Acids and Bases Flashcards
Arrhenius acid-base theory
● Arrhenius acids give hydronium ions, H3O
+ (aq) [or hydrogen ions,
H+ (aq)] after dissolving in water.
Examples: HCl (aq), HNO3 (aq), CH3COOH (aq) and H2SO4 (aq)
● Arrhenius bases give hydroxide ions, OH-(aq) in water.
Examples: NaOH (aq), KOH (aq), Ca(OH)2 (aq)
Acid dissociation
● When an acid dissolves in water, the acid molecule react with H2O molecules by
giving out a hydrogen ion, to form hydronium ions, H3O
+
and a negative ion.
● For example: HCl (aq) + H2O (l) → H3O
+ (aq) + Cl- (aq)
what does proticity mean
number of moles of hydrogen ions/ protons produced
per mole of acid molecules when dissolved in water.
monoprotic acids
Monoprotic acids produce one mole of H+ (aq) ions per mole of acid molecules.
Examples: HCl (aq), HNO3
(aq) and CH3COOH (aq)
■ CH3COOH (aq) + H2O (l) ⇌ CH3COO- (aq) + H+(aq)
polyprotic
- diprotic acids produce two moles of H+ (aq) ions and triprotic
acids produce three moles of H+ (aq) ions per mole of molecules.
Examples: H2SO4
(diprotic) and H3PO4
(triprotic)
Brønsted-Lowry theory
In 1923, Johannes Nicolaus Brønsted (from Denmark) and Thomas Martin
Lowry (from England) published essentially the same theory about how acids
and bases behave independently. Both names have been used for the theory.
● The theory classifies a substance as an acid or a
base in terms of whether it donates or receives
H+ (aq) ions/ protons when dissolved in water or
in a reaction.
In this theory, acid and base must work together so that the protons donated
by the acid can be received by the base.
Brønsted-Lowry acid
donates H+ (aq) ions → proton donors
○ Examples: HCl (aq), HNO3 (aq), CH3COOH (aq) and H2SO4 (aq)
Brønsted-Lowry base:
receives H+ (aq) ions → proton acceptors
Examples: NaOH (aq), KOH (aq), NH3 (aq), Na2CO3 (aq), NaHCO3 (aq)
Brønsted-Lowry theory thinking
● Consider the following generic equation for a reaction between an acid, HA (aq) and a base, B (aq):
HA (aq) + B (aq) → A- (aq) + BH+ (aq)
● When looking at the backward process, BH+
donates H+ ion to A- to form B and at the same time, A-receives H+ ion from BH+ to form HA.
○ BH+ (aq) is a Brønsted-Lowry acid and A- (aq) is a Brønsted-Lowry base (for the
backward reaction).
Conjugate acid-base pairs
● In conclusion, two conjugate acid-base pairs can be identified in each
acid-base reaction.
● The identities of ‘acid’ or ‘base’ mainly describe the role of the chemical species as proton donors/ acceptors.
A Brønsted-Lowry acid and its conjugate base always differ by one proton.
A Brønsted-Lowry base and its conjugate acid always differ by one proton.
Dissociation of water (1)
● Water has a tendency to dissociate (break up) into ions when in solution:
H2O (l) + H2O (l) ⇌ H3O
+ (aq) + OH- (aq)
● Pure water dissociates only slightly - about one water molecule out of every 10
million dissociates and the rest remain undissociated (in molecular form).
Dissociation of water (2)
● Water molecule ionizes (separates into ions) to form a hydronium/ hydrogen ion
and a hydroxide ion:
H2O (l) + H2O (l) ⇌ H3O
+ (aq) + OH- (aq)
● As seen in the equation, water can either → Identify the conjugate acid-base pair in this reaction. donate or receive H+ (aq) ion, which means it can be both a proton donor and a proton acceptor.
○ Water can act as both Brønsted-Lowry acid and base.
○ It is described as being amphoteric
Amphoteric species
● Apart from water, the conjugate base of many polyprotic acids are also
amphoteric.
Amphoteric species (1)
Example 1: HPO4
2- (aq)
○ It is the conjugate base of
H2PO4- (aq)
Amphoteric species (2)
● Example 2: HCO3
- (aq)
○ It is the conjugate base of H2CO3
Amphoteric species (3)
HS-
Conjugate base of H2S
Properties of acids and bases
● Acids have sour taste, while bases have bitter taste
and slippery feeling.
● They can be distinguished using universal indicator,
pH paper, litmus solution or pH meter.
● They can only show these properties when dissolved
in water.
Indicators (1)
Litmus solution
It is just used to identify whether an
aqueous solution is acidic or basic.
Acidic: Red
Neutral: Purple
Basic: Blue
Indicators (2)
Universal indicator
It can measure the acidity and the
alkalinity of an aqueous solution by
providing an approximate pH value.
Indicators (3)
pH paper, pH meter
Indicators (4)
● Methyl orange and phenolphthalein are
indicators mainly used in titration
experiments to identify the end-point and
hence determine the concentration of a
solution.
● These indicators show a sharp colour
change when the acid/ base is neutralised.
Acidity and alkalinity of solutions
● In any type of aqueous solutions (acidic, neutral or basic), they all contain both H+
(aq) ions and OH- (aq) ions.
● This is due to the self-ionization / dissociation of water molecules, although the
extent of dissociation is very small.
H2O (l) + H2O (l) ⇌ H3O
+ (aq) + OH- (aq)
what is the acidity of a solution dependent on
● The acidity of a solution depends on the relative concentration of H
+ (aq) ions and
OH- (aq) ions in an aqueous solution.
what is concentration
● Concentration refers to the amount (number of moles) of solute dissolved in a
unit volume (dm3
) of solution.
● In 1 dm3
solution, if solution A contains larger amount of solute than solution B,
we will describe them as either:
○ solution A is more concentrated than solution B, or
○ solution B is more dilute than solution A.
Note: 1 dm3
= 1000 cm3
how is concentration calculated
number of moles of solute (mol)/volume of solution (dm3)
● Unit: mol dm-3 or M
● Example:
Concentration of solution when 0.05 moles of HCl dissolved in 100 cm3
of water
= 0.05 mol ÷ 0.1 dm3
= 0.5 mol dm-3
what is pH
pH = -log10 [H3O
+ (aq)] or pH = -log10 [H+ (aq)]
● In chemistry, “p” in pH represents ‘-log10’ and “H” represents [H+ (aq)]. pH can
hence be expressed mathematically.
● On the other hand, as pH = -log10 [H+ (aq)], making [H+ (aq)] the subject,
[H+ (aq)] = 10-pH
Relationship between [H+ (aq)],
[OH- (aq)] and pH value
● As pH decreases (down the chart), the acidity of
solution increases, and the solution becomes
more acidic.
● When the pH value of an aqueous solution
decreases by one unit, [H+ (aq)] increases by
10-fold while [OH- (aq)] decreases by 10-fold.
When pH increases by 2 units,
[H+ (aq)] decreases by 10^2
times.
[OH-(aq)] increases by 10^2 times
when pH increases by x units,
[H+ (aq)] would decrease by
10^x times.
when pH decreases by x units
[H+ (aq)] would increase by
10^x times
Reaction of acids with metals
● Dilute acids such as HCl (aq) and H2SO4 (aq) react
with reactive metals (those above hydrogen in the
Reactivity Series) to produce a salt and hydrogen
gas, H2 (g).
● The H2 (g) produced can be tested by a burning
splint - a “squeaky pop” sound can be heard.
reactivity series
potassium K
sodium Na
calcium Ca
magnesium Mg
aluminium Al
carbon C
zinc Zn
iron Fe
tin Sn
lead Pb
hydrogen H
copper Cu
silver Ag
gold Au
platnium Pt
what is the salt
A salt is a compound when the ionizable hydrogen atom(s) of an acid is replaced
by metal ion(s) or other cation(s).
give the Reaction of acids with metals (Zn and Mg with HCl and HNO3)
● Reaction:
metal + acid → salt + hydrogen
● Examples:
zinc + hydrochloric acid → zinc chloride + hydrogen
Zn (s) + 2 HCl (aq) → ZnCl2
(aq) + H2 (g)
magnesium + nitric acid → magnesium nitrate + hydrogen
Mg (s) + 2 HNO3 (aq) → Mg(NO3
)2 (aq) + H2 (g)
Reaction with metal hydroxides or oxides
● This reaction is known as neutralisation.
● The reaction is exothermic, which means heat is released during the process.
● When an acid reacts with a base, salt and water are produced:
acid + metal oxide/ hydroxide → salt + water
● Example: sulphuric acid + sodium oxide → sodium sulphate + water
H2SO4 (aq) + Na2O (s) → Na2SO4
(aq) + H2O (l)
Reaction with metal carbonates and hydrogencarbonates
● Dilute acids react with metal carbonate or hydrogencarbonate to give salt, water
and carbon dioxide gas, CO2 (g).
acid + metal carbonate/ hydrogencarbonate → salt + water + carbon dioxide
● Example:
hydrochloric acid + calcium carbonate → calcium chloride + water + carbon dioxide
2 HCl (aq) + CaCO3 (s) → CaCl2
(aq) + H2O (l) + CO2 (g)
Reaction with metal carbonates and hydrogencarbonates (testing of CO2)
● The produced CO2 (g) can be tested with limewater which is
saturated calcium hydroxide solution, Ca(OH)2 (aq).
The colourless solution turns milky.
● The observation can be explained using the following equation:
Ca(OH)2 (aq) + CO2 (g) → CaCO3 (s) + H2O (l)
● The product in this reaction - calcium carbonate, CaCO3 (s) - is
an insoluble white compound
Reaction of acids with ammonia
● Acids react with ammonia to produce an ammonium salt:
acid + ammonia → ammonium salt
● Example:
(1) hydrochloric acid + ammonia → ammonium chloride
HCl (aq) + NH3
(aq) → NH4Cl (aq)
(2) sulphuric acid + ammonia → ammonium sulphate
H2SO4 (aq) + 2 NH3 (aq) → (NH4)2SO4(aq)
