Self-Ionisation of Water and pH Concepts

The process where water molecules react with each other to produce hydrogen and hydroxide ions.

The equilibrium constant for the self-ionization of water, at a specific temperature.

A measure of the hydrogen ion concentration in a solution, expressed as the negative logarithm (base 10) of the hydrogen ion concentration.

A solution with a pH value of 7 at 25 °C

A scale used to measure the acidity or basicity of a solution based on the concentration of hydrogen ions.

A measure of the hydrogen ion concentration in a solution, calculated as -log₁₀[H⁺]

An acid that completely dissociates in water, meaning it readily donates a proton (H⁺).

An acid that doesn't completely dissociate in water, meaning it is a poor proton donor.

Two chemical species related by the transfer of a proton (H⁺).

A measure of the strength of an acid, calculating the ratio of products (H+ and A-) and reactants (HA) at equilibrium.

Calculate the pH of a 0.04 M sulfuric acid solution. Sulfuric acid is a strong acid that completely dissociates in water, releasing 2 H+ ions per 1 molecule of H₂SO₄.

Calculate the pH of a 0.15 M sodium hydroxide solution. Sodium hydroxide is a strong base, completely dissociating in water to release one OH⁻ ion for each molecule of NaOH.

The pH scale is typically useful for measuring acidity and basicity of dilute aqueous solutions. It's limited to a narrow range (0.1 - 14) under specific conditions (25°C, dilute solutions).

To find the pH of a weak acid, determine the dissociation constant (Ka) for the acid and the acid's concentration. The pH is calculated based on the equilibrium expression and the Ka value

Use the formula [H⁺] = √(Ka * M acid), where Ka is the dissociation constant and M acid is the concentration (in moles/liter) of the weak acid

A substance that changes color depending on the pH of the solution it's in. They act as weak acids or weak bases.

The principle that adding acid (H+) or base (OH-) to an indicator shifts the equilibrium reactions, causing color change.

A measure of the extent to which a weak acid dissociates into ions in aqueous solution.

A graph that shows the change in pH of a solution as a strong base is added to a solution of acid. The equivalence point is where moles of base = moles of acid.

The point in a titration at which the moles of base added equals the moles of acid present.

A substance that changes color when the pH of a solution reaches a certain value. Used to signal the endpoint in a titration.

Titration of a strong acid with a strong base. The pH jumps sharply at the equivalence point.

Titration of a weak acid with a strong base. The pH jump at the equivalence point is less dramatic. The pH at the equivalence point will be above 7.

Titration of a weak acid with a weak base. The pH change is gradual with no sharp jump. Difficult to detect the endpoint using an indicator.

A pair of acids and bases that differ by only one proton. For example, the HIn and In^- pair.

A theory that defines acids as proton (H+) donors and bases as proton acceptors.

An indicator is a weak base that can change color depending on the pH of the solution. In its acidic form, it has one color, and in its basic form, it has another color. It's a weak base because it doesn't completely dissociate in water like a strong base.

The pH range of an indicator is the specific pH interval where it noticeably changes color. This range helps identify whether a solution is acidic or basic by observing the color change.

A pH titration is an experiment where we monitor the pH changes of a solution as titrant (base) is gradually added from a burette to an analyte (acid). This allows us to track the neutralization process and determine the equivalence point.

In a typical pH titration, we use a burette to add the base, a beaker to hold the acid, a magnetic stir bar to mix the solution, a pH sensor to measure the pH, and software to collect and plot the data.