Fundamentals of chemisty: Acid and bases and pH

What is pH?

• The measure of the acidity of a solution (pH stands for “power of hydrogen”)

• a measure of the activity of dissolved hydrogen ions

• Hydrogen ions occur as various of cations including protons and hydronium ions (H3O+) in solution

What happens to the conc of hydrogen ions in pure water?

• In pure water at 25 °C the concentration of hydrogen ions

(H+) equals the concentration of hydroxide ions (OH-)

• "neutral" corresponds to a pH level of 7.0

• ACID Solutions are where concentration of [H+] exceeds

that of [OH- ]. They have a pH value lower than 7.0

• BASIC Solutions are where [OH- ]exceeds [H+]. They

have a pH value greater than 7.0

pH of everyday solutions

What is low pH?

• low pH values solutions with high conc of hydrogen ions

What is high pH?

• high pH values solutions with low conc of hydrogen ions

What is pure water?

• Pure water has pH of 7.0, other solutions are often described with reference to this value

Acid and bases definitions

• Acids solutions that have a pH less than 7 (I.e. more hydrogen ions than water)

• Bases a pH greater than 7 (I.e less hydrogen ions than water)

• The definition of weak and strong acids OR weak and strong bases does not refer to pH value

• It describes how well an acid or bases ionizes in solution

Bronsted-Lowry Theory:

• Acid is a proton donor

• Bases is proton acceptor

NOTE: this definition is independent of water

What is the pH of a solution defined as?

pH = -log[H3O+]

• pH scale - a scale that indicates the acidity or basic nature of a solution

• Measure of number of H+ ions (or equivalent) in solution

• Ranges from 0 (very acidic) to 14 (very basic)

pH = -log[H+]

• A measure of the effective concentration of hydrogen ions (rather than the actual concentration)

• In practice hydrogen ions could be shielded/hidden so not available to participate in chemical reactions e.g. inside a protein molecule

• Having solutions at the correct pH is critical to living organisms

The case of water

• pure water almost 100% molecular (covalently bonded)

• Very small amount is ionised (dissociated)

  > H20 + H20 → H3O+ + OH-

  • In pure water at room temperature

    [H3O+] = 1 × 10-7 M

    [OH-] = 1 × 10-7 Mu

• [ ] = mean the concentration of a substance

• pH = - log [H3O+]

• pH = - (-7)

• pH = 7

However, when pure water is exposed to the atmosphere

CO2 will be absorbed and react with water to form carbonic acid (HCO3- and H+) so pH lowered to approx 5.7

pH is a logarithmic scale

• pH 7 → 0.0000001 M

• pH 6 → 0.0000001M

• pH 5 → 0.00001M

• pH 4 → 0.0001M

• pH 3 → 0.001M

• pH 2 → 0.01M

• pH 1 → 0.1M

• pH 0 → 1M

The difference between each value is 10-fold

What are examples of strong and weak acids?

Strong acids:

• HCl Hydrochloric acid

• HNO3 Nitric acid

• H2SO4 Sulphuric acid

• HI Hydroiodic acid

Weak acids:

• HClO Hypochlorous acid

• HNO2 Nitrous acid

• H2SO3 Sulphurous acid

• HF Hydrofluoric acid

• CH3COOH Ethanoic acid

• H2CO3 Carbonic acid

Strong Acids

Strong acid:

• HCl + H20 → H+ + Cl-

• HCl + H2O → Cl- + H3O+

• strong acids; dissociated reaction goes to completion - no unrelated acid remains in solution

• Reaction is: HX + H2O → H3O+ + X- usually simplified to: HX → H+ + X-

• To calculate pH on;y need to know conc of the acid (HX) present

Weak Acids

Weak acids:

• CH3COOH + H+ => CH3COO- + H2O

• CH3COOH(aq) + H2O(l) => CH3COO-(aq) + H3O+(aq)

  > H3O+ is an hydronium ion

• Weak acids: dissociation reaction does not go to completion - unrequited acid remains in solution

• An equilibrium is reached between the hydrogen ions and the conjugate base

• Reaction is: HX => H+ + X-

• For methanoic acid: HCOOH(aq) => H+ + HCOO-

• To calculate pH need to know balance of equilibrium reaction (equilibrium constant) - called Aciditiy contact (Ka)

=> is reversible arrows

Bases

• A strong base is a base which hydrolyzes completely, raising the pH of the solution towards 14

• Arrhenius bases are water-soluble and donate hydroxide ions (OH-)

• Alkalis are bases. However, not all bases are alkalis

• Alkalis are Arrhenius bases and hydroxides of the alkali

metals e.g. sodium, potassium

• Strong bases: NaOH, KOH, Ca(OH)2

• Weak bases: NH3 (ammonia)

Bases and pH

pH = -log[H3O+]

• The strong base sodium hydroxide ionizes into hydroxide

and sodium ions in solution:

NaOH → Na+ + OH−

• Pure water dissociates :

2H2O(l) → H3O+(aq) + OH−(aq)

• When these are mixed the H3O+ and OH− ions combine to

form water molecules:

H3O+ + OH− → 2 H2O

• The hydroxide ion ‘removes’ the available hydronium /

hydrogen ions, lowering their concentration, so pH value

goes up

Neutralisation

• the reaction between an acid and a base

• will produce a salt and neutralized base

• hydrochloric acid and sodium hydroxide form sodium

chloride and water:

• HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)

H+ (aq) + OH- (aq) =>  H2O (l)

• Neutralization will not always give a solution with pH 7.0

– Only if similar strength acids and bases reacted

– a strong acid and a weak base will give a weakly acidic salt and vice-versa

Buffers

• • Buffer solution - solution which resists large changes in pH when small volumes of acids or bases are added.

  • Buffers consist of either:

  • a weak acid and its salt or a weak base and its salt

  • Buffers are essential to life and biochemical reaction

• Buffer solution: an aqueous solution consisting of a

mixture of a weak acid and its conjugate base or a weak

base and its conjugate acid

• Buffers mean that pH of solution changes very little when a

small amount of acid or base is added to it

• Buffer solutions are used as a means of keeping pH at a

nearly constant value in biochemical applications

• Buffers are in cells and blood to maintain physiological pH

– E.g. blood plasma is at pH 7.4 via bicarbonate-carbonic acid

Henderson-Hasselbalch equation

• Useful for estimating the pH of a buffer solution and finding the equilibrium pH in acid-base reactions

Buffers in biology

• Resistance to changes in pH make buffer solutions

essential for many biochemical processes

• An ideal buffer for a particular pH has a pKa equal to that

pH. Thus, the solution has maximum buffer capacity

• Buffer solutions are necessary to keep correct pH for

enzymes to work - many enzymes work only under very

precise conditions

• A buffer of carbonic acid (H2CO3) and bicarbonate

(HCO3−) is present in blood plasma - maintains a pH

between 7.35 and 7.45