Acid and Base Equilibrium

Acid and Base Equilibrium SCH4U0

Arrhenius Theory • In grade 11 we learned about the Arrhenius theory of acids and bases; • Arrhenius states that; • Acids are substances that dissociate in water, producing hydrogen ions • Bases are substances that dissociate in water, producing hydroxide ions

Arrhenius Neutralization • Arrhenius also concludes that acids and bases react in neutralization reactions; • The hydrogen ions combine with the hydroxide ions to make water Water Salt Base Acid

Problems with Arrhenius’ Theory • However, this theory has many problems and does not explain the behaviour of several compounds. • First, there are several compounds that do not contain hydroxide ions, but will produce them in aqueous solutions • Thus, the solutions they produce are basic (have high pH) • Ex: ammonia

Problems with Arrhenius’ Theory • Second, Arrhenius’ theory only applies to aqueous solutions • However, acids and bases will react with one another in other solvents or in other states like gas or solid • So compounds behave acidic/basic without having to dissociate • Third, Arrhenius’ theory states that hydrogen ions will dissociate and float freely in water • Though this will certainly not occur • Hydrogen ions are free protons and will react with the first negatively charged atom they touch

Problems with Arrhenius’ Theory • In water, every free proton will almost immediately collide with a water molecule • Which has very negatively charged oxygen atoms • These will undoubtedly bond • So there are no free protons floating around • Only hydronium ions hydronium

Problems with Arrhenius’ Theory • Lastly, Arrhenius can’t properly account for the fact that different acids create different pH, even when they have the same concentration • Eg: 1M HCl(aq) 1M CH3COOH(aq) pH= 0 pH= 2.37

Brønsted-Lowry Theory • Since the Arrhenius theory is flawed, a new theory has been developed (independently) by Johannes Nicolaus Brønsted and Thomas Martin Lowry in 1923 • This theory is much more general and applies to compounds in any solvent or state • It also addresses the fundamental question of what occurs during an acid-base reaction

Acid-Base Reactions • To understand the Brønsted-Lowry Theory, we must first answer the question of what fundamentally occurs during a neutralization reaction • What bonds have been made/broken • What particles have moved or been exchanged • So an acid-base reaction is essentially just the transfer of a proton from one compound to another A proton has been exchanged

Brønsted-Lowry Theory • This means that several reactions can be classified as acid-base reactions under this theory that would not have been under the Arrhenius Theory • Eg:

Acid/Base Labels • How can we classify acids/bases with this new theory? • Dissociation is not as important, only proton transfer • Acid: A substance that loses a proton in a chemical reaction. • Base: a substance that gains a proton in a chemical reaction. • We also label the compounds for the reverse reaction Base Base Acid Acid

Practice • Which of the following compounds are acids/bases in the reactions below? Base Base Acid Acid Acid Base Acid Base

Conjugate Pairs • It is often useful for us to discuss which product is created from a particular reactant • We use the term “conjugate” to do this • A conjugate pair are two compounds in a reaction (one a reactant, and the other a product) that are identical except for one proton • Meaning that the reactant became the product by losing/gaining a proton • Conjugates always have the opposite behaviour; if one is an acid, the other is a base

Practice • Which of the following substances are conjugate pairs in the reactions below?

Amphoteric Substances • One of the interesting things about the Brønsted-Lowry Theory is that it classifies acids/bases for specific reactions only • So some substances are acids in some reactions but bases in others • We call these substances “amphoteric” • Example: Water Base Acid

Dissociation • One of the other implications of the Brønsted-Lowry Theory is that we must re-define dissociation • At least for acids and bases • Acid dissociation typically suggests that a proton separates from the acid to float freely in solution • But we know that is not very likely • So we define dissociation as the reaction between a substance and water

Dissociation • An acid-dissociation is the dissociation reaction where the substance behaves like an acid • The base-dissociation is where it behaves like a base • Eg: • The acid dissociation of acetic acid • The base dissociation of ammonia

Dissociation Constants • Just like any other equilibrium, we can define an equilibrium constant for a dissociation equilibrium • For acid dissociation it is Ka • For base dissociation it is Kb • Example:

Practice • For each of the substances below, write an acid and base dissociation reaction and the associated Ka/Kb equation

Acid and Base Strength • The acid/base dissociation constants can be used to analyze how strong a particular acid/base is. • “Strength” refers to how much hydronium/hydroxide is present at equilibrium. • The larger the K value, the more it is shifted to the right at equilibrium • And the more is present

Acid and Base Strength • An acid or base is considered strong if the K value is larger than 1 • In this case we assume that the acid/base has dissociated 100% • Meaning there are no reactants left at equilibrium • Any acid or base that does not dissociate 100% (its K value is less than 1) is called weak • Keep in mind that weak does not equate to bad • An acid could dissociate 50% and be considered weak • Whereas another could dissociate 0.000001% and is also called weak

Strong Acids/Bases • There is a very short list of strong acids/bases • The vast majority of acids/bases not on this list are weak • Strong Acids: • Strong Bases: • Hydroxide containing ionic compounds • Though, some are fairly insoluble

Acid and Base Strength • What makes an acid/base strong or weak? • There are several factors, but the main one to consider is the forward/reverse reaction rate constants • The rate constants will be large/small depending on the activation energy • And how difficult it is to achieve the transition state • So what matters is how difficult it is to ionize the acid, or protonate the base

Acid Strength • For an acid, Ka is increased if; • The forward rate constant is increased; • This happens when the acidic hydrogen is very positively charged, thus attracting the negatively charged base (and lowering EA) • Example: • Electronegative elements pull electron density away from hydrogen • Resonance can pull electron density away from hydrogen

Acid Strength • For an acid, Ka is also increased if; • The reverse rate constant is decreased; • This happens when the conjugate base is not very negatively charged (and thus is weak) • If the conjugate base is small, the charge is more dense • If the conjugate base has resonance, the electron density is more diffuse

Base Strength • As we saw in the previous slide, base strength depends on the amount of charge on the basic atom • This can be reduced through resonance and size • This is also altered by the #of electrons (-2 ions vs -1, etc…) • Arrange the following bases in order of decreasing Kb (strongest to weakest) • Ammonia > Hydrazine due to electronegativity • Hydrazine > Urea due to resonance

Acid and Base Equilibrium