Chapter 7: Apply physical chemistry concepts

Physical chemistry concepts form a crucial foundation for understanding chemical reactions and properties encountered in various professional fields across Kenya. Whether in healthcare, agriculture, or business, professionals benefit from grasping how substances interact at the molecular level. This chapter explores the identification of acids and bases, focusing first on acids, enabling students to apply this knowledge in contexts such as pharmaceutical formulation, soil treatment, or quality control in food processing.

7.1 Identify Acids and bases

A laboratory technician often meets a bottle whose label is faded, a sample whose nature is unknown, or a solution that must be confirmed as acidic or alkaline before it is used. Identifying acids and bases means using safe, reliable tests to decide whether a solution is acidic, alkaline or neutral, and how strong it is. In a water-testing laboratory the result decides how a sample is treated, in a soil-testing laboratory it decides whether the soil is acidic, and in a food-quality laboratory it is one measure of whether a product is fresh. This topic covers the meaning of the terms, the characteristics of acids and bases, and the safe tests used to tell them apart. The individual acids and bases a technician handles are taught in Acids and in Bases.

Meaning of acid, base and alkali

An acid is a substance that releases hydrogen ions (H⁺) when it dissolves in water, or more generally a substance that gives a proton to another substance. A base is a substance that neutralises an acid, or more generally accepts a proton. An alkali is a base that dissolves in water and releases hydroxide ions (OH⁻). Sodium hydroxide and ammonia solution are alkalis. Copper(II) oxide and iron(III) hydroxide are bases that do not dissolve in water to any useful extent, so they are bases but not alkalis, yet they still dissolve in acids.

Acidic properties appear only when the acid is in water: dry hydrogen chloride gas does not turn dry blue litmus paper red, but the same gas dissolved in water does. A neutral solution contains equal concentrations of H⁺ and OH⁻. Pure water at 25 °C is neutral and has a pH of 7.

Characteristics of acids

• pH below 7: the hydrogen ion concentration is greater than that of pure water, so the pH is less than 7.
• Indicators: blue litmus turns red, methyl orange turns red, and phenolphthalein stays colourless.
• Reaction with reactive metals: magnesium, zinc and iron react with dilute acids to give a salt and hydrogen gas. Copper does not react with dilute hydrochloric or sulphuric (VI) acid.
• Reaction with carbonates: carbonates and hydrogencarbonates fizz and give off carbon dioxide, which turns limewater milky.
• Neutralisation: acids react with bases and alkalis to give a salt and water.
• Electrical conduction: acid solutions conduct electricity, because the acid forms ions in water.

Characteristics of bases

• pH above 7: an alkali has a lower hydrogen ion concentration than pure water, so its pH is greater than 7.
• Indicators: red litmus turns blue, methyl orange turns yellow, and phenolphthalein turns pink.
• Neutralisation: bases react with acids to give a salt and water.
• Electrical conduction: alkali solutions conduct electricity, because they contain ions.
• Soapy feel: concentrated alkalis feel soapy because they react with the oils of the skin, which is also why they burn. Touch is never used as a test.
• Insoluble bases: metal oxides and hydroxides that do not dissolve do not change an indicator, but they dissolve in acids to form salts.

Why tasting, touching and smelling are not tests

Acids are sour and alkalis bitter, but tasting a laboratory chemical is never done, even when the substance is thought to be harmless or to be a food ingredient. Acids and alkalis burn the mouth and throat, and unknown solutions can be poisonous. The same applies to touching a solution to feel whether it is slippery and to sniffing a container directly. Vapours of ammonia, hydrochloric acid and nitric (V) acid damage the airways, so any smell test is done only when a procedure asks for it, by wafting a little vapour towards the nose with the hand, never by inhaling from the container. The safe tests in this topic use indicators, a pH meter, conductivity, reaction tests on small samples and titration.

The pH scale

The pH scale measures how acidic or alkaline a solution is. It is defined from the hydrogen ion concentration, written [H⁺] and measured in mol/L:

pH = −log₁₀[H⁺]

At 25 °C the scale runs from 0 to 14: below 7 is acidic, 7 is neutral and above 7 is alkaline. Each step of one pH unit is a tenfold change in hydrogen ion concentration, so a solution of pH 3 has 100 times the hydrogen ion concentration of a solution of pH 5. For an alkali, the pOH is found from the hydroxide ion concentration and the two are linked at 25 °C:

pOH = −log₁₀[OH⁻]

pH + pOH = 14

Worked example: pH of an acid and of an alkali

A technician needs the pH of a solution of a strong acid with [H⁺] = 2.0 × 10⁻⁴ mol/L, and of a solution of sodium hydroxide of concentration 0.0020 mol/L.

• Acid formula: pH = −log₁₀[H⁺]
• Acid substitution: pH = −log₁₀(2.0 × 10⁻⁴)
• Acid working: log₁₀(2.0 × 10⁻⁴) = log₁₀ 2.0 + log₁₀ 10⁻⁴ = 0.301 − 4 = −3.699
• Acid result: pH = 3.70
• Alkali working: sodium hydroxide is a strong alkali, so [OH⁻] = 0.0020 mol/L = 2.0 × 10⁻³ mol/L, and pOH = −log₁₀(2.0 × 10⁻³) = 3 − 0.301 = 2.699
• Alkali result: pH = 14 − 2.699 = 11.30

Litmus paper

Litmus is a dye from lichens, supplied as red and blue paper strips. Acidic solutions turn blue litmus red, alkaline solutions turn red litmus blue, and a neutral solution leaves both colours unchanged. Litmus changes colour at about pH 5 to 8, so it says only whether a solution is acidic or alkaline, not how strong it is. To use it, dip a clean glass rod into the sample and touch the rod on the paper. Never dip the paper into the sample bottle, because that contaminates the stock, and use a clean rod for each solution. Keep litmus paper dry and in a closed container, because damp or old paper gives weak colours.

Indicator solutions

An indicator is a dye that changes colour over a narrow range of pH. A few drops are added to a small sample in a test tube or conical flask. The common laboratory indicators are compared below.

Bromothymol blue is green when the solution is near neutral. An indicator tells only whether the pH is on one side of its range or the other, so the indicator is chosen to suit the work. A coloured or cloudy sample can hide the colour change, and in that case a pH meter is used instead.

Universal indicator and pH paper

Universal indicator is a mixture of several indicators that gives a gradual series of colours across the pH scale, so it shows an approximate pH. Typical colours are red at pH 1 to 3, orange at about pH 4, yellow at pH 5 to 6, green at pH 7, blue at pH 8 to 10 and violet at pH 11 to 14. Makers use slightly different colours, so always compare the result with the colour chart supplied with the product that is being used. Universal indicator paper and wide-range pH paper work in the same way and are matched against a chart on the packet. The result is read to about one pH unit.

Natural indicators

Natural indicators give a low-cost alternative when manufactured indicators are not available. They are suitable for teaching and for rough checks, but not for measurements that must be reported.

• Red cabbage extract: chop red cabbage leaves, steep them in hot distilled water, strain and cool the purple liquid. It turns red or pink in acid, stays purple near neutral, turns blue to green in alkali and yellow in strongly alkaline solutions.
• Turmeric paper: paper soaked in turmeric extract and dried stays yellow in acid and neutral solutions and turns reddish brown in alkali.
• Checking before use: test every natural indicator against a known acid and a known alkali first, because its colours depend on the plant and on how it was prepared.

Reaction with metals and carbonates

Reaction tests confirm that a solution is an acid. They are done on small volumes, after an indicator has already shown the solution to be acidic.

• Test with a metal: add a short piece of magnesium ribbon to a few cubic centimetres of the sample in a test tube. Fizzing shows that a gas is formed, and holding a lighted splint at the mouth of the tube gives a squeaky pop if the gas is hydrogen. Mg + 2HCl → MgCl₂ + H₂
• Test with a carbonate: add a little sodium carbonate to the sample. Fizzing shows carbon dioxide, which turns limewater milky when the gas is passed into it. Na₂CO₃ + 2HCl → 2NaCl + H₂O + CO₂
• Limits of the tests: weak acids react slowly, and dilute nitric (V) acid does not usually give hydrogen with metals, because the nitrate ion is reduced instead. Run these tests in a fume cupboard when the acid is concentrated.

Conductivity test

Acid and alkali solutions conduct electricity because they contain ions, a property explained in Electrolytes. A conductivity meter, or a simple circuit with two graphite electrodes, a low-voltage direct current supply and a lamp or ammeter, shows whether a solution conducts. A strong acid or strong alkali conducts well, a weak acid or weak alkali conducts poorly, and a sugar solution does not conduct. Conductivity alone cannot separate an acid from a base or from a salt solution, because a solution of sodium chloride conducts too, so it supports the indicator results and does not replace them.

Measuring pH with a pH meter

A pH meter measures the voltage between a glass electrode and a reference electrode and displays the result as pH. It reads to a fraction of a pH unit, which is much finer than a colour chart, and it works on coloured and cloudy samples. Follow this procedure:

1. Switch the meter on and allow it to warm up for the time given in its manual.
2. Rinse the electrode with distilled water and blot it dry with a soft tissue. Never rub the glass bulb.
3. Calibrate with buffer solutions of known pH, for example pH 7 first and then pH 4 or pH 10, whichever is nearer to the expected pH of the sample. Use fresh, labelled buffers and record the calibration.
4. Rinse and blot the electrode again so that no buffer is carried into the sample.
5. Place the electrode in the sample so that the bulb is fully covered, stir gently and wait until the reading is steady. Record the pH and the temperature.
6. Rinse and blot the electrode between samples.
7. At the end, check the meter against one buffer. If the reading has drifted, repeat the measurements. Return the electrode to its storage solution, never to dry storage.

The temperature of the sample affects the reading, so many meters have a temperature sensor that corrects for it, and the temperature is always recorded with the pH.

Titration

An acid-base titration finds the concentration of an acid or alkali by adding a solution of known concentration from a burette until the indicator changes colour at the end point. The indicator is chosen to suit the strengths of the acid and base.

• Strong acid and strong base: methyl orange or phenolphthalein is suitable.
• Weak acid and strong base: phenolphthalein, because the solution at the end point is slightly alkaline.
• Strong acid and weak base: methyl orange, because the solution at the end point is slightly acidic.
• Weak acid and weak base: no indicator gives a sharp change, so a pH meter is used.

Worked example: concentration of ethanoic acid

A 25.0 mL sample of ethanoic acid solution needs 18.5 mL of 0.100 mol/L sodium hydroxide to reach the end point with phenolphthalein. The equation is CH₃COOH + NaOH → CH₃COONa + H₂O, so the acid and base react in a 1 : 1 ratio.

• Moles of base: n(NaOH) = 0.100 mol/L × 0.0185 L = 0.00185 mol
• Moles of acid: n(CH₃COOH) = 0.00185 mol
• Concentration: c = 0.00185 mol ÷ 0.0250 L = 0.0740 mol/L
• Result: the ethanoic acid solution has a concentration of 0.0740 mol/L.

Comparing the tests

Start with the quickest, least wasteful test, litmus or universal indicator, and add a second method when the result decides a purchase, a treatment or a report. Record the sample label, date, method, result, temperature and the calibration used.

Safe practice when identifying an unknown solution

• Treat it as hazardous: wear safety goggles, a laboratory coat and chemical-resistant gloves until the solution is identified.
• Use small volumes: test a few cubic centimetres, never the whole sample.
• Never mix unknown solutions: acids mixed with some other chemicals give toxic gases or heat.
• Label everything: every container holds a label with the contents, date and name.
• Never return a used portion to the stock bottle: it contaminates the stock.
• Dispose by the laboratory procedure: collect waste in labelled containers.

Importance of accurate identification

A wrong result can be costly or unsafe. A water-testing laboratory uses pH to decide how a sample is treated, because pH affects how fast metal pipes corrode and how well treatment works. A soil-testing laboratory measures the pH of a soil and water mixture because soil acidity affects the nutrients available to crops and decides whether lime is advised. A food-quality laboratory follows the acidity of a fermented product, because pH falls as acids form. In a research or teaching laboratory, an acid mistaken for an alkali can lead to a violent reaction during neutralisation or to ruined buffer solutions.

Challenges in identification and how to overcome them

• Coloured or cloudy samples: the colour change is hidden, so use a pH meter.
• Interfering substances: a mixture can mask the acid or alkali property, so use two methods.
• Limited equipment: where no pH meter is available, use litmus, universal indicator or a natural indicator checked against known solutions.
• Old indicators and meters out of calibration: check indicators with a known acid and alkali, and calibrate the pH meter with fresh buffers before each session.
• Temperature changes: record the temperature and use a meter with temperature correction.
• Contaminated stock: store chemicals in closed, labelled containers and use clean rods and pipettes.
• Limited training: write standard operating procedures for each test and train every technician in them.

Chapter Summary

This chapter explored the identification of acids and bases, establishing a foundation for understanding their chemical properties. It examined key acids including hydrochloric acid, sulphuric (VI) acid, and nitric (V) acid, detailing their characteristics and common applications. The discussion then shifted to bases such as sodium hydroxide, ammonia solution, and calcium hydroxide, highlighting their roles in various chemical processes. The chapter also covered essential gas laws, explaining Boyle's Law, Charle's Law, Dalton's Law of partial pressures, and Graham's Law of diffusion, which describe the behavior of gases under different conditions. Electrochemistry concepts were introduced, focusing on reduction, oxidation, the process of electrolysis, and Faraday's Laws of electrolysis, which quantify the relationship between electric charge and chemical change. Finally, the chapter addressed the nature and function of electrolytes, underscoring their importance in conducting electricity in solutions. Together, these topics provide a comprehensive understanding of fundamental physical chemistry principles applicable across scientific and industrial contexts.