Have you ever wondered what makes an antacid tablet so effective? How about toothpaste? What’s the cream used to help relieve the pain caused by wasp stings made from? These are all everyday examples of neutralisation reactions between acids and bases.
What are acids and bases?
There are multiple different definitions of acids and bases, depending on who you ask. Take acids, for example.
- The Swedish physicist-turned-chemist Svanthe Arrhenius first defined an acid in 1884 as a substance that dissociates into hydrogen ions in solution.
- In 1923, two scientists called Johannes Nicolaus Brønsted and Thomas Martin Lowry independently arrived at a different theory of acids and bases. They said an acid is a proton donor.
- Also in 1923, Gilbert N. Lewis defined an acid as an electron pair acceptor.
A hydrogen ion is actually just a proton. Hydrogen atoms contain one proton and one electron. Remove the electron in an ionisation reaction and all you are left with is a proton.
In this article, we’re interested in the second definition, the one advanced by Brønsted and Lowry.
An acid is a proton donor.
Monoprotic acids donate just one proton per acid molecule in a solution, whilst diprotic acids donate two.
The word for acid comes from the Latin term ‘acidus’, signifying sour. Acids turn damp blue litmus paper red. In contrast, bases turn red litmus paper blue and have a soapy texture.
A base is a proton acceptor.
Acids and bases dissociate in solution. This means that they split up into ions. For example, acids always split up into protons and a negative ion, whilst bases dissociate into hydroxide ions and a positive ion.
\[HA(aq) \rightarrow H^+(aq) + A^-(aq)\]
\[B(aq) + H_2O(l) \rightarrow BH^+(aq) + OH^- (aq)\]
Conjugate acids and bases
You can also find conjugate acids and bases.
A conjugate acid is a base that has gained a proton, whilst a conjugate base is an acid that has lost a proton.
Acids and bases all come with a paired conjugate acid or base. For example, the conjugate acid of ammonia is ammonium, \(NH_4^+\):
Base Conjugate acid
\[NH_3(aq) + H_2O(l) \rightarrow NH_4^+ (aq) + OH^-(aq)\]
Neutralisation reactions
Acids and bases react together in neutralisation reactions.
Neutralisation reactions form salts. Salts are ionic compounds consisting of positive and negative ions held together in a giant lattice. To name them, we state the cation (positive ion) first followed by the anion (negative ion). For example, sodium chloride, which is actually just common table salt. Another example of a salt is calcium chloride, which is used to de-ice roads.
Fig. 1 - Sodium chloride. The positive sodium ions (purple) and negative chloride ions (green) are arranged in a giant lattice structure
To fully neutralise a solution, you add just enough base to react with all of the acid: there should be neither acid nor base leftover.
In a neutral solution, the concentrations of hydrogen and hydroxide ions are equal.
As we mentioned above, taking antacid tablets, brushing your teeth, and soothing wasp stings all involve neutralisation reactions. Antacid tablets contain bases such as magnesium hydroxide, \(Mg(OH)_2\), which neutralise excess hydrochloric acid produced by the stomach. On the other hand, toothpaste is alkaline and reacts with the acids produced by bacteria living in your mouth. Wasp stings are also alkaline. Thus, creams and balms often contain acids to neutralise the sting and calm the affected area.
An alkali is a base that is soluble in water.
Concentrated and dilute acids and bases
You’ve probably heard of the dangers of acids and images of corrosive warning signs fill your head. Whilst it is true that both acids and bases can be extremely dangerous, you generally only have to worry about concentrated acids and bases.
Concentration refers to the number of acid or base molecules in solution. A concentrated acid or base contains a lot of molecules dissolved in solution whereas a dilute acid or base contains fewer.
Strong and weak acids and bases
Take a look at the following equations:
\[HCl \rightarrow H^+ + Cl^-\]
\[CH_3COOH \rightleftharpoons H^+ + CH_3COO^-\]
Both hydrochloric acid, \(HCl\), and ethanoic acid, \(CH_3COOH\), are acids, just as their names suggest. They both dissociate to donate protons in solution. However, you’ll notice something different about their equations. Whilst the reaction involving hydrochloric acid goes to completion, the one involving ethanoic acid is reversible. This happens because hydrochloric acid is a strong acid whilst ethanoic acid is weak.
Strong acids and bases are acids and bases that fully dissociate in solution. In contrast, weak acids and bases only partially dissociate in solution.
Strong acids include hydrochloric acid, found in gastric juices, and sulfuric acid. Weak acids include ethanoic acid, found in malt vinegar, and citric acid, found in citrus fruits like lemons. Strong bases include sodium hydroxide whilst weak bases include ammonia. You’ll explore weak acids and bases more in Weak Acids and Bases.
Although the difference between a strong and weak acid may seem trivial, it becomes important when calculating pH, as you’ll see later. But before we explore that, we need to define what pH actually is.
What is the pH scale?
\(pH = -\log([H^+])\). It is a measure of hydrogen ion concentration in solution.
The pH scale was invented by a Danish brewer and chemist named Søren Peder Lauritz Sørensen, who was looking to control the acidity of his beer. Solutions with a high hydrogen ion concentration have a low pH and vice versa.
You now know that acids release protons (hydrogen ions) in solution. This means that acids have a low pH. On the other hand, bases have a high pH.
Fig. 2 - The pH scale
Calculating pH
Calculating pH can get a little tricky. There are lots of equations you need to learn, and it is easy to get confused between moles and concentrations. The following table gives you some of the values you need to understand in order to calculate pH, as well as equations linking them.
Fig. 3 - Important values within the topic of acids and bases
In later articles, we’ll explore all these values in greater detail and walk you through the different methods of calculating pH. However, the process can be summarised with the following flow chart:
Fig. 4 - A flow chart used to work out the pH of acids and bases
Measuring pH
If you are carrying out an acid-base reaction such as a neutralisation, which we’ll explore below, you might want to know the pH of a solution at regular intervals. Calculating the pH each time could get a little laborious. Fortunately, we have some different ways of finding pH instantaneously.
- Indicators are substances that show a distinct observable change when the conditions they are in change. This is most commonly a colour change. The universal indicator is a mix of different indicators that spans the whole colour spectrum according to whether a substance is acidic or alkaline.
- pH meters are digital instruments that accurately measure the pH of a solution. They do this by measuring differences in electrical activity.
Fig. 5 - Universal indicator1. Remember that a low pH is acidic and a high pH is alkaline. commons.wikimedia.org
What are buffer solutions?
A buffer solution is a solution that maintains a constant pH when small amounts of acid or alkali are added to it.
In Buffer Solutions, you’ll learn about how these extraordinarily useful solutions work. There are many systems that simply wouldn’t work if their pH fluctuated outside of a narrow range, such as your circulatory system.
The pH of blood is maintained by three systems, most notably by one called the bicarbonate buffer system. A constant pH of around 7.4 is needed to maintain optimum conditions for enzyme activity. When your cells respire, they release \(CO_2\) into the bloodstream. This reacts with water to turn into the bicarbonate ion, \(HCO_3^-\), which exists in equilibrium with carbonic acid, \(H_2CO_3\). Any acids produced by cellular activity, for example, lactic acid, are neutralised by bicarbonate ions whilst any bases are neutralised by carbonic acid. Overall, this maintains a steady pH.
What are titrations?
Suppose you have a solution of hydrochloric acid but you aren’t sure what its concentration is. A common way of finding out the concentration of an unknown acid or base is through a titration reaction. You do this by neutralising a fixed volume of an acid or base of known concentration with your acid or base of unknown concentration, and measuring the volume of unknown substance needed.
To accurately carry out a titration, you need to know its equivalence point.
The equivalence point is the point where just enough base has been added to sufficiently neutralise an acid in solution, or vice versa.
To determine when you have reached the equivalence point, you use an indicator, as we mentioned above. Indicators are useful because they change colour at a specific pH. This is known as the endpoint.
The endpoint of a titration is the point where the indicator just changes colour.
If the end point of a titration is the same as its equivalence point, you can use the indicator’s colour change to tell you when you have added just enough base to neutralise the acid, or vice versa. You can then use the chemical equation for the reaction to work out the concentration of the unknown acid or base. You’ll learn more about this in pH Curves and Titrations.
pH curves and suitable indicators
Plotting the pH change in a neutralisation reaction against volume of acid or base added produces a curved graph known as a pH curve. It has three distinct sections:
- An initial shallow-sloping section, where the pH barely changes.
- A sharply-sloping section, where the pH changes rapidly.
- Another shallow-sloping section, where the pH barely changes.
The equivalence point of a titration lies in the middle of this sharply-sloping section. If an indicator’s endpoint also lies in this section, you can use the indicator in your titration.
Fig. 6 - A pH curve for the addition of a strong base to a weak acid. The endpoints of two indicators are shown. Here, phenolphthalein would be a suitable indicator as its endpoint lies in the sharply-sloping section of the graph, which also contains the equivalence point
Carrying out titrations
A fun practical activity could be carrying out a simple titration or general neutralisation reaction. In fact, you’ll probably do many titrations over the course of your studies. If you are carrying out a titration, make sure you use a suitable indicator, but you could also use a pH meter. Let’s walk through the process using hydrochloric acid and sodium hydroxide.
- Measure out 30 cm3 of HCl using a volumetric flask. Pour the solution into a conical flask.
- Add 2-3 drops of your indicator and swirl the flask.
- Rinse a burette, first with distilled water and then with NaOH. Set up the burette using a stand and clamp so that it is suspended above the conical flask.
- Fill the burette with your titrant. Note the value of this solution shown on the burette. This is your start value.
- Add the titrant to the conical flask in 1 cm3 intervals, swirling after each addition, until the solution in the conical flask changes colour. Note the value shown on the burette. This is your end value.
- To work out the titre, subtract your end value from your start value. This gives you the volume of titrant added to the flask.
- Repeat the experiment again until you have three titre values within 0.1 cm3 of each other. As you reach the end of the titration, near the point of colour change, add the titrant dropwise. The colour change takes place over a very small volume range and so adding the titrant in reduced amounts allows you to be more precise.
If you are using a pH meter, use it to measure the pH of the solution in the conical flask each time you add more of the titrant. As you near the point of colour change, add the titrant in smaller quantities as explained above.
This is the set-up for a typical titration.
Fig. 7 - The setup for a typical titration
Titrations have many useful applications in everyday life. For example, they’re used to determine the degree of contamination of wastewater and to find out the nutritional content of certain foods, such as the proportion of saturated and unsaturated fatty acids present. The cosmetic industry also used titrations to make sure the pH of their products stays within a safe range for human skin.
Fluid Electrolyte and Acid-Base Balance Introduction to Body Fluids
To finish off, let's explore fluid, electrolytes and the acid-base balance of homeostasis. Although this is more important in biology and you most likely won't encounter it in your chemistry exam, it is a very important topic!
In the body of an average adult, there is an average of 40 L of body fluids (figure 3). The intracellular fluid is the fluid inside body cells, and it consists of mostly water and electrolytes like potassium (K+), magnesium (Mg2+), and HPO42-. The extracellular fluid is found outside body cells and contains electrolytes such as Na+, Cl-, HCO3- and Ca2+.
Electrolytes are basically chemical substances that, when dissolved in water, release cations and anions.
One of the many functions of electrolytes is to help maintain acid-base balance in homeostasis. In this case, acids are considered electrolytes that release H+ ions in water, while bases are electrolytes that release OH- ions in water.
Homeostasis is the tendency of our bodies to return to a steady state after an environmental change. A body's ability to maintain homeostasis is essential to life. For instance, if changes in blood pH occur and the body is unable to return the blood pH to its normal range, it can lead to fatal consequences!
If you understand the concepts discussed in this explanation, you'll have a really strong foundation that will help you during your chemistry tests!
Acids and Bases - Key takeaways
There are many different definitions of acids and bases. The Brønsted-Lowry definition defines acids as proton donors and bases as proton acceptors.
Acids and bases dissociate in solution. Acids dissociate into hydrogen ions and bases dissociate into hydroxide ions.
Whilst strong acids and bases fully dissociate in solution, weak acids and bases only partially dissociate.
Concentration is a measure of the number of acid or base molecules in solution.
A neutralisation reaction is a reaction between an acid and base to form a salt. A neutral solution contains equal concentrations of hydrogen and hydroxide ions.
pH is a measure of hydrogen ion concentration in solution. We can calculate the pH of various solutions using values such as Ka, Kb, Kw and pOH.
Titration reactions help you work out the concentration of an acid or base. They use indicators to show when you have reached the equivalence point of the reaction.
Buffer solutions are solutions that maintain a constant pH when small amounts of acids or bases are added to them.
References
- Credits and Attribution- Ariadna.creus, CC BY-SA 4.0(https://creativecommons.org/licenses/by-sa/4.0/) , via Wikimedia Commons
- Theodore Lawrence Brown, Eugene, H., Bursten, B. E., Murphy, C. J., Woodward, P. M., Stoltzfus, M. W., & Lufaso, M. W. (2018). Chemistry : the central science (14th ed.). Pearson.
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