Acid and base are essential topics in the field of Chemistry. They are two major categories of substances which include chemicals used in laboratories and substances that are commonly found in our everyday lives. Additionally, it is important to understand these concepts from an examination perspective. In this article, we will explore what acid and base are and look at their various characteristics and properties.

Robert Boyle’s Definition of an Acid and Base

As per Robert Boyle, he defined acidic and basic substances as:

a) Acid as,

1) Having a sour taste

2) Being corrosive and hazardous to human health

3) Litmus

Based on their occurrence, acids are divided into two types: Natural and Mineral.

Natural Acids: These are obtained from natural sources, such as fruits and animal-derived products.

Examples of acids include: lactic, citric, and tartaric acid, etc.

Mineral Acids: Mineral acids are substances derived from minerals that are capable of producing acidic solutions when dissolved in water.

For example, Hydrochloric acid (HCl), Sulphuric Acid (H2SO4), and Nitric Acid (HNO3) etc.

b) Base as:

1. Having a slick surface

2) Changing the color of litmus from red to blue.

Arrhenius Definition of Acids and Bases

Arrhenius suggested that;

Acid

Acid is a substance that releases H+ ions when dissolved in water.

(\begin{array}{l}HCL_{(g)}\overset{H_{2}O}{\rightarrow} H^{+}_{(aq)} + CL^{-}_{(aq)}\end{array})

Hydrogen chloride ionizes in water to give hydrogen ions (H⁺).

Depending on the number of protons the acid gives, it can be classified as a mono-, di-, or tribasic acid.

Monobasic Acids:

• HCl
• Nitric Acid
• Acetic Acid

Dibasic Acids:

• Sulphuric Acid
• Phosphorous Acid

Phosphoric Acid: Tribasic Acid

(\begin{array}{l}NaOH_{(s)}\overset{H_{2}O}{\longrightarrow} Na^{+}_{(aq)} + OH^{-}_{(aq)}\end{array} )

Base

A base is a substance that produces OH– ions when dissolved in water.

NaOH dissociates in water to give the hydroxide (OH–)

Based on the number of available hydroxide ions, a base can be classified as monoacidic, diacidic, or triacidic.

Mono basic:

• NaOH
• NH4OH

Dibasic:

• Ca(OH)₂
• Zn(OH)₂

Tribasic: Fe(OH)3, Al(OH)3

Acids and bases have similar properties due to the neutralization of bases by acids and vice versa. This process occurs when the H+ ion from the acid reacts with the OH– ion from the base, forming water.

H<sup>+</sup>(aq) + OH<sup>-</sup>(aq) → H<sub>2</sub>O(l)

Strong acids such as HCl, HNO3, H2SO4, and HCIO4 ionize completely in solutions, producing larger hydrogen ions. Similarly, strong bases such as NaOH, KOH, and (CH3)4NOH produce larger hydroxide ions when ionized.

Weak acids and bases are those acids and bases that are only partially dissociated in solution, resulting in a lower concentration of hydrogen or hydroxide ions.

Compounds containing ionizable hydrogen or hydroxide ions can be either an acid or a base. However, CH4 is not an acid. Similarly, CH3OH, C2H5OH, etc., having OH groups are not bases.

Advantages and Limitations of Arrhenius Theory of Acids and Bases

i) Only applicable to aqueous solutions, as acids and bases are defined by their ionization in water.

ii) Nonmetal oxides are acidic due to the release of oxygen ions, while ammonia, sodium carbonate, and metal oxides are basic because they release hydroxide ions.

Relative Strengths of Acids and Bases

The levelling effect states that all strong acids and bases are equally ionized and water is amphoteric, meaning they have the same acidic or basic strength in water.

The strength of acids depends on the solvent. Acetic acid is not able to take up protons and must be forced to do so. Therefore, acids such as HCIO4, H2S04, HCl, and HN03, which have similar strengths in water, follow the order of: HClO4 > H2SO4 > HCl > HNO3 in acetic acid.

The real strength of acids can be judged by solvents. On the basis of proton interaction, solvents can be classified into four types:

Protophilic Solvents: Solvents that have a greater tendency to accept protons, such as water, alcohol, liquid ammonia, etc.

Protogenic Solvents: Solvents that have the tendency to produce protons, such as water, liquid hydrogen chloride, and glacial acetic acid.

Amphiprotic Solvents: Solvents that can act as both a proton donor and proton acceptor, such as water, ammonia, ethyl alcohol, etc.

Aprotic Solvents: Solvents that do not donate or accept protons, such as benzene, carbon tetrachloride, and carbon disulphide.

HCI acts as an acid in the water, a stronger acid in NH3, a weak acid in CH3COOH, neutral in C6H6 and a weak base in HF.

Bronsted-Lowry Theory of Acids and Bases

Bronsted acids are hydrogen-ion donors or proton donors.

Bronsted bases are hydrogen-ion acceptors or proton donors.

HCl donates an H<sup>+</sup> ion to a water molecule to form H<sub>3</sub>O<sup>+</sup>.

(\begin{array}{l}HCl(g) + H_2O(l) \rightleftharpoons H_3O^+(aq) + Cl^-(aq) \end{array})

HCl acts as an acid and water is a base.

Neutralization, as per Bronsted model, involves the transfer of an H⁺ ion from an acid to a base. Acids can be

i) Neutral molecules. \ $$HCl(g) + NH_3(aq) \rightleftharpoons NH_4(aq) + Cl^-(aq)$$

ii) Positive ions $\begin{array}{l}NH_4^+(aq) + OH^-(aq) \rightleftharpoons NH_3(aq) + H_2O(l)\end{array}$

iii) Negative ions. \ $\begin{array}{l} H_2O(l) + H_2PO_4^-(aq) \rightleftharpoons HPO_4^{2-}(aq) + H_3O^-(aq) \end{array}$

Conjugate Acid-Base Pairs

(\begin{array}{l}HCl(aq)+NH_3(aq) \rightleftharpoons NH_{4}^{+}(aq) + H_3O^{-}(aq)\end{array})

HCl and Cl⁻ differ by a proton and so do NH₃ and NH₄⁺.

NH4+, like HCl, can donate a proton and thus is an acid. Since it is derived from the base ammonia, it is referred to as the conjugate acid of base ammonia.

Similarly, Cl^- like ammonia can accept a proton and hence a base. It is considered as the conjugate base of the acid HCl.

The Bronsted–Lowry transfer can be written as:

$$\begin{array}{l}HCl(aq) + NH_3(aq) \rightleftharpoons NH_4^+(aq) + Cl^-(aq)\end{array}$$

Acid + Base ⇒ Conjugate Acid + Conjugate Base

Conjugate Pairs

A conjugate acid contains one more H atom and a positive charge than the base that formed it.

A conjugate base contains one less H atom and one more negative charge than the acid-forming it.

The product side of the reaction will always list the conjugates.

A weak acid forms a strong conjugate base and vice versa.

Reactions always proceed from a strong acid/base to a weak acid/base.

The Acid-Base Nature of Compounds

1. Compounds containing hydrogen bound to a nonmetal are referred to as nonmetal hydrides, and are typically acidic.

(\begin{array}{l}HCl_{(g)}\overset{H_{2}O}{\rightarrow} H^{+}_{(aq)} + Cl^{-}_{(aq)}\end{array} )

(\begin{array}{l}H_{2}S_{(g)} \overset{H_{2}O}{\rightarrow} H^{+}_{(aq)} + HS^{-}_{(aq)}\end{array})

2. Metal hydrides contain hydrogen (H⁻) bound to a metal, which gives the H⁻ (or hydride) ion.

(\begin{array}{l}NaH_{(s)}\rightarrow Na^{+}{(aq)} + H^{-}{(aq)}\end{array})

The H^- ion, with its pair of valence electrons, can abstract an H⁺ ion from a water molecule and increase OH^- ion concentration. Therefore, in a solution, metal hydrides are bases.

(\begin{array}{l}NaH_{s} + H_{2}O_{l} \rightarrow Na^{+}_{aq} + OH^{-}_{aq} + H_{2(g)}\end{array})

(\begin{array}{l}CaH_{2(s)} + 2H_{2}O_{(l)} \rightarrow Ca^{2+}_{(aq)} + 2OH^{-}_{(aq)} + 2H_{2(g)}\end{array} )

3. Nonmetal oxides dissolve in water to form acids, while CO2 dissolves in water to give carbonic acid.

(\begin{array}{l}CO_{2(g)} + H_{2}O_{(l)} \rightarrow H_{2}CO_{3(aq)}\end{array})

$\ce{SO3(g) + H2O(l) -> H2SO4(aq)}$

(\begin{array}{l}P_{4}O_{10(s)} + 6H_{2}O_{(l)} \rightarrow 4H_{3}PO_{4(aq)}\end{array} )

Metal oxides, containing O^2- ions, react with water to produce a pair of OH^- ions and a base.

(\begin{array}{l}O^{2}{(aq)} + H{2}O_{(l)} \rightarrow 2OH^{-}_{(aq)}\end{array})

Metal oxides therefore fit the operational definition of a base.

(\begin{array}{l}CaO_{s(aq)} + H_{2}O_{(l)}\rightarrow Ca^{2+}{(aq)} + 2OH^{-}{(aq)}\end{array})

Metal hydroxides, such as LiOH, NaOH, KOH, and Ca(OH)2, are considered bases.

(\begin{array}{l}NaOH_{s} \rightarrow Na^{+}{(aq)} + OH^{-}{(aq)} \overset{H_2O}{\leftarrow}\end{array})

The large difference in electronegativities between sodium ( EN = 2.5) and oxygen ( EN = 3.5) results in the electrons in the Na O bond being drawn towards the more electronegative oxygen atom, rather than being shared equally. As a result, when NaOH dissolves in water, it dissociates to give Na+ and OH– ions.

(\begin{array}{l}NaOH_{(s)} \overset{H_2O}{\rightarrow} Na^{+}_{(aq)} + OH^{-}_{(aq)}\end{array})

6. Nonmetal hydroxides, such as Hypochlorous Acid (HOCl), HONO2, O2S(OH)2, and OP(OH)3, are

Acids

(\begin{array}{l}HOCl_{(aq)} \rightarrow H^{+}_{(aq)} + OCl^{-}_{(aq)} \end{array})

The difference in electronegativities of the chlorine and oxygen atoms is small ( EN = 0.28), so the electrons in the Cl O bond are shared more or less equally. However, the O H bond is polar ( EN = 1.24)  so the electrons in this bond are drawn towards the more electronegative oxygen atom, resulting in OCl– and H+ ions.

(\begin{array}{l}HOCl_{(aq)} \rightarrow H^{+}{(aq)} + OCl^{-}{(aq)}\end{array})

The acidic hydrogen atoms in the non-metal hydroxides are not bound to the nitrogen, sulfur, or phosphorus atoms, but only to the oxygen atom. These compounds are known as oxyacids.

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Amphoteric compounds, such as Al2O3 and Al(OH)3, lie between metal and nonmetal oxides or metal and nonmetal hydroxides and can act as either acids or bases. For example, Al(OH)3 acts as an acid when it reacts with a base.

(\begin{array}{l}Al(OH){3(s)} + OH^{-aq} \rightarrow AlOH{4}^{- (aq)} \end{array})

Conversely, it acts as a base when it reacts with an acid.

(\begin{array}{l}Al(OH){3(s)} + 3H^{+} \rightarrow Al^{3+}{(aq)} + 3H_{2}O_{(l)}\end{array})

The Lewis Concept of Acid-Base Reactions

Species that can donate electron pairs are Lewis bases and are called acids.

1. Molecules having a central atom with an incomplete octet (less than 8 electrons):

• BF₃
• BCl₃
• AlCl₃
• MgCl₂
• BeCl₂
• etc.

2. Molecules having a central atom with empty d-orbitals:

• SiX₄
• GeX₄
• TiCl₄
• SnX₄
• PX₃
• PF₅
• SF₄
• SeF₄
• TeCl₄, etc.

Molecules with multiple bonds between atoms of different electronegativity, such as CO2, SO2, and SO3, can be attacked by a Lewis base. When this happens, one electron pair is shifted towards the atom with the more negative electronegativity.

4. Simple Cations:

• H⁺
• Ag⁺

Lewis bases are species that can donate a pair of electrons to form a covalent bond.

1. Neutral species having at least one lone pair of electrons:

2. Negatively Charged Species or Anions: For example, Chloride, Cyanide, Hydroxide Ions, etc.

3. It may be noted that all Bronsted bases are also Lewis bases, but not all Bronsted acids are Lewis acids.

The following compounds, for example, contain non-bonding pairs of electrons.

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

pH Scale and Acidity

pH and Solutions

Hydrolysis, Salts, and Types

Buffer Solutions

Frequently Asked Questions on Acid and Base

Examples of Natural Acids:

• Citric Acid
• Lactic Acid
• Acetic Acid
• Tartaric Acid
• Malic Acid

Examples of mineral acids include hydrochloric acid, sulfuric acid, and nitric acid.

An Arrhenius acid is a substance that increases the concentration of hydrogen ions when dissolved in water, resulting in an increase in the acidity of the solution. Examples of Arrhenius acids include nitric acid, hydrochloric acid and sulfuric acid.

An Arrhenius acid is an acid that gives H+ ions in water. For example, Sulfuric acid is a dibasic acid.

Sulfuric acid is a dibasic acid.### The Bronsted-Lowry concept of an acid is a substance that donates a proton, and a base is a substance that accepts a proton.

Bronsted acids are proton donors, whereas Bronsted bases are proton acceptors.