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Oxidation and Reduction

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1 Oxidation and Reduction
Topic 9 Oxidation and Reduction

2 9.1 Introduction to oxidation and reduction
Oxidation originally, the addition of oxygen or loss/ removal of hydrogen Reduction originally, the loss/ removal of oxygen or addition of hydrogen Redox reactions when both oxidation and reduction occur in a chemical reaction

3 History of Chemistry Joseph Preistley discovered oxygen in 1774 and it was called phlogiston until Antoine Lavoisier Antoine Lavoisier determined the role and name of oxygen in the mid- to late- 1700’s. He also termed oxidation for reactions involving oxygen Justus von Liebig proposed that reduction was the removal of oxygen and addition of hydrogen

4 Oxidation and Reduction now
Refer to loss/ gain of electrons in a chemical reaction Many of the reactions considered as original oxidation and reduction (in reference to oxygen) still fall under the category of redox now

5 Remember!! Oxidation Is Loss Reduction Gain

6 2Mg(s) + O2(g) 2MgO(s) Oxidation: Mg Mg2+ + 2e-
equations 2Mg(s) + O2(g) 2MgO(s) Oxidation: Mg Mg2+ + 2e- Electrons in product Reduction: O2 + 4e- 2O2- Electrons as reactants

7 Half- equations Describe only half of the total reaction
Broken into oxidation half and reduction half When the two halves are combined to represent the full reaction, the equation must be balanced such that the electrons on the reactant side equal the electron on the product side

8 Balancing Half-equations
Na Na+ +e- O2+ 4e-  2O2- Have to multiply the whole half-equation for Na by 4 to gain equal electrons before combining equations Sum: 4Na +O2 + 4e-  4Na+ + 2O2- + 4e- Cancel electrons: 4Na +O2 +  4Na+ + 2O2-

9 Oxidation numbers How the oxidized/ reduced element or compound is determined Rules to determine oxidation state Any uncombined element has an oxidation number of 0 (Ex. O2) A simple ion has an oxidation number equal to its charge For a compound, the oxidation number is 0 For an oxoanion, the number is equal to the sum of the oxidation numbers Hydrogen has an oxidation number of +1, unless combined with a metal, then it will be -1 Oxygen has an oxidation number of -2, except in H2O2 (-1) and OF2 (+2)

10 Covalent bonds The more electronegative element gets the negative oxidation number Less electronegative gets a positive oxidation number Less electronegative is named first in the compound Ex. ClF Cl +1 O.N. F -1 O.N.

11 Positive O.N. implies that the element has ‘lost control’ of electrons
Negative oxidation numbers implies that element has ‘gained control’ of electrons Positive O.N. implies that the element has ‘lost control’ of electrons

12 Naming inorganic compounds
O.N. are used to name ionic inorganic compounds Referred to as stock notation Oxidation number inserted right after name of ion and roman numerals after name/ symbol of element Ex. FeCl2 [Fe2+ 2Cl-] iron(II) chloride FeCl3 [Fe3+ 3Cl-] iron(III) chloride Notation only used for transition metals, tin and lead

13 Identifying redox reactions
Recognized by: Deducing all O.N. of atoms in the molecular, ionic, or half- equation Examine all numbers to see if any O.N. on atoms have changed Increase in O.N. is oxidation, decrease is reduction

14 Redox vs. not 2FeCl2(s) + Cl2(g)  2FeCl3(s)
Oxidation numbers of iron, +2 and +3 O.N. of chlorine, 0 and -1 Redox rxn MgO(s) + 2HCl(aq) MgCl2(aq) +H2O(l) O.N. of all remain the same Not a redox rxn

15 Occurs when a single species is oxidized and reduced simultaneously
disproportionation Occurs when a single species is oxidized and reduced simultaneously Ex. Cl2(g) + H2O(l)  HOCl(aq) + HCl(aq) One Cl decreases from 0 to -1, the other increases from 0 to +1

16 9.2 Redox equations Constructing half-equations
1. write down formulae of reactant and product (for one half) 2. Balance with respect to present ion(s) 3. Balance oxygen with water molecule(s) 4. Balance hydrogen with hydrogen ion(s) 5. Determine charges on both sides of the half-equation 6. Balance charges by adding electrons to side that is more positive Work example on pg. 237 on board

17 For oxidation half-equation, electrons will be on left hand side
Con’t For oxidation half-equation, electrons will be on left hand side For reduction half-equation, electrons will be on right hand side Use a similar process to determine each type of half-equation

18 Forming redox equations
Full redox equations are constructed by combining two half-equations One for the oxidizing agent and one for reducing agent Ex: Work example on pg. 238 on board

19 Redox titrations Similar to acid-base titrations, but has electron transfer instead of a hydrogen transfer Overall equation used for stoichiometric calculations is obtained by combining two half-equations such that, the electrons lost equal electrons gained Ex. Pg. 238 on board

20 Common Oxidizing agents for redox titrations
Acidified manganate (VII) ions Manganate (VII) ions are purple, but reduced form (Mn(II)) almost colorless. Potassium manganate (VII) difficult to prepare and reacts slowly Acidified dichromate (VI) ions Orange in color, reduced form is green (Cr(III)). Can be used as primary standard Iron (III) ions or salts Iodine Red brown in color colorless when reduced Starch improves the color change No indicators necessary for these since a color change is involved Acidified hydrogen peroxide

21 Common reducing agents for redox titrations
Iron (II) salts or ions Ethanedioic (oxalic) acid and ethanedioate (oxalate) ions Performed at 80°C to catalyze Able to determine concentration with high accuracy Hydrogen peroxide In presence of a more powerful oxidizing agent Iodide ions Sodium thiosulfate (VI) or thiosulfate (VI) ions

22 Oxidizing agents A substance that brings about the oxidation of substances by accepting electrons from the substance they oxidize Undergo the process of reduction Common oxidizing agents Oxygen Ozone (O3) Chlorine Acidified potassium manganate (VII) Acidified aqueous potassium dichromate (VI) Acidified hydrogen peroxide Metal ions Hydrogen ions Manganese (IV) oxide

23 Reducing agents A substance that brings about reduction of a substance by donating electrons to the reduced substance Common reducing agents Hydrogen atom Carbon Carbon monoxide Metals More reactive metals

24 The terms oxidizing and reducing agents are relative.
extention The terms oxidizing and reducing agents are relative. A weak reducing agent can be ‘forced’ to become an oxidizing agent in the presence of a more powerful reducing agent

25 Reactions of metals with metal ions in solution
9.3 Reactivity Reactions of metals with metal ions in solution Can test reactivity of metals by placing a piece of Mg metal into a solution of aq metal ions to observe whether the Mg changes color to indicate a chemical rxn Procedure is repeated with different types of metals

26 Activity series With the data collected during experimentation with the metals, a reactivity or activity series can be constructed Such that, the metals are arranged in order of increasing reactivity The type of reactions that take place are displacement reactions The more reactive metal ‘pushes out’ the less reactive metal from its salts Can also be performed by heating two sold metals together, or by reacting with water

27 Using reactivity series
Able to deduce the probability of a redox rxn from a reactivity series Carbon and hydrogen can be included in an activity series Metals above H will displace H from dilute acids, those below will not Metals above carbon cannot be produced by reduction of metal ions, those below can

28 Displacement can occur with non-metals also
Especially the halogens A reactivity series for halogens corresponds to their placement on the periodic table

29 9.4 Voltaic cells Background
These attractions of opposite charges produce these ‘simple batteries’ Electric charge Positive and negative charges Measured in coulombs (C). Electron = 1.6E-19C Carried by electrons from negative terminal to positive Electric current Rate at which electric charge flows through a circuit Measured in amperes (A). 1A= 1C s-1 Large current produced by a large amount of slowly moving charge or a small amount of quickly moving charge

30 Potential difference When there is a difference in electric potential between two points in a circuit Gives electrical energy to the charge The charge transfers energy into other forms Heat, light, chemical energy Measured in volts (V)= 1C J The charge simply carries energy around the circuit, is not used

31 Daniell cell is a simple voltaic cell
Voltaic cells Daniell cell is a simple voltaic cell Constructed by placing Zn electrode in zinc sulfate solution and a Cu electrode into a copper (II) sulfate solution The two electrode are connected by wires and a high-resistance voltmeter to allow electrons to flow A salt bridge allows ions to flow and maintain electrical neutrality Ex. Pg. 247

32 What’s happening in a daniell cell
Zn undergoes oxidation (higher in reactivity series) Zn ions dissolve into the water Electrons flow from surface of Zn electrode through external circuit to surface of Cu electrode to reduce it Continues until either all Zn electrode or Cu ions are used up Anode where oxidation occurs in a voltaic cell Cathode where reduction occurs in a voltaic cell

33 Anode and cathode Electric current flows from anode to cathode Difference in electric potential between anode and cathode, measured by a voltmeter, is the cell potential The further apart the two metals are in the activity series, the larger the cell potential will be

34 9.5 electrolytic cells A substance that allows electricity to pass though itself in a conductor Metals Graphite Aqueous solutions of acids, alkalis, ionic compounds Semi-conductors only slightly conduct A substance not allowing electricity to flow through itself is an insulator Common insulators: Non-metallic elements (except graphite) Dry samples of covalent compounds Solid samples of ionic substances

35 Conduction of electricity
Substance must contain electrically charge particles that move freely when a potential difference or voltage are introduced Valence electrons are the charged particles in solid and liquid states of metals Ionic solids do not conduct because their ions cannot move freely

36 Electrolysis of molten salt
When electric current passes through a metal, the metal is unaffected When electric current passes through an ionic substance (molten or in a solution), chemical decomposition occurs Substances that undergo this process are electrolytes Process of electrolyte decomposition electrolysis Especially used in industry to prepare Al, Cl,NaOH and H

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