1 Chapter 20 Oxidation-Reduction Reactions (Redox Reactions)

1

Chapter 20 Oxidation-Reduction Reactions (Redox Reactions)

2

Launch Lab

• Complete the “Penny Chemistry” Lab

• ¼ cup = 57 mL• 1 tsp = 5 mL

3

The chemical changes that occur when electrons are transferred between reactants are called oxidation – reduction reactions

4

-    -principal source of energy on earth-    -combustion of gasoline-    -burning of wood-burning food in your body

oxidation reactions

5

Oxidation reactions are always accompanied by a reduction reaction

Oxidation -    originally meant combining with oxygen-    iron rusting (iron + oxygen)  Reduction -    originally meant the loss of oxygen from a compoundremoving iron from iron ore ( iron II oxide)

6

7

20.2 Electron Transfer in Redox Reactions 

Today OXIDATION means:-    a complete or partial LOSS of ELECTRONS REDUCTION means:-    a complete or partial GAIN of ELECTRONS Memory Device :

LEO the lion says GER or OIL RIG

8

The substance that donates electrons in a redox reaction is the REDUCING AGENT 

The substance that takes electrons in a redox reaction is the OXIDIZING AGENT

9

Oxidation is…

–the loss of electrons

–an increase in oxidation state

–the addition of oxygen

–the loss of hydrogen

2 Mg + O2 2 MgO

notice the magnesium is losing electrons

Reduction is…

–the gain of electrons

–a decrease in oxidation state

–the loss of oxygen

–the addition of hydrogen

MgO + H2 Mg + H2O

notice the Mg2+ in MgO is gaining electrons

10

11

Oxidation states are numbers assigned to atoms that reflect the net charge an atom would have if the electrons in the chemical bonds involving that atom were assigned to the more electronegative atoms.

Oxidation states can be thought of as “imaginary” charges. They are assigned according to the following set of rules:

Oxidation States

20.3 Assigning Oxidation Numbers (ON)

12

#1The ON of a simple ion is equal to its

ionic charge

+1 +2 -3Na + Cu 2+ N3-

13

#2The ON of hydrogen is always +1, except in metal hydrides like NaH

where it is –1

+1 -1 HCl NaH

14

#3 The ON of oxygen is always –2 except in

peroxides like X2O2 where it is –1

-2 -1H2O H2O2

15

#4 The ON of an uncombined element is

always zero

0 0 0Na Cu N2

16

#5For any neutral(zero charge) compound, the sum of the ON’s is always zero

+4-2 CO2

17

#6 For a complex ion, the sum of the ON’s

equals the charge of the complex ion

+7 -2MnO4

1-

18

Examples - assigning oxidation numbers

Assign oxidation states to all elements:

H2 SO3 SO42-

K+ NH3 MnO4-

Cr2O72- CH3OH PO4

3-

ClO3- HSO3

- Cu

19

20

Assignment

1-OBWS 1-5

2-Text 1-3,(pp481-86) 8-15(pp498) Chp 20

3-Worksheet #2 Oxidation Numbers

21

20.4 Oxidation # Changes

an increase in oxidation number of an atom signifies oxidation

a decrease in oxidation number of an atom signifies reduction

+2 to +4

0 to -1

22

Identifying Redox Reactions

Oxidation and reduction always occur together in a chemical reaction. For this reason, these reactions are called “redox” reactions.

Although there are different ways of identifying a redox reaction, the best is to look for a change in oxidation state:

23

SnCl2 + PbCl4 SnCl4 + PbCl2

CuS + H+ + NO3- Cu+2 + S + NO + H2O

-1+2 -1+4 -1+4 -1+2

+2 = LEO

-2 = GER

-2+2 +1 -2+5 +2 0 -2+2 -2+1

-3 = GER

+2 = LEO

RA

OA

RA

OA

24

25

In each reaction, look for changes in oxidation state.

If changes occur, identify the substance being reduced, and the substance being oxidized.

Identify the oxidizing agent and the reducing agent.

Examples - labeling redox reactions

H2 + CuO Cu + H2O0 -2+2 0 -2+1

= +1 (H is oxidized)

= -2 (Cu is reduced)

(reducing agent)

(oxidizing agent)

26

5 Fe2+ + MnO4- + 8 H+ 5 Fe3+ + Mn2+ + 4 H2O

Zn + 2 HCl ZnCl2 + H2

Try These!!

+1 = Fe 2+ is oxidized (reducing agent)

- 5 = Mn 7+ is reduced (oxidizing agent)

+2 = Zn 0 is oxidized (reducing agent)

- 1 = H 1+ is reduced (oxidizing agent)

27

How to write net ionic equations

• 1) write a balanced equation Cu(s) + 2NaCl(aq) 2Na(s) + CuCl2 (aq)

2) Ionize any aqueous substancesCu(s) + 2Na1+

(aq) 2Cl1-(aq) 2Na(s) + Cu2+ (aq) 2Cl

1- (aq)

3) Remove any like substances (spectators)

Cu(s) + 2Na1+(aq) 2Cl1-

(aq) 2Na(s) + Cu2+ (aq) 2Cl 1- (aq)

4) Sum up what’s left

Cu(s) + 2Na1+(aq) 2Na(s) + Cu2+ (aq)

The Net Ionic Equation (the reaction that is really occurring)

28

Table 12.1 Strength of oxidizing and reducing agentsInquiry into Chemistry Chapter 12

Oxidizing Agent Reduction Reducing Agent

Oxidation

Stronger Oxidizing Agent

Cu 2+ Cu

Zn 2+ ZnStronger Reducing Agent

29

Strongest Oxidizing Agent Weakest Reducing AgentBa 2+ (aq) Ba (s)

Ca 2+ (aq) Ca (s)

Mg 2+ (aq) Mg (s)

Al 3+ (aq) Al (s)

Zn 2+ (aq) Zn (s)

Cr 3+ (aq) Cr (s)

Fe 2+ (aq) Fe (s)

Cd 2+ (aq) Cd (s)

Tl + (aq) Tl (s)

Co 2+ (aq) Co (s)

Ni 2+ (aq) Ni (s)

Sn 2+ (aq) Sn (s)

Cu 2+ (aq) Cu (s)

Hg 2+ (aq) Hg (s)

Ag 2+ (aq) Ag (s)

Pt 2+ (aq) Pt (s)

Au 1+ (aq) Au (s)

Weakest Oxidizing Agent Strongest Reducing Agent

Oxidation Reduction Table

12.2Inquiry into Chemistry

30

Spontaneous Reaction

Loses 2 e -

Gains 2 e-

Stronger Reducing

Agent

StrongerOxidizing

Agent

Pt (s) Sn (s)Sn 2+ (aq) Pt 2+ (aq) ++

Compare Reducing Agents

Compare Oxidizing Agents

31

Non Spontaneous Reaction

Loses 2 e -

Gains 2 e-

StrongerOxidizing

Agent

Stronger Reducing

Agent

Compare Reducing Agents

Compare Oxidizing Agents

Mg (s) Fe2+ (aq) Mg 2+ (aq) Fe (s)+ +

32

Assignment

1-OBWS 6,7

2-Text 4, 16, 17 Chp 20

3-Worksheet # 3 Oxidizing Agents and Reducing Agents

4- Investigation 12.A Testing Relative Oxidizing and Reducing Strengths of Metal Atoms and Ions (see table 12.1)

5-Question Sheet for Table 12.2

6-Go back and answer part 4 of each Q on Worksheet #3

33

34

20.5 Balancing Redox Equations

1)the oxidation number change method

There are two methods used to balance redox reactions

2)the half reaction method

35

These methods are based on the fact that the total number of electrons gained in reduction must equal the total number of electrons lost in oxidation

Redox reactions are often quite complicated and difficult to balance. For this reason, you’ll learn a step-by-step method for balancing these types of reactions, when they occur in acidic or in basic solutions.

36

Oxidation Number Change Method

1)Assign ON to all atoms

2)Identify which atoms are oxidized and which are reduced

Balance the following: Fe2O3 + CO Fe + CO2

Fe2O3 + CO Fe + CO2

+3 -2 -2+2 0 -2+4

Fe2O3 + CO Fe + CO2

+3 -2 -2+2 0 -2+4

-3 (Fe reduced)

+2 (C oxidized)

37

3) Make the total increase in oxidation number equal the total decrease in oxidation number by using appropriate coefficientson the reactant side only.

Fe2O3 + CO Fe + CO2

+3 -2 -2+2 0 -2+4

-3

+2

(x 2 atoms) = 6 electrons gained

(X 3 atoms) = 6 electrons lost

3

4) Finally check to be sure that the equation is balanced both foratoms and charge.

Fe2O3 + CO Fe + CO23 2 3

38

39

Assignment

1-OBWS 8

2-Text 5, 18,19

3-Practice Sheet 20A

4-Investigation 12.B Redox Reactions and Balanced Equations

40

Balancing Equations with the Half-Reaction Method

1) First split the original equation into two half-reactions, one “reduction” and the other “oxidation”.

In each half-reaction, follow these steps:

2) Balance all elements except “H” and “O”.

3) Balance the “O’s” by adding water, H2O.

4) Balance the “H’s” by adding hydrogen ions, H+.

If your rxn is taking place in an acidic solution, skip to step 8

If your rxn is taking place in a basic solution proceed to step 5

5) Adjust for basic conditions by adding to both sides the same # of OH- ions as the number of H+ ions already present

6) Simplify the equation by combining H+ and OH- that appear on the same side of the equation into water molecules.

7) Cancel any water molecules present on both sides of the equation

8) Balance the charges by adding electrons

9) Recombine the ½ reactions into a complete balanced equation.

41

Example:

Fe2+ + Cr2O72- Fe3+ + Cr3+ acidic solution

Cr2O72- Cr3+

Fe2+ Fe3+

1( )

6( )

Cr2O72- + 6 Fe2+ + 14 H+ 2 Cr3+ + 6 Fe3+ + 7 H2O

2 7 H2O+14 H+ +6 e- +

1e-+

42

What if the solution was basic?

Notice that the method has assumed the solution was acidic - we added H+ to balance the equation. The [H+] in a basic solution is very small. The [OH-] is much greater.

For this reason, we will add enough OH- ions to both sides of the equation to neutralize the H+ added in the reaction.

The hydrogen and hydroxide ions will combine to make water, and you may have to do some canceling before you’re done.

Cr2O72- + Fe2+ + H2O Cr3+ + Fe3+

Try this in a basic solution!!!

43

Cr2O72- + Fe2+ + H2O Cr3+ + Fe3+ Basic Solution

Cr2O72- Cr3+14OH- + 14OH-

Cr2O72- + 6 Fe2+ + 7 H2O 2 Cr3+ + 6 Fe3+ + 14 OH-

7 H2O14H+ 2( ) 1

Fe2+ Fe3+( )6 + 1e-

+++6 e- + 14 H2O

44

Balancing Redox Equations PracticeBalance in acidic solution:

H2C2O4 + MnO4- Mn2+ + CO2

5 H2C2O4 + 2 MnO4- + 6 H+ 2 Mn2+ + 10 CO2 + 8 H2O

Balance in basic solution:

CN- + MnO4- CNO- + MnO2

3 CN- + 2 MnO4- + H2O 3 CNO- + 2 MnO2 + 2 OH-

45

46

Assignment

1-OBWS 9, 10, 11

2-Worksheet #4 Half Reactions

3-Practice Sheet 20B

47

Redox Reactions - What’s Happening?

Zinc is added to a blue solution of copper(II) sulfate

The blue colour disappears…the zinc metal “dissolves”, and solid copper metal precipitates on the zinc strip

The zinc is oxidized (loses electrons)

The copper ions are reduced (gain electrons)

Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s)

Zn (s) + Cu2+ (aq) Zn2+ (aq) + Cu (s)

48

Copper ions (Cu2+)

collide with the zinc

metal surfaceA zinc atom (Zn)

gives up two of its

electrons to the

copper ion

The result is a

neutral atom of Cu

deposited on the

zinc strip, and a

Zn2+ ion released

into the solution

49

Voltaic Cell

Electrolysis

SIMULATIONS

Redox Titration

50

Assignment :1-Web Quest

Oxidation/Reduction 2-Blue Print Lab

3-Review Worksheet