Oxidation-Reduction Reactions PDF

MED-102 General Chemistry Oxidation-Reduction Reactions  LOBs covered Describe the concepts of oxidation and reduction Calculate oxidation numbers in chemical formulas Identify the oxidizing and reducing agents in a redox reaction Predict reaction reactivity on the...

MED-102 General Chemistry Oxidation-Reduction Reactions  LOBs covered Describe the concepts of oxidation and reduction Calculate oxidation numbers in chemical formulas Identify the oxidizing and reducing agents in a redox reaction Predict reaction reactivity on the basis of the Activity Series  Oxidation-Reduction (Redox) Reactions Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g) This is the reaction where solid magnesium metal is placed in a solution of hydrochloric acid. The magnesium dissolves in the acid, and bubbles of a colorless gas (hydrogen) are evolved. Why is this a redox reaction? It is because magnesium loses two electrons as in: Mg(s) → Mg2+(aq) + 2e- (loss of electrons = oxidation) and the hydrogen ions from hydrochloric acid gain electrons to become hydrogen gas: 2 H+(aq) + 2e- → H2(g) (gain of electrons = reduction) WATCH: https://www.youtube.com/watch?v=RX6rh-eeflM  Assigning oxidation numbers (1) An atom in its elemental form has an oxidation number of 0. (2) An atom in a monatomic ion has an oxidation number equal to its charge. (3) An atom in a polyatomic ion or in a molecular compound has the same oxidation number it would have if it were a monatomic ion Hydrogen can be either +1 or -1. When bonded to a metal, H is -1, when bonded to a nonmetal, H is +1. Oxygen is usually -2, except in peroxides, which contain the O22- anion, where oxygen is -1, or in an O – O covalent bond in a molecular compound. Halogens are usually -1, except when bonded to oxygen or fluorine, in which case they have positive oxidation numbers. (4) The sum of oxidation numbers is 0 for neutral compounds, and is equal to the total charge for a polyatomic ion. Also, watch the entire video from the previous slide  Exercises ACTIVITY Assign oxidation numbers to all elements in the following: (a) CdS (b) AlH3 (c) S2O32- (d) Na2Cr2O7 (e) FeSO4  Exercises - Answers (a) CdS (b) AlH3 (c) S2O32- (d) Na2Cr2O7 (e) FeSO4 (a) CdS – S is in the same group as O, Group 6. Therefore, it forms S2- ions. This means that Cd must be Cd2+. (b) AlH3 – H is bonded to a metal (Al), therefore, it must have an oxidation number of -1. Three H’s make -3, so the Al must be Al3+. (c) S2O32- - The O’s are -2, so they give a total of -6. The total charge on the polyatomic ion is -2, which means that the two S’s must add up to +4. So, each S is +2. (d) Na2Cr2O7 – The Na is always +1, so two of them make +2. This means that the Cr2O7 polyatomic ion must have a charge of -2. Seven O’s give -14. Since the total ionic charge is -2, the two Cr’s must add up to +12. Therefore, each Cr is +6. (e) FeSO4 – Here, we need to know that the sulphate anion, SO4, carries a total charge of -2. Therefore, the iron is +2. Four O’s make -8. Since the total charge in the sulphate anion is -2, the sulphur must have a charge of +6.  Identifying Redox Reactions Using the concept of oxidation numbers we can easily identify which element is oxidized and which element is reduced. Here is an example: H2(g) + ½ O2(g) → H2O(l) 0 0 +1 -2 Thus, hydrogen goes from an oxidation number of 0 to +1, by losing one electron for each H atom (2 electrons lost in total), and oxygen goes from an oxidation number of 0 to -2, a gain of 2 electrons. Hydrogen is therefore oxidized, and oxygen is reduced. More examples  One more example The substance that causes a reduction is called the reducing agent The substance that causes an oxidation is called the oxidizing agent. Here is an example: Also remember: the reducing agent is oxidized, and the oxidizing agent is reduced  Activity Series The higher a metal is in the Activity Series, the more chemically reactive it is. The more reactive a metal is, the easier it loses electrons.  Another example An example of a redox reaction we have already seen is: Mg(s) + 2 H+(aq) → Mg2+(aq) + H2(g) Here, placing magnesium in an acid solution will result in a reaction where the magnesium will all dissolve in the acid. This reaction happens spontaneously. However, if we tried the same type of reaction with gold metal, no reaction will take place.  Activity Series We see that Mg is higher than hydrogen, H2, in the Activity Series. Any metal higher than H2 is the Activity Series will react with an acid. Gold (Au), a very unreactive metal, is lower than H2 in the Activity Series and therefore does not react with an acid.  Summary for Revision In some chemical reactions, there is exchange of electrons between the reactants. These reactions are known as oxidation-reduction reactions, or redox reactions. Oxidation is defined as the loss of electrons. Reduction is defined as the gain of electrons. In covalent compounds, the exchange of electrons is more difficult to visualize. For this purpose, we use the concept of oxidation numbers. There are rules for assigning oxidation numbers. Aside from identifying the elements oxidized and reduced, we can also identify the substances that carry out the oxidation and the reduction. The reducing agent carries out a reduction, and is oxidized in the reaction. The oxidizing agent carries out an oxidation, and is reduced in the reaction. The Activity Series places metals according to their reactivity. At the top we have more reactive metals, and near the bottom we have metals that do not oxidize (rust) easily. Metals near the top of the Activity Series prefer to exist as cations, whereas those lower down prefer to stay in metallic form. Precious metals, such as gold, silver and copper, often used in making jewelry, are near the bottom of the Activity Series, and do not easily rust. This explains their selection in making jewelry, and their relatively high price. Metals that are above H2 is the Activity Series will react spontaneously with an acid. Those below H2 will not.