Chemistry: Acidic: 4

4. Because of the prevalence and importance of acids, they have been used and studied for hundreds of years. Over time, the definitions of acid and base have been refined

Outline the historical development of ideas about acids including those of:
- Lavoisier
- Davy
- Arrhenius
- Lavoisier: In the 1780s, the French chemist, Antoine Lavoisier, found that non-metal oxides reacted with water forming acidic solutions. He concluded that an acid must contain oxygen.
- Davy: In 1815, the English chemist, Humphry Davy, observed that all known acids contained hydrogen that could be replaced by reaction with a metal. He also noted that compounds of metal with oxygen were bases. However, he knew that some compounds containing hydrogen were not acids. He believed that it was the way that a substance was structured that somehow determined its acidity.
(Lavoisier and Davy's definitions were based on observable properties)
- Arrhenius: In 1884, the Swedish chemist, Svante Arrhenius, put forward definitions based on concepts about particles too small to be directly observed. Arrhenius proposed that: an acid produced hydrogen ions H+ when dissolved in water. A base produced hydroxide ions OH- when dissolved in water.
Arrhenius’ model for acids and bases:
- pure acids consist of neutral molecules that dissociate into ions when dissolved in water
- bases are substances that dissociate in water to provide cations and OH- ions
- when an acid or base dissociates, the ions produced are mobile in solution, which is therefore a good conductor
- the behaviour of acids is due to the H+ ion
- the behaviour of bases is due to the OH- ion

Outline the Brönsted-Lowry theory of acids and bases
This theory was independently outlines in 1923 by the Danish chemist, Johannes Bronsted, and the British chemist, Thomas Lowry, which overcame the limitations of the Arrhenius definition. It proposed that an acid is a proton donor and a base is a proton acceptor. The production of hydrogen ions is due to the properties of the acid relative to those of the solvent. If a substance has a greater tendency to give up protons than the solvent, then the substance will be an acid. This theory brought about the idea of conjugates (bases and acids).

Describe the relationship between an acid and its conjugate base and a base and its conjugate acid
When an acid donates a proton, it forms its conjugate base. HCl + H2O Cl- + H3O+ acid conjugate base
When a base accepts a proton, it forms its conjugate acid. HCl + H2O Cl- + H3O+ base conjugate acid

Identify a range of salts which can form acidic, basic or neutral solutions and explain their acidic, neutral or basic nature
Strong acid + strong base = neutral salt eg. NaCl
Strong acid + weak base = acidic salt, eg. NH4Cl
Weak acid + strong base = basic salt eg. CH3COONa
Weak acid + weak base = neutral salt, eg. NH4NO2
Salt ions formed from weak acids or weak bases can react with water to reform the acid or base. In undergoing these hydrolysis reactions, they release OH- or H+, which can produce basic or acidic salt solutions.
Ammonium salt solutions are acidic, because NH4+ + H2O NH3 + H3O+
Sodium chloride solution is neutral, because Na+ and Cl- (ions from the strong base NaOH and the strong acid HCl) do not undergo hydrolysis.
Sodium carbonate solution is basic, because the carbonate ion from the weak acid carbonic acid can hydrolyse. CO32- + H2O HCO3- + OH-
Similarly, potassium acetate solution is basic. CH3COO- + H2O CH3COOH + OH-
If a salt is made up of two ions that hydrolyse to the same extent, the salt solution could be close to neutral, e.g. ammonium acetate NH4CH3COO. NH4+ + H2O NH3 + H3O+ CH3COO- + H2O CH3COOH + OH-The resulting reaction, H3O+ + OH- 2H2O, results in a neutral solution.

Identify conjugate acid/base pairs
Common acids and bases and their conjugates
Acid
Conjugate Base
Base
Conjugate Acid
HCl
Cl-
OH-
H2O
HNO3
NO3-
NH3
NH4+
H2SO4
HSO4-
CN-
HCN
HSO4-
SO42-
CO32-
HCO3-
NH4+
NH3-
HCO3-
H2CO3
CH3COOH
CH3COO-
S2-
HS-
H2O
OH-
H2O
H3O+
H3O+
H2O


Strong acids produce very weak conjugate bases
Strong bases produce very weak conjugate acids
Moderately weak acids produce moderately weak bases (eg. CH3COOH to CH3COO-)

Identify amphiprotic substances and construct equations to describe their behaviour in acidic and basic solutions
A molecule or ion that can behave as a proton donor or acceptor is called amphiprotic. Amphiprotic means protons on both sides. An amphiprotic molecule or ion can donate or accept a proton. Whether an amphiprotic species behaves as an acid or a base depends on the nature of the other species it is reacting with; that is, whether or not the other species is a stronger acid or base than the amphiprotic species itself.
Water is an amphiprotic molecule:
Water as an acid: H2O H+ + OH- Water as a base: H+ + H2O H3O+
The hydrogen carbonate (bicarbonate) ion is an amphiprotic ion: As an acid HCO3- H+ + CO32- As a base H+ + HCO3- H2CO3
Zinc oxide (ZnO), aluminium oxide (Al2O3), HSO3-, H2PO4- and HPO42- are also amphiprotic.

Identify neutralisation as a proton transfer reaction which is exothermic
Neutralisation is the reaction between an acid and a base to form a salt and water. The solutions reacted to demonstrate neutralisation are usually of a strong acid, such as hydrochloric acid, and a strong base, such as sodium hydroxide. acid + base salt + water
HCl + NaOH NaCl + H2O
H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O
The net ionic equation for reaction is: H+ + OH- H2O
The net ionic equation shows that neutralisation is a proton transfer reaction. A proton from the acid transfers to the hydoxide ion of the base.
All neutralisations are exothermic, because the reaction creates covalent bonds. When the heat of neutralisation is measured for a range of strong acids and strong bases, the amount of heat released is always about 57 kJ per mole of water formed. This is the heat change for the following reaction: H+ + OH- H2O DH = - 57 kJ mol-1

Describe the correct technique for conducting titrations and preparation of standard solutions
Titration:
- Rinse entire burette with a portion of the solution to be dispensed
- Fill burette
- Read the starting volume off the burette
- Rinse pipette with the other solution
- Pipette a fixed volume of the other solution into a conical flask
- Add a few drops of a suitable indicator to the liquid in the conical flask. Choose an indicator that changes colour at about the pH f the salt solution formed at the equivalence point
- Add the solution from the burette gradually while swirling the contents of the flask until the desired indicator colour change occurs.
- Repeat at least twice
Preparation of a standard solution
- Calculate the moles of the substance required
- Calculate the required mass
- Accurately weigh a mass close to the required mass in a beaker
- Dissolve all of the measured mass in water, transfer to volumetric flask, wash inside beaker twice with water and transfer washings to volumetric flask
All of the weighed mass must enter the volumetric flask
- Dilute the solution to the exact volume marked on the flask
- Stopper the flask then invert and rotate to fully mix
- Calculate exact concentration
- Label solution with name, and date

For a chemical to be suitable to prepare as a standard solution, it must:
- be a water soluble solid
- have high purity usually Analytical Reagent (A.R.) grade
- have an accurately known formula
- be stable in air, i.e. it does not lose or gain water or react with oxygen or carbon dioxide in air.
The concentration is usually calculated in mol L-1.
At senior high school level equipment such as burettes, pipettes and volumetric flasks give readings to three significant figures. Calculations are carried out to three significant figures.

Qualitatively describe the effect of buffers with reference to a specific example in a natural system
A buffer controls the level of acidity or basicity in a solution. If an acid or a base is added to a buffer solution, there is hardly any change in pH.
A buffer solution is any solution containing either a weak acid and its conjugate base, or a weak base and its conjugate acid, eg. hydrogen carbonate ions, HCO3-, and carbonate ions, CO32-.
If an acid is added to the buffer, the hydrogen ions are removed by H+ + HCO3- H2CO3
If a base is added to the buffer, hydroxide ions are removed by OH- + HCO3- H2O + CO32-
The net effect is that the pH of the solution containing buffer changes only slightly, but will eventually the buffer cannot resist a drastic increase in pH.
Hydrogen carbonate ions are important in maintaining the pH of human blood at about 7.4. The oxygen carrying capacity of the blood haemoglobin and the activity of cell enzymes depend very strongly on the pH of the body fluids. Slight variations in pH are essential for the stimulation of certain physiological functions. However, pronounced changes (extended acidosis or alkalosis) will lead to serious disturbances of normal body functions and even death.

Gather and process information from secondary sources to trace developments in understanding and describing acid/base reactions
Scientist
Acid Definition
Base definition
Notes
Arrhenius
Provides H+ in solution
Provides OH- in solution
- Water solutions only
Bronsted-Lowry
Acid is a proton donor
Base is a proton acceptor
- Acid must contain hydrogen
- Could classify carbonates, etc
Lewis
Acid is an electron pair acceptor
Base is an electron pair acceptor
- Electron pair donor
- Broadest definition of all: applicable to organic and bio- chemistry

Choose equipment and perform a first-hand investigation to identify the pH of a range of salt solutions
Choose the most appropriate equipment available to you. If possible, a pH meter or data logger with probe would be most suitable, otherwise select indicator solution or indicator paper to measure the pH of the salt solutions.
To ensure you perform a valid investigation, prepare solutions of salts of equal concentration.


Perform a first-hand investigation and solve problems using titrations and including the preparation of standard solutions, and use available evidence to quantitatively and qualitatively describe the reaction between selected acids and bases
1. Calculate the moles of standard solution which have reacted:
moles of standard solution = (conc. in mol L-1) x volume of known in L (usu. 0.025)
2. Write a balanced equation for the reaction which occurs with the other reactant, ie. the unknown solution.
3. Use this equation to determine how many moles of unknown solution are required for the neutralisation reaction.
4. Calculate the concentration of the unknown solution using:
conc. of unknown = moles required of unknown / volume of unknown used (from burette) (mL)

Perform a first-hand investigation to determine the concentration of a domestic acidic substance using computer-based technologies
Vinegar contains acetic acid which can be titrated against standardised NaOH(aq) using a pH probe attached to a data logger. The data recorded can be used to draw a graph. The endpoint is where the pH changes most rapidly.




Analyse information from secondary sources to assess the use of neutralisation reactions as a safety measure or to minimise damage in accidents or chemical spills
Neutralisation is an exothermic reaction, so excess heat will be produced. This means that neutralisations shouldn’t be performed to clean up chemical spills on a person, but in the case of a tanker spill, it would stop dangerous chemical leakage and runoff, and excess heat wouldn’t really matter. However, the amount of chemicals needed to neutralise another chemical must be accurately calculated, otherwise the acidity/basicity will go too far over the other way.
A substance containing an amphiprotic ion, such as the hydrogen carbonate ion in NaHCO3, is quite suitable for neutralising chemical spills.
If the chemical spill contains an acid, H+ + HCO3- H2O + CO2 If the spill contains a base, HCO3- + OH- CO32- + H2O.
Sodium hydrogen carbonate is used because:
- it is a stable solid which is safely handled and stored
- if too much is used there is less danger than from excess sodium hydroxide
- it is amphiprotic
- it neutralises chemical spills of acids, bases and of unknown acidity/basicity