Definitions of Acids and Bases Tutorial

Key Concepts

Many substances can be classified as either acids or bases based on their properties.

The definition of what an acid or a base is, is dependent on the chemical system being studied.
For aqueous solutions, acids are often defined by their ability to produce hydrogen ions (H⁺, also known as protons or hydrons)¹.

The definition that is most useful for students of Chemistry working with aqueous solutions² is the definition proposed by Brønsted and Lowry independently in 1923 (also known as the Brønsted-Lowry definition, or Brønsted-Lowry theory, or the Lowry-Brønsted definition, or Lowry-Brønsted theory):
An acid is a substance that can donate a hydrogen ion,
that is, an acid is a proton donor.
A base is a substance that can accept a hydrogen ion, that is,
a base is a proton acceptor.

When a Brønsted-Lowry acid donates (loses) a proton it produces a hydrogen ion (proton) and a conjugate base:

acid → proton + conjugate base

HB → H⁺ + B^-

When a Brønsted-Lowry base accepts (gains) a proton it produces a conjugate acid:

base + proton → conjugate acid

B^- + H⁺ → HB

A monoprotic acid is a Brønsted-Lowry acid that can only donate 1 proton :
HA → H⁺ + A^-

A diprotic acid is a Brønsted-Lowry acid that can donate 2 protons :
H₂A → 2H⁺ + A^-

A triprotic acid is a Brønsted-Lowry acid that can donate 3 protons :
H₃A → 3H⁺ + A^-

The term polyprotic is used to describe any Brønsted-Lowry acid that is capable of donating more than 1 proton.

An amphiprotic substance is one that can either donate or accept a proton.

The definition of an amphiprotic substance is based on the Brønsted-Lowry definition of acids and bases.

HB^- donating a proton Brønsted-Lowry Acid					HB^- accepting a proton Brønsted-Lowry Base
acid	→	proton	+	conjugate base	base	+	proton	→	conjugate acid
HB^-	→	H⁺	+	B^2-	HB^-	+	H⁺	→	H₂B
HB^- is amphiprotic, it can either accept a proton to form H₂B, or, it can donate a proton to form B^2-

The term "amphoteric substance" is derived from an earlier definition of acids and bases proposed by Arrhenius:
An amphoteric substance is one that can act as either an Arrhenius acid, or, as an Arrhenius base:
Arrhenius acids produce protons in aqueous solution :
HA → H⁺_(aq) + A^-_(aq)

Arrhenius bases produce hydroxide ions in aqueous solution:
BOH → OH^-_(aq) + B⁺_(aq)

An Arrhenius acid reacts with an Arrhenius base (OH^-_(aq)) to produce a salt.
An Arrhenius base reacts with an Arrhenius acid (H⁺_(aq)) to produce a salt.

An amphoteric substance can therefore react with either an Arrhenius acid (H⁺_(aq)), or, with an Arrhenius base (OH^-_(aq)) to produce a salt.

Unless otherwise stated, or implied, the terms acid and base as used in High School Chemistry will refer to Brønsted-Lowry acids and bases.
For example, if the definition being used in a particular example is that of Arrhenius, the acid will be stated to be an Arrhenius acid.
If a question refers to an amphoteric substance, the implication is that the Arrhenius definition of acids and bases is being used.

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Defintions of Acids and Bases in Chronological Order

ACIDS

BASES

Lavoisier
(1779)

Definition:
An acid contains oxygen.

Limitations:
Does not account for substances with acidic properties that do not contain oxygen, for example, hydrochloric acid, HCl.

Davy
(1810)

Definition:
An acid contains hydrogen.

Limitations:
Does not account for substances which contain hydrogen but do not display acidic properties.

Arrhenius Definitions
(1884)

Definition:
An acid dissociates (or ionises) in water to produce hydrogen ions, H⁺_(aq) (protons or hydrons).

Definition:
A base dissociates (or ionises) in water to produce hydroxide ions, OH^-_(aq).

Examples:

acid	→	proton	+ anion
hydrochloric acid:
HCl	→	H⁺_(aq)	+ Cl^-_(aq)
nitric acid:
HNO₃	→	H⁺_(aq)	+ NO₃^-_(aq)

Examples:

base	→	hydroxide ion	+ cation
sodium hydroxide:
NaOH	→	OH^-_(aq)	+ Na⁺_(aq)
potassium hydroxide:
KOH	→	OH^-_(aq)	+ K⁺_(aq)

Limitations:
Only accounts for acids that:

are aqueous solutions

have hydrogen in their structure, eg, HCl

does not account for amphoteric substances
(those that can act as an Arrhenius acid or an Arrhenius base)

Limitations:
Only accounts for bases that:

are aqueous solutions (bases that are also known as alkalis)

have OH^- already in their structures, eg, NaOH

does not account for amphoteric substances
(those that can act as an Arrhenius acid or an Arrhenius base)

Definition:
An amphoteric substance can act as either an Arrhenius acid or as an Arrhenius base, that is:

An Arrhenius acid will react with an Arrhenius base.
An Arrhenius base will react with an Arrhenius acid.

Example:
According to the Arrhenius definition of acids and bases, zinc hydroxide, Zn(OH)_2(s), should be a base, that is, it dissolves in water to produce hydroxide ions, OH^-_(aq).³
However, zinc hydroxide can be shown to be amphoteric, it can act as both an Arrhenius acid by reacting with an Arrhenius base, and, it can act as an Arrhenius base by reacting with an Arrhenius acid!

(i) Zinc hydroxide will act like an Arrhenius acid when it dissolves in aqueous Arrhenius base (OH^-_(aq)):

Zn(OH)_2(s) + OH^-_(aq) → Zn(OH)₃^-_(aq)

(ii) Zinc hydroxide will act like an Arrhenius base when it dissolves in aqueous Arrhenius acid (H⁺_(aq)):

Zn(OH)_2(s) + H⁺_(aq) → Zn(OH)⁺_(aq) + H₂O_(l)

Brønsted-Lowry Definitions
(1923)

Definition:
An acid is a species that donates a proton (H⁺).

Definition:
A base is a species that accepts a proton (H⁺).

Examples:

HB H⁺ + B^-

HB is acting as an acid by donating a proton, H⁺.
B^- is the conjugate base of the acid HB .

Examples:

B^- + H⁺ HB

B^- is acting as a base by accepting a proton, H⁺ .
HB is the conjugate acid of the base B^- .

Examples:

acid	→	conjugate base	+ proton
HCl	→	Cl^-	+ H⁺
HNO₃	→	NO₃^-	+ H⁺
H₂SO₄	→	HSO₄^-	+ H⁺
HSO₄^-	→	SO₄^2-	+ H⁺
H₂O	→	OH^-	+ H⁺

Examples:

base	+ proton	→	conjugate acid
OH^-	+ H⁺	→	H₂O
NH₃	+ H⁺	→	NH₄⁺
CO₃^2-	+ H⁺	→	HCO₃^-
HCO₃^-	+ H⁺	→	H₂CO₃
H₂O	+ H⁺	→	H₃O⁺

Definition:
An amphiprotic substance can act as a proton donor and as a proton acceptor.
An amphiprotic substance can act as either a Brønsted-Lowry acid or as a Brønsted-Lowry base.

Example:
hydrogen carbonate (bicarbonate) ion (HCO₃^-) is amphiprotic

(i) hydrogen carbonate (bicarbonate) ion acting as a Brønsted-Lowry acid:

HCO₃^- → CO₃^2- + H⁺

(ii) hydrogen carbonate (bicarbonate) ion acting as a Brønsted-Lowry base:

HCO₃^- + H⁺ → H₂CO₃

Limitations:
The solvent system must be protonic, for example, the solvent system can be water (H₂O_(l)), or liquid ammonia (NH_3(l)).

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Sample Questions with Worked Solutions

Question 1. What is the conjugate base of H₂S ?

Define "conjugate base":
A Brønsted-Lowry acid donates (loses) a proton to form a conjugate base.

Identify the Brønsted-Lowry acid in the question:
H₂S must be the Brønsted-Lowry acid because we are asked to find its conjugate base

Write the balanced chemical equation to represent the Brønsted-Lowry acid donating (losing) a proton (H⁺) to produce its conjugate base:

acid → proton + conjugate base

H₂S → H⁺ + HS^-

State your answer:
HS^- is the conjugate base of the acid H₂S

Question 2: Explain why H₂PO₄^- is said to be amphiprotic.

Define "amphiprotic":
An amphiprotic substance is one that can either:
(i) donate (lose) a proton
or
(ii) accept (gain) a proton

Write a balanced chemical equation to show H₂PO₄^- donating (losing) a proton:
H₂PO₄^- → H⁺ + HPO₄^2-

Write a balanced chemical equation to show H₂PO₄^- accepting (gaining) a proton:
H₂PO₄^- + H⁺ → H₃PO₄

H₂PO₄^- is amphiprotic because :
H₂PO₄^- can either:
(i) lose a proton to form HPO₄^2-
or
(ii) gain a proton to form H₃PO₄

Question 3. Explain why the carbonate ion, CO₃^2-, is a Brønsted-Lowry base but not an Arrhenius base.

Define each term:
A Brønsted-Lowry base accepts (gains) a proton (H⁺) to form a conjugate acid.
An Arrhenius base dissociates in water to produce hydroxide ions (OH^-)

Write a balanced chemical equation to show NH₃ acting as a Brønsted-Lowry base:
CO₃^2- + H⁺ → HCO₃^-

Demonstrate why CO₃^2- is not an Arrhenius base:
CO₃^2- does not contain any hydroxide ions (OH^-).
Therefore CO₃^2- cannot dissociate in water to produce hydroxide ions.

Write an answer to the question:
CO₃^2- is Brønsted-Lowry base because it can accept a proton to form its conjugate acid, HCO₃^-.
CO₃^2- is not an Arrhenius base because it cannot dissociate in water to produce hydroxide ions.

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1. There are lots of other definitions of acids and bases, for example,

Solvent-system definition extends the Arrhenius concept beyond water to other self-dissociating solvents, an acid increases the concentration of cations related to the solvent, a base increases the concentration of anions related to the solvent.

Lewis (1923) acids are electron-pair acceptor, Lewis bases are electron-pair donors.

Usanovich (1939) acids react with bases, give up cations, or accept anions or electrons, and Usanovich bases react with acids, give up anions or electrons, or combines with cations.

Lux-Flood (1947) acids are oxide acceptors, Lux-Flood bases are oxide donors.

2. The solvent does not have to be water, but it does have to be protonic, for example, the solvent could be liquid ammonia or sulfuric acid.

3. Zinc hydroxide, Zn(OH)₂, is only sparingly soluble in water at 25°C so much so that it is usually referred to as insoluble.
K_sp is very, very small, so very little Zn(OH)_2(s) will dissolve in water.
K_sp = 5 × 10^-17