Topic 5A: acid-base equilibria

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Brønsted–Lowry acid: a proton donor; a molecule or ion that donates a hydrogen ion in a reaction.

Brønsted–Lowry base: a proton acceptor; a molecule or ion that accepts a hydrogen ion in a reaction.

conjugate acid-base pair: the pair of the acid and its conjugate base (the form of the acid after it's given up a proton

Acids and bases must always react in pairs, because when an acid gives up a proton there must necessarily be a substance accepting that proton, thereby acting as a base.

acid + base ⇌ conjugate base + conjugate acid

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The pH scale is based on the autoionization process that water undergoes. This is an equilibrium process in which water dissociates into its composite ions, H+ ions and OH- ions.

pH Is defined as the power of the hydrogen ion and is used to indicate the concentration of hydrogen ions in a solution.

the pH = − log [H+] and that pOH = − log [OH−].

Ka, the acid dissociation constant, and Kb, the base dissociation constant are measures of the extent of dissociation of a weak acid or base in equilibrium.

pKa and pKb are the logarithmic scales of Ka and Kb; they are the inverses of each other in terms of value.

pKa + pKb= 14

Kw= Ka x Kb

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Equivalent mass: the mass that yields 1 mole of hydrogen ions in the case of an acid; In the case of a base, the mass of the base that reacts with 1 mole of hydrogen ions.

Acid: a compound that produces H+ ions in water solution. Can be strong or weak.

Aqueous solutions of acids have the following characteristics:

sour taste

changes color litmus from blue to red

react with certain metals to produce gaseous H2

react with bases to form a salt and water

have a pH less than 7

common examples include acetic acid, sulfuric acid, tartaric acid

React with carbonates to create a salt, water, and carbon dioxide

Definitions of acids:

Arrhenius acid: any substances that increase the concentration of hydronium ions in solution

Bronsted-Lowry acid: any substance that can act as a proton donor

Lewis acid: any substance that can accept a pair of electrons

Strong acid: an acid that dissociates fully into two ions in water. These are HCl (hydrochloric acid), HBr (hydrobromic acid), HI (hydroiodic acid), HNO3 (nitric acid), H2SO4 (sulfuric acid; only the first proton is strong), and HClO4 (perchloric acid).

Strong acids are presumed to not involve an equilibrium process because they dissociate completely and are not likely to go back to their acidic form.

HA --> A- + H+

Acidity: measure the concentration of H+ ions in a dilute water solution, measured by the pH scale

The overall strength of an acid is determined by the polarity of the molecule, and the strength of the hydrogen-ion bond.

The more polar the molecule, the greater the electron density within the molecule is drawn away from the proton. This results in a greater partial positive charge on the proton, thereby resulting in a weaker H-A bond, ultimately allowing the proton to dissociate in solution more readily.

The stability of the conjugate base is an indicator of the strength of the acid. Stronger acids have the larger Ka and a more negative pKa than weaker acids.

Base: compound that produces hydroxide ions in water solution. Can be strong or weak.

Definitions of bases:

Arrhenius base: any compound that donates a hydroxide ion in solution

Bronsted-Lowry base: any compound capable of accepting a proton

Lewis base: any compound capable of donating an electron pair

Strong base: fully dissociates in water. These include hydroxides of groups 1A and 2A (LiOH, NaOH, KOH, and Ba(OH)2).

A strong base is a compound that can readily deprotonate from other compounds.

Strong bases are presumed to not involve an equilibrium process because they dissociate completely and are not likely to go back to their basic form.

For bases like Ba(OH)2, the base concentration is 1/2 [OH-]

Common examples include alkaline metal and some alkaline earth metal hydroxides (sodium hydroxide (NaOH), calcium hydroxide (Ca(OH)2), lithium hydroxide (LiOH), barium hydroxide (Ba(OH)2).

There are many weak bases, and many of them are analogs of ammonia.

Amphiprotic compounds: compounds that can serve as an acid or base, such as water.

Polyprotic acids: acids that can lose greater than one proton, therefore resulting in more than one Ka.

Salts of polyprotic acids can act as weak conjugate bases. These salts have their own Kb for each proton that they can accept.

Acidic salts are formed when a strong acid reacts with a weak base.

Basic salts are formed when a strong base reacts with a weak acid.

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weak acids, those that dissociate only to a very small extent; Examples: HF, HCN, H2CO3, H3PO4, acetic acid, and oxalic acid.

the pH of weak acids can be calculated using the equilibrium constant for acids as follows -->

The Ka and the hydrogen ion concentration are directly related.

The majority of acids are weak.

Only ~ 1% of a weak acid solution typically dissociates in water. Therefore the concentration of hydrogen ions in a weak acid solution is always less than the concentration of the undissociated species.

The strength of a weak acid is represented as either an equilibrium constant or a percent dissociation.

In order to calculate the pH of a weak acid using the Ka value, you must assume that [H+] and [A-] are equal. You must also presume that the concentration of the acid is the given concentration due to such low levels of dissociation (i.e the concentration has not decreased due to dissociation).

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Buffer: A solution of weak acid and its conjugate base, or weak base and its conjugate acid, used to resist pH change from an added solute.

Usually composed of approximately equal concentrations of an acid and its conjugate base. Buffers are able to do this because they contain acidic components to neutralize OH- ions and basic component neutralize H+.

They have many chemical applications. Even blood in the human body is a buffer solution. This keeps proteins from denaturing. Buffer solutions are also useful in fermentation processes and in dyeing fabrics.

This resistance to change in pH arises from the presence of an equilibrium between the acid and its conjugate base. In accordance with Le Chatelier’s principle, the equilibrium shifts towards the products (H+, A-) when a base is added, and towards the reactant (HA) when an acid is added. This is with the presumption that the buffer solution is made of a weak acid.

A buffer solution can be created by determining the exact amounts of acid and conjugate base needed to have a solution of a certain pH. This is determined using the Henderson-Hasselbach equation. It is also useful for finding the equilibrium pH in and acid- base reaction.

To make an alkaline buffer, The ratio of base to acid is set up on the basis of the desired pH.

pH can be calculated using an ICE table -->

Titration curves demonstrate the impact of a buffer on a solution and how it resists changes in pH.

Buffers affect the titration of a strong base and weak acid, or strong acid and weak base.

On the titration curves of both of these situations, in the beginning the pH changes very gradually, indicating the formation of a buffer system as the titration approaches the equivalence point.

In the titration of a strong base and a weak acid, the pH at the equivalence point does not equal 7 due to the production of the conjugate base during the titration.

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Acid-base reactions:

A neutral ionic compound is formed when a strong acid reacts with a strong base; The cation is taken from the base and an iron from the acid.

Strong acid + Strong base → Neutral salt pH = 7

HCl + KOH → KCl + H2O

An acidic ionic compound is formed when a strong acid reacts with a weak base.

Strong acid + Weak base → Acidic salt pH < 7

HCl + NH4OH → NH4Cl + H2O

A basic ionic compound is formed when a weak acid reacts with a strong base.

Weak acid + Strong base → Basic salt pH > 7

HF + NaOH → NaF + H2O