ACIDS & BASES

Acids and Bases reactions occur in everyday life and are essential for understanding our world. How does pH value affect our environment?

Why is it important to monitor and maintain the pH of the water in aquariums, soil and our blood? What exactly is pH? How is it measured?

Milk of magnesia is a medicine that usually relieves uncomfortable gastrointestinal symptoms within 30 minutes and constipation within six hours. Why is the milk of magnesia an antacid?

Keywords  Acidity  Basicity (Monoprotic, diprotic, triprotic)  Bronsted-Lowry Theory - Proton donor/acceptor - Acid-base Conjugate pair - Amphiprotic  Lewis Theory - Lone pair electrons - Dative/Coordinate bond

Recall Questions What is an acid? What is a base/alkali? A substance which produces hydrogen ions (protons) when dissolves in water. A base refers to substances like metal oxides and metal hydroxides. A substance which reacts with acid to form salt and water only An alkali is a soluble base which in solution produces hydroxide ions. Most bases are insoluble in water. 3 soluble bases are NaO/NaOH,KO/KOH,CaO/Ca(OH) 2

Recall Questions What causes acidity? It is the hydrogen ions that give an acid its acidic properties when they dissolve in water and dissociate into ions. E.g. HCl gas is a covalent compound. When dissolves in water, it forms HCl acid which dissolciate to form ions.

Recall Questions What is basicity(proticity)? Basicity refers to the no.of H atoms in one molecule of acid that acn be repleced by a metal. refers to the no. of H+ that can be replaced by one molecule of that acid. E.g. HCl (monobasic),H 2 SO 4 (dibasci),H 3 PO 4 (tribasic)

Bronsted-Lowry theory An acid is defined as a molecule or ion that acts as a proton donor (H+). A base is defned as a molecule or ion that acts as a proton acceptor (H+). HCl(g) + H 2 O(l) H 3 O + (aq) + Cl - (aq)

 Acids that have single proton to donate – monoprotic (monobasic). E.g. HCl(aq), HNO 3 (aq), HNO 2 (aq)  Acids that have 2 protons to donate – diprotic E.g. H 2 SO 4 (aq), H 2 SO 3 (aq), H 2 CO 3 (aq)  H 3 PO 4 (aq) is triprotic.

Hydrogen chloride gas dissolved in water HCl(g) + H 2 O(l) H 3 O + (aq) + Cl - (aq) The equation can be split into (i) HCl(aq) Cl - (aq) + H + (aq) acid conjugate base (ii) H 2 O(l) + H + (aq) H 3 O + (aq) base conjugate acid

CH 3 COOH(l) + H 2 O(l) H 3 O + (aq) + CH 3 OOO - (aq) donates H + acidbase acidbase Acid-base conjugate pair conjugate NH 3 (g) + H 2 O(l) NH 4 + (aq) + OH - (aq) Water is sometimes described as amphiprotic because it can accept or donate a proton. donates H +

Competition between acid/base and its conjugate (i) HCl(g) + H 2 O(l) H 3 O + (aq) + Cl - (aq) acid base conjugate acid conjugate base (ii) CH 3 COOH(l) + H 2 O(l) H 3 O + (aq) + CH 3 OOO - (aq) acid base conjugate acid conjugate base (i) Water is a much stronger base than chloride ion and has a stronger tendency to accept proton.The equilibrium shifts more to the right. (ii) Ethanote ion is a much stronger base than water molecule. The equilbrium shifts to the left.

Gas-phase acid-base reaction HCl(g) + NH 3 (g) NH 4 Cl(s)  The Bonsted-Lowry model can be extended to gas- phase acid-base reaction.  It involves the transfer of hydrogen ion from hydrogen chloride to ammonia.

(i) HCl(g) + H 2 O(l) H 3 O + (aq) + Cl - (aq) acid base conjugate acid conjugate base (ii) CH 3 COOH(l) + H 2 O(l) H 3 O + (aq) + CH 3 OOO - (aq) acid base conjugate acid conjugate base  Strong acids produce relatively weak conjugate bases in aqueous solutions.  Weak acids produce relatvely strong conjugate bases in aqueous solutions.

Common acids & conjugate bases in order of strengths

Lewis theory  A Lewis acid is defined as a substance that can accept a pair of electrons from another atom to form a dative (coordinate) covalent bond.  A Lewis base is defined as a substance that can donate a pair of electrons to another atom to form a dative (coordinate) covalent bond. B: H +  + BH

Examples  Reaction between ammonia and proton H 3 N: H +  + NH 4  Reaction between a water molecule and proton H 2 O: H +  H 3 O +

Lewis bonding In complex ions formed by transition metals The 6 water molecules, each donate a lone pair electrons from oxygen of their water molecules to the empty 3d orbitals of iron. What does each water molecule and iron(III) ion act as in the reaction above?

Dative (Coordinate) bond  A dative covalent bond is always formed in a Lewis acid-base reaction.  For a substance to act as a base, it must have space to accept the lone pair of electrons.

Strong and weak acids and bases Strong acid  When strong acid dissolves, virtually all acid molecules react with the water to produce hydronium ions In general for a strong acid HA HA + H 2 O(l)  H 3 O + (aq) + A - (aq) or HA  H + (aq) + A - (aq) 0% 100% 0% 100% Examples : HCl, H 2 SO 4,HNO 3, HClO 4

Strong and weak acids and bases Weak acid  When a weak acid dissolves in water, only a small % of its molecules (typically 1%) react with water molecules to release hydrogen or hydronium ions. The equilibrium lies on the left- hand side of the equation. HA + H 2 O(l) H 3 O + (aq) + A - (aq) or HA H + (aq) + A - (aq) 99% 1% Examples : CH 3 COOH, aqueous carbon dioxide

Distinguish between strong and weak acids Base on the information above, how do we distinguish betwee strong and weak acids of the same concentration (e.g. HCl and CH 3 COOH)? 0.1 mol dm -3 HCl(aq)0.1 mol dm -3 CH 3 COOH (aq) [H + (aq)]0.1 mol dm mol dm -3 pH Electrical conductivityhighlow Relative rate of reaction with magnesium fastslow Relative rate of reaction with calcium carbonate fastslow

How to distinguish between strong and weak acids?  A weak acid has a lower concentration of hydrogen ions and hence a higher pH than a stronger acid of the same concentration.  A weak acid, because of its lower concentration of hydrogen ions, will have much poorer electrical conductor than a stronger acid of the same concentration.  Weak acids react more slowly with reactive metals, metal oxides, metal carbonates and metal hydrogencarbonates than strong acids of the same concentration.  Strong and weak acids can also be distnguished by measuring and comparing their enthalpies of neutralisation. What is the difference between the strength (strong and weak) and the concentrated (concentrated or dilute)?

Strong and weak acids and bases Strong base  A strong base undergoes almost 100% dissociation/ionisation when in dilute aqueous solution. BOH  B + (aq) + OH - (aq) 0% 100% Examples : NaOH, KOH, Ba(OH) 2

Strong and weak acids and bases Weak base  All bases are weak except the hydroxides of groups 1 and 2.  Weak bases are composed of molecules that react with water molecules to release hydoxide ions. In general for a weak molecular base, BOH  The equiibrium lies on the left side of the equation. BOH + (aq) B + (aq) + OH - (aq) Examples : aqueous ammonia, ethylamine, caffeine, bases of nuclei acids

The pH indicator  scale that measures the strength of an acid and alkali.  pH of a substance is measured when it is dissolved in water.  pH stands for “power of hydrogen”  [H + ] = 1 x 10 -n moldm -3 ( n = pH number)

The pH Scale

pH probe and meter An accurate method of measuring pH value. A pH probe is dipped into the solution being tested and the pH value is then read directly from the meter.

pH Calculation  pH is a measure of the concentration of H + ions in a solution.  pH = -log 10 [H + (aq)] Example: If the concentration of H + is 2.50 x moldm -3, what is the pH? pH = -log (2.50 x ) = 2.60

Example: Calculate the concentration of H + of a solution that has a pH = log[H + ] = 3.2 log [H + ] = -3.2 [H + ] = 6.31 x 10 -4

Example: (a) What is the pH pf 10cm 3 of 0.1 moldm -3 HCl? pH = -log (0.1) = 1 (b) If 90cm 3 of water is added to the acid, what happens to the pH? Total volume = 100cm 3 In 10cm 3 solution, concentration of H + is 0.1 moldm -3 In 100cm 3 solution, concentration of H + is 0.01 moldm -3 pH = -log (0.01) = 2 (c) If the solution from (b) is diluted by a factor of 10 5, what is the approximate pH? The pH will increase by 4 to 6.