Ionization of Water and pH Scale

Ionization of Water

• Water is slightly ionized into a hydrogen ion (H+) and a hydroxide ion (OH-).
• At room temperature, only about one out of every 10^9 water molecules is ionized at any given time.
• Pure water mostly consists of H2O molecules.

Concentrations of H+ and OH-

• At 25°C, water dissociates, with [H+] and [OH-] concentrations being 1x10^-7 M.
• The concentration of water is a constant 55.5 M.
• The ion product of water (Kw) is the product of the equilibrium constant and the concentration of water, and is 1x10^-14 M^2.
• In water, the product of the concentrations of [H+] and [OH-] must always be 1x10^-14 M^2.
• A solution is neutral when [H+] = [OH-].
• In most solutions, the concentration of H+ and OH- is not equal, and when the concentration of one increases, the other decreases so that their product equals 1.0 x 10^-14.

pH Scale

• pH is a shorthand way to represent the amount of H+ in a solution.
• pH is defined as the negative logarithm of the hydrogen ion concentration
• pH = -log[H+]
• pH scale ranges from 1 to 14.
• A pH from 1-7 means the solution is acidic, and the smaller the number, the more acidic.
• A pH from 7-14 means the solution is basic, and the larger the number, the more basic.
• A pH of 7, the dissociation of water, is neutral.
• The concentration of [H+] in an aqueous solution is usually quite small, and we therefore express the concentration of hydrogen ions as the negative logarithm of [H+].

Acids and Bases

• An acid is a proton donor, represented as HA → H+ + A-.
• Strong acids completely ionize in water and go 100% to the right, which is an irreversible reaction, for example: HCl → H+ + Cl-.
• Weak acids do not ionize completely into H+ + F-, and this is a reversible reaction, for example: HF ←→ H+ + F-.
• A base is a proton acceptor or a hydroxyl donor, represented as B + H+ → BH+.
• A strong base completely dissociates in solution into its ions, which is an irreversible reaction, for example: NaOH → Na+ + OH-.
• Weak bases do not fully ionize in water and this is a reversible reaction, for example: NH3 + H+ ←→ NH4+.

Ionization Constants of Acids

• HA → H+ + A-
• Ka = [H+][A-] / [HA]
• Ka is a measure of the ability of an acid to dissociate
• Strong acids have high Ka and are ionized 100% in water.
• Weak acids have a low Ka and are ionized partially (~10%) in water.

Physiological pH and Buffers

• It's important for the human body that pH remains nearly constant.
• ~7 inside cells
• 7.4 ± 0.05 in blood.
• This is achieved by the use of buffers.
• Buffers are solutions of a weak acid (HA) with its conjugate base, or a weak base with its conjugate acid.
• Buffers resist changes in pH because they have acid to neutralize bases and bases to neutralize acids.
• Buffers resist changes to the pH of a solution when H+ or OH- is added to the solution.
• Buffers accept hydrogen ions from the solution when they are in excess and donate hydrogen ions when they have been depleted.

Common Ways to Produce Buffers

• Weak acid + its conjugate base (e.g., CH3COOH/CH3COO-)
• Weak base + its conjugate acid (e.g., NH3/NH4+)
• Buffers require an acidic component and a basic component.

Buffer Systems of the Body

• Proteins are important buffers in the body, mainly intracellular, and include haemoglobin.
• Phosphate buffer (H2PO4- : HPO42-) is mainly intracellular.
• H2CO3 is a buffer found in the blood (extracellular).

Henderson-Hasselbalch Equation

• A modification of the equation for the dissociation of a weak acid, used to determine the pH of a buffer solution:
• pH = pKa + log ([base]/[acid])
  • pH is the log of the molar concentration of the hydrogen
  • pKa is the equilibrium dissociation constant for an acid
  • [base] is the molar concentration of the conjugate base
  • [acid] is the molar concentration of an acidic solution