1.1.3 Acids and bases
What are acids and bases? Acids can be described as proton donors; this is known as the Lowry- Bronsted definition of an acid. Any substance which has the ability to produce an H+ ion is capable of acting as an acid. This means that compounds will be capable of acting as acids if they: a) are ionic, and contain hydrogen with a negative counter ion eg HNO3; b) are covalent and contain H bonded to an electronegative atom eg HOCH3. Compounds can exist with hydrogen as a negative ion (hydride); when the counter ion is an electropositive metal eg in NaH. These compounds are not acidic: not all compounds with hydrogen present, act as acids.
The release of a proton in aqueous solution can be shown by the reaction below, for hydrochloric acid.
+ - HCl(aq) + H2O(l) → H3O (l) + Cl (aq)
The proton is actually taken up by a water molecule.
In order to classify substances as bases or alkalis, we need to test their solubility and whether or not they react with an acid
We can summarise the results in the table below.
Substance Solubility in Effect on pH Did it Alkali or water of the HCl neutralise the base (or hydrochloric neither) acid? MgO (s) Insoluble Increases it √ Base NaOH (s) Soluble Increases it √ Alkali NaHCO (s) 3 Soluble Increases it √ Alkali sulphamic acid (s) Lowers (or Soluble doesn’t X Acid change it)
Technical: Requires 8 test tubes, spatula, substances, UI, rack, bung
1 A base can then be described as a proton acceptor; this is known as the Lowry-Bronsted definition of a base. If a substance is capable of accepting a proton, then it is capable of neutralising an acid; or raising the pH of a solution. Common bases are metal oxides, metal hydroxides and ammonia. The release of hydroxide in aqueous solution can be shown by the equation below.
+ - NaOH(s) + aq → Na (aq) + OH (aq)
Activity 1. Match the acids with their formulae
Name of acid Formula hydrochloric HCOOH
sulphuric HNO3
nitric H3CCOOH ethanoic HCl
methanoic H2SO4
phosphoric H3PO4
Activity 2. Match up the formulae with the bases
Name of base Formula sodium hydroxide Ca(OH)2
ammonia H3CNH2 calcium hydroxide NaOH methyl amine MgO
magnesium oxide NH3
You may also see aqueous ammonia described as ammonium hydroxide. This is because the ammonia reacts with water to produce ammonium hydroxide.
Dot and cross diagram for ammonium hydroxide
2 The reactions of acids
Acids commonly react with alkalis, bases and carbonates. An alkali is a soluble base. A base is a compound which can neutralise an acid.
Acids reacting with alkalis
Write balanced equations, with state symbols, to show the following reactions (in pencil). a) hydrochloric acid with sodium hydroxide HCl + NaOH → NaCl + H O ______(aq) (aq) (aq) 2 (l) b) nitric acid with calcium hydroxide
2HNO3(aq) + Ca(OH)2(aq) → CaCl2(aq) + 2H2O(l) ______c) sulphuric acid with ammonia H SO + NH → (NH ) SO ______2 4(aq) 3(aq) 4 2 4(aq)
Acids reacting with bases
Write balanced equations, with state symbols, to show the following reactions. a) hydrochloric acid with magnesium oxide
2HCl(aq) + MgO(s) → MgCl2(aq) + H2O(l) ______b) sulphuric acid with calcium oxide H SO + CaO → CaSO + H O ______2 4(aq) (s) 4(aq) 2 (l) c) phosphoric acid with copper (ii) oxide 2H PO + 3CuO → Cu (PO ) + 3H O ______3 4(aq) (s) 3 4 2(aq) 2 (l)
Acids reacting with carbonates
Write balanced equations, with state symbols, to show the following reactions. a) hydrochloric acid with calcium carbonate
2HCl(aq) + CaCO3(aq) → CaCl2(aq) + H2O(l) + CO2(g) ______
3 b) sulphuric acid with lithium carbonate
H2SO4(aq) + Li2CO3(aq) → Li2SO4(aq) + H2O(l) + CO2(g) ______c) ethanoic acid with sodium carbonate
2H3CCOOH(aq) + Na2CO3(aq) → 2H3CCOONa(aq) + H2O(l) + CO2(g) ______
Ionic equations We can also represent the reaction of an acid with an alkali by using an ionic equation. So the reaction between hydrochloric acid and sodium hydroxide becomes:
+ - H (aq) + OH (aq) → H2O(l)
The spectator ions, Na+ and Cl- have been removed from the original equation; they don’t really do anything or change their oxidation state.
The guidance from the examiners A spectator ion does not change. If the state changes, then the ion is not a spectator. For example, solid sodium carbonate with acid should be shown as an ionic equation as follows:
+ + Na2CO3(s) + 2H (aq) –> 2Na (aq) + CO2(g) + H2O(l)
The sodium ions have clearly changed as they have gone from part of an ionic lattice to hydrated ions.
But aqueous sodium carbonate and acid is:
2– + CO3 (aq) + 2H (aq) –> CO2(aq) + H2O(l)
The sodium ions can be omitted as they are (aq) throughout and do not change.
Write ionic equations, with state symbols, to show the following reactions (in pencil). i) hydrochloric acid with sodium hydroxide
+ + + - H (aq) + NaOH(s) → Na (aq) + H2O(l) or... H (aq) + OH (aq) → H2O(l) ______ii) hydrochloric acid with magnesium oxide
+ 2+ 2H (aq) + MgO(s) → Mg (aq) + H2O(l) ______iii) sulphuric acid with lithium carbonate
+ + 2H (aq) + Li2CO3(s) → 2Li (aq) + H2O(l) + CO2(g) ______
+ 2- or... 2H (aq) + CO3 (aq) → H2O(l) + CO2(g)
4 Strong and weak acids
When a strong acid ionises in water, its molecules fully dissociate fully to form hydrogen ions and chloride ions. The hydrogen ions are also known as + protons. The proton is donated to the water to form an oxonium ion (H3O ).
+ - HCl(aq) + H2O(l) → H3O (aq) + Cl (aq)
The position of equilibrium lies so far to the right hand side that all the reactants have been converted into product.
A weak acid only partially dissociates in water and an equilibrium is established which lies to the left hand side; most of the acid has not been dissociated. This is the case for most organic acids.
+ - CH3COOH(aq) + H2O(l) → H3O (aq) + CH3COO (aq)
Some of the protons are donated to water; there are four species in the reaction mixture.
Activity True or false (write a T or F next to each statement)
T a) Hydrochloric acid is classed as a strong acid because it fully dissociates in aqueous solutions. F b) A strong acid has a high number of moles per volume of solution. T c) A weak acid is capable of reacting with carbonates. T d) Methanoic, ethanoic and citric acids are all weak acids. F e) Aqueous sulphuric acid contains more sulphate ions than oxonium 2- + ions.(H2SO4(aq) + 2H2O(l) → SO4 (aq) + 2H3O (l))
Calculation to find formula of a hydrated salt
0.05 moles of the salt H2C2O4.nH2O weighs 6.3 g
Find the mass of 1 mole of hydrated salt Use this to find the value of n
Mass of 1 mole hydrated salt = (1/0.05) x 6.3 = 126 g Mr H2C2O4 = (2 x 1.0) + (2 x 12.0) + (4 x 16.0) = 90.0 Therefore total mass of H2O in 1 mole hydrated salt = 126 – 90 = 36 g So n = 36/18 = 2
Alternative ratio method (without calculating mass of 1 mole) Mr H2C2O4 = (2 x 1.0) + (2 x 12.0) + (4 x 16.0) = 90.0 Mass of 0.05 moles H2C2O4 = 90.0 x 0.05 = 4.5 g
Assume mass H2O = 6.3 – 4.5 = 1.8 g H2O Moles H2O = 1.8/18.0 = 0.1 moles
Molar ratio H2C2O4 : H2O is 0.05 : 0.1 ≡ 1 : 2 therefore n = 2
5 Salts
A salt is produced when the H+ ion of an acid is replaced by a metal ion or + NH4 . The reactions described above all produce salts; these are normally produced as aqueous solutions.
Activity How could you make the following salts?
a) Ammonium chloride; HCl + NH3 → NH4Cl ______
b) Copper (ii) nitrate; 2HNO3 + CuO → Cu(NO3)2 + H2O ______
c) Sodium iodide; HI + NaOH → NaI + H2O or 2HI + Na2CO3 → 2NaI + H2O + CO2 ______
d) Lithium phosphate;
H3PO4 + 3LiOH → Li3PO4 + 3H2O or 2H3PO4 + 3Li2CO3 → 2Li3PO4 + 3H2O + 3CO2 ______
e) Calcium ethanoate. 2H3CCOOH + Ca(OH)2 → Ca(H3CCOO)2 + 2H2O
or... 2H3CCOOH + CaO → Ca(H3CCOO)2 + H2O ______
or... 2H3CCOOH + CaCO3 → Ca(H3CCOO)2 + H2O + CO2
If the solutions are left to evaporate, at room temperature, then the hydrated salts are formed. A hydrated salt has the metal ion, the counter ion and water of crystallisation.
Example Name hydrated copper (ii) sulphate Metal ion Cu2+ 2- Counter ion SO4 Formulae CuSO4.5H2O
If the blue hydrated copper (ii) sulphate is heated strongly the water of crystallisation is removed to produce the white anhydrous salt, CuSO4.
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