What is pH?
The term pH stands for ‘potential of hydrogen ion’. It was proposed in 1909 by Danish biochemist Soren Sorensen as a more convenient way to describe hydronium and hydroxide ion concentrations in aqueous solutions, as both concentrations are extremely small.
pH of Solutions
Sorensen, a Danish biochemist, devised a scale to quantify acidity in terms of H+ concentrations in a solution. He defines pH as the ‘negative value of a solution logarithm to the base 10 of the concentration of H+ ions present’. Let us see what is the formula of pH and how to find the pH value of a solution?
The pH of Solution Formula:
$\mathrm{pH}=-\log _{10}\left[H^{+}\right]~~\text {or}~~\mathrm{pH}=\dfrac{1}{\log _{10}\left[\mathrm{H}^{+}\right]}$
Just as pH represents the concentration of hydrogen ions, pOH represents the concentration of hydroxyl ions.
$\mathrm{pOH}=-\log _{10}\left[\mathrm{OH}^{-}\right]~~\text {or } \mathrm{pOH}=\dfrac{1}{\log _{10}\left[\mathrm{OH}^{-}\right]}$
In other words, pH is a scale that specifies the acidity or basicity of an aqueous solution. Acidic solutions with higher ion concentrations have lower pH values when measured than basic or alkaline solutions.
The pH of an aqueous solution is determined by the pH scale, which in water typically ranges from 0 to 14. In any case, students should remember these points.
pH Value of Acidic Solution: Low hydroxide concentrations and high hydronium concentrations characterise acidic solutions. Acidic solutions contain more hydrogen ions per litre than neutral solutions.
pH Value of Basic Solution: High hydroxide concentrations and low hydronium concentrations characterise basic solutions. Basic solutions contain more hydroxide ions per litre than neutral solutions.
The pH of Pure Water
Pure water has a pH of 7. The concentration of hydroxide ion and hydrogen ion is the same in pure water, which makes it neutral. One might ask, what is the pH value of a neutral solution?
Neutral solutions have a pH value of 7. It will comprise an equal amount of H+ and OH- ions. Human blood and water are excellent examples of neutral solutions. Acids and bases can be neutralised and adjusted to a pH of 7. Water can dissociate to form H+ and OH ions. When these ions have the same charge, the value is 7 or neutral.
pH Examples
Here is an example of pH in everyday life: In our digestive system, the hydrochloric acid produced in our stomach aids digestion, while causing no harm to the stomach. This pH is required for the enzyme pepsin to be activated, which aids in the digestion of protein in food. Acids cause tooth decay because when we eat sugary foods, the bacteria in our mouth degrade them and acid is formed. Some more examples of natural substances whose pH value is less than 7 are lemon juice, vinegar, soda, rain acid, banana, urine, etc. Examples of substances whose pH value is more than 7 are sea water, baking soda, eggs, Great salt lake, milk of magnesia, ammonia solution, soapy water, liquid drain cleaner, etc.
pH of Salts
As salts are formed by the neutralisation reaction of acid and base, their pH should be equal to 7, indicating that they are neutral.
However, the pH of the majority of salts is less than or greater than 7. The pH of salts can be equal to 7, greater than 7, or less than 7. As a result, salts can be neutral, acidic, or basic.
Salts formed by the reaction of the same strength of acid and base are neutral and have a pH of 7. These salts are known as ‘neutral salts’.
Sodium chloride is neutral with a pH of 7. Sodium carbonate has a pH of 9 and is basic while ammonium chloride is acidic with a pH of 6.
Examples of neutral salt formation with a pH value of 7
$\mathrm{NaOH}+\mathrm{HCl} \rightleftharpoons \mathrm{NaCl}+\mathrm{H}_{2} \mathrm{O}$
An acidic salt is formed when a strong acid reacts with a weak or mild base. This occurs because the weak base was unable to completely neutralise the strong acid, resulting in only partial neutralisation. For instance, ammonium sulphate, ammonium chloride and so on.
$\mathrm{NH}_{4} \mathrm{OH}+\mathrm{HCl} \rightleftharpoons \mathrm{NH}_{4} \mathrm{Cl}+\mathrm{H}_{2} \mathrm{O}$
When a strong base reacts with a weaker acid, basic salts are formed. This occurs because a weaker acid cannot completely neutralise a strong base. A basic salt is formed as a result of the partial neutralisation of the base. For instance, sodium carbonate, calcium carbonate, magnesium carbonate and so on.
Conclusion
pH indicates how much H+ has dissociated from molecules in a solution. The higher the concentration of H+ ions in the solution and the stronger the acid, the lower the pH value. Similarly, as the pH increases, the concentration of H+ ions in the solution decreases and the acid becomes weaker.
FAQs on pH of Solutions for JEE
1. What causes water to self-ionise?
The chemical reaction in which two water molecules react to produce a hydronium (H3O+) and a hydroxide ion (OH-) is known as ‘self-ionisation of water’ (also autoionization of water and auto association of water). It is an example of autoprotolysis that takes advantage of the amphoteric nature of water. Hydronium ions have the same or equal concentration as hydroxide ions in pure water. The concentrations of hydronium and hydroxide ions are equal at $25^{\circ}$C. Water's ion product is the equilibrium condition for self-ionisation and is expressed as follows:
$\begin{align} &\mathrm{K}_{\mathrm{w}}=\left[\mathrm{H}_{3} \mathrm{O}^{+}\right]\left[\mathrm{OH}^{-}\right]=1.0 \times 10^{-14} \\ &\mathrm{pK}{ }_{\mathrm{w}}=\mathrm{pH}^{+}+\mathrm{pOH}^{-}=14 \end{align}$
2. Explain what factors influence the pH of a solution.
Many factors influence the pH of a solution such as the temperature, pressure, etc. The concentration of both H+ ions and OH- ions is 107 M at a neutral pH of 7 (pure water). The pH of basic solutions rises with increasing pressure (for example, NH3 in water), while it falls in acidic solutions (for example SO2 in water). A solution's pH is inversely proportional to the temperature. When the temperature of a solution rises, so do its molecular vibrations, allowing ionisation and the formation of H+ ions. More H+ ions induce acidic behaviour.
3. What are the drawbacks of using a universal indicator to determine pH?
First, let us look at the benefits. They are inexpensive, produce a good result to about a pH unit and can be discarded after use. To about a pH unit, universal indicators are adequate. If you want more precise numbers, that's your problem. They are also single-use only, making them unsuitable for continuous measurement.
A universal indicator would only provide slow and gradual colour changes that are inadequate to accurately determine an endpoint. When possible, the industry has long preferred potentiometric detection via automated instruments over indicator endpoints in titration.