SCH4U · Unit 5 test review

Electrochemistry: Test Review

Everything for the Unit 5 test on one page: the key facts from all 4 lessons, the mistakes that cost the most marks, and a practice test that points you back to the lessons you need.

Key facts

Oxidation Numbers and Redox Reactions

Oxidation is loss of electrons (oxidation number goes up); reduction is gain of electrons (oxidation number goes down).
The oxidizing agent is reduced; the reducing agent is oxidized.
Sum of oxidation numbers = 0 for a neutral compound, or = the charge for a polyatomic ion.
Usual values: O is −2 (−1 in peroxides), H is +1 (−1 in metal hydrides), F is −1, uncombined elements are 0.
In a balanced redox equation, total electrons lost = total electrons gained.

Balancing Redox with Half-Reactions

Order in acidic solution: other atoms, then O with $H_{2} O$ , then H with $H^{+}$ , then charge with $e^{-}$ .
Oxidation half-reaction: electrons are products. Reduction half-reaction: electrons are reactants.
Multiply the half-reactions so the electrons lost equal the electrons gained, then add and cancel.
In basic solution, add $OH^{-}$ to both sides to match each $H^{+}$ , form water, and cancel.
A correctly balanced redox equation balances in both atoms and total charge.

Galvanic Cells and Cell Potential

Anode: oxidation (negative electrode in a galvanic cell). Cathode: reduction (positive electrode).
Electrons flow through the wire from anode to cathode; in the salt bridge, anions move toward the anode and cations toward the cathode.
$E_{cell}^{\circ} = E_{cathode}^{\circ} - E_{anode}^{\circ},$ using reduction potentials for both.
$E_{cell}^{\circ} > 0$ : spontaneous as written. $E_{cell}^{\circ} < 0$ : not spontaneous (the reverse is).
Standard hydrogen electrode: $2 H^{+} + 2 e^{-} \to H_{2}$ , $E^{\circ} = 0.00 V$ by definition.
Cell notation: anode $∣$ anode solution $∥$ cathode solution $∣$ cathode.

Electrolytic Cells, Corrosion and Batteries

Electrolytic cell: electrical energy drives a non-spontaneous reaction ( $E_{cell}^{\circ} < 0$ ). Anode is positive, cathode is negative, but oxidation is still at the anode.
Minimum voltage needed for electrolysis is at least the size of the negative $E_{cell}^{\circ}$ (e.g. about 4.07 V for molten NaCl).
In electroplating, the object to be plated is the cathode and the plating metal is the anode.
Corrosion of iron: $Fe \to Fe^{2 +} + 2 e^{-}$ and $O_{2} + 2 H_{2} O + 4 e^{-} \to 4 OH^{-}$ ; needs both water and oxygen.
Galvanizing and sacrificial anodes protect iron by giving it a more active metal to be oxidized instead.
$Q = I t$ and $n (e^{-}) = \frac{I t}{F}$ , where $F \approx 96485 C/mol$ .

Where marks get lost

Mixing up the agents: saying the oxidizing agent is the substance oxidized. The oxidizing agent takes electrons, so it is the one that gets reduced. (Oxidation Numbers and Redox Reactions)
Forgetting to multiply by the number of atoms, e.g. solving $x + 7 (- 2) = - 2$ for dichromate instead of $2 x + 7 (- 2) = - 2$ , which gives Cr = +12 instead of +6. (Oxidation Numbers and Redox Reactions)
Assuming oxygen is always −2 and missing peroxides such as $H_{2} O_{2}$ , where O is −1. (Oxidation Numbers and Redox Reactions)
Treating an oxidation number like a real ionic charge and writing S in $SO_{4}^{2 -}$ as a $S^{6 +}$ ion. (Oxidation Numbers and Redox Reactions)
Balancing electrons per atom but forgetting that a molecule like $O_{2}$ or $Cr_{2} O_{7}^{2 -}$ contains two atoms that each change. (Oxidation Numbers and Redox Reactions)
Adding electrons to the wrong side. Check by charge: the electrons always go on the more positive side of the half-reaction. (Balancing Redox with Half-Reactions)
Balancing O with $O_{2}$ or with free O atoms instead of with $H_{2} O$ . (Balancing Redox with Half-Reactions)
Stopping once the atoms balance and never checking total charge, which is the only way to catch a wrong number of electrons. (Balancing Redox with Half-Reactions)
In basic solution, adding $OH^{-}$ to only one side, which unbalances the equation. (Balancing Redox with Half-Reactions)
Leaving electrons in the final overall equation because the half-reactions were not multiplied to the same number of electrons. (Balancing Redox with Half-Reactions)
Flipping the sign of the anode's $E^{\circ}$ AND subtracting it, which counts the reversal twice. Use reduction potentials from the table in $E_{cathode}^{\circ} - E_{anode}^{\circ}$ , or flip one sign and add, but never both. (Galvanic Cells and Cell Potential)
Multiplying $E^{\circ}$ by the coefficient used to balance electrons, e.g. using $2 \times 0.80 V$ for silver. (Galvanic Cells and Cell Potential)
Drawing electrons moving through the salt bridge or solution. Electrons travel only through the wire; ions carry the charge in solution. (Galvanic Cells and Cell Potential)
Getting salt-bridge ion flow backward. Anions head toward the anode, which is producing positive ions; cations head toward the cathode, which is using positive ions up. (Galvanic Cells and Cell Potential)
Writing cell notation with the cathode on the left. (Galvanic Cells and Cell Potential)
Thinking the anode is always negative. It is negative in a galvanic cell but positive in an electrolytic cell; what stays the same is that oxidation happens at the anode. (Electrolytic Cells, Corrosion and Batteries)
Expecting sodium metal from electrolysis of aqueous NaCl. Water is reduced more easily than $Na^{+}$ , so hydrogen forms; sodium metal needs molten NaCl. (Electrolytic Cells, Corrosion and Batteries)
Making the object to be electroplated the anode. It must be the cathode, where metal ions are reduced onto it. (Electrolytic Cells, Corrosion and Batteries)
Believing tin plating protects scratched steel the way zinc does. Tin is less active than iron, so exposed iron becomes the anode and corrodes faster. (Electrolytic Cells, Corrosion and Batteries)
Forgetting to convert minutes to seconds in $Q = I t$ , or forgetting the electron ratio for ions like $Cu^{2 +}$ that need 2 electrons per atom. (Electrolytic Cells, Corrosion and Batteries)

Practice test

24 questions from across the unit. Each one shows the lesson it comes from.

Oxidation Numbers and Redox Reactions
1
What is the oxidation number of chlorine in the chlorate ion, $ClO_{3}^{-}$ ? Include the sign.
Oxidation Numbers and Redox Reactions
2
In $Fe_{2} O_{3} + 3 CO \to 2 Fe + 3 CO_{2}$ , which substance is the reducing agent?
Oxidation Numbers and Redox Reactions
3
Which reaction is NOT a redox reaction?
Oxidation Numbers and Redox Reactions
4
When $H_{2} S + HNO_{3} \to S + NO + H_{2} O$ is balanced with the lowest whole-number coefficients, what is the coefficient of $H_{2} O$ ?
Oxidation Numbers and Redox Reactions
5
In the balanced equation $MnO_{4}^{-} + 5 Fe^{2 +} + 8 H^{+} \to Mn^{2 +} + 5 Fe^{3 +} + 4 H_{2} O,$ how many electrons are transferred in total?
Oxidation Numbers and Redox Reactions
6
For $Zn (s) + Cu^{2 +} (aq) \to Zn^{2 +} (aq) + Cu (s),$ match each term to the correct species.
Substance oxidized
Substance reduced
Product of oxidation
Product of reduction
Balancing Redox with Half-Reactions
7
Which is the correctly balanced half-reaction for $NO_{3}^{-} \to NO_{2}$ in acidic solution?
Balancing Redox with Half-Reactions
8
How many electrons must be added to balance $Cr_{2} O_{7}^{2 -} + 14 H^{+} \to 2 Cr^{3 +} + 7 H_{2} O ?$
electrons
Balancing Redox with Half-Reactions
9
A student balances a reaction in acidic solution and gets $\dots + 6 H^{+} \to \dots$ . To convert it to basic solution, what should they do next?
Balancing Redox with Half-Reactions
10
Balance $MnO_{4}^{-} + SO_{3}^{2 -} \to MnO_{2} + SO_{4}^{2 -}$ in basic solution with the lowest whole-number coefficients. What is the coefficient of $OH^{-}$ (it appears as a product)?
Balancing Redox with Half-Reactions
11
Put the steps for balancing a half-reaction in acidic solution in order.
1. 1Balance O by adding H2O
2. 2Balance charge by adding electrons
3. 3Balance H by adding H+
4. 4Balance atoms other than O and H
Balancing Redox with Half-Reactions
12
When balancing a half-reaction in acidic solution, what species do you add to balance oxygen atoms? Give its formula.
Galvanic Cells and Cell Potential
13
Using $E^{\circ} (Ag^{+} / Ag) = + 0.80 V$ and $E^{\circ} (Mg^{2 +} / Mg) = - 2.37 V$ , calculate $E_{cell}^{\circ}$ for the spontaneous magnesium-silver cell. Give your answer to 2 decimal places.
V
Galvanic Cells and Cell Potential
14
In the cell $Zn (s) ∣ Zn^{2 +} (aq) ∥ Cu^{2 +} (aq) ∣ Cu (s),$ which way do the nitrate ions in a $KNO_{3}$ salt bridge move?
Galvanic Cells and Cell Potential
15
Using approximate values $E^{\circ} (Fe^{3 +} / Fe^{2 +}) = + 0.77 V$ and $E^{\circ} (Cu^{2 +} / Cu) = + 0.34 V$ , is $2 Fe^{3 +} + Cu \to 2 Fe^{2 +} + Cu^{2 +}$ spontaneous?
Galvanic Cells and Cell Potential
16
At which electrode of a galvanic cell does reduction occur? (one word)
Galvanic Cells and Cell Potential
17
Write the chemical symbol of the metal that is the anode in a spontaneous cell made from nickel ( $E^{\circ} = - 0.26 V$ ) and silver ( $E^{\circ} = + 0.80 V$ ) half-cells.
Galvanic Cells and Cell Potential
18
Match each part of the zinc-copper galvanic cell to its role.
Zinc electrode
Copper electrode
Wire
Salt bridge
Electrolytic Cells, Corrosion and Batteries
19
In an electrolytic cell, which statement is true?
Electrolytic Cells, Corrosion and Batteries
20
Copper is refined using a current of 1.50 A for 1.00 h. What mass of copper, in grams, is deposited at the cathode? Use $F = 96485 C/mol$ and $M (Cu) = 63.55 g/mol$ . Give 3 significant figures.
g
Electrolytic Cells, Corrosion and Batteries
21
Why does a scratch in a galvanized (zinc-coated) steel bucket not cause the exposed steel to rust?
Electrolytic Cells, Corrosion and Batteries
22
A student wants to copper-plate a steel key. Which setup is correct?
Electrolytic Cells, Corrosion and Batteries
23
Put the events of iron corrosion under a water droplet in order.
1. 1Further oxidation by air produces hydrated iron(III) oxide (rust)
2. 2Fe2+ and OH− ions meet and form iron(II) hydroxide
3. 3Oxygen and water gain electrons to form OH−
4. 4Iron at the anode region loses electrons to form Fe2+
5. 5Electrons travel through the metal to the cathode region
Electrolytic Cells, Corrosion and Batteries
24
Match each technology to the electrochemical idea behind it.
Magnesium rod in a water heater tank
Recharging a lithium-ion battery
Hydrogen fuel cell
Refining copper
Painting a steel bridge

Endless drill

Fresh numbers every time, for the calculation questions in this unit.

Endless drillnew numbers every time

Generating a problem…