Brønstead-Lowry Acids and Bases

The Arrhenius definition of an acid and a base:

• Acids - substances that when dissolved in water release H⁺ ions
• Bases - substances that when dissolved in water release OH^- ions
• The definition of an Arrhenius acid and base emphasizes the H⁺ and OH^- ions in water

HCl is an Arrhenius acid:

• When HCl dissolves in water it actually transfers a proton to a water molecule:

HCl(g) + H₂O(l) ® H₃O⁺(aq) + Cl^-(aq)

• The recognition that the release of H⁺ ions by acids involves H⁺ transfer led to a new proposal for the definition of what is an acid, and what is a base

The Brønstead-Lowry definition of an acid and a base:

• Acids - a substance that can transfer a proton to another substance
• Base - a substance that can accept a proton from another substance
• These definitions emphasize proton transfer, and can include solvents other than water (aqueous solutions are not part of the definition, proton transfer is the key feature)
• The definition was developed independently in 1923 by Johannes Brønstead and Thomas Lowry

HCl(g) + H₂O(l) ® H₃O⁺(aq) + Cl^-(aq)

• In the reaction of HCl with H₂O, HCl is a Brønstead-Lowry acid (donates a proton to H₂O), and the H₂O (in this particular reaction) is a Brønstead-Lowry base (accepts a proton from the HCl)

HCl(g) + NH₃(g) ® Cl^-(g) + NH₄⁺(g)

• In the reaction of HCl with NH₃, a proton is transferred from the HCl to the NH₃.
• The HCl is the Brønstead-Lowry acid
• NH₃ is the Brønstead-Lowry base
• H₂O is not involved in the reaction (or definition) of the acid or base in this reaction

NH₃(aq) + H₂O(l) ó NH₄⁺(aq) + OH^-(aq)

• In the above reaction of ammonia with water, the ammonia is a Brønstead-Lowry base, and the H₂O is acting as a Brønstead-Lowry acid

• The ammonia is also an Arrhenius base (it increases the concentration of OH^- ions in water)

• A substance can only work as a Brønstead-Lowry acid (i.e. donates a proton), if another substance simultaneously acts as a Brønstead-Lowry base (i.e. accepts the proton)
• To be a Brønstead-Lowry acid, a molecule must have a H atom that can it can lose as an H⁺ ion
• To be a Brønstead-Lowry base, a molecule must have a non-bonding pair of electrons that it can use to bind the H⁺ atom (this non-bonding pair will form the basis of the "shared" electrons in the new covalent bond with the H atom)

In any acid-base equilibrium both the forward and reverse reactions involve proton transfer reactions. For example, the general reaction of a Brønstead-Lowry acid with water proceeds as follows:

HA(aq) + H₂O(l) ó A^-(aq) + H₃O⁺(aq)

acid + base ó base + acid

• In the forward reaction, HA donates a proton to H₂O
• HA is the Brønstead-Lowry acid
• H₂O is the Brønstead-Lowry base
• In the reverse reaction, H₃O⁺ donates a proton to A^-
• H₃O⁺ is the Brønstead-Lowry acid
• A^- is the Brønstead-Lowry base

• HA and A^- are termed a conjugate acid-base pair. They differ only in the presence or absence of a proton.
• A^- is the conjugate base of HA
• H₃O⁺ and H₂O are also a conjugate acid-base pair.
• H₃O⁺ is the conjugate acid of H₂O
• In any acid-base proton transfer reaction there will be conjugate acid-base pairs

A strong acid is a molecule that has a strong preference to donate a proton.

• Thus, its conjugate base will have a weak tendency to accept a proton

A strong base is a molecule that has a strong preference to accept a proton

• Thus, its conjugate acid will have a weak tendency to donate a proton

All molecules of a strong acid in water will donate their protons (to H₂O)

• The conjugate base has no tendency to be protonated.

Weak acids have conjugate bases that have a moderate tendency to be protonated.

• Thus, in solution, only a fraction of the molecules of a weak acid will donate a proton. There will be a significant concentration of both the acid and conjugate base forms in solution

What about molecules that contain H atoms, e.g. methane (CH₄), but do not appear to have any acidic character?

• Their conjugate bases must be quite strong!
• CH₃^- is the conjugate base if CH₄ behaves as an acid
• CH₃^- is such a strong base that it strips H⁺ from H₂O with such vigor that NO CH₃^- will remain in solution
• Thus, CH₄ (the conjugate acid to CH₃^-) is technically an incredibly weak acid - absolutely NO molecules will donate a proton
• If we were to consider the "reaction" of methane with water we would have:

CH₄(g) + H₂O(l) ¬ CH₃^-(aq) + H₃O⁺(aq)

acid + base ó base + acid

• The reaction proceeds entirely to the left
• When comparing CH₃^- and H₂O (both are bases in this reaction), the CH₃^- is the stronger base
• When comparing CH₄ and H₃O⁺, the H₃O⁺ is the stronger acid

CH₄(g) + H₂O(l) ¬ CH₃^-(aq) + H₃O⁺(aq)

acid + base ó stronger base + stronger acid

Strong acids and bases are strong electrolytes and exist in solution entirely as ions.

• There are relatively few common strong acids and bases

The monoprotic (one proton) strong acids:

• HCl (hydrochloric acid)
• HBr (hydrobromic acid)
• HI (hydroiodic acid)
• HNO₃ (nitric acid)
• HClO₃ (chloric acid)
• HClO₄ (perchloric acid)

The diprotic (two protons) strong acid:

• H₂SO₄ (sulfuric acid)

These acids are completely ionized in H₂O(l). Thus, the reaction is represented with a single arrow in the direction of proton donation:

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

Or an equally valid equation for an aqueous solution of HCl:

HCl(aq) ® Cl^-(aq) + H⁺(aq)

The protons from an aqueous solution of a strong acid are typically in such vast excess to the natural ionization of water that their concentration determines the [H⁺] of an aqueous solution

There are few common strong bases. They are typically the ionic hydroxides of group 1A metals, and some of the heavier group 2A metals:

• LiOH, NaOH, KOH, etc.
• Ca(OH)₂, Sr(OH)₂, Ba(OH)₂

These molecules ionize completely in aqueous solution.

• The [OH^-] of an aqueous solution of a strong base is determined by the concentration of the base - it is in vast excess to the natural ionization of water

Strongly basic solutions can also be produced by certain substances that react with water to form OH^- ions

• Oxide ion (O^2-) is an example of this (in the form of metal oxides, e.g. CaO)

O^2-(aq) + H₂O(l) ® 2OH^-(aq)

CaO(aq) + H₂O(l) ® 2OH^-(aq) + Ca²⁺(aq)

• O^2- will react completely with H₂O to produce OH^- ion (i.e. it is a strong base)
• Other strong bases like O^2- include H^- and N^3-

Bimolecular compounds of the form X-H (i.e. group 7 hydrides) will have different acid strengths depending on the ease with which the proton can be released (i.e. the ease with which the X-H bond can be broken).

• The further the bonding electrons are from the nucleus, the less energy is required to ionize them completely.
• In this regard, as you move down the periodic table the atomic radii increases and valence electrons are further away from the nucleus.
• Thus, bond strengths generally decrease as you move down the periodic table, and it requires less energy to ionize a proton in an X-H compound.
• The easier it is to release a proton, the more acidic the compound. Thus, group 7 hydrides are stronger acids as you move down the periodic table:

Similarly, the electronegativity of the elements generally increases as you move to the right of the periodic table.

• For any two elements that are bonded together, it is easier to break the bond and ionize the two atoms if there is a large difference in electronegativities.
• The extreme case would be a metal halide - the large electronegativity difference means that they don't share bonding electrons, rather they ionize readily (since the metal wants to give up its valence electron and the halide wants to get one).
• Thus, the greater the electronegativity difference between a hydrogen and the other atom in a H-X bond, the easier to ionize and thus the stronger the acid. The electronegativity of H is about the same as carbon.
• So, as you move right in the periodic table beyond beyond carbon, the stronger the acid for X-H bonds (i.e. CH₄ < NH₃ < H₂O < HF as far as acid strength):

Another effect upon the ability to ionize an X-H bond is the effect of an electron-withdrawing group (i.e. a group with a large electronegativity) in the proximity of an X-H bond. This is most often observed for halides bound to various locations within carboxylic acid compounds.

• The closer they are bonded to the carboxylic acid ionizable proton (-O-H bond) the stronger the carboxylic acid. The reason for this is that their high electronegativity tends to delocalize electrons and withdraw them towards the halide.
• This effect can be distributed over the molecule and influence the electrons comprising the X-H bond.
• If they are withdrawn away from the H, then it is easier to ionize the H (and the acid will be stronger). Here are examples: