Unit 9- Redox/Electrochemistry

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* Reminder: when an atom has a positive charge, there are less electrons due to them being negatively charged. Same for when atoms have negative charges, there it has more electrons.

* Any code blocks () will represent subscripts. Ex: Carbon Dioxide= CO2

Oxidation States/Numbers:

→ The charge of elements. (The ones on the upper right hand corner of the periodic table for each element)

• Since most elements have more than one oxidation state, you can calculate the exact ones for each element in a specific reaction or compound.
  • An element could have an oxidation state not written in the periodic table.

Rule for calculating Oxidation states:

1. Oxygen is always 2-, with two exceptions

   1. H2O2, O is 1-
   2. OF2, O is 2+

2. Hydrogen and most group 1 ions are almost always 1+ (Can also be written as +)

3. Halogens (group 7 ions) will almost always be 1- (Can also be written as -)

4. Elements on their own (in a reaction) are always 0, unless a charge is written on the element symbol.

5. Diatomic elements are always 0.

6. In a compound the sum of the oxidation states must equal zero

   1. If it’s a polyatomic ion, the sum must equal the charge of the compound.

Examples of calculations:

• BeO
  • Since O is always 2-, the oxidation state of Be must be 2+, due to the fact that the sum of the numbers should be zero in neutral compounds

Redox reactions:

→ Redox is short for “Oxidation Reduction”

Oxidation and Reduction:

• Oxidation → The lose of electrons by an atom (an atom becoming more positive)
• Reduction →The gaining of electrons by an atom (an atom becoming more negative)
  • Think of it as the oxidation number is reducing, rather than the amount of electrons.
• A way to remember is LEO the lion says GER
  • LEO stand for “Loosing electrons is Oxidation”
  • GER stand for “Gaining electrons is Reduction”

The Reactions:

• In redox reactions one element is losing electrons (Oxidizing), while another is gaining those electrons (Reducing). These two thing happen simultaneously
  • In these reaction the element which is oxidized (lost electrons), is known as the Reducing Agent.
  • The Element which is reduced (gained electrons), is known as the Oxidizing Agent.
• The oxidation states/charges of elements will change from the reactants to the products.

For Example:

• Na + Cl → NaCl, is a redox reaction, since Na and Cl are both neutral on the reactants side, but have charges on the products.
  • Na on the products side has a charge of 1+, and Cl has a charge of 1-.
  • Na lost an electrons so it was oxidized (was the reducing agent)
  • Cl gained an electron so it was reduced (was the oxidizing agent)
• If a reaction does not have both oxidation and reduction happening it is not a redox reaction.

Half Reactions:

→ Redox reactions must be balanced for BOTH mass and CHARGE. To do this we use half reactions to show both processes (Oxidation and Reduction).

• A redox reaction can be split in to two half reactions, one for the element being oxidized, and the other for the element being reducded.
  • If there are any ions or elements where the charge doesn’t change, they are left out of these half reactions.
  • These are known as spectator ions.
  • In the half reaction we also show the electrons, represented by "e-”.
• Before writing half reaction make sure the mass is balanced (same amount of each element on both sides)
  • Then when the half reaction are written, there should be the same amount of “e-” in both halves.
  • If not then you need to manipulate them so they do.
• When the two half reaction are added together you can cancel out the “e-”, and you should be left with a balanced version of the original reaction
  • Some reaction are already balanced for both.

Examples:

• Na + Cl → NaCl, same as above.
  • The Oxidation half reaction would be: Na→ Na^1+ + e-
  • The Reduction half reaction would be: Cl + e- → Cl^1-
  • Charges are usually written as exponents/superscripts.
  • Notice the e- lost by Na, is the same e- gained by Cl. (This reaction is already balanced)
• (Insert picture of 3. on pg 12)

Simultaneous Reactions and Table J:

• On Table J, metals that are higher up will be oxidized
  • And the lower ones will be reduced.
• So if K and Ca were in a redox reaction, K should oxidize since it is higher up than Ca.
  • If the metal that gets oxidized is lower, than the metal being reduced than the reaction will NOT be spontaneous. (There are scenarios where this could happen)
  • If the metal that is oxidized is higher than the one being reduced, the reaction WILL be spontaneous.

Cells/Batteries:

→ There are two types of cells/Batteries: Galvanic/Voltaic cells, and Electrolytic cells. Though they work differently they do have some of the same parts. (They are both referred to as electrochemical cells)

• There are two electrodes in each battery. The electrons are where the electrons are oxidized and reduced.
• Anode → The electrode where oxidation will occur in the cell.
  • Similar to an anion, because it is the negative end of a batterie.
• Cathode → The electrode where reduction will occur in the cell.
  • Like a Cation, it is the positive end.
• You can remember these by using "An Ox”, and "Red Cat”
  • An Ox, stands for Anode is Oxidation
  • Red Cat, stands for Cathode is Reduction.
• The Anode and the Cathode are connected by a wire, where the electrons will flow through.

Voltaic/Galvanic Cells:

→ Both names mean the same thing.

• These cells produce energy
• The Anode and Cathode are separated, either in two separate containers filled with salt solutions, or by a membrane in the same container.
  • There is still a wire for connection.
  • Electrons will flow from the Anode to the Cathode.
  • From the electrode which loses electrons to the one which gains electrons.
• A salt bridge also connects the two solutions in the beakers. This helps maintain equilibrium by moving the icons around.
  • The ions will flow the opposite way as the electrons. from the cathode side to the anode side.
• As the cell goes through Oxidation and Reduction the Anode will lose mass and the Cathode will gain mass.
  • This is because the Anode loses atoms to the solution as they become ions from losing their electrons. (They fall off of the Anode)
  • The Cathode gains those electrons, which the ions in the solution will accept and they will then become neutral atoms and attach to the Cathode.
• It converts chemical energy to electrical energy.

Example: (Picture)

• Zn is the anode because it is higher up on Table J, than Cu, so it will oxidize.
  • The electrons are flowing from Zn (anode) to the Cu (Cathode)
• The salt bridge is carrying ions from the Cu side to the Zn side.
• Over time the anode will lose Zn atoms as they become ions, and the cathode will gain Cu atoms as they become neutral.

Electrolytic Cells:

• These cells require external energy to work. (Usually a battery)
  • This is because it is forcing a spontaneous reaction to happen.
• The Anode and Cathode are switched. So the one being oxidized is the one which is usually reduced in a Galvanic cell.
• They convert electrical energy into chemical energy
• There is no salt bridge since the entire cell is in one container.
• These cells are usually used to cat items so the cathode is usually an item like a key or some type of cutlery
  • This process is known as electroplating.

Example: (Picture)

• The Silver anode is losing electrons and those electrons are being transferred to the spoon.
  • The spoon accepts them and then the Ag+ ions in the solution take the electrons and attach to the spoon.
  • This will plate the spoon in silver. Increasing the mass of the spoon.
• The voltage source gives electricity to force the nonspontaneous transfer of electrons

Next Unit: Organic Chemistry