Chemistry

Chemistry

Grade 12

Based on the Nelson Chemistry 12 textbook

Product Constant for Water, Kw

We can use the product constant for water, Kw , to calculate the hydrogen ion concentration or the hydroxide ion concentration in an aqueous solution of a strong or weak acid or base at SATP, if the other concentration is known.

Since Kw = [H+(aq)][OH-

(aq)] Then Kw = [H+

(aq)] [OH-

(aq)] And Kw = [OH-

(aq)] [H+

(aq)] In neutral solutions: [H+

(aq)] = [OH-(aq)]

In acidic solutions: [H+(aq)] > [OH-

(aq)] In basic solutions: [H+

(aq)] < [OH-(aq)]

Strong Acid

For strong acids we can use the concepts that strong acids ionize quantitatively in solution and use the value of Kw to calculate [H+(aq)]and [OH-(aq)] in acidic solutions.

EXAMPLE A 0.20 mol/L solution of hydrobromic acid, HBr(aq),

at SATP is found to have a hydrogen ion concentration of 0.20 mol/L. Calculate the [OH-(aq) ].

0.20 mol/L HBr (aq) H+

(aq) + Br -(aq)

0.20 mol/L 0.20 mol/L H2O(l) H+

(aq) + OH-(aq)

We can ignore the contribution of H+(aq) from autoionization

The OH-(aq) can also be ignore.

Therefore the major entities in the solution are: (which are the major entities affecting acid-base characteristics) H+

(aq) and Br -(aq)

but the Br -(aq) is the conjugate base of a strong acid and

therefore can be ignored as weak

So since HBr(aq) ionizes quantitatively (100%), [H+

(aq)] = 0.20 mol/L

Now we can use the Kw expression to calculate the concentration of OH-

(aq).

[OH-(aq)] = __Kw____

[H+(aq)]

[OH-(aq)] = 1.0 x 10-14/0.20 = 5.0 x 10-14 mol/L

Strong Base

Just as with acids, we can use the two concepts for solutions of strong bases – that they dissociate quantitatively in solution and the value of Kw to calculate the hydrogen ion concentration or hydroxide ion concentration.

EXAMPLE Calculate hydrogen ion concentration in a

0.20 mol/L solution of magnesium hydroxide, a strong base.

Mg(OH)2(aq) Mg2+

(aq) + 2OH-(aq)

0.20 mol/L 0.20 mol/L 2(0.20 mol/L) = 0.40 mol/L

Major entities: Mg2+(aq), 2OH-(aq), H2O(l)

pH – measure of the hydrogen ion concentration

pH = -log[H+(aq)] if [H+(aq)] = 2.5 x 10-14 pH = -log [H+(aq)]

= -log (2.5 x 10 -14)

pH = 13.60 In water and any neutral solution, the pH is

7.00 (show calculations, as on p. 541).

To build on the previous chart: In neutral solutions: [H+

(aq)] = [OH-(aq)] pH = 7.00

In acidic solutions: [H+(aq)] > [OH-

(aq)] pH < 7.00

In basic solutions: [H+(aq)] < [OH-

(aq)] pH > 7.00

pOH – measure of the hydroxide ion concentration

pOH = -log[OH-(aq)]

[OH-(aq)] = 10-pOH

pH + pOH = pKw = 14.00

EXAMPLE What is the pH of a solution whose pOH is 2.3? [OH-

(aq)] = 10-pOH

= 10-2.3 [OH-

(aq)] = 5.0 x 10-3

The pH of Strong Acids and Bases

The pH of Strong Acids The pH of solutions of strong monoprotic acids is

calculated from the concentration of H+(aq) ions, which is assumed to be the molar concentration of the solute molecule before ionization.

The pH of Strong Bases As with strong acids, the pOH and the pH of strong

bases are determined entirely by the OH-(aq) ion contributed by the dissociation of one of the ionic hydroxide solution.

Weak Acids

weak electrolyte does not ionize completely in water to form

hydrogen ions most common acids are weak acids: HF(aq),,

H2CO3(aq), H2S(aq), H3BO3(aq) A weak acid is an acid that partially ionizes in

solution but exists primarily in the form of molecules.

Weak Bases

Arrhenius theory of bases states that bases are soluble ionic hydroxides that dissociate in water into positive metal ions and negative hydroxide ions.

There are some molecular and ionic compounds, other than hydroxides, which also dissolve in water to produce basic solutions that are called weak bases as they are not as basic as ionic hydroxide solutions of the same concentration.

The Brønsted-Lowry definition of a weak base is a compound that reacts non-quantitatively (incompletely) with water to form an equilibrium that includes hydroxide ions according to the following general equation:

B(aq) + H2O(l) OH-(aq) + HB+(aq)

Percent Ionization of Weak Acids

most weak acids ionize less than 50%, unlike strong acids that ionize close to 100%

Percent Ionization (p) is defined as follows: p = conc. of acid ionized x 100% conc. of acid solute For a general acid ionization reaction: HA(aq) H+

(aq) + A-(aq)

p = [H+(aq)] x 100%

[HA(aq)] [H+

(aq)] = p x [HA(aq)] 100

If we know the pH of a weak solution, we can calculate the percent ionization of the acids.

Ionization Constants for Weak Acids

We can consider equilibrium solutions of a weak acid dissolved in water to be just like the equilibrium systems we looked at in Chapter 7

we can use the equilibrium law expression and calculate equilibrium constants

acid ionization constant: Ka

Ka is usually determined experimentally Percent ionization can be used to calculate the Ka

value (using an ICE table)

EXAMPLE

Calculate the acid ionization constant, Ka, of acetic acid if 0.1000 mol/L solution at equilibrium at SATP has a percent ionization of 1.3%.

HC2H3O2 (aq) H+ (aq) + C2H3O2

-(aq)

Ka = [H+(aq)][C2H3O2

-(aq)]

[HC2H3O2 (aq) ]

Create ice table and substitute the values back into the Ka expression to solve. Ka = 1.7 x 10-5

Ionization Constants for Weak Bases

Weak bases from dynamic equillibria in aqueous solutions

The reaction of weak bases with water may be defined by the equilibria law)

OH-(aq) ions are produced and affect the acid-base characteristics of solutions

Calculated the same way as weak acids

Relationship between Ka and Kb

Kw = KaKb

This equation allows us to convert the Ka values of acids into the Kb values of their conjugate bases, and vice versa, given the value of Kw, which is a constant.

Polyprotic AcidsPolyprotic Acids

Can anyone remember acids with more than one ionizable proton (more than one proton to give away)?

Exs: sulfuric acid, phosphoric acid, and boric acid Donate one proton at a time in a stepwise fashion Each ionization reaction has its own acid ionization

constant, Ka1, Ka2 … The acid in each step is weaker, generally:

Ka1 > Ka2 > Ka3

Acid-Base Properties of Salt SolutionsAcid-Base Properties of Salt Solutions

Salts

Solids Composed of cations and anions Dissolve in water May or may not alter the pH of a solution,

due to the cation, anion, or both

Salts that form Neutral Solutions

Salts of cations from strong bases and anions of strong acids have no effect of the pH of an aqueous solutionExs: NaCl(aq), KCl(aq), NaI(aq), NaNO3(aq)

Cation from Strong Base

Anion from Strong Acid

Salts

NaOH HCl NaCl

KOH HCl KCl

NaOH HI NaI

NaOH HNO3 NaNO3

Salts that form acidic solutions

The salt of a weak base (cation) and strong acid (anion) dissolves in water to form acidic solutions.

The cation reacts with water to liberate H+

The solution has a pH less than 7

NH4Cl dissociates NH4+ & Cl-

NH4+ + H2O H3O+

(aq) + NH3(aq)

Salts that form basic solutions

The salt of a strong base (cation) and weak acid (anion) dissolves in water to form basic solutions.

The anion reacts with water to liberate OH-

The solution has a pH greater than 7

NaC2H3O2 dissociates Na+ & C2H3O2-

C2H3O2- (aq)+ H2O HC2H3O2(aq) + OH-

(aq)

Salts that act as acids and bases

Some salts: contain both the cation of a weak base and the anion of a

weak acid both ions can hydrolyze we can predict whether the solution is acidic, basic or neutral:

Ka > Kb then the solution is acidic

Kb > Ka then the solution is basic

Acid-Base TitrationAcid-Base Titration

Start by matching titration terms

Totally Terrific Titration Terms

Titration: a chemical analysis involving the progressive addition of a solution of known solute concentration into a solution of unknown concentration to determine the amount of a specified chemical

Titrant: solution of known concentration, usually found in a buret during titration

Terms

sample: the solution of unknown concentration being analyzed in a titration

primary standard: chemical that is available in a pure and stable form that can be used to produce an accurate concentration

standard solution: a stock solution that is of known concentration that is used to create the titrant

Terms

endpoint: point in titration when the pH indicator changes colour

equivalence point: point in titration when chemically equivalent amounts of reactants have reacted (point at which equal amounts of H3O+(aq) and OH-(aq) have been added); generally, an equilibrium is established at this point

Terms

Standardization: a titration used to find the concentration of the titrant using a primary standard

indicator: an acid-base indicator that will change colour at a known pH to signify to signify a specific pH in the neutralization reaction

Titration of a strong acid with a strong base

At the equivalence point [H+]=[OH-] or pH=7. Phenolphthalein is a popular indicator because

it is colourless in acidic solutions and pink in basic.

Remember to consider the reaction at the molar level. (convert to moles!)– C=n/V and C1V1 = C2V2

Titration of a strong acid with a strong base

Titration curve – a plot of the pH vs. Volume of titrant added.

pH

Vol. of titrant

Titration Websites

Titrations– http://www.chemguide.co.uk/physical/acidbaseeqi

a/phcurves.html (curves & notes)

– http://www.chem.uoa.gr/applets/AppletTitration/Appl_Titration2.html (show different graphs)

– http://www.vias.org/simulations/simusoft_titration.html (simulation)

Buffers

Website:– http://www.chemguide.co.uk/physical/acidbaseeqi

a/buffers.html#top