Aim Redox 1 – Why is redox so important in your life?

Moving Electrons and You• Cell phones require electricity

• Electricity – the movement of electrons

• Batteries in cell phones are composed of metals and other electrolytes that allow electrons to flow through the cell phone

Oxidation States

• Remember – electron movements are also why we have bonding in many compounds

• If you lose an electron, – you lose negative charge– You become positive or more positive (less negative)– Metals tend to do this– Example: if sodium loses an electron:

Na0 Na+1 + 1 e-

• If you gain an electron, – you gain negative charge– You become negative or more negative (less positive)– Nonmetals tend to do this– Example: if chlorine gains an electron:

Cl0 + 1 e- Cl-1

Oxidation StatesA Review of the Rules of Oxidation States or

Numbers• Free elements are always zero: Fe, Cl2 , Ca • Ions are the charge assigned to them: Fe 3+ Ca 2+

• polyatomic ions (Table E): NO3- SO4

2-

• Some elements have only one oxidation state: Group 1 and 2

• Some elements usually have a particular oxidation state with exceptions:

• oxygen has a -2 oxidation state • Hydrogen has a +1 oxidation state • Compounds’ oxidation numbers always add up to zero

• In chemical reactions, each element in each compound has an oxidation state

• Examples

3H2 + N2 2NH3

H = 0 N = 0 N = -3, H = +1• Who loses electrons? H0 H+1 + 1 e-• Who gains electrons? N0 + 3 e- N-3

Cu(NO3)2 + Mg Mg(NO3)2 + Cu

Cu = +2 Mg = 0 Mg = +2 Cu = 0

N = +4 N = +4

0 = -2 O = -2• Who loses electrons? Mg0 Mg+2 + 2 e-• Who gains electrons? Cu+2 + 2 e- Cu0

Oxidation-Reduction Reactions (Redox)

In a redox reaction: • Some substances are lose electrons

– They are OXIDIZED– Oxidation: the loss of electrons, become more

positive– LEO – Lose Electrons: Oxidation

• Some substances are gain of electrons– They are REDUCED– Reduction: the gain of electrons, become more

negative– GER – Gain Electrons: Reduction

Identifying Redox Rxns• Redox - another type of reaction

– Elements in the compound change oxidation states

– Example: N2 + 3H2 2NH3

0 0 (-3)(+1)

Nitrogen gains e-, becomes more negative

Hydrogen loses e- becomes more positive

– Example: 2H2O O2 + 2H2

(+1)(-2) 0 0

Hydrogen gains e-, becomes more negative

Oxygen loses e- becomes more positive

Half Reactions in Redox Reactions• Half reaction - reactions that show either a gain

or loss of electrons

• Example 1: What are the half reactions?

2Mg(s) + O2(g) 2MgO(s)

– Oxidation half reaction:

LEO – someone loses electrons

2Mg0 2Mg2+ + 4e-

– Reduction half reaction:

GER – someone gains electrons

O20 + 4e- 2O2-

Half Reactions in Redox Reactions

Example 2:

3Cu(s) + 2Al(NO3)3(aq) 3Cu(NO3)2(aq) + 2Al(s)

• Oxidation half reaction:

LEO – someone loses electrons

Cu0 2Cu 2+ + 2 e-

•Reduction half reaction:

GER – someone gains electrons

Al+3 + 3e- Al0

What are the half reactions in each?

1. N2 + 3 H2 2 NH3

Reduction: N0 + 3 e- N -3

Oxidation: H0 H+1 + e-

2. Mg + 2 HCl H2 + MgCl2Oxidation: Mg0 Mg+2 + 2e-

Reduction: H+1 + e- H0

3. 2 H2O 2 H2 + O2

Oxidation: O-2 O0 + 2 e- Reduction: H+1 + e- H0

4. CH4 + 2O2 CO2 + 2 H2OReduction: O0 + 2 e- O -2

Oxidation: C-4 C+4 + 4e-

Aim Redox 2 – What metals make a better battery?

Metal Reactivity • Table J – shows the activity of

metals and nonmetals in terms of moving electrons

• Active metals can give up electrons to less active metals

• Related to the position on Table J– Higher metals oxidize (lose

electrons) more easily – to lower ones (which gain the

electrons and are reduced)

Metal Reactivity

• Examples: which metal is oxidized/reduced in each of the following pairs:

1. Zn oxidized and Cu reduced

2. K oxidized and Mg reduced

3. Cu reduced and Fe oxidized

Acid reactivity can also be

determined from Table J

• Metals above hydrogen will react with acids to produce H2 gas

• Which metals don’t react with acids?

Why is Table J important?

1. It describes corrosion

• Active metals corrode easily

• Corrosion: loss of metallic properties due to action of air, water, and chemicals

• Two examples:

• Rust: 4Fe + 3O2 2Fe2O3

• Aluminum foil: forms a thin coating of aluminum oxide that protects the inner aluminum

Why is Table J important?2. It explains how electrons

can flow in electrochemical cells (aka voltaic cells or batteries)

• metals will react spontaneously to move electrons between them

• More active metals are oxidized• Less active metals are reduced• The movement of electrons

between the two metals creates an electric current

Why is Table J important?

3. It describes how energy can be added to metals in various electrolytic cells

• Electrolysis – the separation of elements in a compound using electricity

• Example: 2NaCl 2Na0 + Cl20

• Oxidation: 2Na+1 +2e- 2Na0

• Reduction: 2Cl-1 Cl20 + 2e-

Why is Table J important?• Hydrolysis – the separation of

hydrogen and oxygen from water using electricity

• Example: 2H2O 2H20 + O2

0

• Oxidation: 2O-2 O20 + 4e-

• Reduction: 4H+1 + 4e- 2H20

• Electroplating – the coating of one metal on to another through electron movement

• Requires a power source

• Example: Ag+1 + 1e- Ag0

• Spontaneous reactions – Movement of electrons from an

active metal to a less active one– Produces electricity and electron

flow– Batteries

• Non-spontaneous reactions– Movement of electrons from a

less active metal to a more active one

– Requires electricity and electron flow

– Battery chargers, electrolytic cells, hydrolysis

• Agents of Redox • Reducing agents

– lose electrons to other atoms– Are themselves oxidized

• Oxidizing agents – gain electrons from other atoms– Are themselves reduced

• In the rxn,

CH4 + 2O2 CO2 + 2 H2OCarbon is being oxidized: C-4 C+4 + 8e-

Therefore, it is the reducing agent (gives up e-)Oxygen is being reduced: O0 + 2 e- O -2

Therefore, it is the oxidizing agent (takes up e-)

Aim Redox 3c – What is electrochemistry and how can it help us build a better battery?

• During a single replacement reaction, more active metals transfer electrons to less active metals– the more active

metal is oxidized– 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

Electrochemical cells

Electrochemical cells - Parts • Half cells - separate containers in which each

half reaction occurs electrodes• Anode – the metal electrode where oxidation

occurs (this metal is oxidized, lose e-) (AN OX) • Cathode – the metal electrode where reduction

occurs (this metal is reduced, gains e-) (RED CAT)

• U-tube or salt bridge - lets ions travel between half cells to complete the circuit

• Electrolyte - Carries the current in its ions• Voltmeter – measures the voltage or flow of

electricity from the elechemical cell

Electrochemical Cell Parts

Electrochemical cells - Function

• In each half cell - one half of the redox reaction occurs– Oxidation of the metal at the anode breaks down

the metal into ions as it loses electrons– Reduction of the metal at the cathode adds

electrons to the metal ions in solution, making them solid metal

– Electrons flow from the oxidation side, across the wire, through the electrical device, and to the reduction side

– This continues until the flow of electrons stops– This occurs when the cells reach equilibrium

Examples of electrochemical cells

• Voltaic Cells (Spontaneous Rxns) or Batteries

– Definition - a system that uses a chemical reaction to produce electricity

– The cathode where reduction occurs is at the positive electrode (CPR)

– The anode where oxidation occurs is at the negative electrode (ANO)

Examples of electrochemical cells

• Automobile battery or lead acid storage battery

• The reaction is as follows:

2H2SO4 + Pb + PbO2 2PbSO4 + 2H2O + energy

Pb0 Pb 4+ Pb 2+

• Half reactions:

– Oxidation at anode: Pb0 Pb 2+ + 2 e-

– Reduction at cathode: Pb 4+ +2e- Pb2+

• The reverse reaction is how we recharge the battery:

2PbSO4 + 2H2O + energy 2H2SO4 + Pb + PbO2

Electrolytic cells (Nonspontaneous Rxns)

• Remember – you cannot find alkali metals and halogens in nature as free elements

• They must be separated by electrolysis from

compounds

• Electrolysis of molten sodium chloride

electricity + 2NaCl 2Na0 + Cl20

Examples of Electrochemical Cells

Electrolytic cells • Electroplating

– Formation of a metal layer on to a surface

– Requires addition of energy

– Electricity is the source of electrons

• Example - SilverplatingElectricity + Ag+

Ag0