Chapter 20: Oxidation -Reduction reactions Section 20.1 - The meaning of oxidation and reduction
Why does this happen to cars? Why is it less of an issue in Arizona compared to other states?
Has the statue of liberty always been green?
Oxidation reactions Oxidation reactions were originally thought to be reactions that involved the combination of an element with oxygen to produce an oxide This happens in a number of everyday reactions Whenever something is burned Whenever something corrodes (rusts) Whenever you bleach a substance When hydrogen peroxide decomposes
Reduction reactions Reduction reactions are the opposite of oxidation Originally, this was believed to signify simply the loss of oxygen from a compound That is a good rule of thumb, but is not always the case A common example is the reduction of iron ore Oxygen is removed, iron ore and carbon dioxide are formed This occurs when iron ore and carbon are heated together
The relationship between oxidation and reduction These two processes always happen together Oxidation does NOT happen without reduction, and reduction does not happen without oxidation In the previous example, iron(iii) oxide is reduced and the carbon is oxidized These complementary oxidation-reduction reactions are commonly known as redox reactions
Does oxidation always require oxygen? No! Although originally scientists thought that oxygen was necessary, now we know that instead a transfer of electrons is key in this style of reaction. This is because oxygen is highly electronegative OILRIG mnemonic
What does this look like in terms of electron transfer? When a metal and a nonmetal are reacted together electrons are transferred from the metal to the nonmetal When you react magnesium and sulfur, two electrons are transferred from a magnesium atom to a sulfur atom The magnesium atoms are more stable by the loss of electrons, and the sulfur is more stable by the gain of electron Magnesium is oxidised and sulfur is reduced
Oxidising and reducing agents A reducing agent is a substance that undergoes oxidation and loses electrons Magnesium is the reducing agent in the prior reaction An oxidising agent is a substance that undergoes reduction and gains electrons Sulfur is the oxidising agent in the prior reaction
Write in your notes and fill in the blanks. In a chemical reaction, when a substance electrons it undergoes and is also called the agent.
Practice identifying oxidizing and reducing agent Page 662 You need to identify the reactants and the products Identify the changes in charges of substances over the course of the reaction Establish which substances have gained electrons, and which substances lost electrons
In covalent compounds, it is not so easy.. When a metal and a nonmetal react and form ions, it is easy to identify the transfer of electrons In covalent compounds electrons do not transfer completely In covalent compounds electrons are shared We can refer to partial loss and partial gains of electrons In polar molecules, such as water, the electrons are pulled towards the oxygen - why? The hydrogen undergoes partial loss of electrons and is oxidized
Definitions of oxidation and reduction Processes Leading to Oxidation and Reduction Oxidation Reduction Complete loss of electrons (ionic reactions) Complete gain of electrons (ionic reactions) Shift of electrons away from an atom in a covalent bond Shift of electrons toward an atom in a covalent bond Gain of oxygen Loss of oxygen Loss of hydrogen by a covalent compound Gain of hydrogen by a covalent compound Increase in oxidation number Decrease in oxidation number
Why is it so important to know about oxidation and reduction reactions? Corrosion is an example of a redox reaction Each year corrosion costs the american economy billions of dollars Iron is a common example Atoms are oxidized by oxygen to form ions of iron, a process that speeds up in the presence of water In this process the oxygen is reduced to oxide ion What happens to these ions? Process speeds up in the presence of salts
Iron can be oxidized to form iron hydroxide or Iron Oxide The oxide ions can react with either Iron ions, (formed from oxidation) to form Fe2O3, or can react with hydroxide ions to form Iron Hydroxide 2Fe(s) + O2(g) + 2H2O(l) 2Fe(OH)2(s) 4Fe(OH)2(s) + O2(g) + 2H2O(l) 4Fe(OH)3(s) This process speeds up in the presence of salts and acids, as they speed up the conductive process and facilitate electron transfer
Do all metals corrode at the same rate? Some metals have an inherent resistance to corrosion Their electrons are tightly held, meaning it is hard for them to be oxidized Some metals oxidize very quickly - and form an oxide coating Example - gold and platinum, these are called noble metals This protects the aluminium object from further corrosion Iron can also form a coating, but this is not tightly packed together, and the metal can still be damaged
Iron(III) oxide Oxygen Water Aluminum oxide Oxygen Water
Why are zinc blocks attached to the hulls of ships The zinc is used a sacrificial metal on the ships to save the iron on the ship When oxygen and water attack the iron, the iron atom loses electrons Zinc and magnesium are better reducing agents than iron, and immediately transfer electrons back to the iron atoms Can also be used for bridges, pipeline, storage tanks...
How else can corrosion be prevented? Coating the surface This could be with oil, paint, plastic, or another metal Air and water are kept away from the metal If the coating is scratched, or removed at all, the exposed metal will begin to corrode
Key Points Oxidation and reduction are complementary reactions involving the loss/gain of electrons An oxidizing agent gains electrons and is reduced A reducing agent loses electrons and is oxidized Corrosion is a common consequence of redox reactions Corrosion can be prevented through coating the surface, or the use of sacrificial metals
Oxidation numbers Section 20.2
How do you express the degree of oxidation or reduction An oxidation number is a positive or negative number assigned to an atom to indicate the degree of oxidation or reduction A number of rules exist surrounding the determination of oxidation numbers A bonded atoms oxidation number is the charge that it would have if the electrons in the bond were assigned to the atom of the more electronegative element
Oxidation numbers and ionic compounds In ionic compounds the oxidation numbers will equal their ionic charge Remember the overall charge of an ionic compound is 0 Metals will always have a positive charge Non metals will always have a negative charge Polyatomic ions can have a positive or negative charge
Oxidation numbers and molecular compounds In molecular compounds, due to the covalent bond there is no ionic charges associated with atoms But, oxygen is reduced when hydrogen and oxygen form water Oxygen is a highly electronegative element, much more so than in the hydrogen The two shared electrons are shifted towards the oxygen In terms of oxidation number, assume that the electrons have been transferred Hydrogen would be +1, Oxygen would be +2
Further rules for oxidation numbers Many elements can have several oxidation numbers In particular, you can use neutral compounds and polyatomic ions Some elements can also have multiple oxidation numbers, such as Cr Chromium metal - oxidation number 0 Potassium dichromate = K2Cr2O7 - oxidation number +6 Chromium (III) potassium sulfate decahydrate - CrK(SO4)2.12H2O - oxidation np. 3
Summary of rules
What is the oxidation number of each kind of atom in the following ions and compounds? a. SO2 c. Na2SO4 b. CO32 d. (NH4)2S
How does this all relate to chemical reactions During redox chemical reactions, the oxidation number of elements may change If copper reacts with silver nitrate, the oxidation number of silver decreases from +1 to 0 Copper ions are also oxidized Oxygen and reduction can be defined in terms of changes in oxidation number
The relationship between oxidation number and gemstones Gemstones get their color from impurities Eg. Iron, chromium, copper The oxidation number of the metal can also change the color of the gemstone Example - iron in beryl
Describing redox reactions Section 20.3
Identifying redox reactions Reactions are either a REDOX reaction or not All reactions fall into one of these two categories REDOX: Single replacement, combination, decomposition, and combustion NON redox - double replacement or acid - base reactiosn
Using oxidation numbers to identify redox reactions In a lightning storm, oxygen and nitrogen molecules react to form nitrogen monoxide N2(g) + O2(g) -> 2NO(g) Is this a redox reaction? Yes! The oxidation number of nitrogen increase to 2, the number of oxygen decreases to -2
Other indicators of redox reactions Color changes often signify that a redox reaction is taking place
Balancing redox equations Redox reactions can be complex, and often cannot be balanced by trial and error The fact that the number of electrons gained in reduction is equal to the number of electrons lost in oxidation can be used to balance an equation Two methods: Oxidation number changes Using half equations
Using oxidation number changes In this method you compare the changes in oxidation number 1) 2) 3) 4) 5) 6) Start with an unbalanced equation Assign oxidation numbers to all atoms in the equation Identify which atoms are being oxidized and which are being reduced Use bracketing lines to connect atoms that undergo oxidation and those that undergo reduction Use coefficients to make the total increase in oxidation number equal the total decrease in oxidation number Make sure the atom is balanced for both atoms and charge Worked example - 1) Fe2O3, 2) sample problem 20.5
Half reactions A half reaction is an equation that shows just the oxidation or just the reduction that takes place in a redox reaction In this method you write and balance the oxidation half reaction, then you write and balance the reduction half reaction In the example: S(s) + HNO3 -> SO2(g) + NO(g) + H2O (l) S(s) + -> SO2(g) Oxidation half reaction NO3-(aq) -> NO(g) Reduction half reaction
Balancing half reactions 1) 2) S(s) + -> SO2(g) Oxidation half reaction NO3-(aq) -> NO(g) Reduction half reaction In equation 1 S is being oxidised as 4 electrons are being removed In equation 2 N is being reduced - charge changes from +5 to +2 Atoms must be balanced in each half equation, and then the charges so that electrons gained equal electrons lost When each half reaction is balanced you combine them to form a balanced chemical equation