REDOX REACTIONS REACTIONS These reactions are all around you and they directly affect and influence your everyday life!

What are the uses of redox reactions? Extraction of metals from their ores. Welding (Acetylene torches). Driving our cars (burning gas). Keeping the swimming pool clean. Bleaching unwanted colors out of our clothes or hair. Burning food in our bodies. Antioxidants, eg. Vitamin E – prevents fats from going rancid. Photosynthesis

When discussing REDOX reactions it is useful to assign oxidation numbers to elements undergoing a transfer of electrons. Therefore, you must understand oxidation states (oxidation numbers).

Oxidation Numbers The oxidation number is a bookkeeping device of the numbers of electrons gained or lost in an atom when it combines with other atoms to form a compound. The possible oxidation numbers of an atom can be derived from its ground state electron configuration. The actual oxidation number of an atom depends on the other atoms in the compound.

Assigning Oxidation Numbers 1. Any free (unattached) element with no charge has the oxidation number of zero. Diatomic gases such as O 2 and H 2 are also in this category. Ag 0 Cu 0 Cl 2 0 Na 0 2. All compounds have a net oxidation state of zero. The oxidation number of all of the atoms add up to zero NaCl =0

3. Any ion has the oxidation state that is the charge of that ion. The ions of elements in Group I (Alkali metals), II (alkaline earth metals), and VII (halogens) and some other element have only one likely oxidation state. Ca +2 F -1 K Polyatomic ions have an oxidation state for the whole ion which is the charge on that ion. The sum of the oxidation numbers of all the atoms that make up the ion equal the charge on the polyatomic ion ( SO 4 ) = -2

5. Oxygen in compound has an oxidation state of minus two, except for oxygen as peroxide, which is minus one. sodium oxide sodium peroxide Na 2 O NaO = = 0

6. Hydrogen in compound has an oxidation state of plus one, except for hydrogen as hydride, which is minus one. Hydrogen sulfide Calcium hydride H 2 S Ca H = = 0

Let’s Assign Oxidation Numbers! Potassium perchlorate KClO 4 Potassium in a Group I ion K +1 Oxygen in a compound is -2 O -2 You must determine the oxidation state of Cl. The sum of oxidation numbers in a compound equal zero, therefore +1 x -2 KClO (x) - 8 =0 The oxidation state of Cl is +7. K +1 Cl +7 O 4 -2

One more time. P 2 O 5 x -2 P 2 O 5 2(x) -10 = 0 x= +5 P 2 +5 O 5 -2

Your turn! Complete the worksheet of assigning oxidation numbers. Check your answers The blue colors are a positive oxidation state, the red colors are a negative oxidation state, relevant to the atoms underneath them.

1. N 2 O6. Ca 3 (PO 4 ) 2 2. LiH7. Ag 3. H 2 S8. CuCl 2 4. FeSO 4 9. H 2 5. NaBrO LiMnO 4

1. N 2 O6. Ca 3 (PO 4 ) 2 2. LiH7. Ag 3. H 2 S8. CuCl 2 4. FeSO 4 9. H 2 5. NaBrO LiMnO 4

N 2 O6. Ca 3 (PO 4 ) LiH7. Ag H 2 S8. CuCl FeSO 4 9. H NaBrO LiMnO 4

What are Redox Reactions? Redox reactions are a family of reactions that are concerned with the transfer of electrons between species Redox reactions are a matched set -- you don't have an oxidation reaction without a reduction reaction happening at the same time. Oxidation refers to the loss of electrons, while reduction refers to the gain of electrons.

Try this! Place a piece of copper wire, shaped like a Christmas tree, in a solution of 0.2M silver nitrate Cu + AgNO 3  What happens?

Notice the crystals on the copper wire and the blue solution. The Cu (s) loses electrons to become Cu +2 (aq) ions (causing the blue color in the solution) and the Ag +1 (aq) ions gain electrons to become Ag (s) (the crystals on the wire).

An increase in oxidation number is Oxidation. This is due to the loss of electrons 0 +2 Cu(s)  Cu (aq) + 2e - A decrease in oxidation number is called reduction. This is due to gaining of electrons Ag (aq) + e 1  Ag (s)

Oxidation reactions occur when a species loses electrons and the charge is more positive. Reduction reactions occur when a species gains electrons and the charge is more negative. A mnemonic to remember this is: OIL RIG Oxidation Is Loss of electrons. Reduction Is Gain of electrons.

Example: Production of Iron Fe 2 O 3 (s) + 3CO(g)  Fe(l) + 3CO 2 (g) Fe went from +3 to 0 C went from +2 to +4 Iron was reduced Carbon was oxidized

Sparkler Burning Mg(s) + O 2 (g) → 2MgO(s) Mg went from 0 to +2 O went from 0 to -2 Oxygen was reduced Magnesium was oxidized

The species that is being oxidized reduces the other reactant. It is called the reducing agent. The species that is being reduced oxidizes the other reactant. It is called the oxidizing agent or oxidant.

When is it a REDOX reaction? A single replacement reaction is always asingle replacement redox reaction because it involves an element that becomes incorporated into a compound and an element in the compound being released as a free element. A double replacement reaction usually is not a redox reaction.double replacement If a compound divides into elements in a decomposition, a decomposition reaction could be a redox reaction.decomposition reaction Synthesis reactions are also redox reactions if there is an exchange of electrons to make an ionic bond. Synthesis

How do you identify a REDOX reaction? 1. Assign oxidation numbers to all atoms. 2. Look for any element that changed oxidation states. The element that became more negative was reduced (reduction) and the element that became more positive was oxidized (oxidation). 3. The oxidized reactant is the reducing agent and the reduced reactant is the oxidizing agent. 4. Write half reactions for both the oxidation and reduction reactions.

The reaction between magnesium metal and oxygen to form magnesium oxide. Mg (s) + Cl 2 (g)  MgCl 2 1. Assign oxidation numbers to all atoms in the equation Mg (s) + Cl 2 (g)  MgCl 2 Oxidation Reduction 2. Identify both the reduction and oxidation

3. Write the half reaction for the reduction reaction showing the reactant gaining electrons to give the new charge. Cl e -  2 Cl -1 Reduction Oxidizing agent 4. Write the half reaction for the oxidation reaction showing the electrons lost as products. Mg o  Mg e - Oxidation reducing agent The number of electrons lost must equal the number of electrons gained.

Now its your turn. Identify oxidation, reduction, the oxidizing agent and the reducing agent in the following Al + 3 Br 2  2AlBr Al + Fe 2 O 3  Al 2 O Fe 3. 2 Na + MgCl 2  2 NaCl + Mg 4. CuO(s) + H 2 (g)  Cu(s) + H 2 O(g) 5. CoCl 2  Co + Cl 2

1. 2 Al Br 2 0  2Al +3 Br 3 -1 Al went from 0 to +3: it is Oxidized and the 2 Al 0  2 Al e - reducing agent Br 2 went from 0 to -1: it is Reduced and the 3 Br e -  2 Br 3 -1 oxidizing agent

2. 2 Al 0 + Fe 2 +3 O 3 -2  Al 2 +3 O Fe 0 Al went from 0 to +3: It is Oxidized and the 2 Al 0 +  2 Al e - reducing agent Fe went form +3 to 0: It is Reduced and the 2 Fe e -  2 Fe 0 oxidizing agent

3. 2 Na 0 + Mg +2 Cl 2 -1  2 Na +1 Cl -1 + Mg 0 Na went from 0 to +1: It is Oxidized and the 2 Na 0  2 Na e -1 reducing agent Mg went from +2 to 0: It is Reduced Mg e -  Mg 0 and the oxidizing agent 4. Cu +2 O -2 + H 2 0  Cu 0 + H 2 +1 O -2 Cu went from +2 to 0: It is Reduced and the Cu e -  Cu 0 oxidizing agent H 2 went from 0 to +1: It is Oxidized and the H 2 0  2 H e - reducing agent

5. Co +2 Cl 2 -1  Co 0 + Cl 2 0 Co went from +2 to 0: It is Reduced and the Co e -  Co 0+ oxidizing agent Cl went from -1 to 0: It is Oxidized and the 2 Cl -1  Cl e - reducing agent

Corrosion Corrosion is a reaction that involves action of an oxidizing agent on a metal. The oxidizing agent is often oxygen dissolved in water. Iron rusting is an example of corrosion Fe (s) + 2 H 2 O(l)  H 2 (g) + Fe(OH) 2 (s) 4 Fe(OH) 2 (s) + O 2 (g)  2 Fe 2 O H 2 O red-brown Rust

Corrosion The Statue of Liberty is a prime example of the corrosion caused by reduction/oxidation. The metal oxidizes which makes it the reducing agent.

1% of the world’s economy is used in the prevention of corrosion. What is done? 1. Paint or grease. This prevents water or oxygen reaching the iron. However, this is only a temporary step. It is the cheapest method or prevention. 2. Plastic. Coating the iron with plastic in a method used for garden chairs. 3. Galvanizing. Coating iron with a metal that oxidizes more readily. Zinc is used to galvanize iron. 4. Chromium plating. Chromium is a more reactive metal than iron and old model cars have chrome plated bumpers.

Electrochemistry Now we will take these concepts a step further and discuss how various chemical species vary in their ability to "pull" electrons—that is, to be reduced. The implications of the different pulling powers of various species are far reaching. As a result of this phenomenon, we are able to power batteries, produce aluminum, extract metals from their salts, and protect metals through electroplating.

Batteries: A voltaic or galvanic cell When a stick of zinc (Zn) is inserted in a salt solution, there is a tendency for Zn to loose electron according to the reaction, Zn  Zn e -. The arrangement of a Zn electrode in a solution containing Zn +2 ions is a half cell, which is usually represented by the notation: Zn | Zn +2

Similarly, when a stick of copper (Cu) is inserted in a copper salt solution, there is also a tendency for Cu to loose electron according to the reaction, Cu  Cu e - This is another half cell: Cu | Cu +2.

The tendency for Zn to loose electron is stronger than that for copper. When the two cells are connected by a salt bridge and an electric conductor as shown to form a closed circuit for electrons and ions to flow, copper ions (Cu +2 ) actually gains electron to become copper metal. This arrangement is called a galvanic cell or battery as shown here. In a text form, this battery is represented by, Zn | Zn +2 || Cu +2 | Cu Oxidation Reduction

Electroplating Electroplating is widely used to coat one metal with a thin layer of another metal. This is carried out either to make an object appear more decorative, as when gold or silver is plated onto nickel, or to protect an object from corrosion as when chromium is plated on to iron. The object to be plated is made the cathode and the anode is usually made from the metal that will be used to coat the cathode. The electrolyte must contain a salt of the coating metal.