What are Acids and Bases ?

There are three main theories to describe what acidsand bases are

In general each theory widens what chemical reactionscan be considered “ Acid – Base” reactions, than theprevious theory

Let us now examine the three separate theories in detail

Theory 1 : The Arrhenius Theory of Acids and Bases

Came in 1884 from his PhD work on proving that ions exist inwater

Work was not widely accepted at the time but he won the 1903Nobel Prize in Chemistry for it

Svante Arrhenius ( 1859 – 1927 )

According to the Arrhenius Theory, Acids and Bases are definedas follows,

• An Acid is a substance which produces hydrogen ions ( H+) in aqueous solution

HCl(aq) H+(aq) + Cl-

(aq)

Example: Hydrochloric Acid ( HCl )

• A Base is a substance which produces hydroxide ions ( OH- ) in aqueous solution

Example: Sodium Hydoxide ( NaOH )

NaOH(aq) Na+(aq) + OH-

(aq)

NeutralisationWhen an Acid and a Base meet they can “cancel” each otherout, forming a salt and water

Acid + Base Salt + WaterHCl(aq) + NaOH(aq) NaCl + H2O

There are 2 separate reactions combined in this equation

1. Na+ + Cl- NaCl

2. H+ + OH- H2O

According to Arrhenius`s Theory the charged hydrogen ion and thehydroxide ion bond together to form “Neutral” water

The Hydronium Ion ( H3O+)

Developments in Chemistry after Arrhenius showed that it wasimpossible for the hydrogen ion ( H+) to exist freely in solution

1. Too Small - proton diameter 100,000 less than an atom

2. Too Reactive – “naked” + charge has to bind to something

In water the H+ ion immediately binds to a water molecule toform the Hydronium Ion as follows,

H+ + H2O H3O+

This fact does not really change our understanding of Acid-BaseChemistry too much. Just remember when you see H+ in anequation it is really the hydronium ion H3O+

Evidence for the Hydronium Ion ( H3O+)Conductivity: Positive ions in water move towards an anode at a rate consistent with them being the same size as a water molecule. Being 100,000 times smaller, protons would move much faster

Let us now re-examine the reaction between HCl and NaOH inthe context of the hydronium ion ( H3O+ )

HCl (aq) Cl- + H+

H+ + H2O H3O+

Overall equation: HCl(aq) Cl- + H3O+

NaOH(aq) Na+ + OH-

HCl(aq) + NaOH(aq)

NaCl + 2H2O

One H2O molecule was taken from solution to carry the H+ chargeso the original reaction on Slide 5 is still correct as only one newH2O molecule is actually produced

Limitations of the Arrhenius Definition1. Many Acid-Base reactions happen in solutions other than water and some don`t require a solution at all

2. Some Acid- Base reactions don` t produce Hydronium ions ( H3O+) or Hydroxide ions ( OH-)

Example:Ammonia (g) + Hydrogen Chloride (g) Ammonium Chloride (s )

NH3 (g) + HCl(g) NH4Cl(s)

Here Ammonia gas ( Base ) reacts with Hydrogen Chloride gas ( Acid )to give Ammonium Chloride (solid)

No solution + no Hydronium or Hydroxide ions = Need for new Theory!!

The Bronsted-Lowry definition of Acids and Bases

Formulated independently in 1923 by the Danish chemist Johannes Nicolaus Bronsted and the English chemist Thomas Martin Lowry

Both recognised the limitations of the Arrhenius definition and pro-posed a more general definition of Acids and Bases that addressedthe problems on the previous slide

Bronsted Lowry

According to the Bronsted-Lowry theory Acids and Bases are defined as following,

• an acid is a proton donor

• a base is a proton acceptor

This definition solves all the problems Arrhenius had with this reaction

NH3(g) + HCl(g) NH4Cl(s)

The Hydrogen Chloride gas (HCl ) donates a proton to the Ammoniagas (NH3) leaving Cl- . This makes it an acid

The Ammonia gas (NH3) accepts a proton from the HCl becoming NH4+.

This makes it a base

Finally the Cl- ion binds ionically to the NH4+ ion to form the salt

ammonium chloride NH4Cl

Acid/Base Equilibria

So far we have only discussed Acid-Base reactions in a “one way”,reactants just make products manner. A + B C + D

Unfortunately things are not that straight forward and sometimesthe products react together to form the reactants A + B C + D

When we combine the above two equations we get a system whichis said to be in Equilibrium. Reactants making products and productsmaking reactants. A + B C + D

We will now examine the concept of chemical equilibrium in relationto the Bronsted – Lowry theory of acids and bases

Conjugate Acids and Bases

Let us examine the general case of an acid ( HA ) dissociating in water

HA(aq) + H2O(l) H3O+ (aq) + A-

(aq)

In the forward reaction HA acts as an acid by donating a proton tothe H2O becoming A-

In the backwards reaction A- acts as a base by accepting a proton from the H3O+ ion becoming H2O

We say that HA is the conjugate acid of A- and that A- is the conjugatebase of HA

Together they are called a conjugate acid-base pair

HA(aq) + H2O(l) H3O+ (aq) + A-

(aq)

The same can be said of the water molecule involved in the reaction

In the forward reaction it acts as a base by accepting a proton fromHA and in the backwards reaction it acts as an acid by H3O+ donatinga proton to A-

It is also a conjugate acid-base pair

conjugate pair

conjugate pair

Because of the reversibility of Acid- Base reactions it is always possibleto find two substances which act as an acid and two substances which act as bases

Example 2: A Base dissolving in Water

NH3(aq) + H2O(l) NH4+

(aq) + OH-(aq)

Here in the forward reaction H2O acts as an acid by donating a protonto NH3

In the previous example ( an acid dissolving in water ) H2O acted as abase by accepting a proton from the acid

Amphoteric Substance: A substance which can act as either an acid or a base

conjugate pair

conjugate pair

Strong Acids / Weak Acids : Strong Bases / Weak Bases

Strong Acid: Eg: Hydrochloric Acid (HCl)

HCl(aq) + H2O(l) H3O+ (aq) + Cl-

(aq)

Cl- is a very weak conjugate base. It cannot even accept a proton fromH3O+ which is an excellent proton donor

• Strong Bronsted Acids have weak conjugate bases

Although theoretically possible, in reality the backwards reaction does not happen so we can ignore it and rewrite the equation as;

HCl(aq) + H2O(l) H3O+ (aq) + Cl-

(aq)

Strong Acid: is one that virtually 100% ionizes in solution

Weak Acid: Eg. Ethanoic Acid (CH3COOH)

CH3COOH + H2O CH3COO- + H3O+

CH3COO- (Acetate) is a strong conjugate base which readily accepts aproton from H3O+ to form H2O

As a result the backward reaction is far more successful than the forward reaction with only around 1% of the CH3COOH ionised toCH3COO- in solution

• Weak Bronsted Acids have strong conjugate bases

Weak Acid: is one which only partially ionizes in solution

Strong Base: Eg: Sodium Hydroxide NaOH

Exactly the same as for a strong acid, the backwards reaction is ignoredand the same definition is used

Strong Base: is one that virtually 100% ionizes in water__________________________________________________________

Weak Base: Eg: Ammonia NH3

NH3(aq) + H2O(l) NH4+

(aq) + OH-(aq)

NH4+ (Ammonium) is a strong conjugate acid which readily donates a

proton to OH- to form H2O in the backwards reaction

• Weak Bronsted Bases have strong conjugate acids

Weak Base: is one that only partially ionizes in solution

Theory 3 : The Lewis Theory of Acids and Bases

Proposed by American Chemist Gilbert Newton Lewis in 1923,the same year as Bronsted-Lowry theory

Looked more at the chemical bonding involved in Acids and Bases

Gilbert N. Lewis

According to the Lewis Theory of Acids and Bases,

• an acid is an electron pair acceptor

• a base is an electron pair donor