Bronsted-Lowry Acid – Base Reactions Chemistry

Bronsted – Lowry Acid Defined as a molecule or ion that is a hydrogen ion donor Defined as a molecule or ion that is a hydrogen ion donor Also known as a proton donor because H + is a proton. Also known as a proton donor because H + is a proton. Examples Examples HCl – a “monoprotic” acid HCl – a “monoprotic” acid H 2 SO 4 – a “diprotic” acid H 2 SO 4 – a “diprotic” acid H 3 PO 4 – a “triprotic” acid H 3 PO 4 – a “triprotic” acid These all dissociate in water to produce H + ions which can then react with a Bronsted-Lowry Base These all dissociate in water to produce H + ions which can then react with a Bronsted-Lowry Base

Bronsted-Lowry Base Defined as a hydrogen ion acceptor. In an acid-base reaction the base “accepts” the hydrogen ion from the acid. In an acid-base reaction the base “accepts” the hydrogen ion from the acid. NH 3 + H +  NH 4 + NH 3 + H +  NH 4 + NH 3 is Bronsted-Lowry Base NH 3 is Bronsted-Lowry Base

Conjugate Acid The compound that is formed when the B/L base gains a proton (H + ). The compound that is formed when the B/L base gains a proton (H + ). NH 4 + is the conjugate acid of NH 3 ( a base) NH 4 + is the conjugate acid of NH 3 ( a base) NH 4 + acts as an acid in the reverse reaction. NH 4 + acts as an acid in the reverse reaction.

Conjugate Base The compound that is formed when a B/L acid gives away it’s proton (H+) The compound that is formed when a B/L acid gives away it’s proton (H+) Cl- is the conjugate base of HCl (an acid) Cl- is the conjugate base of HCl (an acid) Cl- acts as a base in the reverse reaction Cl- acts as a base in the reverse reaction

The conjugate acid and bases are the acid and bases for the reverse reaction!

Amphoteric - Definition: A substance that can act as an acid (proton donor) or a base (proton acceptor) Definition: A substance that can act as an acid (proton donor) or a base (proton acceptor) Water (H 2 O) is amphoteric. Water (H 2 O) is amphoteric. HBr + H 2 O  Br - + H 3 O + HBr + H 2 O  Br - + H 3 O + NH 3 + H 2 O  NH OH - NH 3 + H 2 O  NH OH - Water acts as a base! Water acts as an acid!

Strong vs. Weak Acids and Bases Strong - Dissociate 100% Strong - Dissociate 100% The following are the strong acids, all the rest are considered weak. (Memorize!) The following are the strong acids, all the rest are considered weak. (Memorize!) HCl – Hydrochloric acid HCl – Hydrochloric acid HBr – Hydrobromic acid HBr – Hydrobromic acid HI – Hydroiodic acid HI – Hydroiodic acid HNO 3 – nitric acid HNO 3 – nitric acid H 2 SO 4 – sulfuric acid H 2 SO 4 – sulfuric acid HClO 4 – perchloric acid HClO 4 – perchloric acid

Strong Bases Strong bases are soluble hydroxides such as NaOH, KOH, and LiOH Strong bases are soluble hydroxides such as NaOH, KOH, and LiOH All the rest are considered weak. All the rest are considered weak.

Weak – Only dissociate < 10%, They are equilibrium reactions. Weak – Only dissociate < 10%, They are equilibrium reactions. We use Equilibrium constants to compare weak acid strength. We use Equilibrium constants to compare weak acid strength. Strong Acids and Bases have very large equilibrium constants Keq Strong Acids and Bases have very large equilibrium constants Keq Weak Acids and Bases have small equilibrium constants (<< 1) Weak Acids and Bases have small equilibrium constants (<< 1)

The equilibrium constant for an acid is called Ka HA H + + A - HA H + + A -

The equilibrium constant for a base is called the Kb NH 3 + H 2 O NH OH - NH 3 + H 2 O NH OH -

The larger the Ka value – the stronger the acid!

Examples AcidConjugate Base AcidConjugate Base HNO 3 _____________ HNO 3 _____________ HCO 3 -1 _____________ HCO 3 -1 _____________ HPO 4 2- _____________ HPO 4 2- _____________ NO 3 -1 CO 3 2- PO 4 3- *Conjugate base is what is left after acid donates it’s H + ion

BaseConjugate Acid BaseConjugate Acid HCO 3 -1 ______________ HCO 3 -1 ______________ F -1 ______________ F -1 ______________ H 2 O ______________ H 2 O ______________ HPO 4 2- ______________ HPO 4 2- ______________ H 2 CO 3 HF H3O+H3O+ H 2 PO 4 1- * Conjugate acid is what you get after the base accepts the H + ion.

Ka’s are used to predict which species will act as an acid and which as a base in an acid base reaction. It also allows you to predict which way an equilibrium reaction is favored. (with the products or reactants) Ka’s are used to predict which species will act as an acid and which as a base in an acid base reaction. It also allows you to predict which way an equilibrium reaction is favored. (with the products or reactants)Remember: The reactant with the higher Ka acts as the acid. The reactant with the higher Ka acts as the acid. Compare the acid and the conjugate acid Ka’s, the one with the higher Ka wants to dissociate more, so the equilibrium will favor the reaction direction that allows it to dissociate! Compare the acid and the conjugate acid Ka’s, the one with the higher Ka wants to dissociate more, so the equilibrium will favor the reaction direction that allows it to dissociate!

Example 1 NH 3 + H 2 O NH 3 + H 2 O Compare Ka values to determine which acts as an acid. Higher Ka value on chart is acid! Compare Ka values to determine which acts as an acid. Higher Ka value on chart is acid! Write the products (remove H+ from acid and add it to the base) Write the products (remove H+ from acid and add it to the base) Compare the Ka values of the acid and conjugate acid. Compare the Ka values of the acid and conjugate acid. Higher Ka wants to dissociate more, so the equilibrium will be favored away from that acid. Higher Ka wants to dissociate more, so the equilibrium will be favored away from that acid. NH 4 + has higher Ka so it wants to dissociate more, so Reactant side is favored! NH 4 + has higher Ka so it wants to dissociate more, so Reactant side is favored! NH OH -

Example 2 HCO HPO 4 2- HCO HPO 4 2-

Example 3 HS- + H 2 CO 3 HS- + H 2 CO 3