Rules for determining Oxidation states

• The oxidation state of an uncombined element is zero. this applies regardless of the structure of the element

  • Example: Cl₂, Xe, S₈

• The sum of the oxidation states of all the atoms or ions in a neutral compound is zero

• The sum of the oxidation states of all the atoms in an ion is equal to the charge of the ion

• The more electronegative element in a substance is assigned a negative oxidation state. The less electronegative element is assigned a positive oxidation state

  • Electronegativity is greatest at the top-right of the periodic table and decreases toward the bottom-left

• Group 1 metals always have a +1 oxidation state

• Group 2 metals always have a +2 oxidation state

• Oxygen always has a -2 oxidation state UNLESS it is in a peroxide or F₂O

  • In a peroxide, the oxidation state of oxygen is -1

  • In F₂O, the oxidation state of oxygen is +2

• Hydrogen always has an oxidation state of +1 UNLESS it is in a metal hydride

  • When hydrogen is in a metal hydride, it has an oxidation state of -1

• Fluorine always has an oxidation state of -1

• Chlorine always has an oxidation state of -1 UNLESS it is in a compound with O or F

  • Chlorine adopts a wide variety of oxidation states so it is safer to assume the oxidation state is not -1

Acid and Base half-reactions

Step 1. Determine whether or not it is an oxidation-reduction reaction

Step 2. Separate the reaction into half-reactions

• The substance being reduced will have electrons as reactants

• The substance being oxidized will have electrons as products

In some cases, it is important to note which half-reaction is being oxidized and which is being reduced

Acidic conditions

• Acidic conditions usually imply a solution with an excess of H⁺ concentration, making the solution acidic

• the balancing starts by separating the reaction into half-reactions

• However, instead of immediately balancing the electrons, balance all the elements in the half-reactions that are not hydrogen and oxygen

• Then, add H₂O molecules to balance any oxygen atoms

• Next, balance the hydrogen atoms by adding protons

• Now, balance the charge by adding electrons and scale the electrons (multiply by the lowest common multiple) so that they will cancel out when added together

• Finally, add the two half-reactions and cancel out common terms

Half-reaction method: acidic

Each reaction is balanced by adjusting coefficients and adding H₂O, H⁺, and e^- in this order:

1. Determine, using oxidation states, what is being oxidized and reduced

2. Balance elements in the equation other than O and H

3. Balance the oxygen atoms by adding the appropriate number of water (H₂O) molecules to the side that has fewer oxygen atoms

4. Balance the hydrogen atoms (including those added in step 2 to balance the oxygen atoms) by adding H⁺ ions to the opposite side of the equation

5. Add up the charges on each side. Make them equal by adding enough electrons (e^-) to the more positive side.

   (rule of thumb: e- and H are almost always on the same side)

6. The e^- in each half-reaction must be made equal; if they are not equal, they must be multiplied by the appropriate integers (the lowest common multiple) to be made the same

7. The half-equations are added together, canceling out the electrons to form one balanced equation. Common terms should also be canceled out

8. The equation can now be checked to make sure that it is balanced

Basic conditions

• Bases dissolve into OH^- ions in solution; hence, balancing redox reactions in basic conditions requires OH^-

• Follow the same steps as for acidic conditions. The only difference is adding hydroxide ions to each side of the net reactions to balance any H⁺

• OH^- and H⁺ ions on the same side of a reaction should be added together to form water

• Again, any common terms can be canceled out

Half reaction Method: Basic

The same steps with the addition of 6b. Each reaction is balanced by adjusting coefficients & adding H₂O, H⁺, e^- and OH^- in this order:

1. Determine, using oxidation states, what is being oxidized and reduced

2. Balance elements in the equation other than O and H

3. Balance the oxygen atoms by adding the appropriate number of water (H₂O) molecules to the side that has fewer oxygen atoms

4. Balance the hydrogen atoms (including those added in step 2 to balance the oxygen atom) by adding H⁺ ions to the opposite side of the equation

5. Add up the charges on each side. Make them equal by adding enough electrons (e^-) to the more positive side (e^- and H⁺ are almost always on the same side)

6. The e^- in each half-reaction must be made equal; if they are not equal, they must be multiplied by appropriate integers (the lowest common multiple) to be made the same

7. The half-equations are added together, canceling out the electrons to form one balanced equation

   • Add the appropriate number of OH^- to neutralize all H⁺ and convert into water molecules

8. Common terms should also be canceled out

9. The equation can now be checked to make sure that it is balanced

Summary

• Acidic half-reactions are balanced by adjusting coefficients & adding H₂O, H⁺, and e^- in that order

• Alkaline half-reactions are balanced by adjusting coefficients & adding H₂O, H⁺, e^-, and OH^- in that order

Electrolytic cells

The electrical conductivity of ionic compounds when molten or dissolved in aqueous solutions is explained by the fact that mobile ions move a particular direction in an electric field.

Electrolysis - The process by which a compound is broken down into its constituent elements using electricity.

• An electrolytic cell is composed of a molten or aqueous electrolyte, a battery, and two electrodes, the anode and the cathode.

• An electric current enters and leaves via the two electrodes, which are usually made of graphite (carbon) or an unreactive metal such as platinum

  • Such metals are know as being inert as they do not take part in the reaction

• The anode is connected to the positive terminal of the battery

  • Negative ions, or anions, are attracted to the anode

• The cathode is connected to the negative terminal of the battery

  • Positive ions, or cations, are attracted to the cathode

• When the ions reach the surface of the electrodes they undergo either oxidation or reduction and are discharged

In an electrolytic cell, the anode is positively charged and the cathode is negatively charged.

Voltaic vs. Electrolytic cells

Voltaic cells

• Oxidation occurs at the negative anode

• Reduction occurs at the positive cathode

• Involves an exothermic spontaneous redox reaction

• The cell converts chemical energy into electrical energy

• The cathode is positive and the anode is negative during discharge

• The cell uses two separate aqueous solutions connected by a salt bridge and an external circuit

• Electric current is conducted by the electrons in the external circuit and the movement of ions in the salt bridge

Electrolytic cells

• Oxidation occurs at the positive anode

• Reduction occurs at the negative cathode

• Involves an endothermic non-spontaneous redox reaction

• The cell converts electrical energy into chemical energy

• The cathode is negative and the anode is positive during electrolysis

• The electrolyte is a molten liquid (or an aqueous solution)

• Electric current is conducted by the electrons in the external circuit and the movement of ions in the electrolyte