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Regents Chemistry Review: Electrochemistry And Redox Reactions

Redox refers to reduction and oxidation. Let's start with a couple of mnemonics to help us memorize necessary information for the chapter.

OIL = Oxidation Is Loss of Electrons

RIG = Reduction Is Gain of Electrons

The substance that is oxidized is called reducing agent and what is reduced is called an oxidizing agent.

Rules for determining oxidation numbers:

  • Free, neutral elements have an oxidation number of 0.

  • Group 1 metals have an oxidation number of +1 (can be found on the Periodic Table)

  • Group 2 metals have an oxidation number of +2

  • Oxygen is -2 in all compounds other than peroxide (H2O2)

  • Hydrogen is +1 unless it is bonded to a metal

  • The sum of oxidation numbers in a neutral molecule is 0

  • The sum of oxidation numbers in a charged molecule (ion) is equal to the charge

Let's practice a problem: Determine the oxidation state of chromium in K2CrO4.

We can look at the periodic table to find oxidation numbers for K and O. K is +1 and O is -2.

Next we can make an equation to find Cr oxidation number. The overall charge of the molecule is 0 since the charge is not shown.

2(+1) + X + 4*(-2) = 0

X is the oxidation number for Cr

We get, 2+x-8 = 0



The oxidation number of Cr is +6

Redox reaction has an oxidation half reaction and a reduction half reaction. In this reaction, electrons are transferred from one atom to another.

An easy way to recognize a redox reaction is if you see an element by itself on one side and as a part of compound on another side. For example: Fe + 2HCl --> FeCl2 + H2. Here, we see Fe on its own on the left and together with Cl on the right. Element in its free, neutral form will have an oxidation number of zero. When it is a part of the compound, it will have an oxidation number other than zero, signifying a transfer of electrons. For example, in FeCl2, the oxidation number of Fe is +2. It is calculated the following way (Cl oxidation number is -1 from the Periodic Table): x+(2*-1) =0, x=+2

Oxidation half reaction will show electrons to the RIGHT of the arrow, since it loses electrons.

Reduction half reaction will show electron to the LEFT of the arrow since it gains electrons.


There are two types of cells: voltaic and electrolytic

Voltaic cell produces electricity by converting chemical energy into electrical. It has spontaneous reactions. Battery is an example of voltaic cell.

An electrolytic cell needs energy to run. It has nonpontanous reactions. Electroplating is an example of an electrolytic cell.

For both of these cells we can memorize

AN OX = Anode is Oxidation

Red Cat = Reduction is Cathode

For both of these cells, electrons flow from the anode to the cathode.

Each cell is made up of the anode, where oxidation half reaction takes place and cathode, where reduction half reaction takes place. The anode and cathode are connected by a wire, through which electrons flow. They are also connected by a salt bridge, which allows ions to travel and prevent charge build up.

Table J presents most active and least active metals on the left side. Most active metals are located at the top of the table. Metals that are more active will be able to replace metals below them from compounds. For example, Li is more active than K and can replace it, Li+ KCl --> LiCl + K.

Most active metals are also most easily oxidized.

For nonmetals, the ones that are most active and more easily reduced.

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