HS Chemistry - Redox Reactions

Unit Summary

Ionic Equations & Redox Reactions

• In an ionic equation, spectator ions are ions that don't change in the equation.

• A net ionic equation is an ionic equation that shows only the ions in a chemical reaction that react together - that is, it removes all the spectator ions.

• To find out what elements will ionize in an ionic equation, we can use the following set of rules $most of the time$:

  • Ionic compounds that dissolve in water will ionize. Almost all ionic compounds will dissolve in water. The major exceptions to this are carbonate compounds, hydroxide compounds, and oxide compounds.
    • There are other exceptions, and exceptions to these exceptions, but they are not notable for now.
  • Molecules, like O₂, will not ionize.
  • Polyatomic ions are not ionized further. For example, if we had H₂SO₄, that would ionize into 2H⁺ and SO₄^2-, but the SO₄^2- wouldn't further ionize.

  • Additionally, if we're given the state symbols:

    • Solids and gases won't ionize, and liquids won't ionize unless they're ionic compounds in the liquid state $like liquid NaCl$.
    • If something is dissolved, it's usually ionized. The exception to this is when it's a molecule that's dissolved $like Cl~2~ dissolved in water. It won't split up into Cl^\+^ and Cl^\-^$.

• A redox reaction is a chemical reaction that involves some elements getting reduced and some getting oxidized:

  • Reduction is gain of electrons $for example, Cl + e^\-^ → Cl^\-^$
  • Oxidation is loss of electrons $for example, Li - e^\-^ → Li^\+^, but this is more commonly written as Li → Li^\+^ + e^\-^$

• Not all equations with ionic compounds have net ionic equations, and not all net ionic equations are redox reactions.

Balancing Redox Reactions

• When balancing redox reactions, not only must the amount of matter on both sides be balanced, but the total charge of both sides must also be balanced.

• The oxidation number is how many electrons an atom can use in bonding - that is, how many electrons can that atom lose, gain, or share in bonding?

  • Oxidation numbers have signs. If the electron is lost, there's a + sign, and if it's gained, there's a - sign.
  • If the electrons are being shared, we act as if the more electronegative atom completely steals the electron$s$ away from the less electronegative atom, and then calculate the oxidation number.

• Oxidation numbers aren't always the same for an atom. They change depending on the circumstance.

• There are a few rules regarding oxidation numbers:

  • In any compound, the oxidation numbers of all the elements will add up to the net charge of the compound.
  • Group 1 & 2 elements will always have oxidation numbers of 1 & 2 respectively.
  • Elements bonded with each other $like O~2~, Au, etc.$ have an overall charge and oxidation number of 0.
  • Monatomic ions $like Br^\-^, K^\+^, Sc^3\+^$ have the same oxidation number as their charge.

  • Some elements will always have the same oxidation number:

    • Fluorine will always have an oxidation number of -1.
    • Hydrogen will always have an oxidation number of either +1 or -1.
      • If Hydrogen is bonded to another metal in ionic bonding $NaH, CaH~2~, etc.$ then its oxidation number is -1. Otherwise, it is +1.
    • Oxygen will almost always have an oxidation number of -2.
      • The exceptions are when it is a preoxide $\-1$ or in F₂O $\+2$.

• If an element's oxidation number in a molecule is unknown, it can be calculated by $net charge of molecule$ - $combined charges of every other element in the molecule$.

• A change in an element's oxidation number $a change in its oxidation state$ over the course of a reaction shows that that element is bieng reduced/oxidized in the reaction. The amount that the oxidation number changes by is the amount by which it is reduced/oxidized. Thus, by calculating elements' oxidation numbers on both sides of a reaction, we can easily tell which elements are being reduced/oxidized in a reaction, and by how much.

• The atom which undergoes reduction is called the oxidizing agent, and the atom which undergoes oxidation is called the reducing agent.

Displacement Reactions

• Metal elements become more reactive as you go further to the left of the Periodic Table and further down the Periodic Table.

• A metal displacement reaction is a special type of redox reaction where a more reactive metal will displace a less reactive metal from its compounds.

  • For example, in the metal displacement reaction 6Na + Al₂O₃ → 3Na₂O + 2Al, Sodium $Na$ displaces Aluminum $Al$ from Oxygen $O$.

• Halogen displacement reactions work similar to metal displacement reactions, but instead of the Group 17 elements becoming more reactive as you go down the Group, they become less reactive as you go down the Group.