Redox and Electrochemistry. Redox Reactions Reduction – Oxidation reactions Involve the transfer...

Redox and

Electrochemistry

Redox Reactions• Reduction – Oxidation reactions• Involve the transfer of electrons from one

substance to another

The oxidation numbers of the atoms will change…. one goes up (oxidation) and one goes down (reduction)

+

Oxidation Number (Oxidation State)• Used to keep track of the transfer of

electrons

• Number is assigned to every atom in a chemical formula, in accordance with certain rules

• NOT an ionic charge, but is often the same as the ionic charge– Possible oxidation states are given on the

periodic table (upper right hand corner)

Rules for assigning Oxidation Numbers

1. For a neutral compound, the sum of the oxidation states must be zero

2. The oxidation state of any atom in an uncombined element is zero

• Element not in chemical combination with another element

• Examples: Na, Mg, H2, Cl2

Rules for assigning Oxidation Numbers

3. The oxidation state of a monatomic ion is equal to its charge

– Examples: Na+ = +1

CaCl2

4. For a polyatomic ion, the sum of the oxidation states must equal the overall charge

– Example: SO42-

Rules for assigning Oxidation Numbers

5. Metals of group 1 always have an oxidation number of +1

6. Metals of groups 2 always have an oxidation number of +2

7. Fluorine is always -1, other halogens are usually -1

8. Aluminum is always +3

Rules for assigning Oxidation Numbers

9. Oxygen is usually -2

Exceptions: – When paired with F (OF2), oxygen will be +2

– Peroxides (H2O2), oxygen will be -1

10. Hydrogen is usually +1

Exceptions: – Metal hydrides (Group 1 or 2 metals paired

with hydrogen), LiH, CaH2, hydrogen will be -1

Examples

Assign an oxidation state to each element in the following:

1. H2SO4

2. SO32-

3. K2CrO4

4. CrCl3

• Reduction– Reduction of charge by gaining electrons

Na+ + e- Na

O + 2e- → O2-

• Oxidation– Increase in charge by loss of electrons

Fe Fe3+ + 3e-

Cl- Cl + e-

LEO the lion says GER

Losing Electrons Oxidation

Gaining Electrons Reduction

Conservation of Matter/Conservation of Charge

• Mass must be conserved – Mass on both sides must be the same

(balanced)

• Charge must be conserved– Net charge on both sides must be the same

(balanced) – add electrons to the higher side

• Reduction and Oxidation reactions must occur together (REDOX reactions)

Half Reactions

• Every Redox reaction consists of a reduction and oxidation reaction

• Each reaction is called a ½ reaction

• A separate equation can be written for each ½ reaction

Examples

1. H2 + Cl2 2HCl

2. Fe + ZnO Zn + FeO

Half Reactions

• Net charge and mass must be the same on both sides of the equation

• The number of electrons must balance out, electrons do not appear in the net equation

• One ½ reaction is reduction and the other is oxidation

Spectator Ion

• Does not change oxidation states in the reaction, same oxidation state on both sides of the equation

• Not every species in an equation is oxidized or reduced, some are spectator ions

Reducing Agent

• Substance which is oxidized– Serves as a source of electrons to make

the reduction reaction occur

– Good reducing agents are substances that lose (donate) electrons easily – elements with low ionization energies

Examples: group 1 and 2 metals

Oxidizing Agent

• Substance which is reduced– Accepts (gains electrons)

– Good oxidizing agents are substances that gain electrons (highly electronegative elements)

Examples: Group 17 elements

Balancing Equations

1. Assign oxidation numbers to all substances in the equation

2. Identify the substances that are being oxidized and reduced (there will be 2 substances that have a change oxidation numbers)

3. Write the oxidation and reduction ½ reactions

Balancing Equations

4. Balance (cancel out) the electrons in the ½ reaction

5. Balance the rest of the equation

6. Check

Examples

1. Fe + Cl2 → FeCl3

2. Fe + CuSO4 → FeSO4 + Cu

3. KMnO4 + HCl → KCl + MnCl2 + H2O + Cl2

Activity Series

Reference Table J

Metals

• The most reactive metals are listed at the top

• An uncombined metal will replace (in a chemical compound) any other metal ion on the table that is below that uncombined metal

• Reactive metals lose electrons easily (low ionization energy)

Examples

Ba + ZnCl2 → Zn + BaCl2• Ba will replace Zn because Ba is

above Zn– Ba is more reactive than Zn

• More reactive means that it loses electrons easier

Nonmetals

• For the halogen nonmetals listed in Table J, the most reactive ones are at the top

• For nonmetals, high reactivity means that they are likely to gain electrons (high electronegativity)

Example:

F2 will replace any other halogen (it is the most reactive)

Examples

1. Which metal is most reactive?

a. Fe b. Zn Cu

2. Will Ba react with Mn2+?

3. Will Na+ react with Cr?

4. Will this reaction occur spontaneously? Mg + Co(NO3)2 →

5. If this reaction does occur, what products would be made?