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 States A 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- SO42- 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? N2 + 3 H2  2 NH3 Reduction: N0 + 3 e-  N -3 Oxidation: H0  H+1 + e- Mg + 2 HCl  H2 + MgCl2 Oxidation: Mg0  Mg+2 + 2e- Reduction: H+1 + e-  H0 2 H2O  2 H2 + O2 Oxidation: O-2  O0 + 2 e- Reduction: H+1 + e-  H0 CH4 + 2O2  CO2 + 2 H2O Reduction: 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)

Zn oxidized and Cu reduced K oxidized and Mg reduced Metal Reactivity Examples: which metal is oxidized/reduced in each of the following pairs: Zn oxidized and Cu reduced K oxidized and Mg reduced 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 + O20 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 H2O Carbon 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?

Electrochemical cells 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 - 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 - Silverplating Electricity + Ag+  Ag0