Chapter 15Acids and Bases.

*Acid-Base TheoriesIn defining what is considered to be an acid and what is considered to be a base, three theories have been proposed:

Arrhenius acid-base theoryBrnsted-Lowry acid-base theoryLewis acid-base theory

We will see that each subsequent theory builds upon what was stated in the previous theory.

*The Arrhenius TheoryIn the Arrhenius theory of acids, an acid dissolved in water increases the concentration of hydronium ions H3O+ in the solution:

Arrhenius acid HA:HA (aq) + H2O (l) H3O+ (aq) + A- (aq)

In this reaction we see all Arrhenius acids contain protons (H+) that are donated to water

*Arrhenius Theory Acid StrengthIn the Arrhenius theory of acids, a strong acid COMPLETELY reacts with water, so there is no HA left at the end of the reaction:

HA (aq) + H2O (l) H3O+ (aq) + A- (aq)

*Arrhenius Theory Acid StrengthIn the Arrhenius theory of acids, a weak acid reacts with water until an equilibrium is reached where HA is still present in the equilibrium mixture:HA (aq) + H2O (l) H3O+ (aq) + A- (aq)

*The Arrhenius TheoryIn the Arrhenius theory of bases, a base dissolved in water increases the concentration of hydroxide ions OH- in the solution:

Arrhenius base M(OH)x:

M(OH)x (aq) Mx+ (aq) + x OH- (aq)

In this reaction we see all Arrhenius bases contain hydroxide (OH-)

*Arrhenius Theory Base StrengthIn the Arrhenius theory of acids, a strong base COMPLETELY dissociates in water, so there is no M(OH)x left at the end of the reaction:

M(OH)x (aq) Mx+ (aq) + x OH- (aq)

*Arrhenius Theory Acid StrengthIn the Arrhenius theory of bases, a weak base only partially dissociates in water until an equilibrium is reached where M(OH)x is still present in the equilibrium mixture:M(OH)x (aq) Mx+ (aq) + x OH- (aq)

*Common strong acids and bases.

*Why do we need to improve on Arrhenius theory?The Arrhenius theory has a drawback!

Certain compounds that DO NOT contain hydroxide can still increase the hydroxide concentration when placed in water.Arrhenius theory does not explain this!

*The Brnsted-Lowry TheoryThe Brnsted-Lowry Theory: an acid Brnsted-Lowry Theory: an acid is any substance that donates protons (H+) while a base is any substance that can accept protons.

This means that Brnsted-Lowry acid-base reactions are proton transfer reactions.

*Proton transfer reactionsPairs of compounds are related to each other through Brnsted-Lowry acid-base reactions. These are conjugate acid-base pairs.

*Proton transfer reactionsGenerally, an acid HA has a conjugate base A- (an H+ has transferred away from the acid). Conversely, a base B has a conjugate acid BH+ (an H+ has transferred toward the base).

*Water in BL acid-base reactionsWhen a Brnsted-Lowry acid is placed in water, it donates a proton to the water (which acts as a base) and establishes an acid-base equilibrium.

*Water in BL acid-base reactions

In the reverse reaction of the equilibrium, the acid H3O+ donates a proton to the base A- to give back water and HA.

*Water in BL acid-base reactionsWhen a Brnsted-Lowry base is placed in water, it accepts a proton from water (which acts as an acid) and establishes an acid-base equilibrium.

*Water in BL acid-base reactions

In the reverse reaction of the equilibrium, the acid BH+ donates a proton to the base OH- to give back water and B.

*Brnsted-Lowry BasesTo accept a proton (to act as a B-L base) requires a molecule to have an unshared pair of electrons which can then be used to create a bond to the H+. All Brnsted-Lowry bases have at least one lone pair of electrons.

*Brnsted-Lowry BasesIn the previous reactions weve seen NH3 has a lone pair of electrons and can act as a B-L base. Also, water has two lone pairs of electrons, and can act as a B-L base.

*Amphiprotic substancesSome substances, like water, have protons that can be donated (BL acid), and lone pairs of electrons that can accept protons (BL base). This is why it can act like an acid OR a base DEPENDING on the other species present.

Such substances are said to beamphiprotic.

*ProblemWrite a balanced equation for the dissociation of each of the following Brnsted-Lowry acids in water:

a) H2SO4 b) HSO4-c) H3O+ d) NH4+

*ProblemWhat is the conjugate acid of each of the following Brnsted- Lowry bases?

a) HCO3-b) CO32-c) OH- d) H2PO4-

*Why do we need to improve on Brnsted-Lowry theory?There are many reactions that behave VERY MUCH LIKE proton transfer reactions that DO NOT involve protons!

*Lewis Acids and Bases

A Lewis acid is an electron pair acceptor, while a Lewis base is an electron pair donor.

These definitions are more general than the Brnsted-Lowry definitions because protons DO NOT need to be involved in Lewis acid-base reactions.

*In general LA + :LB LA-LBBF3 is a Lewis acid where the B atom can accept an electron pair. The N of the NH3 has a lone pair that can be donated, making NH3 a Lewis base.

*Metal ions as Lewis acidsMany metal ions have the ability to act as Lewis acids. The ions are willing to accept electron pairs from LIGANDS (which act as Lewis bases) because this often stabilizes the ion in solution. The result is often called a complex ion.

*Problem.

*Comparing theoriesSince Arrhenius acids must contain protons, then ALL Arrhenius acids ARE ALSO Brnsted-Lowry acids.

Weve already seen that NOT ALL Brnsted-Lowry bases are Arrhenius bases.

*Comparing theories

There are Lewis acids (like metal ions) that ARE NOT Brnsted-Lowry acids.

ALL Brnsted-Lowry bases must all have at least one lone pair of electrons, so ALL Brnsted-Lowry bases MUST ALSO BE Lewis bases

*BL acid and base strengthBrnsted-Lowry acid-base equilibria are competitions!

The equilibrium is the result of a tug-of war between the two bases in the system as they fight for protons given away by the two acids.

*Acid Strength and Base StrengthThe acid that is better at donating protons OR the base that is better at accepting protons will be found in lesser amounts at equilibrium compared to the other acid (or base).

*Strong BL acids in waterA strong acid (HA) is one that almost completely dissociates in water (which acts as a base). The conjugate base A- will be a very weak base.

*Strong BL acids in water

At equilibrium, there will be very little to no HA present in the system, and the concentration of A- will essentially be the same as the initial concentration of HA.

*Weak BL acids in waterA weak acid (HA) is one that partially dissociates in water (which acts as a base). The acid is not as good at donating protons to the water. The conjugate base (A-) will be a weak base. Overall

*Weak BL acids in water

At equilibrium, there will be some A- and H3O+ present in the system. However, the concentration of HA will still be significant at equilibrium.

*.

*.Notice that the strongest acids have the weakest conjugate bases, and the strongest bases have the weakest conjugate acids!

*Hydrated Protons and Hydronium IonsThe ultimate proton-donor is a proton itself!

In water there is no such thing as H+.

Often more than one water molecule will crowd around the proton to give hydrates with the formula H(H2O)n+ where n is 1 to 4.

*Hydrated Protons and Hydronium Ions

*Dissociation of Water

It is possible for one water molecule to act as an acid while another water molecule acts as a base at the same time. This leads to the self-ionization of water equilibrium:

H2O (l) + H2O (l) H3O+ (aq) + OH- (aq)

The equilibrium constant for this reaction is called theion-product constant for water, Kw.

Kw = [H3O+][OH-]

*At 25 C, Kw = 1.0 x 10-14

so [H3O+] = [OH-] = 1.0 x 10-7 mol/L

Relatively few water molecules are dissociated at equilibrium at room temperature!

We will always assume that [H3O+] [OH-] = 1.0 x 10-14 at 25 C.

*Acidic[H3O+] > 1.0 x 10-7 M or [OH-] < 1.0 x 10-7 MBasic[OH-] > 1.0 x 10-7 M or [H3O+] < 1.0 x 10-7 MNeutral[H3O+] = [OH-] = 1.0 x 10-7 M

At 25 C

*We also find, since

[H3O+] [OH-] = 1.0 x 10-14 = Kw

then

[H3O+] = 1.0 x 10-14 / [OH-] and [OH-] = 1.0 x 10-14 / [H3O+]

At 25 C

*.At 25 C

*ProblemThe concentration of OH- in a sample of seawater is 5.0 x 10-6 mol/L. Calculate the concentration of H3O+ ions, and classify the solution as acidic, neutral, or basic.

*ProblemAt 50 C the value of Kw is 5.5 x 10-14.What are the [H3O+] and [OH-]in a neutral solution at 50 C?

*The pH Scale[H3O+] in water can range from very small (strongly basic) to very large (strongly acidic) it is sometimes easier to use a negative logarithmic (power of 10) scale to express [H3O+] with a term we call the pH of a solution.

pH = - log [H3O+]

*pH and acidity

pH > 7 is basic pH < 7 is acidic pH = 7 is neutral

*pOH and aciditypOH < 7 is basic pOH > 7 is acidic pOH = 7 is neutralpOH = - log [OH-] Or [OH-] = 10-pOH

*pH scale

*.Kw = 1.0 x 10-14 = [H3O+] [OH-]

pKw = - log (1.0 x 10-14) = 14.00 (2 sigfigs! The 14 is not significant!)

pKw = - log ([H3O+] [OH-]) = (- log [H3O+]) + (- log [OH-]) = pH + pOH

so 14.00 = pH + pOH (at 25 C)!

*ProblemCalculate the pH of each of the following solutions:

a) A sample of seawater that has an OH- concentration of 1.58 x 10-6 mol/L

b) A sample of acid rain that has an H3O+ concentration of 6.0 x 10-5 mol/L

*ProblemCalculate [H3O+] and [OH-] in each of the following solutions:

a) Human blood (pH 7.40)b) A cola beverage (pH 2.8)

*Measuring pHWe often measure the pH of a solution with a chemical acid-base indicator. Indicators are B-L acids (symbolized HIn) where the acid form has a different colour than the conjugate base form (In-)

HIn (aq) + H2O (l) H3O+ (aq) + In- (aq)colour A colour B

*.Indicators tend to change colour only in small pH ranges of about 2 units.

*Measuring pHTo make a universal indicator that covers the pH range from about 1 to 12, a mixture of several different indicators with different pH ranges is used.

*.pH 1 2 3 4 5 6 7 8 9 10 11Methyl violet

Phenolphthalein

Bromothymol blue

Bromocresol green

Universal indicator

Methyl orange

*Universal indicator

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