The equation to use here is:

Here,  is the molarity of the acid,  is the volume of the acid,  is the molarity of the base, and  is the volume of the base. Don't forget to convert the volume to liters!

Below is a generic acid-base neutralization reaction:

The products are always water, and a salt. This salt is produced from the resulting ions  and . The  from the acid replaces the  from the base, and the  from the acid replaces the  from the base. Since there are two replacements, acid-base neutralizations are classified as double-replacement reactions.

Arrhenius acid/base theory was created by Swedish chemist Svante Arrhenius, and is the oldest acid/base classification. According to his classification, acids are compounds that increase the concentration of  ions in a solution, while bases are compounds or elements that either decrease the concentration of  ions in solution or increase the concentration of  ions in a solution. The other two answers describe the Brønsted–Lowry theory of acids and bases.

For every acidic or basic solution, the product of the hydroxide ion concentration and the hydronium ion concentration will be equal to , the dissociation constant for water. In other words:

Since we are given the hydroxide ion concentration, we can determine the hydronium ion concentration using this equation.

Acetic acid is a weak acid, and its dissociation reaction is written as:

Using an ICE table, we can find the pH of the solution when the sodium acetate is added:

I: There is initially a 0.50M concentration of acetic acid. Because sodium acetate will dissolve completely, there will be a 0.050M concentration of acetate ions in the solution.

C: As acetic acid dissociates, the concentrations of acetate ions and hydronium ions will increase by an unknown concentration . Conversely, the acetic acid concentration will decrease by the same amount.

E: By setting the equilibrium expression equal to the acid equilibrium constant, we can solve for the value of .

*Because the value for  will be so much lower than the initial concentration, we can omit it from the acetate expression as well as the acetic acid concentration:

Keep in mind that  is equal to the concentration of hydronium ions now in the solution.

Use this value in the equation for pH:

The acid dissociation constant (Ka) is used to distinguish strong acids from weak acids. Strong acids have exceptionally high Ka values.

The Ka value is found by looking at the equilibrium constant for the dissociation of the acid. The higher the Ka, the more the acid dissociates. Thus, strong acids must dissociate more in water. In contrast, a weak acid is less likely to ionize and release a hydrogen ion, thus resulting in a less acidic solution.

This equation makes it clear that the more the acid converts from its original form to its ionzied, dissociated form, the higher the Ka value will be.

The Ka is the acid dissociation constant, and thus it is what determines how strong the acid is. Stronger acids dissociate to a greater extent and produce lower pH values.

The definition of  for  is:

Set up an "ICE" table:

Plug in values:

Use the quadratic formula to solve:

Set up an "ICE" table

Plug in values:

Use the quadratic formula:

In order to find the base dissociation constant for the conjugate base, we can start by finding the acid dissociation constant for the acid. Since a 1M solution of the acid has a pH of 4.6, we can find the proton concentration of the solution.

Since the acid is monoprotic, we can set the following equilibrium expression equal to its acid dissociation constant.

We can see that, since the acid is monoprotic, the concntration of protons will be equal to the concentration of the acid anion. The final concentration of the acid molecule will be equal to the initial concentration, minus the amount of protons formed. Using these values, we can solve for the equilibrium constant for the acid.

Now that we have the acid dissociation constant, we can find the conjugate base's dissociation constant by setting the product of the two values equal to the autoionization of water.