acids, bases and salts

acid → definition

an acid is a substance that can donate a proton (H^+) to another substance

base → definition

a base is a substance that can accept a proton (H^+) from another substance

examples of acids

1. hydrochloric acid (HCl) → H+ and Cl-

2. sulfuric acid (H_2SO_4) → H+ and SO_4^{2-}

3. nitric acid (HNO_3) → H+ and NO_3^-

4. phosphoric acid (H_3PO_4) → H+ and PO_4^{3-}

5. carbonic acid (H_2CO_3) → H+ and CO_3^{2-}

6. ethanoic acid (CH_3CO_2H) → H+ and CH_3CO_2^-

examples of alkalis

1. sodium hydroxide (NaOH) → OH- and Na+

2. potassium hydroxide (KOH) → OH- and K^+

3. calcium hydroxide (Ca(OH)_2) → OH- and Ca^{2+}

4. aqueous ammonia (NH_3) → OH- and NH_4+

aqueous ammonia as a base

NH_3 (aq) + H_2O (l) ⇌ NH_4+ (aq) + OH- (aq)

• aqueous ammonia can dissolve in water and take in protons

strong acid → definition

a strong acid is one which dissociates fully in solution to give protons(H^+)

strong acid

• degree of dissociation: approximately 100%

• type of arrow used in dissociation equation: → (single arrow)

• monobasic: HNO_3

• dibasic: H_2SO_4

monobasic → definition

one mole of H^+ is dissociated per mole of an acid

dibasic → definition

2 moles of H^+ is dissociated per mole of an acid

strong base → definition

a strong base is one which dissociates fully in solution to give hydroxide ions (OH^-)

strong base

• degree of dissociation: approximately 100%

• type of arrow used in dissociation equation: → (single arrow)

• monobasic: KOH

• dibasic: Ca(OH)_2

weak acid → definition

a weak acid is one which dissociates partially in solution to give protons (H^+)

weak acid

• degree of dissociation: <<< 100%

• type of arrow used in dissociation equation: ⇌ (half arrow)

• examples: organic acids

weak base → definition

a weak base is one which dissociates partially in solution to give hydroxide ions (OH^-)

weak base

• degree of dissociation: <<< 100%

• type of arrow used in dissociation equation: ⇌ (half arrow)

• examples: ammonia

strength → definition

strength of an acid is a measure of the extent/percentage of dissociation of an acid in solution

concentration → definition

concentration of an acid is the number of moles of undissociated acid per unit volume

power of hydrogen (pH) → definition

the pH of a solution is defined as the negative logarithm to the base 10 of the concentration of hydrogen ions [H^+] in solution in mol dm^{-3}

formula for pH

• pH = -log_{10}[H^+]

• [H^+] = 10 ^{-pH}

pH value and concentration of hydrogen ions [H^+]

• the larger the pH value, the smaller the [$H^+$]

  • acidic: more H^+ than OH^- → lower pH

  • alkali: more OH^- than H^+ → higher pH

pH scale

at 25°C:

• neutral solution: [H^+] = [OH^-] → pH = 7

• acidic solution: [H^+] > [OH^-] → pH < 7 (1-6)

• alkaline solution: [H^+] < [OH^-] → pH > 7 (8-14)

indicators → definition

indicators are dyes or a mixture of dyes which change colour when added to acids or alkalis

physical properties of acids

• have a sour taste

• have pH less than 7

• turn damp blue litmus paper red

• may be solids (e.g. citric acid), liquids (e.g. concentrated sulfuric acid) or gases (e.g. hydrogen chloride gas)

• can conduct electricity

  • can act as charge carriers as they exist as free mobile ions in aqueous solutions

physical properties of alkalis

• have a bitter taste and feel soapy

• have pH greater than 7

• turn damp red litmus paper blue, turn universal indicator blue or violet

• can conduct electricity

  • can act as charge carriers as they exist as free mobile ions in aqueous solution

neutralisation → definition

neutralisation refers to the reaction between an acid and a base to form salt and water

common operating conditions for Haber process

1. high pressure of 250 atm → force the reaction in a certain direction

2. moderate temperature of 450°C → reaction itself produces a lot of heat

3. iron catalyst

4. molar ratio of N_2 : H_2 is 1 : 3

acid-base nature of period 3 oxide

trend: the nature of the oxides change from basic (metal oxides) to amphoteric to acidic (non-metal oxides) across the periods

group 1 and group 2 → metal oxides

• elements: Na (group 1) and Mg (group 2)

• formula: Na_2O and MgO → giant ionic lattice (structure)

• acid-base nature: basic

  • dissolves to form alkalis

  • behaves as bases which can react with acids

group 13 → metal oxide

• elements: Al

• formula: Al_2O_3 → giant ionic lattice (structure)

• acid-base nature: amphoteric

  • properties of both acids and base → reacts with both acids and bases

group 14, group 15 and group 16 → non-metal oxide

• elements: Si (group 14), P (group 15) and S (group 16)

• formula: SiO_2; P_4O_6, P_4O_10; SO_2, SO_3 → covalent

  • P_4O_6: phosphorus (III) oxide

  • P_4O_{10}: phosphorus (V) oxide

• acid-base nature: acidic

  • dissolves in water to form an acid

  • behaves as acids which can react with bases

basic oxides

• can react with acids to form salt and water

• usually insoluble in water (some can dissolve to form alkalis)

basic oxides and reaction with water

1. Na_2O + H_2O → 2NaOH

2. K_2O + H_2O → 2KOH

basic oxides and reaction with acids

1. Na_2O + 2HCl → 2NaCl + H_2O

2. K_2O + 2HCl → 2KCl + H_2O

3. MgO + 2HCl → MgCl_2 + H_2O

4. CaO + 2HCl → CaCl_2 + H_2O

acidic oxides

• can react with bases to form salt and water

• some can dissolve in water to form acidic solutions

  • SiO_2 not soluble in water due to its giant molecular structure and numerous strong covalent bonds

acidic oxides and reaction with water

• SO_2 + H_2O → H_2SO_3 (sulfurous acid)

• SO_3 + H_2O → H_2SO_4(sulfuric acid)

• P_4O_6 + 6H_2O → 4H_3PO_3 (phosphorous acid)

• P_4O_{10} + 6H_2O → 4H_3PO_4 (phosphoric acid)

acidic oxides and reaction with bases

• P_4O_{10} (s) + 12NaOH (aq) → 4Na_3PO_4 (aq) + 6H_2O (l)

• P_4O_6 + 8NaOH → 4Na_2HPO_3 + 2H_2O

• SO_2 (g) + 2NaOH (aq)→ Na_2SO_3 (aq) + H_2O (l)

• SO_3 (g) + 2NaOH (aq) → Na_2SO_4 (aq) + H_2O (l)

• SiO_2 (s) + 2NaOH (l) - 350°C → Na_2SiO_3 (l) + H_2O (g)

amphoteric oxides

• react with acids and bases to form salt and water

• do not react with water

amphoteric examples

1. aluminium oxide (Al_2O_3)

2. lead (II) oxide (PbO)

3. zinc oxide (ZnO)

amphoteric oxides and reaction with acid

Al_2O_3 + 6HCl → 2AlCl_3 + 3H_2O

amphoteric oxides and reaction with base

Al_2O_3 + 2NaOH + 3H_2O → 2NaAl(OH)_4

neutral oxides

• will not react with acids/bases → no reaction at all

• examples: carbon monoxide (CO), water ($H_2O$), nitric oxide (NO) and nitrous oxide ($N_2O$)

• reaction with acids and bases:

  • neutral oxides do not react with acids or bases

salts

• ionic compounds

• to determine type of reagents that could be used to form a salt → analyse the cations and anions of the salt

• water of crystallisation in many hydrated salts can be driven off by heating the crystals of the salt strongly to form anhydrous salts

water of crystallisation

water bonded chemically within the crystals

hydrated salts

salts that contain water of crystallisation

anhydrous salts

salts that do not contain water of crystallisation

to name hydrated salts

• the left hand side denotes the salt

• the dot (.) denotes the presence of water of crystallisation

• the right hand side denotes the water of crystallisation

formation of salts

1. acid + metal → salt + hydrogen gas

2. acid + carbonate → salt + water + carbon dioxide gas

3. acid + base → salt + water (neutralisation)

4. ammonium salt + base → salt + water + ammonia gas

   • need heat to dry off ammonia gas

5. 2 soluble (aqueous) reagent → insoluble salt

reactivity series

• K (most reactive) → Na → Ca → Mg → Al → Zn → Fe → Pb → H (H^+ acid) → Cu → Ag → Au (least reactive)

  • K and Na: highly reactive → explosive

  • K, Na, Ca, Mg, Al, Zn, Fe, Pb: more reactive than H → reacts with acids

  • Cu, Ag, Au: less reactive than H → will not react with acids

reaction of an acid and a metal

• when a reactive metal (above hydrogen in reactivity series) reacts with a dilute acid, salt and hydrogen gas are formed

• metal + acid → salt + hydrogen gas

• extremely reactive metals in the reactivity series (group 1 metals: sodium and potassium) react violently and explosively

reaction of an acid and a carbonate/hydrogen carbonate

• when a carbonate/hydrogen carbonate reacts with an acid, salt, water and carbon dioxide gas are formed

• carbonate ({CO_3}^{2-}) + acid → salt + water + carbon dioxide gas

• hydrogen carbonate ({HCO_3}^-) + acid → salt + water + carbon dioxide

reaction of an acid and a base

• when an acid reacts with a base, salt and water are formed

  • acid + base → salt + water

• when ammonia reacts with an a acid, only a salt is formed

  • acid + ammonia (NH_3) → salt (cation: {NH_4}^+)

reaction of an alkali and an ammonium salt

• when an ammonium salt is warmed in the presence of an alkali, salt, water and ammonia gas are formed

  • ammonium salt ({NH_4}^+) + base - warm → salt + water + ammonia gas (NH_3)

precipitation reactions

• insoluble salts can be formed by a precipitation reaction through the mixing of two soluble (aqueous) reagents

• when an aqueous solution that contains the anion of an insoluble salt is mixed with an aqueous solution that contains the cation of that salt → the insoluble salt will precipitate out of the mixture

preparation of salt → guidelines

• soluble salt: usually prepared from reactions with acids

• insoluble salt: usually prepared by the precipitation method (mixing of two soluble salts)

reaction with acids

• if salt to be prepared is soluble in water while reactants are soluble:

  1. reaction of acid with excess metal

  2. reaction of acid with excess base (insoluble in water)

  3. reaction of acid with excess carbonate (insoluble in water)

reaction of an acid and excess metal

• to produce a salt that is soluble in water, excess metal is used such that the acid is completely used up → limiting reagent

• hydrogen gas will escape to the surroundings → only the soluble salt and the excess metal will remain in the reaction container at the end of the reaction

  • the soluble salt and excess metal can be easily separated through filtration

• this method is not suitable for:

  • very reactive metals such as sodium and potassium → react violently with acids

  • metals below hydrogen in the reactivity series such as copper, silver and gold → do not react with dilute acids

reaction of an acid and excess insoluble base

• a salt can be prepared by a reaction between a base (metal oxide or metal hydroxide) and an acid

• the procedure for preparation of salt from recall of insoluble base and acid is similar to the procedure to prepare salt from acid and excess metal

reaction of an acid with excess insoluble carbonate

• carbonates and hydrogen carbonates can be soluble/insoluble in water

• insoluble carbonate or hydrogen carbonate can be added in excess to an acid to produce a salt

reaction of an acid with a soluble base or carbonate

• to overcome problem, an experiment has to be carried out to determine the exact volume of alkali (soluble base) that reacts with a know amount of acid, in the presence of an indicator → titration

  • the reaction is complete when the indicator changes colour. the volumes of acids and alkalis used are noted

• to prepare the salt, the experiment is repeated with the same amount of acid and alkali used, in absence of indicator → can cause contamination to the salt prepared

preparation of insoluble salts via ionic precipitation reaction

• criteria

  1. choose 2 suitable ionic compounds that when combined will result in a precipitate of the desired solution

  2. necessary to dissolve any solid starting material to get an aqueous solution first

sample question: why PbO(s) cannot be added directly to HI (aq) to prepare PbI_2 (s)

• PbO + 2HI → PbI_2 + H_2O

• reaction of PbO with HI forms insoluble PbI_2, which coats the surface of PbO solid. this protective layer prevents PbO from reacting with HI, leading to low yield of PbI_2