Bronsted-Lowry Acids and Bases

• Bronsted-Lowry focused on Acid and Base Equilibria and solubility Equilibria

Lecture Targets: Students should be able to:

• Define acids and bases based on Bronsted-Lowry theory
• Classify a compound as an acid or base based on Bronsted-Lowry theory
• Identify the conjugate base of an acid and the conjugate acid of a base
• Describe the hydrolytic reaction of acids and bases
• Predict if a conjugate will undergo hydrolysis
• Explain the acid-base behavior of amphoteric molecules

Ammonia (NH3):

• An important chemical with uses from agricultural practices to industrial processes
• Ammonia is not an Arrhenius base because does not have hydroxide ion (OH⁻) to give off
• Ammonia is not an Arrhenius acid because it does not give off H⁺ when dissolved in water
• pH paper shows that the pH of ammonia is around 11
• A new definition of base and acids is needed to account for this deviation from the Arrhenius definition

Johannes Nicolaus Brønsted:

• Born February 22, 1879, in Varde, Denmark, and died December 17, 1947, in Copenhagen
• A Danish physical chemist known for a widely applicable acid-base concept identical to that of Thomas Martin Lowry of England
• Both men introduced their definitions simultaneously in 1923, independently
• Brønsted was also an authority on the catalytic properties and strengths of acids and bases
• His chief interest was thermodynamic studies, but he also did important work with electrolyte solutions.

Thomas Martin Lowry:

• Born October 26, 1874 – November 2, 1936
• An English physical chemist
• Elected a Fellow of the Royal Society in 1914
• In 1920, he became the first holder of a chair of physical chemistry at Cambridge University
• In 1923, studied changes in optical rotation caused by acid- and base-catalyzed reactions of camphor derivatives, leading to formulation of the protonic definition of acids and bases, independently of Brønsted

Brønsted-Lowry Theory:

• Defines an acid as a substance that donates a H⁺ ion (or proton)
• Defines a base as a substance that accepts a H⁺ ion (or proton)
• NH₃(aq) + HCl(aq) → NH₄Cl(aq) can be used to show how H⁺ moves in the reaction
• From the Ammonia(NH₃) + Hydrochloric acid(HCl) reaction, HCl donated H⁺ to NH₃
• HCl is classified as an acid and NH₃ is classified as a base
• With this definition, even ions can now be classified into acids and bases
• Does not necessarily produce water

Brønsted-Lowry Theory: Acids

• Question: Are Arrhenius acids Brønsted-Lowry acids?
• An acid must contain H in its formula (e.g., HNO₃ and H₂PO₄ are two of many examples)
• All Arrhenius acids are Brønsted-Lowry acids

Brønsted-Lowry Theory: Bases

• Question: Are Arrhenius bases Brønsted-Lowry bases?
• Must contain a lone pair of electrons to bind the H⁺ ion; a few examples are NH₃, CO₃²⁻, and F⁻, as well as OH⁻ itself
• Brønsted-Lowry bases are not Arrhenius bases, but all Arrhenius bases contain the Brønsted-Lowry base OH⁻
• An example reaction: Benzoic acid (C₆H₅COOH) reacting with bicarbonate ions (HCO₃⁻) to form benzoate ions (C₆H₅COO⁻) and carbonic acid (H₂CO₃): C₆H₅COOH(aq) + HCO₃(aq) → C₆H₅COO(aq) + H₂CO₃(aq)

Hydrolysis:

• Reaction in which sulfide ions (S²⁻) and water (H₂O) react in equilibrium with hydrosulfide ion (HS⁻) and hydroxide ion (OH⁻): S²⁻(aq) + H₂O(l) ⇌ HS⁻(aq) + OH⁻(aq)
• Are reversible reactions involving water
• Another reaction that results in hydrolysis is cyanic acid reacting with water, to form cyanide and hydronium ions: HCN(aq) + H₂O(l) ⇌ CN⁻(aq) + H₃O⁺(aq)

Amphoterism:

• A phenomenon exhibited by water, where it can act as both an acid and a base depending on the reaction
• Amphoterism is not exclusive to water and can also be observed with some ions such as the bicarbonate ion (HCO₃⁻)
• H₂O donates the H⁺ and NH₃ accepts it.
• Water is amphiprotic because it acts as a base (accepts an H⁺) in one case and as an acid (donates an H⁺) in the other

Acid-Base Reactions (Bronsted-Lowry): Reactants and Products

• Brønsted-Lowry definition provides a new way to look at acid-base reactions
• H₂S and HS⁻ are a conjugate acid-base pair: HS⁻ is the conjugate base of the acid H₂S
• NH₃ and NH₄⁺ form a conjugate acid-base pair: NH₄⁺ is the conjugate acid of the base NH₃
• Every acid has a conjugate base, and every base has a conjugate acid

Conjugate Acid-Base Pairs:

• For any conjugate acid-base pair: acid₁ + base₂ ⇌ base₁ + acid₂
• The conjugate base has one fewer H and one more minus charge than the acid
• The conjugate acid has one more H and one fewer minus charge than the base

Conjugates Summary

• NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq): NH₃ is the base, H₂O is the acid, NH₄⁺ is the conjugate acid, and OH⁻ is the conjugate base
• C₆H₅COOH(aq) + HCO₃(aq) → C₆H₅COO(aq) + H₂CO₃(aq): C₆H₅COOH is the acid, HCO₃ is the base, C₆H₅COO⁻ is the conjugate base, and H₂CO₃ is the conjugate acid

Conjugates Classification

• Similar to acids and bases, conjugates may be classified as strong or weak
• Strong conjugates can undergo hydrolysis, whereas weak conjugates don't Strong conjugates arise from weak acids or bases
• On the other hand, weak conjugates originate from strong acids or bases

Strong Acids and Bases Information

• The dissociation of strong acids or bases is not reversible _ Every other acid or base not included in the table is weak
• Means that the strong acids and bases can undergo hydrolysis
• Common weak acids and bases are organic compounds such as acetic acid and methylamine
• The knowledge of strong acids and bases can be used to determine relative strengths of conjugates
• The conjugate base of a strong acid is weak
• The conjugate base of a weak acid is strong