ACID BASE EQUILIBRIA

ACID BASE EQUILIBRIACHAPTER 14 1ACID/ BASE DEFINITIONSARRHENIUS

BRONSTED-LOWRY

LEWIS21. ARRHENIUS

Most limited definition of acids and bases.

Acids supply H+ in aqueous solution.Bases supply OH-1 in aqueous solution.

Limited since many bases do not contain OH-1.

Ex. NH3

3 2. BRONSTED-LOWRY

More general definition.

Acids are proton (H+) donors.

When an acid loses a proton, it becomes a conjugate base.

Bases are proton acceptors.

Once the base accepts a proton, it becomes a conjugate acid.

General form of a BL acid base reaction:

HA + H2O A-1 + H3O+

4EXAMPLES:Formic acid dissociates (HCOOH)

Perchloric acid dissociates

Acetic acid dissociates

5Relationship between acids and their conjugates:

The stronger the acid, the weaker its conjugate base.

63. LEWIS

Most general definition.

Acids are electron pair acceptors.

Ex. BF3

Bases are electron pair donors. Ex. NH3

7ACID DISSOCIATION EQUILIBRIUM

Ka expressions can be written for an acid dissociation reaction.

Ka = [ products]power [reactants]power

Write the Ka expressions for the sample BLreactions:

8Higher Ka values, more dissociation, stronger acids.Strong acids do not have a Ka value s(extremely large) since the equilibrium lies so far to the right due to complete dissociation of a strong acid.The strong acids are:

SulfuricNitricPerchloricHydrochloricHydrobromicHydroiodic

Other Ka values are given in the appendix.

Take note of acetic acids Ka value and memorize it for the AP exam.

9WATERWaters Ka value is 1.0 x 10-14 How?

Consider two waters reactingthis is called auto-ionization of water.

What is the concentration of H+ and OH- in a sample of water.

10MEASURING ACID and BASE STRENGTH

pH scale ranges from 0-14

pH = - log [H+]

pOH = -log [OH-1]

pH + pOH = 14

11For water @ 25 C

pH = 7

[H+] = 10-pH

pOH = 7

[OH-1] = 10-pH

Kw = [H+][OH-1] and pKw = pH + pOH = 1412SolutionpHpOH[H+1] [OH-1]Acid/base/neutralA6.88B8.4 E -14C3.11D1.0 E -713ACID STRENGTH and CHEMICAL STRUCTURE

When a substance is dissolved in water, it may behave as an ACID, a BASE, or produce a NEUTRAL solution.

How does the chemical structure of a substance determine such behaviors?

14FACTORS AFFECTING ACID STRENGTH

Strength measured by:

Ka valueRecall Ka meaning.

pH

pH = -log [H+] measures how much an acid dissociates

151. BOND POLARITY

a proton is transferred only when the H is the positive pole of the compound

H + X

in hydrides, ex. NaH

the H is negatively charged, so no H+ could result.in non-polar molecule, ex. CH4

the electrons are not being pulled from the H, so no H+ could result.

162. BOND STRENGTH also affects acid strength

strong bond less likely to dissociate, weaker acid

weak bond stronger acid

173. CONJUGATE STRENGTH

stable conjugates come from a strong acid

18TYPES of ACIDS and STRENGTH TRENDS

Acids that can donate more than one proton are called POLYPROTIC acids

2 protons, ex. H2SO4 = diprotic

3 protons, ex. H3PO4 = triprotic

191. BINARY ACIDS made of H and one other element. General form H-X.H-X bond strength determines strength of the acid.

Bond strength DECREASES as the size of X increases.

Changes more drastically down a group.

20Going across a period, look at ELECTRONEGATIVITY.

IVVVIVII2CH4NH3H2OHF3SiH4PH3H2SHCl212. OXYACIDS

contain OH groups bound to a central atom.

OH groups acting as bases:

When OH is bound to a group with extremely low ELECTRONEGATIVITY.

Ex. metals

Ca(OH)2KOH

22OH groups acting as acids:

Bound to a NON-METAL.

Does not readily lose OH.

As the ELECTRONEGATIVITY of Y increases, so will the acidity.

The O-H bond becomes more POLAR and loss of H+ is favored.

When an additional OXYGEN is bound to the central Y, further increases the POLARITY of the OH bond favoring loss of the H+.

23Rules for comparing oxyacid strength:

1. for oxyacids that have the same number of OH groups and the same number of O atoms, acid strength increases with increasing electronegativity of the central atom

2. for oxyacids with the same central atom, acid strength increases as the number of oxygens attached to Y increases.

24Example:Place the following oxyacids in order of increasing strength:HClOHClO2HClO4HClO3HBrO

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3. CARBOXYLLIC ACIDScontain a carboxyl group

additional oxygens attached to the carboxyl group draws electron density from the OH group, increasing its polarity.

strength of the carboxylic acid increases when the number of electronegative atoms increases.

26Example:CH3OH is this an acid?

CH3COOH what acid is this? How does its strength compare to CH3OH?

27MORE PRACTICEArrange the compounds in each of the following series in order of increasing acid strength:AsH3HINaHH2O

28Arrange the compounds in each of the following series in order of increasing acid strength:H2SeO3H2SeO4H2O

29Explain whya. HCl is stronger than H2S as an acid.

b. benzoic acid (C6H5COOH) is a stronger acid than phenol (C6H5OH)

c. H2SO4 is stronger than HSO4-1

30Calculating pH of strong acidsLarge Ka, so products are favored.Assume the [acid]0 = [ H+ ] due to complete dissociation.

What else contributes [ H+ ] in solution?

Auto- ionization of water

Assumed to be negligible for strong acids (unless they are dilute >10-6 M)

So, pH = - log [acid]o

31pH example (strong)Calculate the pH and [OH-1] of a 5 x 10_3 M solution of HClO4.32pH of weak acid solutionsKa is small

Consider all sources of H+ in solution

ICE table

Ka expression

Solve, approximate whenever possible (check > Ka2 > Ka3 , so many times it is possible to ignore the contribution of H+ from the second (and third)dissociation.

Sulfuric acid is the exception to this rule!!!44pH example (polyprotic acid)Calculate the pH, [PO4-3] and [OH-] of a 6.0 M phosphoric acid solution.

1. Write all relevant dissociation reactions and find Ka values for each.452. ICE, assumptions, Ka1

3. Ka2 to find [HPO4-2]464. Ka3 to find [PO4-3]

5. [OH-], pOH, pH47ACID/ BASE properties of SALTSIonic compounds dissociate in water and the resulting ions can make the solution acidic, basic, or neutral.

Na+ and other alkali and alkaline earth metals do NOT exhibit acid or base properties.What about Na (s)?Conjugate of strong acids or bases do NOT exhibit acid or base properties.

Kw = Ka x Kb = 1x10 -14If Ka> Kb, acidicIf Kb> Ka, basic

48Salt examplesPredict whether each of the following will create an acid, base, or neutral aqueous solution.Na3PO4

KI

NH4F49pH example (of a salt) Calculate the pH of a 0.500 M NaNO2 solution (Ka = 4.0 x 10-4)Reaction

Kb

ICE50ACID-BASE PROPERTIES of SALT SOLUTIONS

When most salts are dissolved in water they will form acidic, basic, or neutral solutions.

When both the cation and the anion have an effect, the ion with the weakest conjugate acid or base will have the greatest influence on pH.

51ANIONS

Reaction of anion in water:A-1 + H2O HA + OH-1

If a strong acid is predicted as a product, the equilibrium will lie to the left. No considerable acidic/ basic effects will be noted.

If a weak acid is predicted as a product, the equilibrium will lie to the right. Due to the production of hydroxide ion, the pH of the solution will increase (become more basic).

52CATIONS

Polyatomic cations are considered to be conjugates of weak bases. Ex. NH4+ + H2O hydroxide ions are formed and the pH is increased.

Most metal ions can also react with water to decrease the pH of an aqueous solution.

Some exceptions include the alkali metals and some alkaline earth metals (usually produce bases in H20).

Metal ions are considered Lewis acids.

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LEWIS DEFINITION of ACIDS and BASES

Lewis base electron pair donor

Ex. NH3

Lewis acid electron pair acceptor

54Metal ionsPositively charged metal ions attract the oxygens from the water molecules.This reaction is known as hydration.usually coordinate with # waters = 2 x metal ions charge

Example: Fe+3 in water: (coordinates with 6 water molecules)

Acid dissociation constants for hydrolysis reactions incerasewith increasing charge and decreasing radius on the metal ion.

Ex. Which would form a more acidic solution Cu+2 or Fe+3?

55COMMON ION EFFECT (add on)Adding a salt to an acid base equilibrium that contributes to the initial conjugate concentration.Ex. What is the pH of a solution that is 0.5M in acetic acid and 2.5M in sodium acetate, NaCH3COO? Consider both dissociation rxns. What ions does the salt contribute to the acids equlibrium? (common ion) What does LeChatelier predict? Would the acid pH increase or decrease due to the common ion?BUFFERS

Solutions that contain weak conjugate acid-base pairs.

Resist changes in pH when small amounts of strong acid or base are added.

Examples of common buffer systems:BloodHydrogen carbonate/ carbonic acid

Contain an acid species to neutralize hydroxide and a basic species to neutralize hydrogen ions.

57WEAK ACID BUFFER

HA H+ + A-

With equilibrium expression:

Ka = [H+] [A-] / [HA]Adding OH-1 to the solution:OH-1 + HA H2O + A-HA = weak acid in the buffer used to absorb excess base

Adding H+ to the solution:H++A- HA A- = conjugate base of the weak acid in the buffer used to absorb excess acid58BUFFER CAPACITY and PH

Buffer capacity the amount of acid or base a buffer can neutralize before pH changes to a considerable degree.

PH depends on the Ka for the acid and the relative concentrations of the acid and base that comprise the buffer.

Calcualting pH and relating to KaUses the Henderson - Hasselbalch equation.

pH = pKa + log ([A-] / [HA])

59Henderson Hasselbalch (*hoff)Because weak acids and bases only slightly ionize, it is possible to use the initial amounts of acid and base when using the Henderson Hasselbalch relationship.

pH = pKa + log ([A-] / [ HA])

Example: Calculate the pH of a buffer composed of 0.12 M lactic acid (HC3H3O3) and 0.10 M sodium lactate. Ka for lactic acid = 1.4 x 10-4.

60Example: What is the ratio of HCO3-1 to H2CO3 in blood of pH 7.4?

61Weak base/ conjugate acid bufferpOH = pKb + log ([BH+]/[B])

ADDITION of A STRONG ACID or BASE to BUFFER SYSTEM

Reactions between STRONG ACIDS and WEAK BASES or a STRONG BASE and WEAK ACID occur essentially to completion. So the buffer should completely consume the additional acid or base.

Calculating pH after the addition of an acid or a base:1. consider the acid-base neutralization reaction and determine its effect on [HA] and [A-]. (Use stoichiometry).2. Use Ka and the new concentration of [HA] and [A-] to calculate [H+] using an ICE table or the Henderson Hasselbalch equation.

63H+ + A- HAOH- + HA H2O + A-Stoichiometry determines new concentrations of the conjugatesadding H+ or OH- to HA, A- at equilibriumpKa + log of the ratio of the [conjugate][original acid or base]Neutralization of excess acidNeutralization of excess base64Example: A buffer is made by adding 0.300 mole acetic acid to 0.3000 mole sodium acetate to make a 1.00L solution. The pH of the buffer is 4.74.calculate the pH of the solution after 0.020 moles of NaOH is added

65calculate the pH of the solution after .020 moles of HCl is added.

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