# Solve problems and perform a first-hand investigation to use pH meters/probes and indicators to distinguish between acidic, basic and neutral chemicals

This experiment is relatively simple. Carefully use a pH meter/probe in a solution in order to determine if it is acidic, basic, or neutral. Don’t forget the pH ranges for each of these classifications, which are stated again below for your convenience.

An acid has a pH less than 7, whereas a base has a pH greater than 7. Substances with a pH of 7 are neutral substances.

# Define acids as proton donors and describe the ionisation of acids in water

Although this dotpoint focuses on acids as proton donors, also keep in mind that this makes a base a proton acceptor. Once you understand that a hydronium ion is equivalent to a proton, you will find acid-base reactions quite simple.

The hydrogen atom consists of one proton and one electron. Thus, when this atom ionises and loses an electron to form H+, only a proton remains.

When an acid molecule reacts to release the hydrogen ion H+, it is thus donating a proton. In the presence of hydrogen ions, water will convert to a cation known as the hydronium ion ($H_3O^+$).

An example of the ionisation of an acid in water using acetic acid- an acid known for its presence in vinegar – is shown below:

$CH_3COOH_{(aq)}$ + $H_2O_{(aq)}$ ↽⇀ $H_3O^+_{(aq)}$ + $CH_3COO^-_{(aq)}$

In the above equation, the donated proton, or hydrogen ion, ionised the water molecule.

It is also useful to note that a monoprotic acid is one that releases one proton. Similarly, a diprotic or triprotic acid releases two or three protons respectively. For example, sulfuric acid is a diprotic acid, and citric acid is a triprotic acid.

Remember- If a substance can ‘disassociate’ into the hydronium ion and an anion, then we will often recognise it as an acid. This dotpoint simply seeks to illustrate this point, noting that the hydrogen ion is simply a proton. Thus an acid is a proton donor.

# Identify acids including acetic (ethanoic), citric (2-hydroxypropane-1,2,3-tricarboxylic), hydrochloric and sulfuric acid

Historical names will not appear often, but systematic names will appear from time to time in examinations. It may also be useful to remember one use for each of these acids, but I would not recommend anything further.

Acetic acid, also known as ethanoic acid, occurs naturally, but is manufactured on an industrial level for the production of items such as soft drink bottles, glue, and vinegar, and is also used as a food additive. Formic acid is methanoic acid, and can be found in ant stings.

Citric acid, also referred to by its systematic name of 2-hydroxypropane-1,2,3-tricarboxylic acid, occurs naturally in citrus fruits. Citric acid is used predominantly as a food additive.

Hydrochloric was historically known as muriatic acid, and finds a wide-scale use industrially due to its nature as a strong acid. One such example is the production of vinyl chloride for its polymer PVC.

Sulfuric acid was historically known as oil of vitriol, and also finds a wide-scale use industrially. A large proportion of sulfuric acid is used in the fertiliser industry, and sulfuric acid is also used as a catalyst in many industrial reactions.

Remember- Acetic acid is ethanoic acid, and formic acid is methanoic acid.

# Identify data, gather and process information from secondary sources to identify examples of naturally occurring acids and bases and their chemical composition

Acids:

• Acetic acid ($CH_3CO_2H$)- Acetic acid is produced naturally through bacterial fermentation, and is used in the production of vinegar and as a
• Citric acid ($C_6H_8O_7$)- Citric acid is produced naturally in citrus fruits, and is used in food both as a preservative, and for
• Formic acid ($HCOOH$)- Formic acid can be found in the stings of ants and bees, and is used for a variety of purposes such as esterification and as a preservative.
• Lactic acid ($C_3H_6O_3$)- Lactic acid is produced within the body by muscles during exercise. It can also be found within certain milk products such as yoghurt and various

Bases:

• Ammonia ($NH_3$)- Certain organisms can produce ammonia from atmospheric nitrogen. Am- monia base is used within household cleaners, and in the production of fertilisers such as ammonium
• Potassium hydroxide ($KOH$)- The leaching (The passing through of water resulting in the dissolving of soluble substances) of the certain wood ashes was known historically to produce potash, or potassium Potash can be used to produce soap as well as fertiliser.
• Sodium bicarbonate ($NaHCO_3$)- Produced naturally in a mineral known as nahcolite, sodium bicarbonate is used for various pest control purposes and Sodium bicarbonate is also used in the kitchen to make dough rise.

Take care not to confuse ammonia ($NH_3$) with ammonium ($NH^+_4$).

Remember- Acetic acid and formic acid are also known as ethanoic acid and methanoic acid respectively.

# Describe the use of the pH scale in comparing acids and bases

The pH scale is used to compare the acidity and basicity of substances, making use of a logarithmic scale so as to take into account the great variances of concentration.

For example, a hundredfold change will only result in a change in pH of 2.

This scale measures how acidic or basic a substance is, by assigning it a number to indicate its nature. An acid has a pH less than 7, whereas a base has a pH greater than 7. Substances with a pH of 7 are neutral substances.

# Describe acids and their solutions with the appropriate use of the terms strong, weak, concentrated and dilute

The use of these terms poses some confusion for most students, so take the time to read and understand these definitions.

The terms ‘strong’ and ‘weak’ are used to describe the degree of ionisation of an acid. If an acid ionises to completion within solution, then the acid is what is termed a strong acid. However, if neutral acid molecules still remain, meaning only some of the acid molecules have ionised, then the acid is a weak acid. Where an acid reaches equilibrium will determine how we classify it.

Furthermore, an acid can be termed ‘concentrated’ or ‘dilute’, referring to the concentration of the solution (mol/L). If the total concentration of the solute is high, then the acid is concentrated. If the total concentration is low, then it is called a weak acid.

Remember- A strong acid is not necessarily concentrated, just as a weak acid is not necessarily dilute. These categories are independent of one another, and a strong acid can be dilute just as a weak acid can be concentrated.

# Plan and perform a first-hand investigation to measure the pH of identical concentrations of strong and weak acids

The approach to this dotpoint will greatly vary depending on your school, but the use of pH probes presents the simplest method. If not, a variety of indicators (Including universal indicator) can be used to narrow down the possible range of pH.

If you do not conduct this experiment, then simply note that the result is essentially ’Equal concentrations of acids do not necessarily yield equal pH measurements.’ This is a practical application of the differing degrees of ionisation between acids. The stronger acid will have a lower pH, reflecting its higher degree of ionisation.

# Gather and process information from secondary sources to write ionic equations to represent the ionisation of acids

You will find these equations easy to do once you get used to them. The acid will always be ionising in water, hence these will be the reactants. The products will then be an anion, as well as a hydrogen ion (Or proton as an acid is a proton donor).

The ionisation of hydrochloric acid (complete):

$HCl_{(l)}$ + $H_2O_{(l)}$$H_3O_{(aq)}^+$ + $Cl_(aq)^-$

The ionisation of sulfuric acid (complete):

$H_2SO_{4 (l)}$ + $2 H_2O_{(l)}$$2 H_3O_{(aq)}^+$ + $SO_{4(aq)}^2-$

The ionisation of acetic/ethanoic acid (partial):

$CH_3COOH^+_{(l)}$ + $H_2O_{(l)}$ ↽⇀ $H_3O^+_{aq}$ + $CH_3$ $COO_{(aq)} ^-$

# Use available evidence to model the molecular nature of acids and simulate the ionisation of strong and weak acids

The diagram introduced in the later dotpoint may aid with your understanding of the distinctions between acids which are strong or weak, and concentrated or dilute.

# Identify pH as $-log_10[H^+]$ and explain that a change in pH of 1 means a ten-fold change in $[H^+]$

Simply feed this function through your calculator to obtain the pH, when given the hydrogen ion concentration. Don’t forget this function uses a base of 10, and appears as ‘log’ on your calculator, as opposed to the natural logarithm, ‘ln’, which uses a base of the constant e.

The relative strength of acids and bases can be measured through the concentration of hydrogen ions present within the molecules. As this concentration is subject to large variances, a logarithmic function is used.

$pH = -log_10[H^+]$, where the square brackets mean ’concentration’. As such, this can simply be read: ‘pH is the negative logarithm of the concentration of hydrogen ions.’

A logarithm is used to effectively provide smaller numbers to work with, as a change in 1 pH represents a tenfold change in hydrogen ion concentration. Similarly, a change in 2 pH represents a hundredfold change.

Thus, an acid with a pH of 2 has a hydrogen ion concentration 10 times more than an acid with a pH of 3.

You also need to be able to work backwards using the logarithmic formula, i.e. work out the hydrogen ion concentration when given the pH. A simple manipulation of the formula provides the result $[H^+]$ = $10^{-pH}$

A point which may be of use which does appear occasionally is the fact that: pH + pOH = 14. Calculations often appear which require this formula, and if you take the time to think about it, it makes sense that adding a measure of hydrogen ion concentration with a measure of hydroxide ion concentration gives you the upper limit of the pH scale. Practice converting from pH to [H+] and calculation questions within this area should prove no difficulty.

# Compare the relative strengths of equal concentrations of citric, acetic and hydrochloric acids and explain in terms of the degree of ionisation of their molecules

This dotpoint is linked quite closely to the terms ’strong’ and ’weak’ acid, so re-read the definition for these terms if they have slipped your mind.

Although three different acids may have equal concentrations, they may have different relative strengths due to the degree of ionisation of their molecules.

Hydrochloric acid is a strong acid, meaning that virtually 100% of the acid molecules ionise when in solution. In contrast, citric acid and acetic acid only ionise around 1% of their acid molecules, making these weak acids. However, note that citric acid can potentially have each acid molecule dissociate into three protons, as it is a triprotic acid. This makes citric acid the stronger of the two.

This relationship can be represented through the use of the following equations. Full-ionisation:

Partial-ionisation:

The use of the symbol → demonstrates that the ionisation runs until completion, whereas ↽⇀ indicates that the reaction is in a state of equilibrium.

Remember- A weak acid is one where neutral acid molecules are prevalent within the acid, meaning that not all the acid molecules have ionised.

# Describe the difference between a strong and a weak acid in terms of an equilibrium between the intact molecule and its ions

The main purpose behind this dotpoint is to point out that the degree of ionisation of an acid is what differentiates a strong acid from a weak acid. Although I would expect that a student have a sound understanding of this difference, the actual definition and calculation of an acid’s degree of ionisation is unlikely to appear in an exam.

The degree of ionisation of an acid is the percentage of molecules which are ionised in solution. This is simply the concentration of hydrogen ions divided by the concentration of the acid, as a percentage. For example, a 0.2 mol/L solution of an acid with a pH of 2 has a degree of ionisation of $\frac {10^{-2}}{0.2} = 5%$

$Degree of Ionisation = \frac {[H+]}{Concentration}$

Remember- A strong acid ionises to completion. A weak acid enters a state of equilibrium between the intact molecule and its ions.

# Process information from secondary sources to calculate pH of strong acids given appropriate hydrogen ion concentrations

You will be expected to be able to calculate pH using information about hydrogen ion concentration, as well as calculating hydrogen ion concentration from pH. Simply use the given formula and your calculator to do so. Don’t forget to use the ‘log’ button on your calculator instead of ‘ln’.

or

Using the above equation and its variation, you will be able to switch between hydrogen ion concentration and pH relatively easily.

# 2.2.14 Gather and process information from secondary sources to explain the use of acids as food additives

Acids find a myriad of uses as food additives. Below are three which I believe would be useful to learn (Also be prepared to give an example of each).

Acids have a wide range of application in the culinary industry, as they can be used as food additives for a variety of purposes. Acids are often used as:

• Preservative- Acids such as acetic acid prevent the growth of microorganisms such as salmonella within food, preserving the food so that it has a longer shelf-life. Phosphoric acid is similarly used in soft drinks to kill
• Antioxidant- Some acids such as ascorbic acid are used as antioxidants, preventing spoilage of food due to
• Flavouring- Acids are also used to add Citric acid is commonly used to add tartness.