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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

• All acids contain hydrogen
• All bases have a lone pair of electrons to bond to hydrogen

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 ${\mathrm{NH}}_4^{+}$ is NH₃

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

• OH^-
• HCO₃⁻
• ^H₂PO₄−

^{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:}

• ^{Contains hydrogen}
• ^{Has a lone pair of electrons to accept a hydrogen ion.}

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	${{\mathrm{HNO}}_3}_{\left(\mathrm{aq}\right)}+{\mathrm{Mg}}_{\left(\mathrm{s}\right)}\to {{\mathrm{MgNO}}_3}_{\left(\mathrm{aq}\right)}+{\mathrm{H}}_{2\left(\mathrm{g}\right)}$
Acid-metal hydroxide	Acid + metal hydroxide = Salt + water	(a)H + (b)OH = (a)(b) + H2O	$Mg(OH{{)}_2}_{\left(\mathrm{s}\right)}+{\mathrm{H}}_2{\mathrm{SO}}_4{ }_{\left(\mathrm{aq}\right)}\to Mg{{\mathrm{SO}}_4}_{ \left(\mathrm{aq}\right)}+2{\mathrm{H}}_2{\mathrm{O}}_{ \left(\mathrm{l}\right)}$
Acid-metal/hydrogen carbonate	Acid + metal/hydrogen carbonate = Salt + water + carbon dioxide	(a)H + (b)CO3 = (a)(b) + H2O + CO2	$NaHC{O_3}_{ \left(s\right)}+HCl{ }_{\left(aq\right)}\to NaC{l}_{\left(aq\right) }+H_2{O}_{ \left(l\right) }+C{O}_2{ }_{\left(g\right)}$

Acid-Base neutralization reactions

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

• This neutralization always occurs when the pH of the solution is 7. 
• The solution forms a salt, dependent on the acid and base.
• The reaction is always exothermic and thus the enthalpy change  (ΔH)  is negative

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: 

$$\begin{gathered} NaOH + HCl \to \mathrm{NaCl} + {\mathrm{H}}_2\mathrm{O} \\ {\mathrm{Na}}^{+}+{\mathrm{OH}}^{-}+{\mathrm{H}}^{+}+{\mathrm{Cl}}^{-}\to \mathrm{NaCl}+{\mathrm{H}}_2\mathrm{O} \end{gathered}$$

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.

• $H_{2}S{O}_4+2N{H}_3\to (N{H}_4{)}_{2}S{O}_4$
• ${\mathrm{CuO}}_{\left(\mathrm{s}\right)}+{\mathrm{H}}_2{\mathrm{SO}}_4{ }_{\left(\mathrm{aq}\right)}\to {\mathrm{CuSO}}_4{ }_{\left(\mathrm{aq}\right)} + {\mathrm{H}}_2\mathrm{O}{ }_{\left(\mathrm{l}\right)}$

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.

Procedure

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

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.

Indicators

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

$pH = -\log \left(\left[H^{+}\right]\right)$

Concentration of hydrogen ions = [H⁺]

Concentration of hydroxide ions = [OH^-]

Example hydrogen ion conc. to pH question

$$\begin{gathered} \left[H^{+}\right] = 0.01 mol \\ pH = -{\log }_{10}\left(0.01\right) = 2pH \end{gathered}$$

Using pH to distinguish between acid, basic and neutral

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

• pH<7  → acidic
• pH = 0 → neutral
• pH>7 → basic

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.

$H_{2}O\iff H^{+}+O{H}^{-}$

The equilibrium value is given by

 $K=\frac{\left[H\right]\left[OH\right]}{\left[H_{2}O\right]}$

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

Buffers

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. 

CO₂ (g) + H₂O (l) ⇌ H₂CO₃ (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) + 3O_2 (g) → 2SO_3 (g)

SO_3 (g) + H₂O _(l) → H₂SO_4 (aq)

H₂SO_{4 (aq)} → H⁺_(aq) + HSO₄^-_(aq)

 

 

S (s) + O₂ (g) → SO₂ (g)

SO₂ (g) + H₂O (l) → H₂SO_3 (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:

N_{2 (g)} + 2O_2 (g) → 2NO_2 (g)

2NO_2 (g) + H₂O _(l) → HNO_3 (aq) + HNO_2 (aq)

HNO_3 (aq) → H⁺_(aq) + NO₃^-_(aq)

HNO_2 (aq) ⇌  H⁺_(aq) + NO₂^-_(aq)

 

Ways to stop/prevent acid rain

• Wash coal to remove sulphur
• Catalytic converter on cars, prevent release of sulphur dioxide.
• Alternative cleaner fuels
• Instal "scrubbers" in power plant to clean emissions of sulphur dioxide before it leaves the plant

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