Chapter 7: Acids, bases and salts (copy)

acid

acid

• proton donors

• pH below 7

• H+ ions make the acidic

strong acid

an acid that completely dissociates in an aqueous solution, produce lots of H+ ions

example of strong acid, why is it a strong acid

hydrochloric acid

HCl (aq) gives H+ (aq) + Cl- (aq)

weak acid

an acid that partially dissociates in an aqueous solution, produce few H+ ions

example of weak acid, why is it a weak acid

ethanoic acid

CH3COOH (aq) reversibly gives H+ (aq) + CH3COO- (aq)

characteristic properties of acids: reaction with metals

acid + metal gives salt + hydrogen gas

characteristic properties of acids: reaction with bases

acid + base gives salt + water

characteristic properties of acids: reaction with carbonates

acid + carbonate gives salt + carbon dioxide + water

acids: effect on litmus

blue to red

acids: effect on thymolphthalein

stays colourless

acids: effect on methyl orange

orange to red

base

• proton acceptors

• oxides or hydroxides of metals

• pH above 7

• OH- ions make them basic

alkali

soluble bases

characteristic properties of bases: reaction with ammonium salts

base + ammonium salt gives salt + water + ammonia gas

bases: effect on litmus

red to blue

bases: effect on thymolphthalein

colourless to blue

bases: effect on methyl orange

orange to yellow

acid (aq) contains

H+ ions

base (aq) contains

OH- ions

neutralisation reaction

neutralisation reaction occurs when an acid reacts with an alkali

H⁺ ions react with the OH^– ions to produce water

e.g H+ (aq) + OH- (aq) gives H2O (l)

hydrogen ion concentration and universal indicator (colours and pH)

the higher the concentration of hydrogen ions, the stronger the acid, but the lower the pH

the lower the concentration of hydrogen ions in a solution, the weaker the acid, the higher the pH

(applied to hydroxide ions as well)

pH

pH is a measure of the concentration of H⁺ ions in solution, but they have an inverse relationship

• lower the pH then the more acidic the solution is

• The higher the pH then the more alkaline the solution is

• A solution with a pH of 7, such as water or distilled water, is described as being neutral 

pH scale is logarithmic, meaning that each change of 1 on the scale represents a change in concentration by a factor of 10

oxide, how are they classified

oxides are compounds made from one or more atoms of oxygen combined with one other element

oxides can be classified based on their acid-base characteristics

acidic and basic oxides; in terms of pH and metallic character

acidic and basic oxides have different properties and values of pH

the difference in their pH stems from whether they are bonded to a metal or a non-metal element

The metallic character of the element influences the acidic or basic behaviour of the molecule

acidic oxides

acidic oxides are formed when a non-metal element combines with oxygen

react with bases

when dissolved in water they produce an acidic solution with a low pH

e.g. CO2 and SO2

basic oxides

basic oxides are formed when a metal element combines with oxygen

react with acids to form a salt and water

when dissolved in water they produce a basic solution with a high pH

e.g. CuO and CaO

neutral oxides

some oxides do not react with either acids or bases and thus are said to be neutral

e.g. CO, NO and N2O

amphoteric oxides

curious group of oxides that can behave as both acidic and basic, depending on whether the other reactant is an acid or a base, in both cases salt and water are formed

e.g. Al2O3 and ZnO

(applies to their hydroxides as well)

aluminium oxide behaving as a base

Al₂O₃ + 6HCl gives 2AlCl₃ + 3H₂O

aluminium oxide behaving as an acid

Al₂O₃ + 2NaOH gives 2NaAlO₂ + H₂O

salt

a compound that is formed when the hydrogen atom in an acid is replaced by a metal

Naming salts involves 2 parts; the name of the metal and the end of the acid

eg. calcium + hydrochloric acid gives calcium chloride

soluble salt preparation: dilute acid and alkali (soluble base) by titration

1. Use a pipette to measure the alkali into a conical flask and add a few drops of indicator (thymolphthalein or methyl orange)

2. Add the acid into the burette and note the starting volume

3. Add the acid very slowly from the burette to the conical flask until the indicator changes to the appropriate colour

4. Note and record the final volume of acid in the burette and calculate the volume of acid added (starting volume of acid - final volume of acid)

5. Add this same volume of acid into the same volume of alkali without the indicator

6. Heat the resulting solution in an evaporating basin to partially evaporate, leaving a saturated solution (crystals just forming on the sides of the basin or on a glass rod dipped in and then removed)

7. Leave to crystallise, decant excess solution and allow crystals to dry

soluble salt preparation: acid to a solid metal, insoluble base or insoluble carbonate

1. Add dilute acid into a beaker and heat using a bunsen burner flame

2. Add the insoluble metal, base or carbonate, a little at a time, to the warm dilute acid and stir until the base is in excess (i.e. until the base stops disappearing and a suspension of the base forms in the acid)

3. Filter the mixture into an evaporating basin to remove the excess base

4. Heat the solution to evaporate water and to make the solution saturated. Check the solution is saturated by dipping a cold, glass rod into the solution and seeing if crystals form on the end

5. Leave the filtrate in a warm place to dry and crystallize

6. Decant excess solution and allow crystals to dry or blot to dry with filter paper

example of soluble salt preparation: acid (dilute sulfuric acid) to insoluble base (copper(II) oxide)

copper(II) oxide + sulfuric acid gives copper(II) sulphate + water

CuO (s) + H₂SO₄ (aq) gives CuSO₄ (aq) + H₂O (l)

insoluble salt preparation: two soluble salts by precipitation

1. Dissolve soluble salts in water and mix together using a stirring rod in a beaker

2. Filter to remove precipitate from mixture

3. Wash filtrate with  distilled water to remove traces of other solutions

4. Leave in an oven to dry

example of insoluble salt preparation: two soluble salts (lead(II) nitrate and potassium sulfate) by precipitation

lead(II) nitrate + potassium sulfate gives lead(II) sulfate + potassium nitrate

Pb(NO₃)₂ (aq) + K₂SO₄ (aq) gives PbSO₄ (s) + 2KNO₃ (aq)

solubility rules

sodium, potassium, and ammonium salts: all are soluble

nitrates: mostly soluble (except Ag, lead(II) Pb(NO3)2)

chlorides: mostly soluble (except Ba, Ca, lead(II) Pb(NO3)2)

carbonates: mostly insoluble (except Na, K, NH4)

hydroxides: mostly insoluble (except Na, K, NH4, Ca is partially soluble)

hydrated substance

a substance that is chemically combined with water

anhydrous substance

a substance containing no water

example of copper(II) sulfate becoming a hydrated and anhydrous salt

copper(II) sulfate crystallises forming the salt hydrated copper(II) sulfate, which is blue

when heated, the water from its structure is removed, forming anhydrous copper(II) sulfate, which is white

can be reversed by adding water to anhydrous copper(II) sulfate:

hydrated copper(II) sulfate reversibly gives anhydrous copper(II) sulfate + water

water of crystallisation

water molecules present in hydrated crystals

e.g. cobalt(II) chloride and copper(II) sulfate

anhydrous to hydrated salt:

CuSO4 + 5H2O gives CuSO4∙5H2O

CoCl2 + H2O gives CoCl2∙6H2O

hydrated to anhydrous salt (by heating):

CuSO4∙5H2O gives CuSO4 + 5H2O

 CoCl2∙6H2O gives CoCl2 + H2O