Introduction to Redox

Mrs. Kay

Chemistry 12

Chapter 18 Pages:713-729

Redox Reactions

• combustion of gas in a car

• rusting of metals• bleaching hair• reactions in batteries• Cut apples turning

brown

What do you think Redox stands for?

Redox = oxidation and reduction reactions

• Oxidation: loss of electrons

• Reduction: gain of electrons

Hint:

Leo the lion says Ger

• redox reactions are a family of reactions that are concerned with the transfer of electrons between species

• Like acid-base reactions, redox reactions are a matched set -- you don't have an oxidation reaction without a reduction reaction happening at the same time

Single Displacement redox reaction

• Cu(s) + 2AgNO3(aq) ---> Cu(NO3)2(aq) + 2Ag(s)

• The silver nitrate solution is transparent, when the copper wire is placed in it, the reaction begins slowly. The wire is coated with silver, while the copper is broken down into ions

Ex: Mg(s) + O2(g) MgO(s)

• Mg loses 2 electrons to become Mg+2

• O2 gains 2 electrons to become O-2

• The total reaction = redox reaction

• The half reactions = oxidation and reduction reactions

half reactions for the overall redox reaction

• Mg Mg+2 + 2e- (Magnesium loses electrons, so it is oxidized)

• O2 + 4e- 2O2- (Oxygen gains electrons, so it is reduced)

We’re Not Finished yet!

• When combining half reactions, you must make sure that the electrons gained = the electrons lost. This ensures balancing of the redox reaction. (so we multiplied the first half reaction by 2, so that the 4e- balanced out)

• 2Mg 2Mg+2 + 4e-

• O2 + 4e- 2O2-

• Total redox reaction is: 2Mg + O2 2MgO

Let’s see that again…

• The unbalanced reaction is as follows:

• Look at each half reaction separately: aluminum metal being oxidized to form an aluminum ion with a +3 charge and oxygen being reduced to form two (2) oxygen ions, each with a charge of -2.

• If we combine those two (2) half-reactions, we must make the number of electrons equal on both sides. The number 12 is a common

Taking care of the number of atoms, you should end up with:

Practice Writing the half reactions and balance the following

equations.

1. Fe + Br2 FeBr3

2. Ni + HgCl2 Hg + NiCl2

3. Sn+4 + Cu Sn+2 + Cu+2

4. CO+ I2O5 CO2 + I2 (this one’s tricky!)

Vocabulary:

• Examine : 2Mg + O2 2MgO

• Mg oxidation # increased from 0 to +2, therefore it has oxidized.

• O2 oxidation # decreased from 0 to –2, therefore it has reduced.

• Oxygen has been reduced because of magnesium, so magnesium is the reducing agent (the atom who gave its electrons away)

• Magnesium has been oxidized because of oxygen, so the oxygen is the oxidizing agent (the atom who accepted electrons)

• In redox: the reducing agent is oxidized and the oxidizing agent is reduced.

Assigning oxidation numbersPg 721-726

• Actual or hypothetical charges assigned using a set of rules.– Used to describe redox reactions, identify

redox reactions and to identify oxidizing agents and reducing agents.

Oxidation Numbers from Lewis structures

• Using electronegativities• In molecules, the more electronegative atom will

have the electrons found closer to its nucleus.• Molecules don’t have CHARGES, but you can

assign Ox # based on which atom “owns” the electrons more

• So the Oxygen = -2 and the Hydrogen each = +1

Oxidation Numbers

• Molecules with the same atoms will have no difference in electronegativity, so they share electrons evenly.

• Ex: Cl2 occurs where Cl-Cl, then pair of electrons are shared equally, so there is NO OWNING of electrons.

• Cl will have an oxidation number of zero.

Oxidation numbers for ionic compounds

• The oxidation number for an ionic compound is equal to its charge

• Example– MgO is made of Mg 2+and O2-

– So, Magnesium has an ox# of +2 and oxygen has an ox# of -2.

Oxidation Number Rules: Pg 724

Rules Example

1. A pure element has an ox # of zero.

Ne, Na, Cl2 = 0

2. A monatomic ion ox # equals its charge

Ca2+ = +2

3. Hydrogen with a metal = -1, Hydrogen with a non-metal =+1

Ex: H2S : H=+1

NaH: H=-1

Oxidation Number rules

Rules Examples

4. Oxygen is usually -2, unless its in a peroxide (H2O2 or OF2)

Li2O : O = -2

H2O2: O= -1

OF2: O=+2

5. In molecules, the more electronegative atom usually has the same charge that it would have had as an ion

PCl3: Cl=-1, so P=+3

CS2: S=-2, so C=+4

Oxidation number rules

Rules Examples

6. The sum of all ox# in a compound or molecule equal zero

CF4: C=+4, because 4 F=-1 each

7. The sum of ox# in a polyatomic ion equal the charge of the ion

NO3-

Exercise: For the following reactions

1. draw a diagram showing the loss and gain of electrons

2. identify the substance oxidized, the substance reduced, the oxidizing agent and the reducing agent.

3. Write the oxidation and reduction half reactions

Practice together:

• Example:

1,2. 2Mg + O2 2MgO

3. 2Mg 2Mg+2 + 4e-

O2 + 4e- 2O2-

0 0 +2 -2

Gain of e-

Substance reduced

Oxidizing agent

Loss of e-

Substance oxidized

Reducing agent

Work on the following in class, finish for homework:

1.Ni + HgCl2 Hg + NiCl2

2.2Na + 2H2O 2NaOH + H2

3.Cl2 +2NaBr Br2 + 2NaCl

4.4NH3 + 7O2 4NO2 + 6H2O

Practice:

• Page 715 # 1-4

• Page 716 # 6 & 7

• Page 726 # 9-12

• http://staff.prairiesouth.ca/~chemistry/chem30/6_redox/redox1_2.htm