MODULE 6
IQ: What is an acid and what is a base?
Properties of acids and bases
Properties of acids
Properties of bases
● Turn litmus indicator red
● Tend to be corrosive
● Taste sour
● React with bases
● Solution have a pH below 7.0
● Turn litmus indicator blue
● Are caustic and feel slippery
● Taste bitter
● React with acids
● Solutions have a pH above 7.0
● An alkali is a soluble base (generally
containing hydroxide)
Conjugate acid-base pairs
+
● A conjugate acid-base pair are two molecules or ions that differ by one proton (𝐻 )
+
−
● HCl(aq) + H2O(l) → 𝐻3𝑂 +(aq) + 𝐶𝑙 (aq)
+
- HCl and Cl- differ by one proton (𝐻 ) hence are a conjugate acid-base pair
+
- 𝐻3𝑂 and H2O also differ by one proton hence are a conjugate acid-base pair
- Cl- is formed by HCl losing a proton, hence is called the conjugate base of HCl
−
- Similarly, HCl is the conjugate acid of 𝐶𝑙
● When writing these conjugate pairs, we must remove species which make the pairing incorrect,
for example:
- NaOH + HCl → NaCl + H2O
−
- The acid/conj base pairing is: HCl/𝐶𝑙 and NOT HCl/NaCl, as these two species do NOT
differ by a single proton, rather differ by a proton and a sodium atom.
→ Hydrochloric acid (HCl) can only exist in an aqueous solution (aq) or as a gas (g) → but as a gas is not
called hydrochloric acid, rather is called hydrogen chloride gas
Amphiprotic substances
● Substances which can either donate or accept protons (depending on the substance they are
reacting with) are called amphiprotic
+
● Example is water, when acids react with water hydronium (𝐻3𝑂 ) ions are produced - water gains
−
a proton. When bases react with water, hydroxide (𝑂𝐻 ) ions are produced - water donates a proton.
Monoprotic and polyprotic acids
● Monoprotic acids can only donate one proton. E.g HCL, HF, HNO3, CH3COOH
● Polyprotic acids can donate more than one proton. This does not happen all in one go, however it occurs
in a number of steps depending on the acid.
Acid/base reactions
- Diprotic acids donate two protons (E.g H2SO4, H2CO3) and do so in two steps
- Triprotic acids can donate three protons. E.g H3PO4, H3BO3 and do so in three steps
Bonding
● Hydronium ions form when the hydrogen ion bonds to a water molecule by sharing one of the
electron pairs of the oxygen atom. This type of bond is a type of covalent bond called a
coordinate bond
- Coordinate bond: A type of covalent bond in which one atom provides the shared
electron bonding pair
● Acid-base indicators are natural or synthetic chemical dyes that change colour as the acidity of
the surroundings change
● Indicators are organic molecules that can change colour as the concentration of acid or base
Changes
Indicator
High acidity
Low acidity
Neutral
Low alkalinity
High alkalinity
Methyl orange
Red
Orange to
yellow
Yellow
Yellow
Yellow
Bromothymol
Blue
Yellow
Yellow
Green
Blue
Blue
Phenolphthalein
colourless
colourless
colourless
pink
crimson
Litmus
Red
Red
Purple
Blue
Blue
● Acidity in water solutions is caused by hydronium ions whereas alkalinity is caused by hydroxide
ions
● Universal indicator (UI) is a mixture of different indicators. Different brands of universal
indicators vary in their colour produced at different acid-base levels and hence a colour card
should be used
UI colour
Strongly
acidic
pH 0-3
Moderately
acidic
pH 4-5
Weakly
acidic
pH 6
Neutral
pH 7
Weakly
alkaline
pH 8
Moderately
alkaline
pH 9-10
Strongly
alkaline
pH 11-14
Red
Orange
Yellow
Green
Greenish
-blue
Bluish-violet Violet
+
−
Bromothymol blue reversible reaction: HBB + H2O ⇌ 𝐻3𝑂 + 𝐵𝐵
+
● 𝐻3𝑂 represents an acid, therefore if we add HCl (an acid) the system will shift to left as there is
+
an increase in 𝐻3𝑂
- Think of equilibrium and the way it will shift
+
● If NaOH (a base) is added, then the OH- ions from the NaOH will react with 𝐻3𝑂 to create H2O,
therefore the equilibrium will shift right
Predict the products of acid reactions and write balanced equations to represent
- Acids and bases
● Acid + Base → Salt + Water
Reaction 1:
- Whole equation: MgO (s) + H2SO4 (aq) → MgSO4 (aq) + H2O (l)
+
+
- Net ionic: MgO (s) + 2𝐻3𝑂 (aq) → 𝑀𝑔 (aq) + 2H2O (l)
Reaction 2:
- Whole equation: Fe(OH)3 (s) + 3HCl (aq) → FeCl3 (aq) + 3H2O (l)
+
+
- Net ionic: Fe(OH)3 (s) + 3𝐻3𝑂 (aq) → 𝐹𝑒 (aq) + 6H2O (l)
Note: Some acid-base reactions won’t form water, however whenever a metal hydroxide reacts with an
acid, water will always be a product.
- Acids and carbonates
● Acid + Carbonate → Salt + Water + Carbon dioxide
Reaction 1:
- Whole equation: NiCO3 (s) + 2HNO3 (aq) → Ni(NO3)2 (aq) + H2O (l) + CO2 (g)
+
+
- Net ionic: NiCO3 (s) + 2𝐻3𝑂 (aq) → 𝑁𝑖2 + CO2 (g) + 3H2O (l)
- Acids and metals
● Acid + Metal → Salt + Hydrogen
- Oxidant: A substance that accepts electrons from another substance called the
reductant
- Reductant: A substance that donates electrons to another substance called the oxidant
Reaction 1:
- Whole equation (redox equation): Mg (s) + H2SO4 (aq) → MgSO4 (aq) + H2 (g)
+
+
- Net ionic: Mg (s) + 2𝐻3𝑂 (aq) → 𝑀𝑔2 (aq) + H2 (g) + 2H2O (l)
Note: In general dilute nitric acid reacts with metals and nitric oxide gas rather than hydrogen gas i
released. In these cases the nitrate ion is the oxidant rather than the hydronium ions.
● A solution is neutralised when the concentration of hydronium and hydroxide ions become equal
Everyday and industrial uses of neutralisation
● Vinegar used to cure bites that are alkaline (wasp stings)
● Baking powder used to cure bites that are acidic (bee stings, ant bites)
● Toothpaste contains bases that neutralise acids in our mouth
● Antacid (weak base) used to neutralise stomach acid
● Acidic byproducts from industries are neutralised before they escape into the environment
Enthalpy of neutralisation
● Neutralisation: The proton transfer reaction between an acid and a base to produce a salt and
water
● For example the reaction of HCl and NaOH to form NaCl and water
● Represented by the symbol ΔHneut.
● These reactions are exothermic, thus ΔH is negative
● Measuring the enthalpy of neutralisation can be done using calorimetry
- Measure the temperature change of the salt solution (K)
- Know the specific heat capacity of water (4.18 Jg-1K-1)
- Measure the mass of the solution (g)
- Multiply these values to find the amount of energy released by the reaction (J)
q= m x c x ΔT which is simplified to q=mcΔT
● From the equation q=mcat, the q is in J, which can be converted to kJ by dividing by 1000
● This leaves a value in kJ, which can then be converted into kJ/mol to get the change in enthalpy,
whereby the moles which is being divided by the kJ is the number of moles of water, which can
be calculated by using the limiting reactant and mole ratio
→ The amount of heat liberated in a neutralisation reaction can be measured using a polystyrene
calorimeter, which is an excellent heat insulator. The heat liberated by the neutralisation reaction is
absorbed by the solution, which rises in temperature
Year
Individual
Acid Theory
1766
Lavoisier
Acids were substances that contained oxygen
1810
Davy
Acids were substances that contained hydrogen that could be
replaced totally or partially by metals
1884
Arrhenius
Acids were substances that ionised in solution to produce 𝐻
ions
1923
Bronsted and Lowry
Acids are proton donors and bases are proton acceptors
1923
Lewis
Acids accept an electron pair and bases donate an electron pair
+
→ Strong Acid: Acid that completely dissociates in water (same for strong bases)
→ Weak acid: Partially dissociates in water (same for weak bases)
● In 1923, Thomas Martin Lowry (English) and Johannes Nicolaus Bronsted (Danish) independently
proposed new definitions for acids and bases
- Acids: Proton donors
- Bases: Proton acceptors
● If a substance has a greater tendency to donate protons than a particular solvent, then in that
solvent it will be an acid
● In 1923, Gilbert Lewis proposed the electron pair acceptor/electron pair donor theory of acids
and bases
- Lewis acid: An electron-pair acceptor
- Lewis base: An electron-pair donor
Ionic equations that represent the dissociation of acids and bases in water, pairs in solution
and amphiprotic nature of some salts
- Sodium hydrogen carbonate
● NaHCO3
- Acting as a base: NaHCO3 + HCl → CO2 + H2O + NaCl
- Acting as an acid: NaHCO3 + NaOH → H2O + Na2CO3
- Potassium dihydrogen phosphate
● KH2PO4
- Acting as a base: KH2PO4 + HCl → KCl + H3PO4
- Acting as an acid: KH2PO4 + NaOH → KHPO4 + H2O + Na
IQ: What is the role of water in solutions of acids and bases?
Conduct a practical investigation to measure the pH of a range of acids and bases
General Method (using pH probe):
1. Place 50mL of the HCl in a clean beaker and measure its pH using the calibrated pH probe
2. Rinse the probe with deionised water and measure the pH of the acetic acid and citric acid
solutions
3. Re-calibrate the pH probe
4. Repeat procedure for the basic solutions
● pH is affected by concentration
+
Calculate pH, pOH, hydrogen ion concentration ([𝐻 ]) and hydroxide ions concentration ([𝑂𝐻
])
for a range of solutions
+
● pH = -log10[𝐻3𝑂 ]
−
● pOH = -log10[𝑂𝐻 ]
+
−𝑝𝐻
● To go from pH to [𝐻3𝑂 ] = 10
● To work out pH using pOH
1. Find out pOH
2. Subtract this value from 14
3. The resultant value is the pH
−
● Recall that pH + pOH = 14 (at 298 kelvin)
● Understand the interrelationship between all these value
● Autoionisation of water is when to water molecules react with each other to form hydronium
and hydroxide ions through proton transfer, given by the reversible reaction below
+
- 2H2O(l) ⇌ 𝐻3𝑂
−
(aq) + 𝑂𝐻
(aq)
● At 298 kelvin, in pure water, hydronium ions and hydroxide ions exist in equal concentrations,
which have been measured to be 1.0x10-7 mol/L. The ionic product for this equilibrium is given
the symbol Kw and is called the water ionisation constant
+
−
- Kw = [𝐻3𝑂 ][𝑂𝐻 ] … in pure water at 298 kelvin, Kw can be worked out to be:
→ Kw = (1.0x10-7)(1.0x10-7) = 1.0x10-14
● Think about the water ionisation constant in terms of equilibrium. When acids are added to
water the hydronium ion concentration increases and hence equilibrium shifts left according to
LCP consequently reduces the hydroxide ions. Kw is only affected by temperature, just like any
another equilibrium constant.
Conduct an investigation to demonstrate the use of pH to indicate the differences between the
strength of acids and bases
General Method (using pH probe):
5. Place 50mL of the HCl in a clean beaker and measure its pH using the calibrated pH probe
6. Rinse the probe with distilled water and measure the pH of the acetic acid and citric acid
solutions
7. Re-calibrate the pH probe
8. Repeat procedure for the basic solutions
Discussion would talk about why different acids/bases have different pH levels.
Answer = due to their strengths. Strong acid/base completely dissociate therefore have higher pH when
compared to weak acid/base of the same molarity. E.g 1M HCl vs 1M CH3COOH (strong vs weak).
Construct models and/or animations to communicate the differences between strong, weak,
concentrated and dilute acids and bases
Calculate the pH of the resultant solution when solutions of acids and/or bases are diluted or
mixed
● Recall the equation for a neutralisation reaction
- Acid + Base → Salt + Water
● Steps
1. Write a chemical equation (balanced)
2. Determine the number of moles of acid/base reacting, using n = c x v
3. Determine which reactant is in excess (limiting would be completely neutralised)
4. If equal then pH = 7, if not determine how much is in excess, in moles
5. Calculate the concentration of excess reactant using the total volume
−
6. Determine pH of solution using pH = -log10[𝐻3𝑂 ]
IQ: How are solutions of acids and bases analysed?
Conduct practical investigations to analyse the concentration of an unknown acid or base by
titration
● Primary standard: prepared using chemicals that are pure
- Known composition, not hygroscopic, stable, soluble in water, high molecular weight,
undergoes complete reaction
- Benzoic acid (C7H5O2)
- Oxalic acid (H2C2O4•2H2O)
- Sodium carbonate (Na2CO3)
- Sodium hydrogen carbonate (NaHCO3)
● Secondary standard: the concentration is determined using primary standard
- HCl, H2SO4, NaOH
● Titrant: Solution added from burette
● Aliquot: Known volume of a liquid
● Titre: Volume of solution delivered from the burette that achieves an end point
● End-point: the point in the titration where the indicator just changes colour (i.e end of titration
→ observable colour change)
● Equivalence point: the point where the acid has stoichiometrically reacted with base → when
exact neutralisation has occurred
To calculate mass of oxalic acid required to make primary:
1. Calculate molar mass of oxalic acid (126.068g/mol)
2. Convert 250mL to L (0.250L)
3. Calculate moles of oxalic acid required to make 0.0500 M oxalic acid solution using n = c x v
(0.0125 mol)
4. Calculate mass of oxalic acid that should be weighed out using m = n x MM (1.57585g)
To make oxalic acid primary standard:
1. Weight out 1.57585g of oxalic acid using electronic balance and transfer to clean beaker
2. Dissolve oxalic acid by adding just enough deionised water to beaker
3. Transfer oxalic acid from beaker into volumetric flask using funnel, ensuring deionised water is
used to completely transfer all traces of oxalic acid
4. Fill volumetric flask to 250mL mark with deionised water, use pipette as meniscus gets closer to
250mL mark
5. Stopper flask, invert and mix thoroughly
To titrate oxalic acid primary with sodium hydroxide to create secondary standard:
1. Clean burette using deionised water and then with sodium hydroxide
2. Clean 25mL pipette a with deionised water and then with oxalic acid
3. Transfer 25mL oxalic acid from volumetric flask to conical flask using pipette and then add 3
drops of phenolphthalein indicator
4. Add sodium hydroxide to the burette above the zero mark and then open the nozzle to allow
excess NaOH to flow into the chemical waste bucket until the meniscus reaches zero mark. This
ensures no air bubbles in burette
5. Place conical flask containing oxalic acid and indicator beneath burette
6. Allow the sodium hydroxide to flow into the conical flask and stop when a colour change occurs,
note this as a rough titration
7. Repeat this three more times using a new solution of 25 mL oxalic acid and indicator, however
slow down as the end point is reached
8. Note the volume of sodium hydroxide used to induce a permanent colour change
● Know the correct chemical to clean certain equipment, for example
- Pipette used to extract oxalic acid from volumetric flask should be rinsed first with
deionised water and then with oxalic acid solution
- The burette which will contain the sodium hydroxide unknown concentration solution
will firstly be cleaned with deionised water and then with sodium hydroxide solution
- The conical flask where the neutralisation reaction will be taking place should only be
cleaned with deionised water as this will not affect the neutralisation reaction, however
should still be dried to remove any possible errors
● Indicator chosen must have a range within equivalence point in order to observe colour change
Titration curves and conductivity graphs analyse data to indicate characteristic reaction profiles
● Wording → strong acid and weak base (or strong acid titrated against strong base) = strong acid
in conical flask and strong base is to be added from the burette
● Electrolyte → an electrolyte is a substance that contains free ions and behaves as an electrically
conductive medium (i.e conducts electricity)
- The more free ions, the higher the electrical conductivity
- Some ions conduct electricity better than other ions
+
+
→ e.g 𝐻 is a better electrical conductor than 𝑁𝐻4
● Aqueous conductivity → Conductivity in aqueous solutions, is a measure of the ability of water
to conduct an electric current
● The warmer the solution, the higher the solubility of the substance being dissolved and
therefore the higher the conductivity as well
- Strong acid/strong base
Strong acid with strong base conductivity graphs
→ high initial conductivity due to the presence
+
of 𝐻 ions
+
→ conductivity decreases as 𝐻 reacts with OH-,
+
creating 𝑁𝑎 ions which have lower conductivity
→ equivalence point occurs at the point on the
graph with the lowest conductivity
→ the solution at the equivalence point contains
+
−
𝑁𝑎 and 𝐶𝑙 ions and H2O molecules
−
→ conductivity increases as excess 𝑂𝐻 ions are
added
- Strong acid/weak base
Strong acid with weak base conductivity graphs
+
→ high initial conductivity due to high [𝐻 ]
+
→ conductivity decreases as [𝐻 ] decreases,due to
+
reaction with NH3 (for e.g) to yield 𝑁𝐻4 , has less
+
conductivity than 𝐻
→ equivalence point occurs at the lowest point on the
graph - when it is first reached
→ conductivity does not change as excess NH3 is added
as it is a weak base and only dissociates to a small extent
- Weak acid/strong base
Weak acid with strong base conductivity graphs
+
→ initial conductivity is low because [𝐻 ] is
low due to partial dissociation of weak acid
→ conductivity increases as strong base
dissociates into ions, OH- ions from base
react with weak acid molecule and leave
−
−
the acid anion i.e 𝑂𝐻 + HA → H2O + 𝐴
→ equivalence point occurs at the sudden
change in conductivity
Conductivity increases after equivalence
point due to excess strong base dissociating