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Topic 9

Topic 9. Oxidation and Reduction. 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

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Topic 9

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  1. 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 • Is • Gain

  6. 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. -/+ O.N. • 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. 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

  17. 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:

  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

  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. 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. 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. Non-metals • 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. 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

  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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