Transcript Chapter Eighteen

1
Chapter Eighteen
Electrochemistry
Prentice Hall © 2005
General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Eighteen
2
Today…
• Turn in:
– Nothing
• Our Plan:
– Test Results
– Notes – Redox Equations
– Worksheet #1
• Homework (Write in Planner):
– Redox Equation WS due Friday
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General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Eighteen
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6% Curve
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Average
High Score
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4
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90% (x2)
General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Eighteen
4
Oxidation–Reduction:
The Transfer of Electrons
Silver metal is formed,
and the solution turns
blue from copper(II)
ions formed.
Electrons from
copper metal are
transferred to
silver ions.
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Chapter Eighteen
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Half-Reactions (Review)
• In any oxidation–reduction reaction, there are two halfreactions:
– Oxidation: a species loses electrons to another species.
(LEO)
– Reduction: a species gains electrons from another
species. (GER)
• Both oxidation and reduction must occur
simultaneously.
– A species that loses electrons must lose them to
something else (something that gains them).
– A species that gains electrons must gain them from
something else (something that loses them).
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Chapter Eighteen
6
Oxidation Numbers
• An oxidation number is the charge on an ion, or a
hypothetical charge assigned to an atom in a
molecule or polyatomic ion.
• Examples: in NaCl, the oxidation number of Na is
+1, that of Cl is –1 (the actual charge).
• In CO2 (a molecular compound, no ions) the
oxidation number of oxygen is –2, because oxygen
as an ion would be expected to have a 2– charge.
• What is the oxidation number on carbon in CO2?
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Chapter Eighteen
Rules for Assigning Oxidation Numbers
7
1. For the atoms in a neutral species—an isolated atom, a molecule,
or a formula unit—the sum of all the oxidation numbers is 0.
– This includes elements in their standard state (Cu (s), Cl2 (g), etc.)
2. For the atoms in an ion, the sum of the oxidation numbers is equal
to the charge on the ion.
3. In compounds, the group 1A metals all have an oxidation number
of +1 and the group 2A metals all have an oxidation number of +2.
4. In compounds, the oxidation number of fluorine is –1.
5. In compounds, hydrogen has an oxidation number of +1.
6. In most compounds, oxygen has an oxidation number of –2.
7. In binary compounds with metals, group 7A elements have an
oxidation number of –1, group 6A elements have an oxidation
number of –2, and group 5A elements have an oxidation number
of –3.
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Chapter Eighteen
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Let’s Try It…
• Assign oxidation numbers to
the elements in each compound:
-1
1. NO3
2. SO2
3. Fe2O3
4. Cu (s)
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Chapter Eighteen
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Identifying Oxidation–Reduction
Reactions
• In a redox reaction, the oxidation number of a species
changes during the reaction.
• Oxidation occurs when the oxidation number
increases (species loses electrons). LEO
• Reduction occurs when the oxidation number
decreases (species gains electrons). GER
• Another mnemonic is OIL RIG!
• If any species is oxidized or reduced in a reaction, that
reaction is a redox reaction.
• Examples of redox reactions: displacement of an
element by another element; combustion;
incorporation of an element into a compound, etc.
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Chapter Eighteen
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Review
• In the reactions below, write oxidation numbers
for each substance and identify which substance is
oxidized and which is reduced?
• N2O + H2
-->
H2O
+
NH3
•
K +
KNO3 --> N2
• Fe2O3 + S
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+
K2O
--> Fe + SO2
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Chapter Eighteen
11
The Half-Reaction Method of
Balancing Redox Equations
1.
2.
3.
4.
5.
6.
Separate a redox equation into two half-equations, one
for oxidation and one for reduction.
Balance the number of atoms of each element in each
half-equation. Usually we balance O and H atoms last.
Balance each half-reaction for charge by adding
electrons to the left in the reduction half-equation and to
the right in the oxidation half-equation.
Adjust the coefficients in the half-equations so that the
same number of electrons appears in each half-equation.
Add together the two adjusted half-equations to obtain
an overall redox equation.
Simplify the overall redox equation as necessary.
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Chapter Eighteen
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Redox Reactions in Acidic and in
Basic Solution
• Redox reactions in acidic solution and in basic solution
may be very different from one another.
• If acidic solution is specified, we must add H2O and/or H+
as needed when we balance the number of atoms.
• If basic solution is specified, the final equation may have
OH– and/or water molecules in it.
• A simple way to balance an equation in basic solution:
– Balance the equation as though it were in acidic solution.
– Add as many OH– ions to each side as there are H+ ions in the
equation.
– Combine the H+ and OH– ions to give water molecules on one side,
and simplify the equation as necessary.
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Chapter Eighteen
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Example 18.1
• Balance the equation in acidic solution:
MnO4-1 + S2O3-2 --> Mn+2 + SO4-2
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Example 18.1 B
• Write a balanced equation for the oxidation
of phosphorus by nitric acid, which is
described by
P4 (s) + H+1 (aq) + NO3-1 (aq) --> H2PO4-1 (aq) + NO (g)
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Chapter Eighteen
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Example 18.2
• Balance the following in basic solution:
Br2 (l) --> Br-1 (aq) + BrO3-1 (aq)
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Chapter Eighteen
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Example 30a
Balance the following in basic solution:
CrO4-2 (aq) + AsH3 (g) --> Cr(OH)3 (s) + As (s)
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Chapter Eighteen
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WS Example 1
HCl + K2Cr2O7 --> KCl + CrCl3 + H2O + Cl2
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WS Example 2
FeCl2 + KMnO4 + HCl --> FeCl3 + KCl + MnCl2 + H2O
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WS Example 3
S-2 + MnO4-1 --> S + MnO2 (basic solution)
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WS Example 4
CuS + NO3-1 --> Cu+2 + S + NO (acidic)
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Chapter Eighteen
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Stop!
• Do the Redox Equations Worksheet.
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Chapter Eighteen
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Today…
• Turn in:
– Get out Redox WS
• Our Plan:
– Questions on Redox WS
– Quiz
– Notes
– Galvanic Cells WS
• Homework (Write in Planner):
– Worksheet Due Wednesday
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Chapter Eighteen
23
A Qualitative Description of
Voltaic Cells
• A voltaic cell uses a spontaneous redox reaction to produce
electricity.
• A half-cell consists of an electrode (strip of metal or other
conductor) immersed in a solution of ions.
This Zn2+ becomes
a Zn atom.
Both oxidation and
reduction occur at the
electrode surface, and
equilibrium is reached.
This Zn atom leaves the
surface to become a Zn2+ ion.
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Chapter Eighteen
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Important Electrochemical Terms
• An electrochemical cell consists of two half-cells
with the appropriate connections between
electrodes and solutions.
• Two half-cells may be joined by a salt bridge (U
shaped tube suspended in gel) that permits
migration of ions, without completely mixing the
solutions.
• The anode is the electrode at which oxidation
occurs.
• The cathode is the electrode at which reduction
occurs.
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Chapter Eighteen
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Helpful Mnemonic
• AUTO (anode = oxidation)
• CAR (cathode = reduction)
OR:
• To remember the charge: Ca+ions are
attracted to the Ca+hode (the t is a plus
sign)
• To remember which reaction occurs at
which terminal: An Ox and Red Cat Anode Oxidation, Reduction Cathode
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Chapter Eighteen
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Important Electrochemical Terms
• In a voltaic cell, a spontaneous redox
reaction occurs and current (electricity) is
generated.
• Cell potential (Ecell) is the potential
difference in volts between anode and
cathode.
• Ecell is the driving force that moves
electrons and ions.
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Chapter Eighteen
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A Zinc–Copper
Voltaic Cell
Positive and negative
ions move through the
salt bridge to equalize
the charge.
… the electrons
produced move
through the wire …
Zn(s) is oxidized to
Zn2+ ions, and …
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… to the Cu(s) electrode,
where they are accepted
by Cu2+ ions to form
more Cu(s).
Reaction: Zn(s) + Cu2+(aq) --> Cu(s) + Zn2+(aq)
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Chapter Eighteen
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Cell Diagrams
•
•
•
•
A cell diagram is “shorthand” for an electrochemical cell.
The anode is placed on the left side of the diagram.
The cathode is placed on the right side.
A single vertical line ( | ) represents a boundary between
phases, such as between an electrode and a solution.
• A double vertical line ( || ) represents a salt bridge or
porous barrier separating two half-cells.
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Chapter Eighteen
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Reaction: Zn(s) + Cu2+(aq) --> Cu(s) + Zn2+(aq)
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Chapter Eighteen
Another look at Cell
Diagrams
30
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Chapter Eighteen
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Example 18.3
Describe the half-reactions and the overall reaction that
occur in the voltaic cell represented by the cell diagram:
Pt(s) | Fe2+(aq), Fe3+(aq) || Cl–(aq) | Cl2(g) | Pt(s)
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Chapter Eighteen
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Standard Electrode Potentials
• Since an electrode represents
only a half-reaction, it is not
possible to measure the
absolute potential of an
electrode.
• The standard hydrogen
electrode (SHE) provides a
reference for measurement of
other electrode potentials.
• The SHE is arbitrarily assigned
a potential of 0.000 V.
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Standard Electrode Potentials
• The standard electrode potential, E°, is based on
the tendency for reduction to occur at an
electrode. (Sometimes it is referred to as standard
reduction potential)
• All standard reduction potentials are for 25 ᵒC and
1 M solutions.
• E° for the standard hydrogen electrode is
arbitrarily assigned a value of 0.000 V.
• All other values of E° are determined relative to
the standard hydrogen electrode.
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Chapter Eighteen
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Standard Electrode Potentials
• The standard cell potential (E°cell)
is the difference between E° of the
cathode and E° of the anode.
• E°cell = E°(cathode) – E°(anode)
• Remember AUTO CAR!
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Chapter Eighteen
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Measuring the Standard Potential
of the Cu2+/Cu Electrode
The voltmeter reading
and the direction of
electron flow tell us
that …
… Cu2+ is more easily
reduced than H+, by
0.340 volts.
Standard
hydrogen
electrode
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Cu2+ + 2e --> Cu
E° = +0.340 V
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Chapter Eighteen
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Measuring the Standard Potential
of the Zn2+/Zn Electrode
The voltmeter reading
and the direction of
electron flow tell us
that …
… Zn2+ is harder to
reduce than H+, by
0.763 volts.
Standard
hydrogen
electrode
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Zn2+ + 2e --> Zn
E° = – 0.763 V
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Chapter Eighteen
F2 is the strongest
oxidizing agent
F– is the weakest
reducing agent
37
Li is the
strongest
reducing
agent
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Chapter Eighteen
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Important Note
• The formula
E°cell = E°(cathode) – E°(anode)
allows us to simply use the
numbers in the table, but if you
were asked for the oxidation value
of a reaction, you would have to
reverse the sign since the values
are all reduction potentials.
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Chapter Eighteen
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Important Points about
Electrode and Cell Potentials
• Standard electrode potentials and cell voltages are
intensive properties; they do not depend on the
total amounts of the species present.
• A “flashlight battery” (D-cell) and a “penlight
battery” (AA cell) produce the same potential—
1.5 volts.
• E° does depend on the particular species in the
reaction (or half-reaction).
• As we shall learn later, cell and electrode
potentials can depend on concentration of the
species present.
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Chapter Eighteen
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Example 18.4
Determine E° for the reduction half-reaction
Ce4+(aq) + e– --> Ce3+(aq), given that the cell voltage for the
voltaic cell
Co(s) | Co2+(1 M) || Ce4+(1 M), Ce3+(1 M) | Pt(s)
is E°cell = 1.887 V.
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Chapter Eighteen
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Example 18.5
Balance the following oxidation–reduction equation, and
determine E°cell for the reaction.
O2(g) + H+(aq) + I–(aq) --> H2O(l) + I2(s)
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Electrode Potentials, Spontaneous
Change, and Equilibrium
• An electrochemical cell does work.
welec = nFEcell
n = number of electrons in the balanced equation
F = 96,485 coulombs per mole.
• The amount of electrical work is also equal to –DG:
DG = –nFEcell
• Under standard conditions:
DG° = –nFE°cell
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Chapter Eighteen
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Criteria for Spontaneous
Change in Redox Reactions
• If Ecell is positive, the forward reaction is
spontaneous.
• If Ecell is negative, the forward reaction is
nonspontaneous (the reverse reaction is
spontaneous).
• If Ecell = 0, the system is at equilibrium.
• When a cell reaction is reversed, Ecell and
DG change signs.
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Chapter Eighteen
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Example 18.6
Will copper metal displace silver ion from aqueous
solution? That is, does the reaction
Cu(s) + 2 Ag+(1 M) --> Cu2+(1 M) + 2 Ag(s)
occur spontaneously from left to right?
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Chapter Eighteen
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The Activity Series Revisited
• In the activity series of metals
(Section 4.4), any metal in the
series will displace a metal
below it from a solution of that
metal’s ions.
• Theoretical basis: The activity
series lists metals in order of
their standard potentials.
• Displacement of a metal from a
solution of its ions by a metal
higher in the series corresponds
to a positive value of Ecell and a
spontaneous reaction.
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Chapter Eighteen
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Visual Example
• http://intro.chem.okstate.edu/1515F01/Labo
ratory/ActivityofMetals/home.html
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Chapter Eighteen
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Equilibrium Constants in Redox
Reactions
• Whereas potential and free energy are related, and free
energy and equilibrium are related, equilibrium and
potential must be related to one another.
DG° = –nFE°cell
and
DG° = –RT ln Keq
therefore
–RT ln Keq = –nFEocell
RT ln Keq
RT
E°cell = ––––––––– = –––– ln Keq
nF
nF
R and F are constant,
therefore at 298 K:
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0.025693 V
E°cell = –––––––– ln Keq
n
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Chapter Eighteen
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Example 18.8
Calculate the values of ΔG° and Keq at 25 °C for the reaction
Cu(s) + 2 Ag+(1 M) --> Cu2+(1 M) + 2 Ag(s)
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Chapter Eighteen
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Thermodynamics, Equilibrium, and
Electrochemistry: A Summary
From any one of the three
quantities Keq, ΔG°, E°cell, we
can determine the others.
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Chapter Eighteen
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Stop!
• Begin working on Worksheet #2!
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Chapter Eighteen
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Today…
• Turn in:
– Nothing
• Our Plan:
– Quick Review
– Notes – Concentration & Electrochemical
Cells
– Finish Electrochemical Cells WS
• Homework (Write in Planner):
– Worksheet Due Wednesday
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Chapter Eighteen
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52
Effect of Concentrations on Cell
Voltage
• A nonstandard cell differs in potential from a standard cell
(1 M concentrations, 1 atm partial pressures).
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Chapter Eighteen
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Effect of Concentrations on Cell
Voltage
From the previous relationships
we can show that:
At 25 °C, and converting to
common logarithms:
Ecell = E°cell
RT
– –––– ln Q
nF
Ecell = E°cell
0.0592 V
– ––––––– log Q
n
• This Nernst equation relates a cell voltage for nonstandard
conditions, Ecell, to a standard cell voltage, E°cell, and to the
concentrations of reactants and products expressed through
the reaction quotient, Q.
• We can use the Nernst equation to find cell potential from
concentrations, or we can measure Ecell and determine the
concentration of a species in the cell.
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Chapter Eighteen
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Example 18.9
Calculate the expected voltmeter reading for the voltaic
cell pictured in Figure 18.13.
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Chapter Eighteen
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Another Nernst Equation Example
18.9 A
• Use the Nernst equation to determine Ecell at
25ᵒC for the following voltaic cells.
a) Zn(s) | Zn2+(2.0 M) || Cu2+(0.050 M) | Cu (s)
b) Zn(s) | Zn2+(0.050 M) || Cu2+(2.0 M) | Cu (s)
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Batteries: Using Chemical
Reactions to Make Electricity
• We often call any device that stores chemical
energy for later release as electricity a battery.
• Technically, a D, C, or AA “battery” is actually a
single electrochemical cell.
• A battery consists of several cells joined together
to produce higher current or higher voltage.
• A 9-volt “transistor” battery, an automobile
battery, and a rechargable “battery pack” are all
true batteries.
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Chapter Eighteen
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The Dry Cell
• A primary cell employs an
irreversible chemical reaction.
• When the reactants inside the
cell are largely used up, the cell
is “dead.”
• The LeClanché cell or dry cell
(right) is the “ordinary” type of
flashlight “battery.”
• Alkaline cells cost more than
the LeClanché cell but they
have a longer shelf life and
longer service life.
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Chapter Eighteen
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The Lead–Acid Storage Battery
• A lead–acid storage battery
used in an automobile uses
secondary cells; they are
rechargeable.
• By connecting the cell to an
external electric energy
source, the discharge
reaction is reversed.
Cell reaction:
Pb(s) + PbO2(s) + 2 H+(aq) + 2 HSO4–(aq) --> 2 PbSO4(s) + 2 H2O(l)
Charging reaction: 2 PbSO4(s) + 2 H2O(l) --> Pb(s) + PbO2(s) + 2 H+(aq) + 2 HSO4–(aq)
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Other Secondary Cells
• The nickel–cadmium (NiCd) cell uses a cadmium
anode and a cathode containing Ni(OH)2.
• A NiCd cell can be recharged hundreds of times. It
produces 1.2 V (a Leclanché cell produces 1.5 V).
• Nickel–metal hydride cells (NiMH) use a metal alloy
anode that contains hydrogen.
• In use, the anode releases the hydrogen, forming
water. Like the NiCd cell, a NiMH cell produces 1.2 V.
• Lithium-ion cells use a lithium–cobalt oxide or lithium–
manganese oxide material as the anode. The
electrolyte is an organic solvent containing a dissolved
lithium salt.
• Many modern laptop computers and cellular phones
use lithium-ion cells.
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Chapter Eighteen
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Fuel Cells
• In a fuel cell, the cell reaction is
equivalent to a combustion reaction.
• The reactants are supplied
externally; the cell does not “go
dead” as long as the oxidizing
and reducing agents are provided.
• Fuel cells are generally operated
under nonstandard conditions and
at temperatures considerably
higher than 25 °C.
• H2–O2 fuel cells are seeing use in
some automobiles.
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Chapter Eighteen
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Corrosion: Metal Loss
Through Voltaic Cells
• In moist air, iron is oxidized to Fe2+, particularly at
scratches, nicks, or dents. These areas are referred to as
anodic areas.
• Other regions of the iron serve as cathodic areas, where
the electrons from the anodic areas reduce O2 to OH–.
• Iron(II) ions migrate from the anodic areas to the cathodic
areas where they combine with the hydroxide ions.
• The iron(II) is then further oxidized to iron(III) by
atmospheric oxygen.
• Common rust is Fe2O3 · x H2O.
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Corrosion of an Iron Piling
One way to minimize
rusting is to provide a
different anode reaction.
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Protecting Iron from Corrosion
• The simplest defense against corrosion of iron is to coat it
(with paint or metal) to exclude oxygen from the surface.
• An entirely different approach is to protect iron with a
more active metal.
• Galvanized iron has been coated with zinc.
• The zinc provides an alternative anode reaction; the zinc
corrodes, protecting the iron.
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Cathodic Protection
• In cathodic protection, an iron object to be protected is
connected to a chunk of an active metal.
• The iron serves as the
reduction electrode and
remains metallic. The
active metal is oxidized.
• Water heaters often
employ a magnesium
anode for cathodic
protection.
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Stop!
• Finish Worksheet #2 –
Electrochemical Cells
Worksheet
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Today…
• Turn in:
– Get out Electrochemical Cells WS
• Our Plan:
– Questions on WS
– Electrochemical Cell Review
– Notes
– Electrolysis WS
• Homework (Write in Planner):
– Worksheet Due Friday (LAST HOMEWORK)
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Chapter Eighteen
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Electrolytic Cells
• A voltaic cell corresponds to a spontaneous cell reaction.
• An electrolytic cell corresponds to a nonspontaneous cell
reaction. The reaction is called electrolysis.
• An electrolytic cell is the opposite of a galvanic cell.
• The external source of electricity acts like an “electron
pump.” It pulls electrons away from the anode, where
oxidation takes place, and pushes them toward the
cathode, where reduction takes place.
• The polarities of the electrodes are reversed from those in
the voltaic cell, because now the external source controls
the flow of electrons.
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Chapter Eighteen
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Electrolysis of Molten
Sodium Chloride
2 NaCl(l) --> 2 Na(l) + Cl2(g)
The nonspontaneous
reaction is driven by
external potential.
Molten NaCl
(around 1000 °C)
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Chapter Eighteen
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Predicting Electrolysis Reactions
• In an electrolytic cell, all combinations of cathode and
anode half-reactions give negative values of E°cell.
• The reaction most likely to occur is the one with the least
negative value of E°cell (requires the lowest applied voltage
from the external electricity source). HOWEVER …
• In many half-reactions, particularly those involving gases,
various interactions at electrode surfaces make the required
voltage for electrolysis higher than the voltage calculated
from E° data.
• Overvoltage is the excess voltage above the voltage
calculated from E° values that is required in electrolysis.
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Chapter Eighteen
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Quantitative Electrolysis
• The unit of electric charge is the
coulomb (C), and the charge carried
by one electron is –1.6022 x 10–19 C.
– The conversion factor we’ll use is
96,485 C/1 mole electrons
• Electric current, expressed in
amperes (A), is the rate of flow of
electric charge (C/s).
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General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Eighteen
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Quantitative Electrolysis
• To calculate the quantitative outcome of an
electrolysis reaction:
1. Determine the amount of charge (C)—the
product of current and time.
2. Convert the amount of charge to moles of
electrons.
3. Use a half-equation to relate moles of
electrons to moles of a reactant or a product.
4. Convert from moles of reactant or product to
the final quantity desired.
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General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Eighteen
72
Example 18.13
We can use electrolysis to determine the gold content of a sample.
The sample is dissolved, and all the gold is converted to Au3+(aq),
which is then reduced back to Au(s) on an electrode of known
mass. The reduction half-reaction is Au3+(aq) + 3e–  Au(s).
What mass of gold will be deposited at the cathode in 1.00 hour
by a current of 1.50 A?
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General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Eighteen
73
Try It Out!
• 18.13 A: For how many minutes must the
electrolysis of a solution of CuSO4 (aq) be
carried out, at a current of 2.25 A, to deposit
1.00 g of Cu (s) at the cathode?
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General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Eighteen
74
Example 18.14
An Estimation Example
Without doing detailed calculations, determine which of the
following solutions will yield the greatest mass of metal at a
platinum cathode during electrolysis by a 1.50-A electric
current for 30.2 min: CuSO4(aq), AgNO3(aq), or AuCl3(aq).
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General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Eighteen
75
Producing Chemicals
by Electrolysis
Electrolysis plays an
important role in the
manufacture and purification
of many substances, including
chlorine, copper, silver,
magnesium, aluminum, lead,
zinc, sodium, fluorine,
titanium, sodium hydroxide,
hydrogen …
Electrolysis of NaCl(aq) is
used to produce H2, Cl2, and
NaOH, all of which have
important industrial uses.
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Chapter Eighteen
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Electroplating
• Electrolysis can be used to coat
one metal onto another, a process
called electroplating.
• Usually, the object to be
electroplated, such as a spoon, is
cast of an inexpensive metal. It is
then coated with a thin layer of a
more attractive, corrosionresistant, and expensive metal,
such as silver or gold.
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General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Eighteen
77
Stop
•Complete the
Electrolysis
Worksheet
Prentice Hall © 2005
General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Eighteen
78
Today…
• Turn in:
– Get out Electrolysis WS
• Our Plan:
– Questions on WS
– Electrochemical Cells Online Lab
• Homework (Write in Planner):
– Take Home Test Due Monday OR
prepare for the AP Exam!
Prentice Hall © 2005
General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Eighteen
79
A Little Redox Humor
Prentice Hall © 2005
General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Eighteen
80
Today…
• Turn in:
– Take Home Test
• Our Plan:
– Begin Qualitative Analysis Lab (read
the pre-lab and procedure FIRST)
• Homework (Write in Planner):
– Work on Lab Report
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General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Eighteen
81
Lab Preparation
• Read the Packet, including the entire
procedure.
• Answer the Pre-Lab Questions
• Begin working on your lab report
– I recommend recording data in the handout that
I gave you and then transferring it to your final
copy.
– You could begin working on the title, purpose,
pre-lab, and procedure, though.
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General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Eighteen
82
Reminders
• Safety – Goggles, gloves, apron at ALL TIMES!
• Clean all glassware and lab supplies before and after
use.
• Label things well (include your initials)!
• Seal containers to store until next class period. Place
them on the back bookshelf.
• All solids go in the trash and liquids down the drain
EXCEPT barium (anion step 2 &3). Put them in the
waste container in the hood!
• Run tests until you get clear and reproducible results
(you can repeat steps if you need to).
• Report is due 5/14 at the beginning of class!
Prentice Hall © 2005
General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Eighteen
83
Today…
• Turn in:
– Nothing
• Our Plan:
– Qualitative Analysis Lab
• Homework (Write in Planner):
– Report Due on 8/14
Prentice Hall © 2005
General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Eighteen