1/21/2012
Chemistry Lab 2011-2012
Presenter:
John R Kiser
Hickory Regional Director
State Supervisor, Chemistry Lab
Introductions
• Topics for 2011-2012: Oxidation/Reduction and
Periodicity
• Regional vs State Topics
• Safety Requirements – Long Sleeve Shirt under lab
apron. Lab Coat would be easier.
• Must bring calculator!
Non-programmable, non-graphing
Need to know topics
• Formula Writing/Nomenclature
• Mole & Stoichiometry Calculations
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Oxidation Reduction (Redox)
Fundamental Concepts
• Oxidation is
Loss of electrons, gain of O, loss of H
• Reduction is
Gain of electrons, loss of O, gain of H
Mnemonic Devices:
LEO the lion goes GER!
OIL RIG
The species being reduced is the oxidizing agent
The species being oxidized is the reducing agent
Activity Series of Elements
• Usually lists the best reducing agent (most easily
oxidized) at top.
• Metal above reacts with ion below.
• Sample lab activity: The metal that reacts with
everything else goes at the top; the ion that reacts
with everything goes (as element, not ion) at bottom
Demo – Hollow Penny
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Determining if a species is oxidized or
reduced: Oxidation Numbers
Rules above take precedence over rules below!
• An atom in its elemental state has an oxidation number of 0.
Oxidation number of H in H2 = 0
• An atom in a monatomic ion has an oxidation number
identical to its charge.
Oxidation number of Fe3+ = +3
• The total sum of oxidation numbers in a polyatomic ion is
equal to the charge. Sum of oxidation numbers in a neutral
molecule is 0. Pay attention to subscripts!
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•
In a compound or polyatomic ion:
1A metal, ox #= +1
2A metal, ox # = +2
Al or B, ox# = +3
H, ox # = -1 if bonded to metal or B
H, ox # = +1 if bonded to another nonmetal
Oxygen, ox# = -2 (EXCEPT IN PEROXIDES, -1)
Fluorine, ox # = -1.
Other halogens (if written to right in formula), ox # = -1
Hypochlorite, can be written as OCl- OR ClO-, ox # of Cl is +1
• Other atoms not on this list can be deduced by following the
rules above
Example
Finding oxidation number of Cl in ClO4 –
Oxidation number of O = -2
Cl + 4*-2 = -1
Cl – 8 = -1 Cl = +7
If oxidation number increases, oxidation is
taking place
If oxidation number decreases, reduction is
taking place
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Redox Half Reactions
In many complex redox reactions, H+ (or OH-) and H2O
are involved in the reaction and may not be obvious
at first.
Therefore, there is a systematic method to balancing
complex redox reactions.
First, separate into two half reactions, one for
oxidation, one for reduction.
Total Reaction:
Cl- + Cr2O72- →
Cl2 + Cr3+
Half Reactions:
Cl- → Cl2
and Cr2O72- → Cr3+
Balancing Half Reactions
•
•
•
•
•
Using Coefficients, balance all atoms BUT H and O: Cr2O72- →2 Cr3+
Balance O by adding H2O
Cr2O72- → 2 Cr3+ + 7 H2O
Balance H by adding H+
14 H+ + Cr2O72- → 2 Cr3+ + 7 H2O
Balance charge by adding e- 14 H+ + Cr2O72- + 6e- →2Cr3+ + 7 H2O
Number of electrons should correspond to change in oxidation
number (keeping number of atoms in mind too!)
• No electrons in final answer, so multiply one (or both) half reaction(s)
by a number so that electrons are equal.
2 Cl- → Cl2 + 2 e2 Cl- → Cl2 + 2 ebecomes 6 Cl- → 3 Cl2 + 6 e-
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Adding Half Reactions Together
6 Cl→
14 H+ + Cr2O72- + 6e- →
3 Cl2 + 6 e2Cr3+ + 7 H2O
Be sure to cancel out electrons, water, and H+ that appears on
both sides
6 Cl- + 14 H+ + Cr2O72- → 3 Cl2 + 2 Cr3+ + 7 H2O
Be sure to double check charges and numbers of atoms!
Electrons should not be left over!
If you are balancing a half-reaction in basic solution, add OH- to
both sides, convert H+ to H2O, and cancel out.
A different approach that starts with oxidation numbers is given
as a pdf file.
Example Problem
Balance the following half reaction in basic solution:
Oxidation or reduction?
OCl - → Cl-
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Galvanic Cells
What happens when a redox reaction is spread out
over 2 beakers?
Oxidation at anode (-), reduction at cathode (+)
Electrons flow from anode to cathode
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Standard Reduction Potential (Eored)
These Eo values are technically Eoreduction. Eooxidation can
also be written for reverse reactions
O Standard
Conditions, 298 K, 1 atm, 1 M solutions
All these values are measured with respect to 2H+/H2
Eocell = Eocathode – Eoanode OR
Eocell = Eored + Eoox
Think of adding the two half reactions together.
Electrons must be equal, but values for Eo are
unchanged if a reaction is multiplied by a value
Shorthand Notation for Galvanic
Cells
Anode half-reaction:
Cathode half-reaction:
Overall cell reaction:
Zn2+(aq) + 2e-
Zn(s)
Cu2+(aq) + 2e-
Cu(s)
Zn(s) + Cu2+(aq)
Zn2+(aq) + Cu(s)
Salt bridge
Anode half-cell
Cathode half-cell
Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
Electron flow
Phase boundary
Phase boundary
Copyright © 2010 Pearson
Prentice Hall, Inc.
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Example Problem
• For the following Galvantic cell, determine the cathode,
anode, overall reaction, and Eocell
Ni (s) | Ni 2+ (aq) || Cu2+(aq) | Cu(s)
The Effect of Concentration on Cell E
What happens when conditions are not standard?
Nernst Equation:
E = E° -
RT
ln Q
nF
or
E = E° -
2.303RT
log Q
nF
or
E = E° -
0.0592 V
log Q
in volts, at 25°C
n
Q = Reaction Quotient (from Equilibrium topics)
Setup is just like an equilibrium constant, except the system is
not at equilibrium. Products on top, reactants on bottom.
Coefficients become exponents. Solutions in M, gases in atm,
pure liquids and solids omitted.
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Consider a galvanic cell that uses the reaction:
Cu(s) + 2Fe3+(aq)
Cu2+(aq) + 2Fe2+(aq)
What is the potential of a cell at 25 °C that has the following ion
concentrations?
[Fe3+] = 1.0 x 10-4 M
[Cu2+] = 0.25 M
[Fe2+] = 0.20 M
Batteries
Lead Storage Battery
Anode:
Pb(s) + HSO4-(aq)
Cathode:
PbO2(s) + 3H+(aq) + HSO4-(aq) + 2e-
Overall:
Pb(s) + PbO2(s) + 2H+(aq) + 2HSO4-(aq)
Copyright © 2010
Pearson Prentice Hall,
Inc.
PbSO4(s) + H+(aq) + 2ePbSO4(s) + 2H2O(l)
2PbSO4(s) + 2H2O(l)
Chapter 17/20
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Batteries
Dry-Cell Batteries
Leclanché cell
Copyright © 2010
Pearson Prentice Hall,
Inc.
Chapter 17/21
Fuel Cells
Hydrogen-Oxygen Fuel Cell
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Electrolysis and Electrolytic Cells
Electrolysis: The process of using an electric current to bring about chemical
change.
Electrolysis of Molten Sodium Chloride
Anode:
2Cl-(l)
Cathode:
Overall:
2Na+(l) + 2e-
2Na+(l) + 2Cl-(l)
Cl2(g) + 2e2Na(l)
2Na(l) + Cl2(g)
Copyright © 2010
Pearson Prentice Hall,
Inc.
Electrolysis and Electrolytic Cells
Cathode has negative charge, connected to negative terminal on
battery; anode has positive charge, connected to positive terminal.
Reduction still at cathode, Oxidation still at anode
Electrolysis of Molten Sodium Chloride
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Electrolysis and Electrolytic Cells
Electrolysis of Water
Anode:
2H2O(l)
O2(g) + 4H+(aq) + 4e-
Cathode: 4H2O(l) + 4e-
2H2(g) + 4OH-(aq)
Overall:
6H2O(l)
2H2(g) + O2(g) + 4H+ (aq)+ 4OH-(aq)
6H2O(l)
2H2(g) + O2(g) + 4H2O (l)
2H2O(l)
2H2(g) + O2(g)
Electrolysis of Aqueous Solutions
Electrolysis of NaI
• Consider reduction of water and reduction of cation
Na+ (aq) + e- → Na (s)
Eored = - 2.71 V
2 H2O (l) + 2e- → H2 (g) + 2 OH- (aq) Eored = - 0.83 V
Less negative (or more positive) reduction is preferred!
• Consider oxidation of water and oxidation of anion
2 H2O (l) → O2 (g) + 4 H+ (aq) + 4 e- Eo ox = -1.23 V
2 I- (aq) → I2 (s) + 2 eEoox = -0.54 V
Less negative (more positive) oxidation preferred!
(Note: reversing from Eored table)
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• Overvoltage makes predictions difficult if both
competing reactions have similar Eo values.
• In electrolysis of aqueous NaCl, even though
oxidation of water is less negative, Cl – is oxidized
to Cl2 gas.
• Poorly understood, mainly due to kinetic factors
and concentration effects.
• Prediction is difficult if values are close on table.
• Rules of Thumb for electrolysis of solutions
Look on left side of SRP table. If cation is higher
than water, cation will be reduced.
Look to right side of SRP table. If anion is below
water, anion will be reduced
Quantitative Aspects of
Electrolysis
Charge(C) = Current(A) x Time(s)
1 mol eMoles of
e-
= Charge(A) x
96,500 C
Faraday constant
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Example Problem
How much Ca will be produced in an electrolytic cell of molten CaCl2
if a current of 0.452 A is passed through the cell for 1.5 hours?
19.8
Periodicity
• How (and why) do many chemical and
physical properties of elements relate to
position on periodic table?
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Electron Configuration
Orbital Diagrams
Rules of the Game..
• Building Up Principle – Start at lowest energy
level and sublevel, work your way up. An atom
“builds on” the previous atom before it.
• Pauli Exclusion Principle – No more than 2
electrons in an orbital, one is spin up, the other
spin down.
• Hund’s Rule – If 2 or more orbitals have the same
energy, one electron goes in each until all are half
full. Then, electrons go back and pair up.
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Examples
Outer Shell Configuration of Elements
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Core vs. Valence Electrons
• Valence Electrons = In main group or
representative elements, valence electrons are in
the outmost energy level
• In examples below, both Na and Mg have 10 core
electrons. Na has 1 valence electron, Mg has 2
valence electrons.
• For representative element, column # before A
gives number of valence electrons
Effective Nuclear Charge - Zeff
• What positive charge do valence electrons
“feel” from nucleus?
• Zeff increases from left to right across a
row – number of protons in nucleus
increase
• Zeff slightly increases from top to bottom
down a column – more diffuse inner
electrons slightly reduce shielding.
• Zeff ≈ Column number before A
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Atomic Radius
Ionic Radius
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Electronegativity
• Electronegativity - Ability of an atom in a
molecule to attract shared electrons in a
covalent bond.
Ion Formation
Main Group Metals
Main group metals typically lose all valence electrons when
an ion forms. This allows metal ion to be isoelectric to a
noble gas and have an octet (8 valence electrons)
Na atom: 1s22s22p63s1 - 1 valence electron
When ion forms, 1 valence electron is lost
Ion: 1s22s22p6 8 valence electrons, isoelectric to Ne
Because 1 electron is lost, charge of ion is 1+ (Na+)
Prediction: Column number before “A” is charge on ion.
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Ion Formation
Transition Metals
• While electrons fill first in ns sublevel and then
(n-1)d sublevel, a transition metal usually loses
the ns electrons first when an ion forms, THEN
the (n-1)d electrons.
Fe atom: [Ar]4s23d6
Fe2+ ion: [Ar]4s03d6 or [Ar]3d6
Fe3+ ion: [Ar]4s03d5 or [Ar]3d5
Ion Formation Nonmetals
When a nonmetal atom forms an ion, electrons are gained to
complete an octet in outer energy level.
Chlorine Atom: [Ne]3s23p5 7 valence electrons
When ion forms, 1 more electron is gained to complete octet
Ion: [Ne]3s23p6 – 8 valence electrons – octet!
Ion is isoelectric to Argon.
Since 1 electron is gained, charge is 1- (Cl-)
Rule of Thumb: 8-Column number before A = charge on ion
OR, count number of columns from noble gases.
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Summary of Ion Formation
Ionization Energy
• The amount of energy needed to remove the highest
energy electron from an isolated neutral atom in the gas
state.
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Why Deviations in Ionization
Energy?
Successive Ionization Energies
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Electron Affinity
• Energy change that occurs when an electron is added to
an isolated atom in the gaseous state.
X + e- → X∆E is < 0 kJ, energy released
Predicting Melting Point
• For metals, mp will decrease down a column in the
periodic table.
• For nonmetals, mp will usually increase down a column in
the periodic table.
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Example – Predicting
Melting Point of Fr
Chang, Chemistry, 9th edition, McGraw Hill
Selected Reactivity Examples
1) Metal + Nonmetal → Ionic compound
Be sure to use ionic charges to write formula of
product
Then, go back and balance equation.
Al (s) + O2 (g) → Al2O3 (s)
Al3+ and O24 Al (s) + 3 O2 (g) → 2 Al2O3 (s)
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2) Reaction of metals with water
Reactivity usually decreases from left to right
across a row, increases down a column
• 3) Oxide reactivity
Most metallic oxides react with water to produce a
base (hydroxide)
Li2O (s) + H2O (l) → 2 LiOH (aq)
Most nonmetallic oxides react with water to
produce an acid.
SO3 (g) + H2O (l) → H2SO4 (aq)
Aluminum oxide can produce an acidic or basic
solution!
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Solubility of
Silver Salts Lab Demo
Thank you!
Please email me at jkiser@wpcc.edu
For follow-up questions, concerns, etc..
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