Rules for Assigning Oxidation Numbers



Name ______________________ Date ______________

UNIT #15: Electrochemistry & Redox Reactions

REVIEW: Rules for Assigning Oxidation Numbers

1. Oxidation numbers for atoms of free elements and diatomic molecules are always zero.

Ex. He = _______

Ne = _______

H2 = _______

2. The oxidation numbers of ions are the same as the charge on the ion.

Ex. Li1+ = _______

Ca2+ = _______

N4- = ________

3. Some elements have only one oxidation state.

a. Group 1 metals always form +1 ions & always have a +1 oxidation state.

b. Group 2 metals always form +2 ions & always have a +2 oxidation state.

Find the Oxidation Number for each of the following:

1. Al = ________ 5. I2 = __________

2. H2= ________ 6. Cu1+ = ________

3. B3+ = ________ 7. O2- = _________

4. Xe = ________ 8. Zn = _________

Rules for Calculating Missing Oxidation Numbers in a Compound

1. Some elements usually have a particular oxidation state.

a. Oxygen has a -2 oxidation state except in peroxides (i.e., H2O2) where it is -1 and

with fluorine (i.e., OF2) where it is +1.

b. Hydrogen has a +1 oxidation state except when combined with a metal, in which case

it is -1.

2. The sum of the oxidation numbers in a compound is always ZERO.

3. The sum of the oxidation states in a polyatomic ion (Table E) is equal to the charge on the polyatomic ion.

4. When trying to determine oxidation states for elements in a compound:

a. For neutral compounds: The sum of the oxidation numbers in a must add up to zero.

Find the oxidation numbers for each of the elements in:

|Formula |Oxidation State of the First |Oxidation State of the Last Element|

| |Element | |

|NaCl | | |

|H2O | | |

|BF3 | | |

|CaBr2 | | |

b. For polyatomic ions: The sum of the oxidation numbers must add up to the charge

on the ion.

Find the oxidation numbers for each of the elements in:

|Formula |Oxidation State of the First |Oxidation State of the Middle |Oxidation State of the Last Element|

| |Element |Element | |

|CO32– | | | |

|C2O42– | | | |

|SCN1- | | | |

|PO43– | | | |

Practice Assigning Oxidation States

|Formula |Oxidation State of the First |Oxidation State of the Middle |Oxidation State of the Last Element|

| |Element |Element | |

|K3PO4 | | | |

| | | | |

|Na2Cr2O7 | | | |

| | | | |

|NaClO4 | | | |

| | | | |

|Li2CO3 | | | |

| | | | |

|FeSO3 | | | |

| | | | |

|MnO41– | | | |

| | | | |

|Cu(NO3)2 | | | |

| | | | |

|KClO3 | | | |

| | | | |

|HSO42- | | | |

| | | | |

|KMnO4 | | | |

| | | | |

|LiOH | | | |

| | | | |

Name: ________________ Date: ___________

Oxidation State Practice

1. What is the oxidation number of sodium in sodium chloride?

2. What is the oxidation number of chlorine in calcium chloride?

3. What is the oxidation number of sulfur in sodium sulfate?

4. What is the oxidation number of chromium in potassium dichromate?

5. What is the oxidation number of phosphorus in the phosphate ion?

6. What is the oxidation number of iron in iron (III) chloride?

7. What is the oxidation state of nitrogen in sodium nitrite?

8. What is the oxidation number of chromium in potassium chromate?

Aim:

How are oxidation numbers used to keep track of electrons during chemical changes?

REDOX REACTION:

______________________________________________________

______________________________________________________

Write the balanced equation for the synthesis of sodium chloride.

The Synthesis of Sodium Chloride:

_____________________________________

___ is oxidized, because it ______ electrons and its oxidation number ___.

___ is reduced, it ______ electrons and its oxidation number ___.

During chemical reactions, electrons are often transferred from one species to another. To keep track of the gain and loss of electrons, we use oxidation numbers.

“LEO says GER”

___________________________________

____________________________________

“OIL RIG”

____________________________________

____________________________________

Medial Summary:

|OXIDATION |REDUCTION |

|The ______ of electrons by an atom. |The ______ of electrons by an atom. |

|Oxidized species shows an ___________ in oxidation number. |Reduced species shows a ___________ in oxidation number. |

|_________ are usually oxidized. |_________ are usually reduced. |

|Example: Sodium metal and chlorine gas combine to form sodium chloride. |

| |

|______________________________________ |

| |

|Na ____ electrons and is ___________. Its oxidation state goes from ___ to ____. |

|Cl ____ electrons and is ___________. Its oxidation state goes from ___ to ____. |

| |

|Oxidation Half-Reaction: __________________________________ |

| |

|Reduction Half-Reaction: __________________________________ |

Practice Determining Oxidized and Reduced Species

Balance the following reactions, assign oxidation numbers to each of the species, and identify the species that is oxidized and the species that is reduced. Write the half-reactions.

1.

Zn + AlCl3 ( ZnCl2 + Al

____ is oxidized (reducing agent); its oxidation number goes from ___ to ____.

____ is reduced (oxidizing agent); its oxidation number goes from ___ to ____.

Oxidation Half-Reaction: _____________________________________

Reduction Half-Reaction: _____________________________________

2.

Mg + O2 ( MgO

____ is oxidized (reducing agent); its oxidation number goes from ___ to ____.

____ is reduced (oxidizing agent); its oxidation number goes from ___ to ____.

Oxidation Half-Reaction: __________________________________

Reduction Half-Reaction: __________________________________

3.

Li + ZnCl2 ( LiCl + Zn

____ is oxidized (reducing agent); its oxidation number goes from ___ to ____.

____ is reduced (oxidizing agent); its oxidation number goes from ___ to ____.

Oxidation Half-Reaction: __________________________________

Reduction Half-Reaction: __________________________________

4. Consider the reaction: Sn+2 + 2Fe+3 ( 2Fe+2 + Sn+4

Which species is oxidized? a) Fe+3 b) Sn+2 c) Fe+2 d) Sn+4

5. Consider the reaction: 2Ag+1 + Cu ( Cu+2 + Ag

Which species is reduced? a) Ag+1 b) Cu c) Cu+2 d) Ag

6. Which of the following is a redox reaction?

a) HCl + KOH ( KCl + H2O

b) 4HCl + MnO2 ( MnCl2 + 2H2O + Cl2

c) 2HCl + CaCO3 ( CaCl2 + H2O + CO2

d) 2HCl + FeS ( FeCl2 + H2S

Note: In order for a reaction to be a redox reaction, there must be a change in oxidation numbers. If an element appears in a compound on one side of the equation and alone on the other, the reaction must be a redox.

6. Which of the following is NOT a redox reaction?

a) Al + O2 ( Al2O3

b) H2S + HNO3 ( S + NO + H2O

c) H2O ( H2 + O2

d) 2HCl + FeS ( FeCl2 + H2S

Note: Double replacement reactions are NEVER redox reactions.

7. Given the balanced ionic equation: 2Al(s) + 3Cu2+(aq) --> 2Al3+(aq) + 3Cu(s)

Compared to the total charge of the reactants, the total charge of the products is

(1) less (2) greater (3) the same (4) need more information

8. Which change in oxidation number indicates oxidation?

(1) –1 to +2 (2) –1 to –2 (3) +2 to –3 (4) +3 to +2

Aim:

How Can We Balance Equations Using the Half-Reaction Method?

Do Now: Review Balancing Half-Reactions

Step 1: Assign oxidation numbers to each of the elements on both sides of the equation.

Step 2: Determine which element was oxidized and which was reduced.

Step 3: Write the partial oxidation half (electrons LOST, so electrons will be on the right (product) side of the equation).

Step 4: Write the reduction half (electrons GAINED, so electrons will be on the left (reactant) side of equation).

Step 5: Make sure the total moles of electrons gained and lost is equal among both half-reactions.

|Example: |

| |

|Mg + Br2 ( MgBr2 |

| |

|Oxidation Half-Reaction: _________________________ |

|Reduction Half-Reaction: _________________________ |

| |

| |

|Fe + Zn+2 ( Fe+2 + Zn |

| |

| |

|Oxidation Half-Reaction: _________________________ |

|Reduction Half-Reaction: _________________________ |

| |

|3Fe+2 + 2Al ( 3Fe + 2Al+3 |

| |

| |

|Oxidation Half-Reaction: _________________________ |

|Reduction Half-Reaction: _________________________ |

More Practice Balancing Half-Reactions

1. Balance the half-reaction for mass by adding coefficients to make the number of atoms equal on each side.

2. Calculate the total charge on each side of the equation. Balance for charge by adding electrons to one side of the equation. Hint: Always add the electrons to the more positive side of the equation.

3. Identify the equation as either oxidation (electrons are a product; oxidation number increases), or reduction (electrons are a reactant; oxidation number decreases).

|Half-Reaction |Oxidation or Reducation |

|Ga ( Ga+3 | |

|O3 ( O-2 | |

|2Cr+3 ( Cr+2 | |

|S+6 ( S-2 | |

|F-1 ( F2 | |

|As+3 ( As+5 | |

Balancing EQUATIONS Using the Half-Reaction Method

__ Mg(s)0 + __ Fe3+ ( __ Mg2+ + __ Fe(s)0

Oxidation half-reaction: ______________________________

Reduction half-reaction: ______________________________

Balance the equation using the half-reaction method:

Notice that your net reaction is balanced for both mass and charge.

Practice Balancing Redox Rxns Using Half Rxns

1. The electrons on each side must be made equal; if they are not equal, they must be multiplied by appropriate integers (the lowest common multiple) to be made the same.

2. The half-equations are added together, canceling out the electrons to form one balanced equation.

3. Add the appropriate coefficients to the original redox equation.

1. __ Zn(s)0 + __ Ag+1 ( __ Zn+2 + __ Ag0

Oxidation half-reaction: ______________________________

Reduction half-reaction: ______________________________

Add the half reactions together:

2. __ Al0+ __ Fe+2 ( __ Al+3 + __ Fe0

Oxidation half-reaction: ______________________________

Reduction half-reaction: ______________________________

Add the half reactions together:

3. __ AlBr3(s) + __ Fe(s) ( __ FeBr2 + __ Al0

Oxidation half-reaction: ______________________________

Reduction half-reaction: ______________________________

Add the half reactions together:

4. __ Al(s) + __ Zn2+(aq) ( __ Al3+(aq) + __ Zn(s)

Oxidation half-reaction: ______________________________

Reduction half-reaction: ______________________________

Add the half reactions together:

Voltaic (Galvanic) Cells

(a.k.a. batteries)

Portable electronic devices run on batteries. The electricity generated by a battery comes from a chemical reaction known as an oxidation-reduction reaction. During a single replacement, a type of oxidation-reduction reaction, more active metals transfer electrons to less active metals. As a result, the more active metal is oxidized, and the less active metal is reduced. If the oxidation and reduction half reactions are physically separated and attached by a wire, electrons will flow through the wire during the reaction and can be used to power our portable electronics. This is done by putting electrolytes, usually aqueous acids, bases, or salts, into separate containers. The separate containers are called half cells because the half reactions are isolated in them. They are connected by a salt bridge which lets ions travel between half cells. Electrodes are immersed into the electrolytes. The electrodes are merely metals with differing activity. Completing the circuit by connecting the electrodes with a wire enables electrons to flow through the wire from the more active metal to the less active metal, reducing the less active metal. The electrode where reduction occurs is called the cathode. The electrode where oxidation occurs is called the anode. The device that produces electric current from a chemical reaction is called a voltaic cell. Several voltaic cells attached together form a battery of cells. A battery produces a higher voltage than a single cell.

• The two solid metal bars are called ________________.

o The more active electrode (________ on Table J) is the site of __________ and is called the ____________.

o The less active electrode (________ on Table J) is the site of __________ and is called the ____________.

• Electrons flow across the external circuit (______) from _________ to __________, generating electricity.

• The ______________ maintains neutrality in the two half-cells by allowing the transport of ions between half-cells.

• The anode loses mass as it becomes oxidized and the cathode gains mass as it becomes reduced.

Practice Understanding Voltaic Cells

Ag/Ag+ // Ni/Ni+2

1. Label the electrodes in the diagram as (Ag0(s) and Ni0(s)).

2. Using Table J, identify the anode (oxided) and the cathode (reduced).

Anode: _______________ Cathode: _______________

3. Label the anode and cathode in the diagram. Show the direction of electron flow across the wire.

4. Choose appropriate electrolyte solutions for each half-cell. Remember, the electrolyte solution should contain the same cation as the electrode in the half-cell.

5. Determine the following reactions, & fill them in the appropriate box under the beakers.

Oxidation half reaction: ______________________

Reduction half reaction: ______________________

4. Write out the balanced equation showing the spontaneous redox reaction that

occurs.

__________________________________________________

VOLTAIC CELLS – ON YOUR OWN

Ag/Ag+ // Pb/Pb+2

1. Label the electrodes in the diagram as (Ag0(s) and Pb0(s)).

2. Using Table J, identify the anode (oxided) and the cathode (reduced).

Anode: _______________ Cathode: _______________

3. Label the anode and cathode in the diagram. Show the direction of electron flow across the wire. Which electrode increases in mass? ___ Which electrode decreases in mass? ___

4. Choose appropriate electrolyte solutions for each half-cell. Remember, the electrolyte solution should contain the same cation as the electrode in the half-cell.

5. Determine the following reactions, & fill them in the appropriate box under the beakers.

Oxidation half reaction: ______________________

Reduction half reaction: ______________________

4. Write out the balanced equation showing the spontaneous redox reaction that

occurs.

__________________________________________________

Practice Regents Questions:

1. Use the voltaic cell to answer the following questions.

a. When the switch is closed, in which half-cell does oxidation occur? __________

b. When the switch is closed, state the direction that electrons will flow through the wire. ____________________

c. Based on the given equation, write the balanced half-reaction that occurs

in half-cell 1.

__________________________

2. Use the voltaic cell below to answer the following questions.

a. In which half-cell will reduction occur when switch S is closed? __________

b. Write the balanced half-reaction equation that will occur in half-cell 2 when

switch S is closed.

______________________________

c. As this voltaic cell operates, the mass of the Zn(s) electrode decreases. Explain, in terms of particles, why this decrease in mass occurs.

________________________________________________________________

________________________________________________________________

3. Base your answers to questions a –c on the diagram of a voltaic cell and the balanced ionic equation below.

a. What is the total number of moles of electrons needed to completely reduce

6.0 moles of Ni2+(aq) ions? ____________

b. Identify one metal from Reference Table J that is more easily oxidized than

Mg(s). ____________

c. Explain the function of the salt bridge in the voltaic cell.

_________________________________________________________

4. In a laboratory investigation, a student constructs a voltaic cell with iron and copper electrodes. Another student constructs a voltaic cell with zinc and iron electrodes. Testing the cells during operation enables the students to write the balanced ionic equations below.

Cell with iron and copper electrodes: Cu2+ (aq) + Fe(s) ( Cu(s) + Fe2+(aq)

Cell with zinc and iron electrodes: Fe2+(aq) + Zn(s) ( Fe(s) + Zn2+(aq)

a. State evidence from the balanced equation for the cell with iron and copper electrodes that indicates the reaction in the cell is an oxidation-reduction reaction.

________________________________________________________________

b. Identify the subatomic particles transferred between Fe2+ and Zn during the reaction in the cell with zinc and iron electrodes. ___________________

c. State the relative activity of the three metals used in these two voltaic cells.

________________________________________________________________________________________________________________________________

Forcing Electrons to Move

An Introduction to Electrolytic Cells

The energy to run most cars comes from gasoline. Except in diesel engines, the energy is released from the gasoline by exploding it with a tiny spark from a spark plug. The energy to make the spark comes from the car’s battery. The battery in the car is called a “wet cell.” It contains sulfuric acid [H2SO4(aq)], a liquid electrolyte. The electricity is generated by the following chemical reaction:

PbO2 + Pb + 2H2SO4 ( 2PbSO4 + 2H2O

Car batteries can last for several years. This is because they get recharged. As the engine spins, a moving magnet in the alternator pushes electrons in a direction opposite to the way they normally flow from the battery. These electrons reverse the chemical reaction that generated electricity in the battery. A cell that uses electricity to produce a chemical reaction in this way is called an electrolytic cell. When the car battery is generating electricity it is a voltaic cell. When it is being recharged, it is an electrolytic cell.

Answer the questions below based on the reading above and on your knowledge of chemistry.

1. Write the chemical reaction that occurs when a car battery generates electricity.

____________________________________

a) What is oxidized? _______ What is reduced? ______

b) Write the oxidation half reaction: ________________________________

c) Write the reduction half reaction: ________________________________

2. Write the chemical reaction that occurs when a car battery is (opposite of the equation in 1a).

a) What is oxidized? _______ What is reduced? ______

b) Write the oxidation half reaction: ________________________________

c) Write the reduction half reaction: ________________________________

3. Aluminum is found in the mineral bauxite (Al2O3). To get pure aluminum, the aluminum needs to be separated from oxygen.

a) Imagine that bauxite forms by the following reaction:

4Al + 3O2 ( 2Al2O3.

o What is oxidized? _______ What is reduced? ______

o Write the oxidation half reaction: ________________________________

o Write the reduction half reaction: ________________________________

b) Write the rxn for the purification of aluminum from bauxite (reverse of reaction above).

________________________________

o What is oxidized? _______ What is reduced? ______

o Write the oxidation half reaction: ________________________________

o Write the reduction half reaction: ________________________________

o Metals tend to [ gain / lose ] electrons and become oxidized. Is the purification of aluminum from bauxite a spontaneous reaction? ____ Why? ___________________

________________________________________________________________

4. Iron is often protected from rusting by a process called galvanizing (i.e., galvanized nails). When a metal is galvanized, it is coated with zinc. One way to coat iron with zinc is through a single replacement reaction: Fe + Zn(NO3)2 ( Fe(NO3)3 + Zn. Since the reaction occurs at the surface of the iron, the iron becomes plated with zinc.

a) What is oxidized? _______ What is reduced? ______

b) Write the oxidation half reaction: ________________________________

c) Write the reduction half reaction: ________________________________

d) Consult the activity series on Table J. How likely is this reaction to occur? Explain.

_____________________________________________________________________

5. Based on the questions above, what is an electrolytic cell? How is it different than a voltaic cell? What are two functions of electrolytic cells? __________________________________________________________________________

__________________________________________________________________________

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