We can compare the rates by setting up the following equations:

We can expand the second equation by distributing the exponent:

We can see that the rate is multiplied by a factor of 128 when the concentration of each reactant is doubled.

We can compare the rates by setting up the following equations:

We can expand the second equation by distributing the exponent:

We can see that the rate is multiplied by a factor of 128 when the concentration of each reactant is doubled.

The balanced reaction does not tell you anything about the rate law of a reaction. The rate law must be determined through experiment. The basic format of our rate law will be:

In order to see how the initial concentration of a reactant affects the initial rate of the reaction, we need to find two trials where two reactants are kept constant and only one is changed. By seeing how the initial reaction rate changes, we can determine the reaction order for each particular reactant and fill in the exponent values.

When the concentration of A is doubled, the initial reaction rate is quadrupled. This can be seen by comparing trials 1 and 3. As a result, the reaction is second order with respect to A.

When the concentration of B is tripled, the initial reaction rate is tripled. This can been seen by comparing trials 1 and 2. As a result, the reaction is first order with respect to B.

Finally, the initial reaction rate will not change regardless of the concentration change of C. This can be seen by comparing trials 3 and 4. This makes the reaction zero order with respect to C.

For reactions with multiple steps, the step with the slowest rate will determine the rate law. In this reaction, the slowest step is the first step, so the reactants in that step will be the components of the rate law.

The first step is written as:

Using the reactants, we can determine the rate law to be 

Note that we cannot determine the value of the exponents, since they will need to be found experimentally.

The rate-determining step in a reaction mechanism is a kinetic bottleneck, in that it prevents the overall reaction from proceeding; thus, it is what determines how quickly the overall reaction can proceed.

Galvanic cells always involve spontanous oxidation-reduction reactions. In any electrochemical cell, the electrons always move from anode to cathode. Also, the anode is always the site of oxidation, and the cathode is always the site of reduction. Since the reaction is spontaneous (net release of free energy) it drives the movement of electrons from the anode to the cathode. Remember, oxidation is loss of electrons and reduction is gain of electrons. Since oxidation always occurs at the anode, we are left with an excess of electrons, making it the negative electrode. It should makes sense that the extra electrons from the anode spontaneously travel to the cathode (positive electrode).

To help remember oxidation-reduction processes, consider the mnemonics "OIL RIG" and "An Ox, Red Cat." OIL RIG stands for "oxidation is loss, reduction is gain" in reference to electrons. An Ox, Red Cat tells us that the anode is the site of oxidation, while the cathode is the site of reduction.

Electrolytic cells use non-spontaneous reactions that require an external power source in order to proceed. The values between galvanic and electrolytic cells are opposite of one another. Galvanic cells have positive voltage potentials, while electrolytic voltage potentials are negative. Both types of cell, however, have oxidation occur at the cathode and reduction occur at the anode.

First we must rearrange the reduction potentials so that when added together, they match the reaction that takes place in the electrochemical cell.

In the overall reaction,  is in the reactant side, so the  equation must be inverted.

Use the equation:  to find the . 

 is product, while  is the reactant.

The cell must be galvanic because the  value is positive. This means, this the reaction is a spontaneous reaction occurs without an outside energy source.

Oxidation always takes place at the anode, regardless of the electrical cell type. The charges on the anode and cathode are reversed between galvanic and electrolytic cells. In electrolytic cells, the cathodes are marked negative and the anodes are marked positive. In galvanic cells, the reverse is true: cathodes are marked positive and anodes are marked negative.

Reduction always occurs at the cathode, and oxidation always occurs at the anode. Since reduction is the addition of electrons, electrons must travel toward the site of reduction.

In an electrolytic cell the negative charge is on the cathode, while the positive charge is on the anode. Since an electrolytic cell requires energy to perpetuate the reaction, we are pushing the electrons against their potential gradient. The electrons, which are negatively charged, are traveling towards the cathode, which is also negatively charged.