Acids and Bases

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Theories of acids and bases

Brønsted-Lowry Acid-Base theory

This theory gives a new definition for acids and bases. The fundamental idea is that when acids and bases react the acid forms its conjugate base and the base forms its conjugate acid via the exchange of protons.

Brønsted–Lowry acid: A proton donor

Brønsted–Lowry base: A proton acceptor

Deduce the acid and the base in a reaction

Acid+base examples

These are not given in the data booklet so they will need to be memorized.

Name of acid Formula
Hydrochloric acid HCl
Nitric acid HNO3
Sulfuric acid H2SO4
Ethanoic acid CH3COOH
Carbonic acid  H2CO3
Phosphoric acid  H3PO4
Benzoic acid C6H5COOH

Conjugate acid-base pairs

These are pairs of acid and bases which differ by one hydrogen ion (proton). 

E.g the conjugate base of NH4+ is NH3

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Amphiprotic and amphoteric species

Amphiprotic species: is a compound which is able to gain a  hydrogen ion to form a conjugate acid or lose a hydrogen ion to form a conjugate base. Thus these species act as both a Brønsted-Lowry acid and base.

Amphoteric species: Can act as either a base or an acid. All amphiprotic species are also amphoteric.

Common amphiprotic species

Deduce if a species is amphiprotic

Very often  in multiple choice you will be asked to identify the amphiprotic species from a list of compounds. To identify it check each option one by one to see if it fits both the acid and the base criteria. If it does, then congratulations you've identified the amphiprotic species.

Amphiprotic must fit the following properties:

Properties of acids and bases

Acid reactions

Acids undergo reaction with different elements such as metals, metal hydroxides and metal/hydrogen carbonates. Each of these have different products. These need to be memorized.

Type of reaction Reaction word equation Pseudo-equation Example
Acid-metal Acid + metal = Salt + hydrogen (a)H + (b) = (a)(b) + H2 HNO3(aq)+Mg(s)MgNO3(aq)+H2(g)
Acid-metal hydroxide Acid + metal hydroxide = Salt + water (a)H + (b)OH = (a)(b) + H2O Mg(OH)2(s)+H2SO4 (aq)MgSO4 (aq)+2H2O (l)
Acid-metal/hydrogen carbonate Acid + metal/hydrogen carbonate = Salt + water + carbon dioxide (a)H + (b)CO3 = (a)(b) + H2O + CO2 NaHCO3 (s)+HCl (aq)NaCl(aq) +H2O (l) +CO2 (g)

Acid-Base neutralization reactions

In neutralization reactions an acid and a base react together to form a salt and water.

Ionic formula and spectator ions

The reason it neutralizes is due to the double-displacement reaction which occurs. We can show this by writing the ionic formula

For example: 

NaOH + HCl NaCl + H2ONa++OH-+H++Cl-NaCl+H2O

Determining formula of salt produced

  1. Remove a hydrogen from the acid
  2. Remove the hydroxide/ammonia part of the base
  3. Combine together and balance

Non-hydroxide neutralization examples

The base can be a hydroxide but it is also good to know what happens with ammonia or an oxide.

Acid base titrations

An acid-base titration is used to find the concentration of an unknown acid/base, if the concentraton of one of them is known.


The procedure is to add the acid/base solution with the known concentration (titrant) to the solution with the unknown concentration (Anylate) slowly until the reaction reaches neutralization. This point can be determined using an indicator which changes colour dependent on the pH of the solution. You record the volume of titrant required to neutralize the anylate then calculate the moles and determine the concentration of the anylate.

Titrant: The solution with the known concentration in buret, can be an acid or a base

Analyte: The solution with the unknown concentration in conical flask, can be an acid or a base.

Buret: a graduated glass tube with a tap at one end, for delivering known volumes of a liquid in titrations.

Titration curve

Image result for acid base titration curves

Image result for acid base titration curves

Image result for Acid base titration diagram

Equivalence point: is the point in a titration where the amount of titrant is enough to completely neutralize the analyte. Thus at this point the moles of titrant is equal to the moles of the solution with unknown concentration. The straight vertical line in the titration curve.


A good indicator is bright, colorful and has a distinct change depending on whether acidic or basic. Indicators must also have their changing point within the range of the neutralization. These are themselves either weak acids or weak bases which react in a reversible reaction which depends on the acidity of the surroundings.

Phenolphthalein: A ph indicator which is colourless in the presence of an acid and pink in the presence of a base. It helps determine the point of neutralisation

Buffer region

The region where the acids and bases don't have the same concentration, so there are still products or reactants to accept or donate hydrogen ions

Thermometric titrations

Heat is released in the neutralisation reaction. However, once the reaction has been neutralized it will stop releasing heat. This means you can determine the equivalence point without an indicator.

Enthalpy of neutralisation: Energy released when 1 mol of  water is produced through the neutralisation of an acid and a base.

The pH scale

What is pH?

pH is a measure of the concentration of H+ ions within a solution.  The lower the pH, the lower the concentration; the higher the pH, the greater the concentration. The idea of pH is to map the wide range of concentrations onto a value ranging from 0 to 14. This is done through the use of a logarithm. The formula for pH is given below

pH = -log(H+)

Concentration of hydrogen ions = [H+]

Concentration of hydroxide ions = [OH-]

Example hydrogen ion conc. to pH question

H+ = 0.01 molpH = -log100.01 = 2pH

Using pH to distinguish between acid, basic and neutral

the pH of a solution tells us whether that solution is acid, basic or neutral. 

Comparing pH

As the pH scale is logarithmic a one unit change in pH represents a 10-fold change in [H+]

So going from pH 4 to pH 0, is a 10,000 fold increase in [H+] !

The ionic product of water \(K_w\)

This is a value which allows us to determine the relationship between [H+] and [OH-]  for aqueous solutions.

When an acid and a base reacts, the hydrogen and hydroxides disassociate and then combine together in a double displacement reaction. We can split this up into two partial reactions, one with the salt and one with water. The water one looks like this.


The equilibrium value is given by


As the reaction lies so far to the left the water concentration has a constant value  and we can ignore it. This gives us an expression for what's called the ionic product constant of water

\(K_w\)= \( [H^+][OH^-] \)


The value of \(K_w\) depends on the temperature, however at 298K this value is known to be \(1.0 \times 10^{-14}\)

This means if we know the concentration of either hydrogen or hydroxides we can calculate the concentration of the unknown one.

Strong and weak acids and bases

Definition of strong and weak

Strong acid/base: Ions fully dissociate. pH either very low or very high

Weak acid/base: Ions partially dissociate, in a reversible reaction. pH on either side of 7, but not too extreme.

Some must-know strong and weak

Strong acids: HCl, H2SO4, HNO3

Strong bases: Alkali hydroxides (LiOH, NaOH, KOH, RbOH, CsOH), Ba(OH)2

Weak acids: CH3COOH, H2SO3, HNO2, H3PO4

Weak bases: Mg(OH)2, NH3

Properties of strong vs weak

  Strong Weak
pH At ends of spectrum Towards middle of spectrum
Electrical conductivity High Low
Rate reaction High Low


A solution which resists pH changes upon the addition of small amounts of strong acid or base. It can either accept hydrogen ions or donate them depending on the change in acidity.

Acid deposition

Any precipitation of acid rain

Why is rain acidic

Acid rain contains dissolved CO2 which lowers its pH to around 5.6. 

CO(g) + H2O (l) ⇌ H2CO(aq)

Forms hydrogen carbonate. This process above is NATURAL, this is NOT considered "acid rain".

Acid rain comes from oxides of 2 main elements: Sulfur and Nitrogen.

Sulfuric and sulfurous acid

Sulfur can be turned into gaseous oxides from burning of fossil fuels. The sulfur present in amino acids (ex. cysteine) from decayed dinosaurs and plants is combusted. This then reacts with water to form sulfuric acid (strong acid) or sulfurous acid (weak acid).


S (s) + 3O2 (g) → 2SO3 (g)

SO3 (g) + H2(l) → H2SO4 (aq)

H2SO4 (aq) → H+(aq) + HSO4-(aq)



S (s) + O2 (g) → SO2 (g)

SO2 (g) + H2O (l) → H2SO3 (aq)

H2SO3 (aq) ⇌  H+(aq) + HSO2-(aq)


Nitric and nitrous acid

Nitrogen can be turned into gaseous oxides in the heat of combustion engines. Usually nitrogen gas and oxygen gas will not react, but at high temperatures (ex. in an engine), they have enough energy to overcome the Activation Energy. This then reacts with water to form nitric acid (strong) and nitrous acid (weak).

Forming nitric acid and nitrous acid:

N2 (g) + 2O2 (g) → 2NO2 (g)

2NO2 (g) + H2(l) → HNO3 (aq) + HNO2 (aq)

HNO3 (aq) → H+(aq) + NO3-(aq)

HNO2 (aq) ⇌  H+(aq) + NO2-(aq)


Ways to stop/prevent acid rain


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