The pH scale is a commonly used scale to measure the acidity or basicity of a substance. The possible values on the pH scale range from 0 to 14. Acidic substances have pH values ranging from 1 to 7 (1 being the most acidic point on the pH scale) and alkaline or basic substances have pH values ranging from 7 to 14. A perfectly neutral substance would have a pH of exactly 7.

The pH, an abbreviation of ‘potential for hydrogen’ or ‘power of hydrogen’, of a substance can be expressed as the negative logarithm (with base 10) of the hydrogen ion concentration in that substance. Similarly, the pOH of a substance is the negative logarithm of the hydroxide ion concentration in the substance. These quantities can be expressed using the following formulae:

pH = -log[H+]

pOH = -log[OH⁻]

It is important to note that the pH scale is a logarithmic scale, meaning that a single-unit increase in pH corresponds to a tenfold increase in hydrogen ion concentration and acidity. For example, a solution with a pH of 3 has ten times the acidity of a solution with a pH of 4 and a hundred times the acidity of a solution with a pH of 5. Additionally, pH is a dimensionless quantity.

Also Check Out: pH and Solutions

Introduction to pH Scale

The pH scale is mostly attributed to a set of chosen solutions whose pH is already established by an international committee. Alternatively, the primary pH standard values can be determined using a concentration cell with transference, wherein the potential difference is measured between a hydrogen electrode and a standard electrode. Typically, the pH of an aqueous solution is measured with the help of either a glass electrode or a pH meter. A colour-changing indicator can also be used. In any case, the pH measurement and pH scale are essential in various fields of Chemistry, Medicine, Water treatment and other applications.

Properties of the pH Scale

The pH scale is very convenient to use, as it can convert odd expressions such as 1.24 × 10-4 into a single number of 3.89.

This scale covers a very wide range of [H+][H+].

The division between zero and 1 is expanded to a linear scale, allowing for a more comprehensive comparison. Additionally, the pH scale expands a compact scale into a large scale for comparison purposes.

The pH scale actually has positive values due to the use of a negative log of [H+][H+]. It should be noted that the [H+][H+] will be smaller if the pH is large.

The pH range has no lower or upper bound due to the fact that it is an indication of the concentration of H+ ions.

Unified Absolute pH Scale

In 2010, a new unified absolute pH scale was proposed. This scale would allow for various pH ranges across various solutions to be measured based on a common proton reference standard, which was to be developed on the basis of the absolute chemical potential of the proton. Furthermore, this model was based on the Lewis acid–base definition and can be applied to liquids, gases and solids.

pH Scale and Acidity

Acid solutions typically contain protons, while basic solutions contain hydroxide ions, which have low concentrations (expressed as negative powers of ten). The pH scale provides a convenient way to express these concentrations as numbers between 1 and 14.

The pH is the negative logarithm to the base ten of hydrogen ion concentration in moles per litre.

pH = -log[H⁺]

p(OH) is the negative logarithm to the base ten of hydroxide ion concentration in moles per litre.

pOH = -log[OH⁻]

pH + pOH = 14

The pH scale is based on neutral water, where [H⁺] = [OH^-] = 10^-7.

pH = +7

The pH of a strong acid decreases with a limit of 1, and the pH of a base increases up to 14.

Generally, acids and bases will have a pH between 0 and 14.

But negative and greater than 14 pH values are also possible.

Also Check Out: Study The pH Change

Limitations of the pH Scale

1. pH values do not directly reflect the relative strength of acids or bases. A solution of pH = 1 has a hydrogen ion concentration that is 100 times greater than a solution of pH = 3 (not three times). A 4 x 10^-5N HCl solution is twice as concentrated as a 2 x 10^-5N HCl solution, but the pH values of these solutions are 4.40 and 4.70 (not double).

2. pH value is zero for 1 N solution of strong acid. The concentration of 2 N, 3 N, 10 N, etc. gives negative pH values.

A solution of an acid having a very low concentration, say 10^-8 N, shows a pH of 8 and hence should be basic, but the actual pH value is less than 7.

4. A range of 0 to 14 provides reasonable (but not definitive) measurements for the scale. In some cases, the concentrations of hydronium ions or hydroxide ions can go beyond one molar, resulting in values below zero or above 14.

Periodic Variation of Acidic and Basic Properties

Hydrides of the Elements of the Same Periods

The increase in acidicity along the period is due to the fact that the stability of the conjugate bases of CH4, NH3, H2O and HF increases in that order, making them increasingly more acidic.

CH₃- < NH₂- < OH^- < F^-

Hydrides of the Elements of the Same Group

1. As the column progresses, the acidic nature decreases. Hydrides of V group elements (NH3, PH3, AsH3, SbH3, BiH3) demonstrate basic character, which decreases as the size increases and the electronegativity decreases from N to Bi. This is due to a decrease in electron density in the sp3-hybrid orbital, resulting in a decrease in electron donor capacity.

The increasing basic properties reflect a decreasing trend in the electron donor capacity of OH–, HS–, HSe– or HTe– ions for the hydracids of VI group elements (H20, H2S, H2Se, H2Te), which act as weak acids. The strength increases in the order H20 < H2S < H2Se < H2Te.

The decreasing bond energies of hydracids of VII group elements (HF, HCl, HBr, HI) explain why these compounds show increasingly acidic properties from HF to HI.

The heat of formation (H-F) of H₂O is 135 kcal/mol, HCl is 103 kcal/mol, HBr is 88 kcal/mol and HI is 71 kcal/mol.

Oxyacids

The acidity of oxyacids of the same element in different oxidation states increases with an increase in its oxidation number.

+1 +3 +5 +7

HC1O2 < HC1O3 < HCIO4 < HCIO

4 + 6 + 3 + 5 = 18

H2SO3 < H2SO4; HNO2 < HNO3

But this rule fails in oxyacids of phosphorus.

H3PO2 > H3PO3 > H3PO4

The electronegativity of the oxyacids of different elements which are in the same oxidation state decreases as the atomic number increases. This is due to an increase in size and a decrease in acidic properties.

HBrO4 > HC1O4 > HIO4

H2SO3 < H2SeO3

But there are a number of acid-base reactions in which no proton transfer takes place, such as:

SO2 + SO2 <=> SO2++ S

Acid1 Base2
Acid2 Base1

Therefore, the protonic definition is not applicable to reactions taking place in non-protonic solvents such as COCl2, S02, N2O4, etc.

Water - Amphoteric Weak Electrolyte

1) Water can behave like both an acid and a base, so it is amphoteric.

Water acts as a base and accepts protons from HCl

(\begin{array}{l}NH_{4}^{+}(aq)+OH^{-}(aq)\rightleftharpoons NH_3(aq)+H_2O(l)\end{array})

Molarity of Water

Molarity = Number of moles per litre of solution = 55.55 mol/L

Ionization Constant of Water

(\begin{array}{l}H_2O \leftrightarrow H_{4}^{+} + OH^{-}\end{array})

(\begin{array}{l}Ka = Kb = \frac{[H^+][OH^-]}{[H_2O]} = \frac{10^{-7} \times 10^{-7}}{55.55} = 1.8 \times 10^{-16}\end{array})

Where, Ka is the acid ionization constant and Kb is the base ionization constant.

pKa = pKb = 15.74, where Ka = $1.8\times10^{-16}$

Degree of Ionization of Water

(\begin{array}{l}H_2O \leftrightarrow H^+ + OH^-\end{array})

Initial concentration (moles): 55.55

At equilibrium, moles = 10^-7

Degree of ionization = α = $\frac{number;of;moles;ionized}{initial;number;of;moles}$ = $\frac{10^{-7}}{55.55} = 1.8\times 10^{-9}$

Only about 2 parts per billion (ppb) of the water molecules dissociate into ions at room temperature.

Ionic Product of Water

The concentration of hydrogen and hydroxide ions in water is the product of it.

Ionic product of water (Kw) = [H+][OH–] = 10^-14

pKw = -log[Kw] = -log₁₀(-14) = 14

The values of Ionic product, pKw, pKa and pKb remain constant regardless of whether the solution is acidic, neutral or basic.

Related Topics

Chemical Equilibrium

Ionic Equilibrium – Degree of Ionization and Dissociation

Equilibrium Constant – Characteristics and Applications

Le Chatelier’s Principle on Equilibrium

Solubility and Solubility Product

Acid and Base

Hydrolysis, Salts, and Types

Buffer Solutions