Chemical Equilibrium Lesson 6 AcidBase Theory Arrhenius Theory

Chemical Equilibrium Lesson # 6 Acid-Base Theory

Arrhenius Theory • Swedish chemist Svante Arrhenius performed experiments and he discovered that acids produce hydrogen ions (H+) in aqueous solution, and bases produced hydroxide (OH-) when they dissolved. This is known as the Arrhenius theory of acids and bases. – HCl (g) → – Na. OH (s) → • This worked well to explain the taste (sour versus bitter) and feel (sticky versus slippery) of acids and bases, but had its limitations. – Not all bases contain hydroxide – Not all acid-base reactions happen in aqueous (water) solution

The Brønsted-Lowry Theory • Danish chemist Johannes Brønsted and English chemist Thomas Lowry proposed that an acid is a hydrogen ion donor, and that a base is a hydrogen ion acceptor. Sometimes we say proton donor/acceptor, as a hydrogen ion lacks its only electron, which makes it only a proton. • This theory improved upon the Arrhenius theory as it did not limit bases and it did not reference any solution.

Brønsted-Lowry Acids • This theory also led to the discovery that ionization of acids and bases is not so simple. The hydrogen ion is not stable on its own, and appears to interact with water – the oxygen pulls the hydrogen ion towards it. In fact, it forms a new, more stable ion, called the hydronium ion. For example, when hydrogen fluoride is dissolved in water: – Arrhenius: HF (aq) → – Brønsted-Lowry: H 2 O (l) + HF (aq) ⇌ – In this reaction, hydrogen fluoride acts like an acid as it “donates” its hydrogen ion to water, the base, which “accepts” the hydrogen. This is a reversible reaction – hence the equilibrium unit!

Brønsted-Lowry Bases • Water can also act like a Brønsted-Lowry acid. It is amphiprotic (also sometimes called amphoteric) meaning it can be classified as both an acid and a base. • For example, when ammonia is dissolved in water: – NH 3 (g) + H 2 O (l) ⇌ – In this reaction, ammonia acts like a base as it “accepts” the hydrogen ion from water, which “donates” and hence acts like an acid. – Brønsted-Lowry bases generally contain at least one atom with one or more lone pairs of electrons (most often O, N, P)

Conjugate Acid-Base Pairs – Since acid-base reactions are reversible, there is always a proton transfer is each direction. That also means there is a Brønsted-Lowry acid and Brønsted-Lowry base on each side of the reaction. – Using the hypothetical reaction: HA (aq) + H 2 O (l) ⇌ A- (aq) + H 3 O+ (aq) • In the forward reaction, HA is the acid and H 2 O is the base. • In the reverse reaction, H 3 O+ is the acid (called a conjugate acid and A- is the base (called a conjugate base) • In general, a conjugate acid is what is formed when a base accepts a hydrogen ion, and a conjugate base is what is formed when an acid loses a hydrogen ion.

Conjugate Acid-Base Pairs – Using the hypothetical reaction: B (aq) + H 2 O (l) ⇌ HB+ (aq) + OH- (aq) • In this case, the conjugate acid is HB+ and the conjugate base is OH-. – In any acid-base reaction, there will always be one conjugate acid-base pair made up of an acid and its conjugate base, and another made up of a base and its conjugate acid.

Non-Aqueous Reactions Example: HCl (g) + NH 3 (g) ⇌ NH 4+ (aq) + Cl- (aq) (sometimes we write it as NH 4 Cl – a coordinate covalent bond) • Here HCl is the acid and Cl- is the conjugate base. • NH 3 is the base and NH 4+ is the conjugate acid.

The Acid and Base Ionization Constants, Ka and Kb • In the acidic reaction: HA (aq) + H 2 O (l) ⇌ A- (aq) + H 3 O+ (aq), Ka = • In the basic reaction: B (aq) + H 2 O (l) ⇌ HB+ (aq) + OH- (aq), Kb =

Competition for Protons – In HA (aq) + H 2 O (l) ⇌ A- (aq) + H 3 O+ (aq), there are two bases (H 2 O on the left, A- on the right) – Both bases compete for the hydrogen ion – if H 2 O has a greater affinity for hydrogen than A-, then the equilibrium position will move to the right – most of the dissolved acid will be A- (aq). – If A- is a much stronger base, than the equilibrium will move to the left – most dissolved acid will be as HA.

Strong vs. Weak Acids & Bases – A strong acid or base is one for which the equilibrium position is an aqueous solution lies far to the right. In an acid, that means all of HA forms the ion A-, and in a base, B forms HB+. • Ka/Kb is large • All group 1 and most group 2 bases (ex: Na. OH, Ca(OH)2) are strong bases.

Strong vs. Weak Acids & Bases – A weak acid or base is one for which the equilibrium position is far to the left, meaning most of acid is still in HA form, and most of the base is still in B form. Very little ionization occurs. • Ka/Kb is small • Most organic acids (carboxylic acids) are weak acids. In ethanoic acid, it is often written as HCH 3 CO 2. The H at the beginning ionizes, the others do not. – There is an important connection between the strength of an acid and the strength of its conjugate base (and vice versa): the stronger the acid, the weaker its conjugate base. Conversely, the weaker the acid, the stronger its conjugate base.