Oxidation)&Reduction)(Redox)Notes) - Amazon S3&Reduction)(Redox)Notes)! Chemical)Activity)...

Oxidation & Reduction (Redox) Notes Chemical Activity (or Chemical Reactivity) is the measure of the reactivity of elements. If an element has high activity, then it means that the element is willing to give up its electrons & react. If an element has low activity, then the element wants to hold onto its electrons and not react. (See Table J)

Practice Problems Predicting Single Replacement Reactions a) AgNO3 + Ni à b) NaBr + I2 à c) Magnesium metal reacts with an aqueous solution of

zinc nitrate. d) Solid zinc is combined with a solution of lithium

nitrate. e) Sodium metal is dropped into water. f) Aluminum metal is combined with aqueous iron (III)

bromide.

Redox Determining Oxidation Numbers 1. Every uncombined element has an oxidation number

of zero. e.g. Cl2 à 0; or Na à 0 2. Monotomic ions have an oxidation number equal to

their charge. e.g. F-‐ à -‐1; or Fe3+ à +3 3. When combined, metals of Group 1 always have an

oxidation number of +1 and metals of Group 2 always have an oxidation number of +2.

4. Fluorine always has an oxidation number of -‐1 in compounds.

5. Hydrogen is +1 in compounds, unless it is bonded to a metal, when it is -‐1.

6. Oxygen is usually -‐2 in compounds. There are two exceptions: a) When oxygen is in a compound that contains F, then its oxidation number is +2. b) When oxygen in a peroxide.

7. The sum of the oxidation numbers in all compounds (which are always neutral) must be zero.

8. The sum of the oxidation numbers in a polyatomic ion must be equal to the charge of the ion.

Redox Practice Problems Determine the oxidation number for the given elements: 1. N in N2O3 2. S in H2SO4 3. C 4. C in CO 5. Na in NaCl 6. H in H2O 7. Ba in BaCl2 8. N in NO2-‐ 9. S in Al2S3 10. S in HSO4-‐

11. Cl in Fe(ClO2)3 12. Fe in Fe(ClO2)3 13. N in NO3-‐ 14. Cu2+ 15. C in CH4 16. Mn in MnO2 17. S in SO32-‐ 18. Mn in KMnO4 19. I in Mg(IO3)2 20. C in C2O42-‐

Redox Oxidation: when an element loses electrons Reduction: when an element gains electrons LEO the lions says GER

(Unbalanced Redox Reaction) Zn + HCl à ZnCl2 + H2 Which element lost electrons? Which element gained electrons?

Redox The reactant that loses electrons oxidizes The reactant that gains electrons reduces *****Confusing Part: The reactant that undergoes reduction is called the

reducing agent The reactant that undergoes oxidation is called the

oxidizing agent

Half-‐Reactions: Unbalanced chemical reaction: Zn + HCl à ZnCl2 + H2 0 +1/-‐1 +2/-‐1 0 Oxidation 1/2 Reaction: Zinc oxidizes (loses electrons): Oxid. # 0 à +2 Zn à Zn2+ + 2e-‐ Note: 2e-‐ must be added to the right side in order to

balance mass and charge! Reduction 1/2 Reaction: Hydrogen reduces (gains electrons: Oxid. # +1 à 0 2H+ + 2e-‐ à H2 ****Redox Reactions must be balanced by MASS AND

CHARGE!!!!

Write the oxidation and reduction 1/2 reactions for each redox reaction, then balance the overall reaction. The first one is done for you. Remember that you must balance both mass and charge! 23. Fe3+ + Co ⇄ Co2+ + Fe Oxidation: 3*(Co → Co2+ + 2e-‐) Reduction: 2*(Fe3+ + 3e-‐ → Fe)

2 Fe3+ + 3 Co ⇄ 3 Co2+ + 2 Fe 24. Ag+ + Ni ⇄ Ni3+ + Ag Oxidation: Reduction: 25. Cu1+ + Pb ⇄ Pb2+ + Cu Oxidation: Reduction:

26. O2 + Sn ⇄ O2-‐ + Sn4+ Oxidation: Reduction: 27. Co3+ + F-‐ ⇄ Co + F2 Oxidation: Reduction:

Redox How can we utilize half-‐reactions? Electrochemical Cells -‐ A cell that involves chemical reactions and electron flow. There are two types of electrochemical cells: i) Voltaic Cell (or Galvanic Cell) – an electrochemical cell that spontaneously produces electron flow. An example is a battery. ii) Electrolytic – an electrochemical cell that requires an electric current to force a nonspontaneous chemical reaction. An example is a battery that is charging.

Voltaic Cells

Zn (s) + Cu2+ (aq) à Zn2+ (aq) + Cu (s) Oxidation half-‐reaction: Zn (s) à Zn2+ (aq) + 2e-‐ Reduction half-‐reaction: Cu2+ (aq) + 2e-‐ à Cu (s) The zinc electrode (solid Zn) releases electrons to the aqueous copper ions (Cu2+). Therefore, electricity moves from the zinc to the copper.

The salt bridge: -‐NO ELECTRONS FLOW THROUGH THE SALT BRIDGE. -‐“Completes the circuit”. -‐Releases ions into the half-‐cells in order to neutralize the solutions. -‐Anions (negative ions) move towards the Anode. -‐Cations (positive ions) move towards the Cathode. **** REMEMBER**** Electrons always flow from A to C (Anode to Cathode). Anode:

-‐AN OX ANode OXidation

-‐Always the MOST ACTIVE metal

-‐Negative (produces electrons)

-‐LOSES MASS as electrons flow (s) à (aq)

Cathode:

-‐RED CAT REDuction CAThode

-‐Always the LEAST ACTIVE metal

-‐Positive (attracts electrons)

-‐GAINS MASS as electrons flow (aq) à (s)

Consider a voltaic cell involving zinc and aluminum. a. Sketch the cell. b. Identify the anode and the cathode. c. Write the balanced overall reaction. d. Which electrode will lose mass? e. In the wire, which way do the electrons flow? toward Zn or Al? f. Which way do the cations from the salt bridge move? toward Zn or Al? Consider a voltaic cell involving magnesium and silver. a. Sketch the cell. b. Identify the anode and the cathode. c. Write the balanced overall reaction. d. Which electrode will lose mass? e. In the wire, which way do the electrons flow? toward Mg or Ag? f. Which way do the cations from the salt bridge move? toward Mg or Ag?

Electrolytic Cells – electrochemical cells that need an external power source in order to drive a chemical reaction. Electroplating

An external power supply must be used to force the chemical reaction.

Now the Anode is the LEAST ACTIVE and the Cathode is the MOST ACTIVE (but the Anode is + and Cathode is -‐).

“AN OX” & “RED CAT” still apply.

Electrons still flow from A to C (Anode to Cathode).

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