Lecture 7: Types of Chemical Reactions

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Lecture 7: Types of Chemical Reactions

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Oxidation-Reduction Reactions

Oxidation-Reduction (Redox) Reactions

Reactions involving the transfer of electrons. eg: in the formation of an ionic compound

2 Na_(s) + Cl_2(g) → 2NaCl_(s)

Half reactions:

2Na(s) → 2Na+ _{+ 2e}- _{electrons lost: oxidation}

Cl₂(g) + 2e- → _2Cl- _{electrons gained: reduction}

# of electrons lost = # of electrons gained

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Redox

Oxidation and reduction always occur together. Electrons lost by one species must be gained by another.

In reactions between metals and non metals, generally,

metals are oxidized and nonmetals are reduced.

Another example of an oxidation reduction : C2H6(g) + O2(g) → CO2(g) + H2O(g)

How can we tell if electrons have been transferred?

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Keeping Track of Electrons: Oxidation States

Assigning Formal Oxidation States (oxidation numbers)

-system for keeping track of electrons -assign “formal” charges to the atoms in a compound/molecule

1. The oxidation number of an atom in its elementalstate is

zero:

eg: Ag, H₂, P₄, Cl₂ , He

2. The oxidation state of a monatomic ion is the same as its charge:

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Assigning Oxidation States

3. Oxygen is (usually) assigned an oxidation state -2

eg: in its covalent compounds: H₂O, CO₂

exceptions: -can be positive (bonded to fluorine): OF₂ -1in peroxides (contain O₂2-_{) :}

H2O2 , Na2O2 -0.5 in superoxides (contain O2-)

NaO2

4. Hydrogen is usually assigned an oxidation state of +1

eg: in its covalent compounds: H2O, NH3, CH4..

exceptions: in metal hydrides (H is -1):

NaH, CaH₂

Assigning Oxidation States

5. For atoms in covalent compounds or polyatomic ions:

-assign formal charge as if the most

electronegative element controls both electrons in a shared pair.

Generally: just treat “as if” these compounds were ionic and assign “charges” and oxidation states based on the previous rules

eg: CO2 - oxygen is more electronegative so treat it as

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Assigning Oxidation States

6. The sum of the oxidation states must equal: -zero, for an electrically neutral compound or -the overall charge, for an ionic species.

Examples: assign oxidation states to each atom

CO₃2- _KMnO

4 CH3OH S2O8 S4O62- HCN

Identifying a Reaction as a Redox Reaction

ƒ Transfer of electrons must take place

ƒ Look for changes in oxidation states of atoms involved in the reaction

ƒ Something must gain electrons: this is reduction and will show a decrease in oxidation state

Oxidizing and Reducing Agents

Na_(s) + Cl_2(g) → NaCl_(s)

Na is oxidized

- supplies electrons to Cl -increasein oxidation number from 0 to +1

- Na is a reducing agent

Cl is reduced

- gains the electrons supplied by Na

-decreasein oxidation number from 0 to –1

- Cl₂is an oxidizing agent

ƒ Reducing Agent

-is or contains the element being oxidized

-contains the atom that shows an increasein oxidation number

ƒ Oxidizing Agent

-is or contains the element being reduced

-contains the atom that shows a decreasein oxidation number

Balancing Oxidation Reduction Reactions

ƒ often too complex to balance by inspection ƒ take advantage of fact that

# of electrons lost = # of electrons gained

Two methods

1) Oxidation state method (for all redox equations, best for non aqueous reactions)

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Oxidation State Method

For Aqueous and Non Aqueous reactions 1. assign oxidation states to all atoms

2. determine which element is oxidized and its increase in oxidation state

3. determine which atom is reduced and its decrease in oxidation state

4. choose coefficients for the species containing the atom oxidized and the atom reduced such that the total increase in oxidation state equals the total decrease in oxidation state

5. balance any other atoms by inspection without changing the coefficients established in step 4

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In class Example

C₂H_6(g) + O_2(g) → CO_2(g) + H₂O_(g)

Oxidation states:

C₂H₆: H +1 O₂: O 0 CO₂: O -2 H₂O: H +1 C -3 C +4 O -2 Changes in oxidation #

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In class example

As4O6 + Cl2 + H2O → H3AsO4 + HCl

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Balancing Redox Reactions

ƒ Half Reaction Method

– for redox reactions occurring in aqueous solution – write separate equations for the oxidation process

and the reduction process

– balance these individually and then add them together for the overall balanced equation

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Reactions Occurring in Acidic Solution

ƒ Write the individual oxidation and reduction half reactions

ƒ For each half reaction:

– balance all elements except hydrogen and oxygen – balance oxygen by adding H₂O

– balance hydrogen by adding H+

– balance the charge by adding electrons (e-₎

ƒ Where applicable, multiply one or both balanced half reactions by an appropriate integer so that the # of electrons lost = # of electrons gained

ƒ Add the two half reactions and cancel out equivalent species on both sides of the equation

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Example

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in acidic solution

Br - _{+ MnO}

4- → Br2(l) + Mn2+ -1 +7 -2 0 +2

10 Br -_{+ 2MnO}

4-+ 16H+ +10e-→5 Br2 + 10 e- + 2Mn2+ + 8H2O

10Br - _{+ 2MnO}

4-+ 16H+ →5 Br2 + 2Mn2+ + 8H2O

oxidation half reaction reduction half reaction

Reactions Occurring in Basic Solution

ƒ Balance each half reaction following the same steps as for acidic solutions

ƒ Add the equations and eliminate common species on both sides

ƒ To each sideof the equation add one OH- _{for each H}+

in the equation

ƒ Combine H+ _{and OH}- _{ions, on the same side, to form} H₂O

ƒ Where possible, eliminate equal numbers of H₂O when they appear on both sides

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Example

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in basic solution

MnO4- + S2- → MnS(s) + S(s)

+7 -2 -2 +2 -2 0

5S2-_{+ 2MnO}

4-+ 2S2- + 16H+ + 10 e- → 5S + 10e- + 2MnS + 8H2O

7S2-_{+ 2MnO}

4-+ 16H+ → 5S(s) + 2MnS(s) + 8H2O

oxidation half reaction reduction half reaction

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Basic solution

7S2-_{+ 2MnO}

4-+ 16H+ → 5S(s) + 2MnS(s) + 8H2O

add OH-_:

combine OH- _{and H}+ _{to form water:}

if possible, eliminate equal numbers of waters from both sides

7S2- _{+ 2MnO}

4- + 8 H2O → 5S(s) + 2MnS(s) + 16OH

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When given an equation to balance:

ƒ check to see if the equation represents an oxidation reduction reaction (look for a change in oxidation state) ƒ if it is not a redox equation then balance by inspection ƒ if it is a redox equation use one of the redox balancing

methods

which method? - (these are guidelines only not rules) – if all or most of the reactants and products are in

molecular form use the oxidation state method

– if all or most of the reactants and products are (aq) or in ionic form use the half reaction method

– if it is specified that the reaction occurs in either acidic or basic solution, use the half reaction method

– if oxygen or hydrogen appear on one side but not on the other, use the half reaction method

Redox Titrations

eg: Addition of an oxidizing agent of known concentration to a solution of a reducing agent of unknown concentration

MnO₄- _{+ 8H}+ _{+ 5e}- → _Mn2+ _{+ 4H} disappears with the formation of Mn2+_{. When there is no more} Fe2+_{left in the solution, then one extra drop of KMnO}

4 will not be reduced and the solution will turn slightly pink.