ΔG° = ΣΔGf°(products) - ΣΔGf°(reactants)
For elements, ΔGf° = 0
7
Exam #1 Topics
Midterm #1 Tuesday 7 February, 7-9 PM
Exam #1 covers through today’s lecture: recaps and coverage at
http://bit.ly/20B17recaps (also linked in syllabus)
Define temperature, pressure, state functions, phases
Bond lengths and strengths
Energy scales
Interactions and potentials – intramolecular, intermolecular, solids, liquids
Periodic trends
Ideal gas law and deviations
Partial pressures and concentrations
Thermodynamics, equilibria, and electrochemistry
Spontaneity, enthalpy, entropy, free energy, cell potentials
Phase transitions and equilibria
Typical, water, CO2
Solid, liquid, gas, supercritical fluid
Energy changes with making and breaking bonds, heat
Acids & bases
Electrochemistry
Relation to thermodynamics, energy storage
Heat capacity
Recap of Lecture #11: Thermodynamics,
Electrochemistry, and Concentrations
Le Chatelier’s Principle
Disturb a system from equilibrium and it will move to restore that equilibrium
è  One way to drive a reaction is to remove product
Quantify with concentration dependence of ΔG and E.
Batteries
Lead acid battery
Dry cell, alkaline cell
Rechargeable Ni-Cd battery
To get higher voltages, stack up cells in series (e.g., car battery 6 × 2 V = 12 V)
Electrolysis
Driving non-spontaneous reactions by applying electrical energy
The least unfavorable potential reaction goes first (there can be overlap)
Overpotentials and concentrated reactants are used
Midterm #1 Tuesday 7 February, 7-9 PM
10 AM Lecture (Section 2, all) in Moore 100
12 noon Lecture:
Sections 3A-D in Physics & Astronomy Building 1425
Sections 3E-I Lakretz 110
8
Recap of Lecture #12: Electrolytic Reactions
Electrolysis
Driving non-spontaneous reactions by applying electrical energy
The least unfavorable potential reaction goes first (there can be overlap)
NaCl(l) vs NaCl(aq)
Overpotentials and concentrated reactants are used
Recap of Lecture #14: Electrolysis & Acid/Base Equilibria
Electrolysis:
Quantify the amount of reaction – n, F, and number of moles
Acids and bases
HX(aq)
H+(aq) + X-(aq)
X-(aq) + H2O(l)
+
Ka = [H ][X ]
[HX]
HX (aq) + OH-(aq)
Exam #1 average = 75
Standard dev = 17
4@104 & 21@ ≥100
-]
Kb = [HX][OH
[X ]
Table of initial and equilibrium conditions
Solve problems by following the amount of reaction
Make and test assumptions about relative significance of initial concentrations and
amount of reaction (our limit here will be <3%)
Fractional dissociation and dilution
9
Recap of Lecture #15: Polyprotic Acids & Buffers
Polyprotic Acids
+
H+(aq) + HX-(aq) Ka1 = [H ][HX ]
[H2X]
H2X(aq)
HX-(aq)
Exam #1 average = 75
Standard dev = 17
4@104 & 21@ ≥100
+
2Ka2 = [H ][X- ]
[HX ]
H+(aq) + X2-(aq)
Table of initial, intermediate, and final equilibrium conditions
Solve problems by following the amounts of reactions
Make and test assumptions about relative significance of initial concentrations and
amount of reaction (our limit here will be <3%)
Buffer
pH = pKa – log10 [HA]
[A-]
pH should be within ~1 unit of pKa
How to make a buffer – acid + conjugate base
Buffer capacity
Recap of Lecture #16:
Solubility, Peroxides, Superoxides, Amphoterism
2-
-
Peroxide: O2
Superoxide: O2
Disproportionation reactions
An oxidizing agent oxidizes something else, and is reduced.
A reducing agent reduces something else, and is oxidized.
Hydride HSolubility product
MaXb(s)
aM+(aq) + bX-(aq)
Ksp = [M+]a[X-]b
In solving equilibria, keep track of stoichiometry
Solubility – how much of the solid (in moles/L) dissolves in a given solution
Amphoterism
e.g., Al(OH)3(H2O)3 is more soluble (as an ion) in acid or base than in neutral
solution - Al(OH)3(H2O)3 vs [Al(OH)2(H2O)4]+ or [Al(OH)4(H2O)2]-
10
Recap of Lecture #17:
Complex Ions + Photon Absorption & Emission
Solubility – Ksp
Drive equilibrium by adding or removing ions (products).
Common ion effect – solubility is reduced
Complex ions
Lewis acid – Lewis base adducts
Coordination determined by central element’s orbitals
Important in enzymes, catalysis, metal/salt dissolution
Formation & dissociation constants → KF = 1/KD
Dissolve insoluble salts
Chelation – multiple Lewis base sites on one molecule
Entropic advantage
Metal Ion Buffer – store metal ions as complex ions
Photon Absorption and Emission
Basis of spectroscopy
Review Session
Monday 27 Feb 7-9 PM Dodd 147
Recap of Lecture #18:
Fluorescence + Metals and Semiconductors
Photon Absorption and Emission, Fluorescence
Basis of spectroscopy
Metals
No band gap
HOMO, LUMO are at the same energy
Electronic excitation is small vs. kT
High electrical and thermal conductivity
Semiconductors and Insulators
Have a band gap
HOMO, LUMO are at different energies
More now!
Review Session
Monday 27 Feb 7-9 PM
Dodd 147
11
Lecture 20 (Kris Barr)
Fluorescence, Metals, Semiconductors, and Energy Scale
Midterm notes
CS 50 Sections 2A-2E
CS 24 Sections 2F-2L
Moore 100 Lecture 3
Regrades in pen only
Conflicts NEED TO BE CONFIRMED BY 1:30 TODAY
Review Session
Tonight Feb 7-9 PM
Dodd 147
Paul’s Office Hours
Tues 3:30 to 4:30 PM
Recap of Lecture #20:
Energy Scales, Fluorescence, Metals and Semiconductors
Photon Absorption and Emission, Fluorescence
Metals
No band gap
HOMO, LUMO are at the same energy
Electronic excitation is small vs. kT (thermal energy)
High electrical and thermal conductivity
Midterm exam #2:
CS 50 Sections 2A-2E
CS 24 Sections 2F-2L
Moore 100 Lecture 3
Regrades for exams taken in
pen only
Semiconductors and Insulators
Have a band gap
HOMO, LUMO are at different energies
Dopants in semiconductors
Elemental identity determines energy of dopant level
which then determines p- or n-type
Energy of dopant level and concentration determine conductivity of material
more dopant atoms è higher conductivity
Do not form a “band” – serve as a source or sink for thermally excited electrons
Exam #2 covers through last Friday’s lecture (along with Kris’ explanations Monday):
recaps and coverage at http://bit.ly/20B17recaps (also linked in syllabus)
Data sheet will be nearly identical to exam #1. No calculators, phones, devices, watches, etc.
12
Recap of Lecture #23: Kinetics
Kinetics
Rate laws
Reaction order
Order and stoichiometry are not the same
Clue to mechanism
Reaction dynamics = kinetics + mechanism
Location of reaction barrier determines effectiveness of translation vs.
vibration at promoting reaction.
Recall the mechanism and kinetics do NOT affect the thermodynamics (state
functions) and equilibra.
Reaction Order
Statistical/graphical analysis of kinetics
Concentration
First order reactions result in exponential changes in [reactants], [products]
Time
Recap of Lecture #24: Kinetics
Kinetics
Compare thermal energy to barrier
Reaction Order
Statistical/graphical analysis of kinetics
Early results of midterm #2:
79 ± 14
with all precincts reporting
Catalysis
Lower barrier to accelerate reaction
Recall the mechanism and kinetics do NOT affect the thermodynamics (state
functions) and equilibria
Catalytic Converters – catalyze multiple reactions
Oxidation: Hydrocarbons, CO → CO2 + H2O
Reduction: NO, NO2 → N2
Catalysts: CuO, Cr2O3, Pt, Rh
13
Recap of Lecture #24: Kinetics/Enzymes
Enzymes
Biological catalysts
Better specificity, turnover numbers, and control than artificial catalysts
Most are proteins, and can also include co-factors (vitamins) and metal ions
(There are also some RNA enzymes.)
Biopolymers
DNA – copolymer of four bases
RNA – copolymer of four bases
Proteins –
Copolymer of 20 amino acids, with ~100 other chemical modifications
Function depends on chemical modifications, environment, conformation
Lewis base sites coordinate metals
Polymerase chain reaction – enzyme to copy DNA
Drugs often interfere with enzyme activity (e.g., bind up active site)
Midterm #2: 79 ± 14
Recap of Lecture #26: Direct & Indirect Band Gap
Semiconductors, Semi-Metals, Nuclear Decay
Direct vs. indirect band gap semiconductors
Direct – top of valence band aligns with bottom of conduction band
light is efficiently coupled in & out - photodetectors, photodiodes, LEDs, lasers
e.g., GaAs
Indirect – direct band gap is larger than indirect gap, misaligned bands
inefficient coupling of light
e.g., Si, Ge
Semi-metal – no band gap, but 0 density of states at EFermi
Nuclear Stability
Unstable elements decay so as to move toward the band of stability
Transuranium elements undergo a series of alpha decays
Neutron-rich isotopes undergo beta decays
Neutron-deficient isotopes undergo positron emission or electron capture
Even proton and neutron numbers + closed shells favor stability
Radioactive Dating
First-order (exponential) decay kinetics
Measure elapsed time by following decay of an isotope
Example – 14C decay measures the time since respiration stopped (end of carbon uptake)
14C half-life of 5700 years, 238U half-life of 4.5x109 years
Accessible ages must be on these scales
After n half-lives, (½)n of original amount remains
14
Recap of Lecture #27: Nuclear Reactions II
Mass Defect
ΔE = Δmc2
Energy scales are much higher than chemical energies (KNOW THEM!)
Exothermic Nuclear Reactions
Rate stability as mass deficit per nucleon.
Binding energy per nucleon peaks at 56Fe
Lighter elements can undergo fusion exothermically
Heavier elements can undergo fission exothermically
Nuclear Chain Reactions
More neutrons produced than absorbed
Subcritical mass – not sustained
Critical mass – self-sustaining reaction
Supercritical mass – increasing reaction rate – bomb
Nuclear Reactors
Fission reactors – most common
Heat water to drive steam turbines
Use moderators to determine neutron capture rate and
thus reaction rate
Breeder reactors produce more fuel than they consume
Fusion reactors do not exist due to difficulty in sustaining high temperatures
required
Exam #2 Topics (in order of requests to review)
Materials – semiconductors, insulators, and metals
Semiconductors: Density of states, dopants
Fluorescence, energy scales, spectroscopy
Complex ions & chelates
Solubility
Disproportionation reactions
Amphoterism
Buffers
Acid/base equilibria
Polyprotic acids
Electrolytic reactions
Quantifying the amount of electrochemical reaction
Peroxide, superoxide, hydride
Oxidizing and reducing agents
15
Recap of Lecture #27: Nuclear Reactions II
Mass Defect
ΔE = Δmc2
Energy scales are much higher than chemical energies (KNOW THEM!)
Exothermic Nuclear Reactions
Rate stability as mass deficit per nucleon.
Binding energy per nucleon peaks at 56Fe
Lighter elements can undergo fusion exothermically
Heavier elements can undergo fission exothermically
Nuclear Chain Reactions
More neutrons produced than absorbed
Subcritical mass – not sustained
Critical mass – self-sustaining reaction
Supercritical mass – increasing reaction rate – bomb
Nuclear Reactors
Fission reactors – most common
Heat water to drive steam turbines
Use moderators to determine neutron capture rate and
thus reaction rate
Breeder reactors produce more fuel than they consume
Fusion reactors do not exist due to difficulty in sustaining high temperatures
required
Final Exam
Comprehensive Final Exam Tuesday 21 March 630-930 PM
CS 24 Sections 2E-2L
CS 50 Sections 2A-2D & 3A-C
CS 76 Sections 3D-3I
Those who are taking a make-up exam were set up last week
(Friday 3-6 PM, CS50 – calculators ok on make-up only, be sure to get extra
questions)
TA Review Session:
Saturday 18 March, 2-4 PM in CS76
Homework
Assessments
Two midterm scores
Final exam
25%
10%
40%
25%
Please complete your course evaluations by Saturday 8 AM
Please feel free to send me comments and suggestions directly
Office hours next week: Tuesday 1030-1130 AM, 230-330 PM &Thursday 230-330 PM
1
Nuclear Binding Energy
Sholto Ainslie Design, Wordpress
Data from NASA Goddard
Thermodynamics and Equilibria
Thermodynamics, free energy (ΔG), enthalpy (ΔH), entropy (ΔS), cell potential (E)
Use intuition, rules of forming stronger bonds (ΔH<0)
Entropy changes – gas/solution/liquid/solid
Know sign conventions
Relate thermodynamics, equilibrium constants, and electrochemistry (vs kinetics)
Phase transitions
Acids & bases, buffers
Solubility
Complex formation – Lewis acid – Lewis base interactions (including enzymes)
Periodic trends
Interaction strengths
Chemical bonds – covalent, ionic, metallic
Weaker interactions – hydrogen bonding, ion-dipole, dipole-dipole,
Electrochemistry – LAnOx & GRedCat
Oxidation states, half reactions, spontaneous & non-spontaneous reactions
Batteries, electrolysis
2
Kinetics
Rate laws
Reaction order
Order and stoichiometry are not the same
Clue to mechanism
Reaction dynamics = kinetics + mechanism
Location of reaction barrier determines effectiveness of translation vs.
vibration at promoting reaction
Recall the mechanism and kinetics do NOT affect the thermodynamics (state
functions) and equilibra
Reaction Order
Statistical/graphical analysis of kinetics
Catalysis
Lower barrier to accelerate reaction equilibration
Recall the mechanism and kinetics do NOT affect the thermodynamics (state
functions) and equilibra
Enzymes are biological catalysts with greater specificity and control than synthetic
catalysts
Materials
Semiconductors
Conduction and valence bands
Band gap, Fermi level
Density of states
n- and p-type, semi-insulating
Direct and indirect band gaps
Reactions of Si to make: insulators, metals, and to add dopants
Conductivity increases with increasing temperature
thermal excitation to conduction band or from valence band
Insulators
Conductivity increases with increasing temperature
Energy level diagram looks like semiconductor but bigger gap (compare to kT)
Metals
No band gap
Conductivity decreases with increasing temperature
3
Nuclear Chemistry
Decays - α, β-, β+, γ, electron capture
Balance decay reactions
Band of stability – also, use periodic table (what are average/common masses?)
Nuclear Stability
Unstable elements decay so as to move toward the band of stability
Transuranium elements undergo a series of α decays
Neutron-rich isotopes undergo β- decays
Neutron-deficient isotopes undergo positron emission or electron capture
Even proton and neutron numbers + closed shells favor stability
Radioactive Dating
First-order (exponential) decay kinetics
Measure elapsed time by following decay of an isotope
Example – 14C decay measures the time since respiration stopped (end of carbon uptake)
14C half-life of 5700 years, 238U half-life of 4.5x109 years
Accessible ages must be on these scales
After n half-lives, (½)n of original amount remains
Nuclear Energy
Mass defect, ΔE = Δmc2 – binding energy per nucleon peaks at 56Fe
Fission and chain reactions, fusion
Chemical Measurements
Infrared spectroscopy – vibrations, chemical fingerprint
Optical and ultraviolet spectroscopy – electronic excitation
X-ray spectroscopies – core levels, elemental identification
Fluorescence
Mass spectrometry
Electrochemical – thermodynamics
Balance cells and half-cells, count electrons
Quantitative
Also important in energy harvesting and storage
X-ray diffraction – spacings in and between molecules
Microscopies – real-space measurements
Know energy scales, both for photons and interaction strengths
4
Photon Energies
Infrared spectroscopy – vibrations, chemical fingerprint
Optical and ultraviolet spectroscopy – electronic excitation
X-ray spectroscopies – core levels, elemental identification
X-ray diffraction – bond lengths in crystals
(Microwave spectroscopy – rotations)
NASA, via Earth & Sky
5
Exam #2 Review
Materials
Semiconductors
Conduction and valence bands
Band gap, Fermi level
Density of states
n- and p-type, semi-insulating
Direct and indirect band gaps
Reactions of Si to make: insulators, metals, and to add dopants
Conductivity increases with increasing temperature
thermal excitation to conduction band or from valence band
EFermi
Insulators
Conductivity increases with increasing temperature
Energy level diagram looks
Metals
No band gap
Conductivity decreases with increasing temperature – scattering of carriers
Exam #2 Review
Energy scales
Reaction free energy changes and interaction strengths vs. photon energies
Measurements & spectroscopies
Energies and wavelengths
e.g., for diffraction, we match photon wavelength to bond length
choose X-rays (not yet covered)
Infrared – vibrational transitions è
Useful for chemical identification of functional groups (e.g., -CH3)
Visible & near UV – electronic transitions (not generally used for chemical ID)
X-ray – core level electronic excitations è
Useful for identifying elements, Li and beyond
16
Exam #2 Review
Buffers and acid/base equilibria
HA
H+ + A-
Ka =
[H+][A-]
[HA]
pKa = -log10Ka
pH = pKa – log10 [HA]
[A-]
pH should be within 1 unit of pKa
The more of the acid and base present,
the greater the capacity of the buffer
Exam #2 Review
Solubility
MaXb(s)
aM+(aq) + bX-(aq)
Ksp = [M+]a[X-]b
Solubility of a solid in a specific solution – how much dissolves in moles/L
Equilibria, LeChatelier’s Principle, Common Ion Effect, and è
Complex Ion Formation
Lewis acid-base adducts
Chelation and the effect of entropy in polydentate ligands
Use complex ions to dissolve solids or to supply/buffer metal ions where they are
needed in solution
Relate thermodynamics, equilibrium constants, and electrochemistry
(vs. kinetics)
Amphoterism – charged hydroxides are more soluble than neutral species
e.g., Al(OH)3(H2O)3 vs [Al(OH)2(H2O)4]+ or [Al(OH)4(H2O)2]-
17
Exam #2 Review
2-
-
Peroxide: O2
Superoxide: O2
Disproportionation reactions
An oxidizing agent oxidizes something else, and is reduced.
A reducing agent reduces something else, and is oxidized.
Hydride H-
Additional Slides
Here are a few additional slides for reference
18
Intramolecular and Intermolecular Potentials
PoMC, Oxtoby, Gillis, & Campion
Transformations: Energy and Order
Forming or breaking bonds? Breaking bonds costs energy (ΔH).
Which state is more ordered (ΔS)?
We will come back
to this topic
quantitatively
(Ch. 12)
http://ch302.cm.utexas.edu/physEQ/physical/physical-all.php
19
Real Gases – corrected figure labels on left
Deviations from ideal gas
Repulsive forces
(volume)
PV
RT
For N2
Attractive forces
(when close together)
P (atm)
PV
RT
P (atm)
wps.prenhall.com
Heat Water, Starting from Ice
Add heat to go from ice to water to steam
http://ch302.cm.utexas.edu/physEQ/physical/physical-all.php
20
Periodic Trends Reminder
Periodic Trends
Know which direction across the periodic table determines property.
Based on filling electron shells
Ionization Energy ì
Low if resulting ion has filled shell rare gas configuration.
(or to a lesser extent – has filled or half-filled subshells)
Same rules for higher oxidation states (e.g., Mg+2)
Electron Affinities í
(Negative values for species with stable anions)
Related to electronegativity ì – many ways to define this.
Determine dipoles within molecules.
Atomic & Ionic Sizes í
Size decreases with more positive oxidation state for isoelectronic atoms/ions.
.
Le Chatelier’s Principle
Disturb a system from equilibrium and it will move back to restore that
equilibrium
Increase a concentration è reaction shifts to decrease (consume) what was added
Increase pressure è reaction shifts to consume gas
+ more
This concept is used to drive reactions – by removing one or more of the products
cf. homeostasis
Stay tuned for buffers
21
Solution Compositions
Concentrations
Most common is molarity
M = moles solute / L solution
Water is 55.5 M
K2SO4(s)
H2O
2K+ + SO42-
Since K2SO4 is soluble in water,
1 M K2SO4(aq) is 2 M K+ and 1 M SO42Others used:
Mole fraction = moles of X / total moles
Mass percentage is common in everyday life
Less common are:
Normality = equivalents (usually acid or base) / L
Still used in medicine, but discouraged in science & engineering
Molality = moles solute / kg solvent
.
Balancing Reactions
Same number of atoms of each element on each side of reaction*
Same total charge on each side
In electrochemistry – we will also cover half-cell reactions.
In reduction, electrons will be a reactant (on left)
In oxidation, electrons will be a product (on right)
Total reaction will eliminate electrons from both sides.
Losing electrons at Anode is Oxidation &
Gaining electrons is Reduction at the Cathode
LAnOx & GRedCat
*Except
in nuclear reactions
.
22
Intermolecular and Intramolecular Forces I
All based on Coulomb’s Law (interacting charges)
www.umkc.edu
Intermolecular and Intramolecular Forces II
All based on Coulomb’s Law (interacting charges)
www.umkc.edu
23
Effect of Concentration on E and ΔG
Half Reactions
aA + bB + ne° –
E1/2 = E1/2
cC + dD
[C]c[D]d
2.303RT
log10
[A]a[B]b
nF
Cell Reactions
aA + bB
cC + dD
° –
Ecell = Ecell
[C]c[D]d
2.303RT
log10
[A]a[B]b
nF
ΔG = ΔG° + 2.303 RT log10
[C]c[D]d
[A]a[B]b
At equilibrium: ΔG = 0, so: ΔG° = –2.303 RT log10Keq
Buffers
slideshare.com
24
Metal Complexes
Coordination compounds
Lewis acid – Lewis base adducts
Important in enzymes, catalysis, metal/salt dissolution
Orbitals and oxidation state of central metal ion determine coordination.
Electronic excitation – absorption and emission
Lewis base ligands split electronic energies of metal ions – leading to color and
spin
Lone pair electrons repel and stay farthest away (as compared to ligands)
Spin
High spin vs. low spin compounds
Compare crystal field splitting (Δ) to the spin pairing energy (P)
Spectrochemical series – relative ligand effect on Δ
Paramagnetic – having one of more unpaired spins
Colors
Complementary colors – if a color is absorbed, the absorbing material will
appear as the complementary color
Red-green, orange-blue, yellow-violet
Other means of color: emission, interference
Energy and Excitations in Solids
Metals
No band gap - HOMO, LUMO are at the same energy
Electronic excitation is small vs. kT
High electrical and thermal conductivity
Conductivity decreases with increasing temperature because of electron scattering
Semiconductors and Insulators
Have a band gap between valence (lower) band and conduction (upper) band –
electronic excitation is >kT
This determines insulator vs. semiconductor
Direct band gap can be excited by photons (very little momentum)
Indirect band gap cannot be photoexcited efficiently
Electrons and “holes” can carry charge
Differentiate with magnetic field (Hall effect)
Energies of (dopant) states in the band gap determine conductivity and whether
electrons or holes dominate current
Conductivity increases with increasing temperature because of thermal excitation
of carriers
25
Semiconductor Electronic Structure
Semiconductor
(intrinsic)
Raising temperature increases conductivity
(i.e., decreases resistivity), because carriers
are thermally excited
EFermi
With impurities,
semiconductors have shifted
Fermi energies, from/to which carriers
can be thermally excited
wikipedia.com
Doping in Semiconductors
Thermally excite carriers from/to dopant level in the band gap
Semiconductor
Semiconductor
p-type
n-type
EFermi
EFermi
26
Direct vs. Indirect Band Gap Semiconductors
Energy and momentum are conserved in photoexcitation
Photons carry little momentum
Direct band gap semiconductors absorb and emit light
more efficiently than indirect band gap semiconductors
Indirect Band Gap Semiconductor
Direct Band Gap Semiconductor
Conduction
Band
E
Conduction
Band
E
Direct band gap
Valence Band
Valence Band
Momentum
Momentum
e.g., Si, Ge
e.g., GaAs, GaN
Semimetals
Semimetals have no band gap, but no states at the Fermi energy
e.g., graphite, graphene
Peplow, Nature 503, 327 (2013)
Millie
Dresselhaus
MIT
?
Kostya
Novoselov
Manchester
27
Band of Stability for Nuclei
arpansa.gov.au/radiationprotection
Example Examinations
Note that the syllabus has changed and there are different numbers of exams.
Averages for the exams shown:
Exam #1: 62
Exam #2: 79
Exam #3: 63
Final Exam: 87
28
Name _______________________ Section # ____ Student ID # ___________
Chemistry 20B, Winter 2016, Sections 1 & 3, Midterm #1
20 January 2016
6 questions + 2 small extra credit problems, 10 pages.
Answer on these sheets only. Additional space on last page.
If you need extra sheets, please ask your proctor or TA.
Note: Only these papers can be used; no other notes are allowed.
Please answer each question concisely. Show your calculations.
You may (and in some cases, must) draw explanatory diagrams.
Label all axes and features on graphs and diagrams.
You may not use a calculator, computer, watch, smart device, or electronics of any sort.
Irrelevant material will be ignored. Incorrect material will result in loss of points.
Table of constants and conversions
Speed of light: c = 3 × 108 m/s
Electron charge magnitude: e = 1.6 × 10-19 C
Plank's constant: ℏ = 1.1 × 10-34 J-s
Gas constant: R = 0.08206 L-atm/mol-K = 8.314 J/mol-K = 1.987 cal/mol-K
Boltzmann constant: kB = 1.4 × 10-23 J/K
Electron rest mass: m = 9.1 × 10-31 kg
Proton rest mass: M = 1.7 × 10-27 kg
1 mole = 6.02 × 1023
You will find a periodic table for your reference on the next page.
1
1
2
3
13
5
10.811
GROUP CAS
RELATIVE ATOMIC MASS (1)
GROUP IUPAC
ATOMIC NUMBER
B
1
2
Nonmetal
16 Chalcogens element
Semimetal
Alkali metal
17 Halogens element
Metal
Alkaline earth metal
- gas
58.693
- liquid
VIIIB
9
10
27 58.933 28
Ne
Ga
11
29
Fe
Tc
IB 12
30
63.546
- synthetic
- solid
STANDARD STATE (100 °C; 101 kPa)
18 Noble gas
Lanthanide
Transition metals
55.845
Actinide
VIB 7 VIIB 8
25 54.938 26
51.996
65.39
IIIA 14
6
10.811
IVA 15
7
12.011
VA 16
8
14.007
VIA 17
9
15.999
VIIA
18.998
18 VIIIA
2 4.0026
He
20.180
HELIUM
10
Ne
39.948
F
18
O
35.453
N
17
C
32.065
B
16
NEON
30.974
FLUORINE
15
OXYGEN
28.086
NITROGEN
14
CARBON
26.982
BORON
54
131.29
KRYPTON
Kr
83.80
Ar
126.90
36
Cl
79.904
S
53
BROMINE
Br
35
P
78.96
Si
34
Al
74.922
ARGON
33
CHLORINE
72.64
SULPHUR
32
PHOSPHORUS
69.723
SILICON
31
ALUMINIUM
13
13
5
PERIODIC TABLE OF THE ELEMENTS
IIA
9.0122
GROUP
2
4
http://www.ktf-split.hr/periodni/en/
1.0079
H
6.941
SYMBOL
BORON
ELEMENT NAME
VB 6
24
50.942
IIB
IA
1
1
3
HYDROGEN
Be
24.305
Li
12
BERYLLIUM
22.990
LITHIUM
11
MAGNESIUM
Na Mg
SODIUM
IVB 5
23
Se
47.867
As
IIIB 4
22
44.956
Ge
3
21
Ga
40.078
Zn
20
Cu
39.098
Ni
19
Co
127.60
Fe
52
Cr Mn
121.76
V
51
Ti
118.71
Sc
50
Xe
114.82
I
49
Te
112.41
Sb
48
Sn
RADON
Rn
(222)
At
86
ASTATINE
(210)
Po
85
POLONIUM
(209)
Bi
69
168.93
70
173.04
MENDELEVIUM
102
(259)
174.97
(262)
NOBELIUM LAWRENCIUM
Lr
103
LUTETIUM
Lu
71
Copyright © 1998-2003 EniG. ([email protected])
BISMUTH
(289)
167.26
Pb
114
Uuq
164.93
UNUNQUADIUM
67
68
Tl
162.50
LEAD
66
THALLIUM
84
In
(285)
MERCURY
112
UNUNBIUM
158.93
(258)
YTTERBIUM
Er Tm Yb
101
THULIUM
ERBIUM
(257)
FERMIUM
Fm Md No
100
Ho
(252)
HOLMIUM
99
Dy
(251)
Es
BERKELIUM CALIFORNIUM EINSTEINIUM
Cf
98
DYSPROSIUM
(247)
Tb
97
TERBIUM
65
208.98
Cd
(272)
157.25
(247)
83
Ag
64
Gd
96
207.2
Pd
151.96
(243)
82
Rh
63
95
204.38
XENON
150.36
(244)
CURIUM
81
IODINE
200.59
TELLURIUM
80
ANTIMONY
196.97
TIN
111
GOLD
Au Hg
79
INDIUM
195.08
CADMIUM
78
SILVER
192.22
PALLADIUM
77
RHODIUM
107.87
Ca
190.23
47
K
76
106.42
SELENIUM
186.21
46
ARSENIC
75
102.91
GERMANIUM
183.84
45
GALLIUM
101.07
Ru
44
ZINC
(98)
Tc
43
COPPER
95.94
NICKEL
42
COBALT
92.906
74
MOLYBDENUM TECHNETIUM RUTHENIUM
Pt
180.95
Ir
73
NIOBIUM
Nb Mo
41
IRON
91.224
CHROMIUM MANGANESE
40
VANADIUM
88.906
TITANIUM
39
SCANDIUM
87.62
CALCIUM
Zr
38
Y
85.468
POTASSIUM
Sr
37
Rb
ZIRCONIUM
Os
178.49
57-71
YTTRIUM
Re
(281)
W
110
Ta
(268)
Hf
62
MEITNERIUM UNUNNILIUM UNUNUNIUM
Mt Uun Uuu Uub
109
PLATINUM
(277)
IRIDIUM
108
OSMIUM
(264)
RHENIUM
107
Hs
(266)
TUNGSTEN
HASSIUM
106
Bh
(262)
TANTALUM
BOHRIUM
105
Sg
(145)
SEABORGIUM
61
Db
144.24
DUBNIUM
60
Rf
140.91
94
EUROPIUM GADOLINIUM
Nd Pm Sm Eu
(237)
AMERICIUM
Pu Am Cm Bk
NEPTUNIUM PLUTONIUM
Np
93
PRASEODYMIUM NEODYMIUM PROMETHIUM SAMARIUM
Pr
59
RUTHERFORDIUM
(261)
HAFNIUM
72
137.33
56
STRONTIUM
132.91
Actinide
Ac-Lr
89-103 104
Lanthanide
La-Lu
RUBIDIUM
Ba
55
Cs
(226)
BARIUM
Ra
88
RADIUM
(223)
CAESIUM
Fr
87
FRANCIUM
140.12
Ce
58
CERIUM
138.91
La
57
LANTHANIDE
LANTHANUM
238.03
U
92
URANIUM
231.04
Pa
91
PROTACTINIUM
232.04
Th
90
THORIUM
(227)
Ac
89
ACTINIDE
ACTINIUM
2
4
5
6
7
Editor: Aditya Vardhan ([email protected])
However three such elements (Th, Pa, and U)
do have a characteristic terrestrial isotopic
composition, and for these an atomic weight is
tabulated.
Relative atomic mass is shown with five
significant figures. For elements have no stable
nuclides, the value enclosed in brackets
indicates the mass number of the longest-lived
isotope of the element.
(1) Pure Appl. Chem., 73, No. 4, 667-683 (2001)
PERIOD
Question 1 (10 points):
In Flint, Michigan, the water supply has become contaminated with lead.
If you were worried that other metals might be in the water, how would you determine
which are present?
Describe how that information is obtained from your choice.
Question 2 (15 points):
A 3-m long cylinder has three compartments with two and each compartment has a
volume of 45 L. Initially, the separators are held in place and each compartment is open
to the air, with the pressure in the room at 760 torr (1.0 atm), and the temperature at 0 °C.
Assume that the air is an ideal gas.
a) How many moles of gas are in each compartment? (5 points)
b) The connections from the compartments to the room are sealed off and:
the pressure in the right chamber is increased to 2.0 atm,
the pressure in the left chamber is increased to 3.0 atm, and
then the separators (pistons) are released and move without friction.
What is the pressure in each chamber?
Where do the separators sit (draw a diagram) and why? (10 points)
Question 3 (30 points):
Heat is added at a constant rate to water as the temperature is raised from -20 °C to
120 °C.
Draw the time course of the temperature and density (two different curves) and
indicate the phase(s) at each temperature.
Label and concisely explain each stage of each curve and indicate where intermolecular
bonds are made or broken.
Question 4 (15 points):
How could you differentiate between ethylene (C2H4), nitrogen (N2), and carbon
monoxide (CO) using mass spectrometry?
Question 5 (15 points):
Why are there warning signs indicating “ionizing radiation”?
Briefly explain in terms of energies and other related issues.
Give two examples of ionizing radiation.
Question 6 (15 points):
How would you experimentally determine the atomic radius of Xe?
Briefly explain your choice of technique in terms of length scales, energies, and related
issues.
Extra credit #1 (2 points):
How are the gas constant, R, and the Boltzmann constant, kB, conceptually related
Extra credit #2 (2 points):
How are the energy units Hz and cm-1 conceptually related?
3
Name _______________________ Section # ____ Student ID # ___________
TOTAL = ________
Chemistry 20B, Winter 2016, Sections 1 & 3, Midterm #2
18 February 2016
5 questions + 2 small extra credit problems, 10 pages.
Answer on these sheets only. Additional space on last page.
If you need extra sheets, please ask your proctor or TA.
1)
Note: Only these papers can be used; no other notes are allowed.
4)
Please answer each question concisely. Show your calculations.
You may (and in some cases, must) draw explanatory diagrams.
Label all axes and features on graphs and diagrams.
5)
2)
3)
EC1+2)
You may not use a calculator, computer, watch, smart device, or electronics of any sort.
Irrelevant and/or incorrect material will result in loss of points.
Table of constants and conversions
Speed of light: c = 3 × 108 m/s
Electron charge magnitude: e = 1.6 × 10-19 C
Plank's constant: ℏ = 1.1 × 10-34 J-s
Gas constant: R = 0.08206 L-atm/mol-K = 8.314 J/mol-K = 1.987 cal/mol-K
Boltzmann constant: kB = 1.4 × 10-23 J/K
Electron rest mass: m = 9.1 × 10-31 kg
Proton rest mass: M = 1.7 × 10-27 kg
1 mole = 6.02 × 1023
ΔG° = -nFE° = -2.303 RT log10Keq
pH = pKa – log10 ([HA]/[A-])
You will find a periodic table for your reference on the next page.
1
1
2
3
13
5
10.811
GROUP CAS
RELATIVE ATOMIC MASS (1)
GROUP IUPAC
ATOMIC NUMBER
B
1
2
Nonmetal
16 Chalcogens element
Semimetal
Alkali metal
17 Halogens element
Metal
Alkaline earth metal
- gas
58.693
- liquid
VIIIB
9
10
27 58.933 28
Ne
Ga
11
29
Fe
Tc
IB 12
30
63.546
- synthetic
- solid
STANDARD STATE (100 °C; 101 kPa)
18 Noble gas
Lanthanide
Transition metals
55.845
Actinide
VIB 7 VIIB 8
25 54.938 26
51.996
65.39
IIIA 14
6
10.811
IVA 15
7
12.011
VA 16
8
14.007
VIA 17
9
15.999
VIIA
18.998
18 VIIIA
2 4.0026
He
20.180
HELIUM
10
Ne
39.948
F
18
O
35.453
N
17
C
32.065
B
16
NEON
30.974
FLUORINE
15
OXYGEN
28.086
NITROGEN
14
CARBON
26.982
BORON
54
131.29
KRYPTON
Kr
83.80
Ar
126.90
36
Cl
79.904
S
53
BROMINE
Br
35
P
78.96
Si
34
Al
74.922
ARGON
33
CHLORINE
72.64
SULPHUR
32
PHOSPHORUS
69.723
SILICON
31
ALUMINIUM
13
13
5
PERIODIC TABLE OF THE ELEMENTS
IIA
9.0122
GROUP
2
4
http://www.ktf-split.hr/periodni/en/
1.0079
H
6.941
SYMBOL
BORON
ELEMENT NAME
VB 6
24
50.942
IIB
IA
1
1
3
HYDROGEN
Be
24.305
Li
12
BERYLLIUM
22.990
LITHIUM
11
MAGNESIUM
Na Mg
SODIUM
IVB 5
23
Se
47.867
As
IIIB 4
22
44.956
Ge
3
21
Ga
40.078
Zn
20
Cu
39.098
Ni
19
Co
127.60
Fe
52
Cr Mn
121.76
V
51
Ti
118.71
Sc
50
Xe
114.82
I
49
Te
112.41
Sb
48
Sn
RADON
Rn
(222)
At
86
ASTATINE
(210)
Po
85
POLONIUM
(209)
Bi
69
168.93
70
173.04
MENDELEVIUM
102
(259)
174.97
(262)
NOBELIUM LAWRENCIUM
Lr
103
LUTETIUM
Lu
71
Copyright © 1998-2003 EniG. ([email protected])
BISMUTH
(289)
167.26
Pb
114
Uuq
164.93
UNUNQUADIUM
67
68
Tl
162.50
LEAD
66
THALLIUM
84
In
(285)
MERCURY
112
UNUNBIUM
158.93
(258)
YTTERBIUM
Er Tm Yb
101
THULIUM
ERBIUM
(257)
FERMIUM
Fm Md No
100
Ho
(252)
HOLMIUM
99
Dy
(251)
Es
BERKELIUM CALIFORNIUM EINSTEINIUM
Cf
98
DYSPROSIUM
(247)
Tb
97
TERBIUM
65
208.98
Cd
(272)
157.25
(247)
83
Ag
64
Gd
96
207.2
Pd
151.96
(243)
82
Rh
63
95
204.38
XENON
150.36
(244)
CURIUM
81
IODINE
200.59
TELLURIUM
80
ANTIMONY
196.97
TIN
111
GOLD
Au Hg
79
INDIUM
195.08
CADMIUM
78
SILVER
192.22
PALLADIUM
77
RHODIUM
107.87
Ca
190.23
47
K
76
106.42
SELENIUM
186.21
46
ARSENIC
75
102.91
GERMANIUM
183.84
45
GALLIUM
101.07
Ru
44
ZINC
(98)
Tc
43
COPPER
95.94
NICKEL
42
COBALT
92.906
74
MOLYBDENUM TECHNETIUM RUTHENIUM
Pt
180.95
Ir
73
NIOBIUM
Nb Mo
41
IRON
91.224
CHROMIUM MANGANESE
40
VANADIUM
88.906
TITANIUM
39
SCANDIUM
87.62
CALCIUM
Zr
38
Y
85.468
POTASSIUM
Sr
37
Rb
ZIRCONIUM
Os
178.49
57-71
YTTRIUM
Re
(281)
W
110
Ta
(268)
Hf
62
MEITNERIUM UNUNNILIUM UNUNUNIUM
Mt Uun Uuu Uub
109
PLATINUM
(277)
IRIDIUM
108
OSMIUM
(264)
RHENIUM
107
Hs
(266)
TUNGSTEN
HASSIUM
106
Bh
(262)
TANTALUM
BOHRIUM
105
Sg
(145)
SEABORGIUM
61
Db
144.24
DUBNIUM
60
Rf
140.91
94
EUROPIUM GADOLINIUM
Nd Pm Sm Eu
(237)
AMERICIUM
Pu Am Cm Bk
NEPTUNIUM PLUTONIUM
Np
93
PRASEODYMIUM NEODYMIUM PROMETHIUM SAMARIUM
Pr
59
RUTHERFORDIUM
(261)
HAFNIUM
72
137.33
56
STRONTIUM
132.91
Actinide
Ac-Lr
89-103 104
Lanthanide
La-Lu
RUBIDIUM
Ba
55
Cs
(226)
BARIUM
Ra
88
RADIUM
(223)
CAESIUM
Fr
87
FRANCIUM
140.12
Ce
58
CERIUM
138.91
La
57
LANTHANIDE
LANTHANUM
238.03
U
92
URANIUM
231.04
Pa
91
PROTACTINIUM
232.04
Th
90
THORIUM
(227)
Ac
89
ACTINIDE
ACTINIUM
2
4
5
6
7
Editor: Aditya Vardhan ([email protected])
However three such elements (Th, Pa, and U)
do have a characteristic terrestrial isotopic
composition, and for these an atomic weight is
tabulated.
Relative atomic mass is shown with five
significant figures. For elements have no stable
nuclides, the value enclosed in brackets
indicates the mass number of the longest-lived
isotope of the element.
(1) Pure Appl. Chem., 73, No. 4, 667-683 (2001)
PERIOD
Question 1 (25 points):
Limestone (CaCO3) is heated to produce lime (CaO) as an inexpensive, solid base in the
following reaction:
CaCO3(s) → CaO(s) + CO2(g)
a) What are the signs of ΔH, ΔS, and ΔG for this reaction? (20 points)
Explain your answer for each.
If any of the answers depend on additional information, indicate on what and how.
b) What can you say about the equilibrium constant Keq? (5 points)
Question 2 (15 points):
Referring again to the reaction:
CaCO3(s) → CaO(s) + CO2(g)
a) Write the equilibrium constant for this reaction (in terms of concentrations, etc.).
(5 points)
b) If a sealed container, initially under vacuum, with both CaCO3(s) and CaO(s) inside is
allowed to reach equilibrium, how might you measure the equilibrium constant?
(5 points)
c) When the volume of the container is doubled (by opening a valve to an evacuated
container), what happens? (5 points)
Question 3 (20 points):
On the first exam, you described what happened when heat is added at a constant rate to
water as the temperature is raised from -20 °C to 120 °C.
a) Draw the time course of the temperature again. Indicate the phase(s) present in each
temperature band. Indicate what determines the slope for each section of the plot.
(10 points)
b) Describe briefly what we know (i.e., signs and any other information) about ΔH, ΔS,
and ΔG at each phase transition? (10 points)
3
Question 4 (25 points):
A lithium ion battery, such as found in a cell phone, operates at a high cell potential
based on the following half-cell reactions:
CoO2 + Li+ + e− ↔ LiCoO2
E° ≅ 1 V
Li+ + C6(polymer/graphite) + e− ↔ LiC6(polymer/graphite)
E° ≅ -3 V
Compare to:
O2(g) + 4H+ + 4e− ↔ 2H2O
2H+ + 2e− ↔ H2(g)
Li+ + e− ↔ Li
E° = 1.23 V
E° = 0.00 V
E° ≅ -3 V
a) What is the overall cell reaction for the lithium ion battery shown? (5 points)
b) What is the overall standard cell potential of the reaction you wrote in (a)? (5 points)
c) What is the n (number of electrons) for the overall cell reaction you wrote in (a)?
(5 points)
c) Which elements of which compounds are being oxidized and which are being reduced
in the overall cell reaction you wrote in (a)? Determine the oxidation states of all the
species in the reaction. (5 points)
d) Can the reaction that you wrote in (a) be carried out in water (aqueous solution)? Why
or why not? (5 points)
Question 5 (15 points):
a) Why do metals absorb infrared light (i.e., why are they not transparent in the infrared)?
(5 points)
b) ZnO is a direct band gap semiconductor with a band gap of 3.3 eV. Why is it
colorless? (10 points, and see extra credit problem)
Extra credit #1 (5 points):
Why is ZnO used in sunscreen?
Extra credit #2 (5 points):
Estimate ΔG° for your answer to question 4b and explain your reasoning.
4
Name _______________________ Section # ____ Student ID # ___________
Signature ____________________
TOTAL = ________
Chemistry 20B, Winter 2016, Sections 1 & 3, Midterm #3
29 February 2016
5 questions + 2 small extra credit problems, 10 pages.
Answer on these sheets only. Additional space on last page.
If you need extra sheets, please ask your proctor or TA.
1)
Note: Only these papers can be used; no other notes are allowed.
4)
Please answer each question concisely. Show your calculations.
You may (and in some cases, must) draw explanatory diagrams.
Label all axes and features on graphs and diagrams.
5)
2)
3)
EC1+2)
You may not use a calculator, computer, watch, smart device, or electronics of any sort.
Irrelevant and/or incorrect material will result in loss of points.
Table of constants and conversions
Speed of light: c = 3 × 108 m/s
Electron charge magnitude: e = 1.6 × 10-19 C
Plank's constant: ℏ = 1.1 × 10-34 J-s
Gas constant: R = 0.08206 L-atm/mol-K = 8.314 J/mol-K = 1.987 cal/mol-K
Boltzmann constant: kB = 1.4 × 10-23 J/K
Electron rest mass: m = 9.1 × 10-31 kg
Proton rest mass: M = 1.7 × 10-27 kg
1 mole = 6.02 × 1023
Faraday constant = 96500 coul/mole
ΔG° = -nFE° = -2.303 RT log10Keq
pH = pKa – log10 ([HA]/[A-])
You will find a periodic table for your reference on the next page.
1
1
2
3
13
5
10.811
GROUP CAS
RELATIVE ATOMIC MASS (1)
GROUP IUPAC
ATOMIC NUMBER
B
1
2
Nonmetal
16 Chalcogens element
Semimetal
Alkali metal
17 Halogens element
Metal
Alkaline earth metal
- gas
58.693
- liquid
VIIIB
9
10
27 58.933 28
Ne
Ga
11
29
Fe
Tc
IB 12
30
63.546
- synthetic
- solid
STANDARD STATE (100 °C; 101 kPa)
18 Noble gas
Lanthanide
Transition metals
55.845
Actinide
VIB 7 VIIB 8
25 54.938 26
51.996
65.39
IIIA 14
6
10.811
IVA 15
7
12.011
VA 16
8
14.007
VIA 17
9
15.999
VIIA
18.998
He
18 VIIIA
2 4.0026
20.180
HELIUM
10
Ne
39.948
F
18
O
35.453
N
17
C
32.065
B
16
NEON
30.974
FLUORINE
15
OXYGEN
28.086
NITROGEN
14
CARBON
26.982
BORON
83.80
Ar
36
Cl
79.904
S
35
P
78.96
Si
34
Al
74.922
ARGON
33
CHLORINE
72.64
SULPHUR
32
PHOSPHORUS
69.723
SILICON
31
ALUMINIUM
13
13
5
http://www.ktf-split.hr/periodni/en/
IIA
9.0122
PERIODIC TABLE OF THE ELEMENTS
2
4
IA
1.0079
H
6.941
SYMBOL
BORON
ELEMENT NAME
VB 6
24
50.942
IIB
GROUP
1
1
3
HYDROGEN
Be
24.305
Li
12
BERYLLIUM
22.990
LITHIUM
11
Na Mg
IVB 5
23
Kr
47.867
Br
IIIB 4
22
44.956
Se
3
21
As
40.078
Ge
20
MAGNESIUM
Ga
39.098
SODIUM
Zn
19
Cu
131.29
Ni
54
Co
126.90
Fe
53
Cr Mn
127.60
V
52
Ti
121.76
Sc
51
Ca
118.71
K
50
Xe
114.82
I
49
Te
112.41
Sb
48
Sn
RADON
Rn
(222)
At
86
ASTATINE
(210)
Po
85
POLONIUM
(209)
Bi
69
168.93
70
173.04
MENDELEVIUM
102
(259)
174.97
(262)
NOBELIUM LAWRENCIUM
Lr
103
LUTETIUM
Lu
71
Copyright © 1998-2003 EniG. ([email protected])
BISMUTH
(289)
167.26
Pb
114
Uuq
164.93
UNUNQUADIUM
67
68
Tl
162.50
LEAD
66
THALLIUM
84
In
(285)
MERCURY
112
UNUNBIUM
158.93
(258)
YTTERBIUM
Er Tm Yb
101
THULIUM
ERBIUM
(257)
FERMIUM
Fm Md No
100
Ho
(252)
HOLMIUM
99
Dy
(251)
Es
BERKELIUM CALIFORNIUM EINSTEINIUM
Cf
98
DYSPROSIUM
(247)
Tb
97
TERBIUM
65
208.98
Cd
(272)
157.25
(247)
83
Ag
64
Gd
96
207.2
Pd
151.96
(243)
82
Rh
63
95
204.38
XENON
150.36
(244)
CURIUM
81
IODINE
200.59
TELLURIUM
80
ANTIMONY
196.97
TIN
111
GOLD
Au Hg
79
INDIUM
195.08
CADMIUM
78
SILVER
192.22
PALLADIUM
77
RHODIUM
107.87
KRYPTON
190.23
47
BROMINE
76
106.42
SELENIUM
186.21
46
ARSENIC
75
102.91
GERMANIUM
183.84
45
GALLIUM
101.07
Ru
44
ZINC
(98)
Tc
43
COPPER
95.94
NICKEL
42
COBALT
92.906
74
MOLYBDENUM TECHNETIUM RUTHENIUM
Pt
180.95
Ir
73
NIOBIUM
Nb Mo
41
IRON
91.224
CHROMIUM MANGANESE
40
VANADIUM
88.906
TITANIUM
39
SCANDIUM
87.62
CALCIUM
Zr
38
Y
85.468
POTASSIUM
Sr
37
Rb
ZIRCONIUM
Os
178.49
57-71
YTTRIUM
Re
(281)
W
110
Ta
(268)
Hf
62
MEITNERIUM UNUNNILIUM UNUNUNIUM
Mt Uun Uuu Uub
109
PLATINUM
(277)
IRIDIUM
108
OSMIUM
(264)
RHENIUM
107
Hs
(266)
TUNGSTEN
HASSIUM
106
Bh
(262)
TANTALUM
BOHRIUM
105
Sg
(145)
SEABORGIUM
61
Db
144.24
DUBNIUM
60
Rf
140.91
94
EUROPIUM GADOLINIUM
Nd Pm Sm Eu
(237)
AMERICIUM
Pu Am Cm Bk
NEPTUNIUM PLUTONIUM
Np
93
PRASEODYMIUM NEODYMIUM PROMETHIUM SAMARIUM
Pr
59
RUTHERFORDIUM
(261)
HAFNIUM
72
STRONTIUM
Actinide
Ac-Lr
89-103 104
Lanthanide
La-Lu
RUBIDIUM
56
137.33
132.91
Ba
55
Cs
(226)
BARIUM
Ra
88
RADIUM
(223)
CAESIUM
Fr
87
FRANCIUM
140.12
Ce
58
CERIUM
138.91
La
57
LANTHANIDE
LANTHANUM
238.03
U
92
URANIUM
231.04
Pa
91
PROTACTINIUM
232.04
Th
90
THORIUM
(227)
Ac
89
ACTINIDE
ACTINIUM
2
4
5
6
7
Editor: Aditya Vardhan ([email protected])
However three such elements (Th, Pa, and U)
do have a characteristic terrestrial isotopic
composition, and for these an atomic weight is
tabulated.
Relative atomic mass is shown with five
significant figures. For elements have no stable
nuclides, the value enclosed in brackets
indicates the mass number of the longest-lived
isotope of the element.
(1) Pure Appl. Chem., 73, No. 4, 667-683 (2001)
PERIOD
Name _______________________
Question 1 (15 points):
HF Ka = 7×10-4
CuF2 Ksp = 7×10-2
[Cu(NH3)4]2+ KF = 2×1013
Rank the solubility of CuF2(s) from highest to lowest in water, 0.1 M HF, 0.1 M NH3.
Concisely explain your answer, showing the reactions involved in each case. (15 points)
Question 2 (25 points):
Ethylenediamene (shown below) is a chelate that is typically abbreviated as “en”
The formation constants, KF, of octahedral Ni2+ complex ions at 25 °C are:
[Ni(en)3]2+
2 ×1018
[Ni(NH3)6]2+
4 ×108
a) Write the reactions for the formation and the formation constants (in terms of
concentrations) for each complex ion (10 points)
b) The ΔHf of the two complex ions are the same to two significant figures. How do you
explain the difference in formation constants? (5 points)
c) If NH3 were added to an aqueous solution of [Ni(en)3]2+ such that both concentrations
were initially 1 M after mixing, what would happen?
Write the reaction and discuss on which side the equilibrium would lie, and why.
(5 points)
d) If ethylenediamine were added to an aqueous solution of [Ni(NH3)6]2+ such that both
concentrations were initially 1 M after mixing, what would happen?
Again, write the reaction and discuss on which side the equilibrium would lie, and why.
(5 points)
Question 3 (20 points):
The monomer styrene (shown below)
3
is used to make polystyrene (styrofoam).
a) What is the structure of the polymer? You may answer by showing the repeat unit.
(10 points)
b) Show how this polymerization reaction proceeds, showing both the reactants and the
products as the polymer grows. Explain concisely why this reaction proceeds over other
possible reactions. (10 points)
Question 4 (30 points):
a) Choose three elements from a row of the periodic table and give the order you expect
for increasing oxidation half-cell potential of the neutral element. Concisely explain your
answer. (5 points)
b) Choose two of the above and write one as a balanced oxidation half-cell reaction and
the other as a balanced reduction half-cell reaction. Use realistic oxidation states.
Combine the two such that the overall reaction is spontaneous as written. (10 points)
c) Give the signs of ΔG° and E° that you expect for the reaction as written in b. What
can you say about Keq? (10 points)
d) Choose three elements from a column of the periodic table and give the order you
expect for increasing oxidation half-cell potential of the neutral element. Concisely
explain your answer. (5 points)
Question 5 (10 points):
DNA has special properties as both a polymer and the holder of our genetic code: it can
be synthesized precisely, copied, and measured (sequenced).
One property of a polymer (that we discussed) is rigidity.
A quantitative measure of this rigidity is the “persistence length” – below the persistence
length, the polymer behaves as a rigid rod; at larger scales, it behaves as a flexible chain.
The persistence length of double-stranded, double-helix DNA is 50 nm.
The persistence length of single-stranded DNA is ~4 nm (even lower at high salt
concentration, by the way).
a) Concisely explain this difference between single- and double-stranded DNA. (5 points)
b) What interactions make double-stranded DNA more rigid? (5 points)
Extra credit #1 (5 points):
4
In questions 2c and 2d, what would happen to the equilibria if the temperature were
raised from 25 °C to 50 °C? Explain your answers.
Extra credit #2 (5 points)
In question 5, how might double-stranded DNA be rigidified further (chemically)?
5
Name _______________________ Section # ____ Student ID # ___________
Signature ____________________
TOTAL = ________
Chemistry 20B, Winter 2016, Sections 1 & 3, Final Exam
Saturday 12 March 2016
5 questions + >1 extra credit problem, 12 pages.
Answer on these sheets only. Additional space on last page.
If you need extra sheets, please ask your proctor or TA.
1)
Note: Only these papers can be used; no other notes are allowed.
4)
Please answer each question concisely. Show your calculations.
You may (and in some cases, must) draw explanatory diagrams.
Label all axes and features on graphs and diagrams.
5)
2)
3)
EC)
You may not use a calculator, computer, watch, smart device, or electronics of any sort.
Irrelevant and/or incorrect material will result in loss of points.
Table of constants and conversions
Speed of light: c = 3 × 108 m/s
Electron charge magnitude: e = 1.6 × 10-19 C
Plank's constant: ℏ = 1.1 × 10-34 J-s
Gas constant: R = 0.08206 L-atm/mol-K = 8.314 J/mol-K = 1.987 cal/mol-K
Boltzmann constant: kB = 1.4 × 10-23 J/K
Electron rest mass: m = 9.1 × 10-31 kg
Proton rest mass: M = 1.7 × 10-27 kg
1 mole = 6.02 × 1023
Faraday constant = 96500 coul/mole
ΔG° = -nFE° = -2.303 RT log10Keq
pH = pKa – log10 ([HA]/[A-])
You will find a periodic table for your reference on the next page.
1
1
2
3
13
5
10.811
GROUP CAS
RELATIVE ATOMIC MASS (1)
GROUP IUPAC
ATOMIC NUMBER
B
1
2
Nonmetal
16 Chalcogens element
Semimetal
Alkali metal
17 Halogens element
Metal
Alkaline earth metal
- gas
58.693
- liquid
VIIIB
9
10
27 58.933 28
Ne
Ga
11
29
Fe
Tc
IB 12
30
63.546
- synthetic
- solid
STANDARD STATE (100 °C; 101 kPa)
18 Noble gas
Lanthanide
Transition metals
55.845
Actinide
VIB 7 VIIB 8
25 54.938 26
51.996
65.39
IIIA 14
6
10.811
IVA 15
7
12.011
VA 16
8
14.007
VIA 17
9
15.999
VIIA
18.998
18 VIIIA
2 4.0026
He
20.180
HELIUM
10
Ne
39.948
F
18
O
35.453
N
17
C
32.065
B
16
NEON
30.974
FLUORINE
15
OXYGEN
28.086
NITROGEN
14
CARBON
26.982
BORON
54
131.29
KRYPTON
Kr
83.80
Ar
126.90
36
Cl
79.904
S
53
BROMINE
Br
35
P
78.96
Si
34
Al
74.922
ARGON
33
CHLORINE
72.64
SULPHUR
32
PHOSPHORUS
69.723
SILICON
31
ALUMINIUM
13
13
5
PERIODIC TABLE OF THE ELEMENTS
IIA
9.0122
GROUP
2
4
http://www.ktf-split.hr/periodni/en/
1.0079
H
6.941
SYMBOL
BORON
ELEMENT NAME
VB 6
24
50.942
IIB
IA
1
1
3
HYDROGEN
Be
24.305
Li
12
BERYLLIUM
22.990
LITHIUM
11
MAGNESIUM
Na Mg
SODIUM
IVB 5
23
Se
47.867
As
IIIB 4
22
44.956
Ge
3
21
Ga
40.078
Zn
20
Cu
39.098
Ni
19
Co
127.60
Fe
52
Cr Mn
121.76
V
51
Ti
118.71
Sc
50
Xe
114.82
I
49
Te
112.41
Sb
48
Sn
RADON
Rn
(222)
At
86
ASTATINE
(210)
Po
85
POLONIUM
(209)
Bi
69
168.93
70
173.04
MENDELEVIUM
102
(259)
174.97
(262)
NOBELIUM LAWRENCIUM
Lr
103
LUTETIUM
Lu
71
Copyright © 1998-2003 EniG. ([email protected])
BISMUTH
(289)
167.26
Pb
114
Uuq
164.93
UNUNQUADIUM
67
68
Tl
162.50
LEAD
66
THALLIUM
84
In
(285)
MERCURY
112
UNUNBIUM
158.93
(258)
YTTERBIUM
Er Tm Yb
101
THULIUM
ERBIUM
(257)
FERMIUM
Fm Md No
100
Ho
(252)
HOLMIUM
99
Dy
(251)
Es
BERKELIUM CALIFORNIUM EINSTEINIUM
Cf
98
DYSPROSIUM
(247)
Tb
97
TERBIUM
65
208.98
Cd
(272)
157.25
(247)
83
Ag
64
Gd
96
207.2
Pd
151.96
(243)
82
Rh
63
95
204.38
XENON
150.36
(244)
CURIUM
81
IODINE
200.59
TELLURIUM
80
ANTIMONY
196.97
TIN
111
GOLD
Au Hg
79
INDIUM
195.08
CADMIUM
78
SILVER
192.22
PALLADIUM
77
RHODIUM
107.87
Ca
190.23
47
K
76
106.42
SELENIUM
186.21
46
ARSENIC
75
102.91
GERMANIUM
183.84
45
GALLIUM
101.07
Ru
44
ZINC
(98)
Tc
43
COPPER
95.94
NICKEL
42
COBALT
92.906
74
MOLYBDENUM TECHNETIUM RUTHENIUM
Pt
180.95
Ir
73
NIOBIUM
Nb Mo
41
IRON
91.224
CHROMIUM MANGANESE
40
VANADIUM
88.906
TITANIUM
39
SCANDIUM
87.62
CALCIUM
Zr
38
Y
85.468
POTASSIUM
Sr
37
Rb
ZIRCONIUM
Os
178.49
57-71
YTTRIUM
Re
(281)
W
110
Ta
(268)
Hf
62
MEITNERIUM UNUNNILIUM UNUNUNIUM
Mt Uun Uuu Uub
109
PLATINUM
(277)
IRIDIUM
108
OSMIUM
(264)
RHENIUM
107
Hs
(266)
TUNGSTEN
HASSIUM
106
Bh
(262)
TANTALUM
BOHRIUM
105
Sg
(145)
SEABORGIUM
61
Db
144.24
DUBNIUM
60
Rf
140.91
94
EUROPIUM GADOLINIUM
Nd Pm Sm Eu
(237)
AMERICIUM
Pu Am Cm Bk
NEPTUNIUM PLUTONIUM
Np
93
PRASEODYMIUM NEODYMIUM PROMETHIUM SAMARIUM
Pr
59
RUTHERFORDIUM
(261)
HAFNIUM
72
137.33
56
STRONTIUM
132.91
Actinide
Ac-Lr
89-103 104
Lanthanide
La-Lu
RUBIDIUM
Ba
55
Cs
(226)
BARIUM
Ra
88
RADIUM
(223)
CAESIUM
Fr
87
FRANCIUM
140.12
Ce
58
CERIUM
138.91
La
57
LANTHANIDE
LANTHANUM
238.03
U
92
URANIUM
231.04
Pa
91
PROTACTINIUM
232.04
Th
90
THORIUM
(227)
Ac
89
ACTINIDE
ACTINIUM
2
4
5
6
7
Editor: Aditya Vardhan ([email protected])
However three such elements (Th, Pa, and U)
do have a characteristic terrestrial isotopic
composition, and for these an atomic weight is
tabulated.
Relative atomic mass is shown with five
significant figures. For elements have no stable
nuclides, the value enclosed in brackets
indicates the mass number of the longest-lived
isotope of the element.
(1) Pure Appl. Chem., 73, No. 4, 667-683 (2001)
PERIOD
Question 1 (15 points):
Can an infrared photon break a typical covalent bond? Why or why not? (15 points)
Question 2 (25 points):
For the reaction of Si(s) + O2(g) → SiO2(s)
a) What are the signs of ΔH, ΔS, and ΔG? Explain each of your answers concisely.
(15 points)
[Hint: what would C do, in place of Si?]
b) Petrified wood comes from dehydration (loss of water) of silicic acid, Si(OH)4 (shown
below), ultimately to form quartz, SiO2 (after the silicic acid has replaced the organic
matter in the plant; the colors from metal ions that are left in the quartz crystals).
Show the first step in the polymerization reaction, starting with the reaction of two silicic
acid molecules, including showing the reaction products. (10 points)
Question 3 (10 points):
Recall from the third midterm that the monomer styrene (left) is used to make
polystyrene (middle), the main component of Styrofoam.
Last Wednesday, at the International Conference on Nanostructures on Kish Island, Iran,
in the talk before mine, Prof. Luis Liz-Marzán (right, one of the world’s leading
nanoscientists, the director of the nanocenter CIC biomaGUNE, San Sebastian, Spain,
who has been knighted by the King of Spain) showed that he could make crystalline
lattices of Au nanoparticles that were coated with polystyrene.
In each of the three electron microscope images below, all the 18-nm Au nanoparticles
(which appear as spots) are the same, other than their polystyrene coating.
Each polystyrene chain is chemically attached to the 18-nm Au nanoparticle at one end,
so each particle has the same number of attached chains. The only difference from top to
bottom in the images is the molecular weight (MW) of the polymer chains: 5800, 21,500
and 53,000 amu. Each spot in the electron microscopy image is an 18-nm Au
3
nanoparticle within a small crystal of nanoparticles; the polymer surrounding each
nanoparticle is not observed (because C and H do not interact with electrons as strongly
as Au does).
a) Why are the Au nanoparticle spacings different and in this order, in terms of the
attached polymer molecular weight? (Ignore the points in the plot for the different sizes
of nanoparticles.) (5 points)
b) Estimate the average polymer chain lengths in each case in terms of number of
monomer units. Recall that each carbon has four bonds and the H atoms are not shown in
the molecular structures above. Show your work. (5 points)
4
Question 4 (20 points, extra points are possible):
Fluorescence is a widely used quantitative method of detection.
For example, in flow cytometry, two different dyes can be used and separately
interrogated to label DNA in terms of the total amount of C-G vs A-T content in
individual chromosomes. This is shown schematically below.
a) Explain concisely how we can tell the two dyes apart. (10 points, + bonus points if you
correctly draw, use, and explain energy level diagrams for both
)
b) Explain concisely how we can tell how much of each type of base pair (C-G and A-T)
and dye is present. (5 points)
c) If a piece of a chromosome is missing, what happens to the two signals? (5 points)
Question 5 (30 points):
The half-life of 14C is approximately 5700 years.
The ratio of 14C/12C in the atmosphere is approximately 1.2×10-12.
If a tree was seeded and started to grow 17,100 (=3×5700) years ago and lived for 5700
years (ok, I know that is a long life for a tree, but nobody was carving their initials in it):
a) Draw and concisely explain a plot of the 14C/12C ratio in the tree for the last 17,100
years. (15 points)
b) For the first 5700 years since the tree was seeded and started to grow and for the time
since, give the functional form of the plot. (5 points)
c) What is the reaction order in [14C] for the 11,400 years since the tree died? Explain
concisely. (5 points)
d) What is the 14C/12C ratio in the (dead) tree now? Show your calculation. (5 points)
Extra credit #1 (10 points):
For question 2a, write balanced half-cell redox reactions that could be used to measure
the Ecell and ΔG for this reaction. Label the oxidation and reduction half-cell reactions.
5