May 09, 2013

Acid Base ChemistryAcid Base ChemistryChapter 19

Properties of Acids• Taste sour• Are strong or weak electrolytes• React with bases to form water and salts• React with active metals to produce H2

• Turn litmus (and cabbage) red• Low pH

Acid Rxns with Metals1. Al, Mg and Zn form hydrogen gas

2Al + 6HCl 2AlCl3 + 3H22. Metal carbonates form carbon dioxideCaCO3 + 2HCl CaCl2 + H2O + CO2

Naming Acids Review

Anion Ending Example Acid Name Example

-ide chloride hydro-(stem)-ic

-ite sulfite (stem)-ous

-ate nitrate (stem)-ic

Properties of Bases

• Taste bitter• Are strong or weak electrolytes• React with acids to form water and salts• Feel slippery• Turn litmus (and cabbage) blue• High pH

Three Acid and Base Theories1. Arrhenius Theory

2. Bronsted-Lowry Theory

3. Lewis Theory

1) Arrhenius Theory

Simplest definition and most restrictive

• Acids ionized to produce H+ and• Bases ionized to produce hydroxide

ions OH-

May 09, 2013

Examples of ArrheniusAn acid in water ionizing to form a H+ ion

HCl(aq) H+(aq) + Cl-(aq)

A base in water dissociating to form hydroxide ion, OH-

NaOH Na+(aq) + OH-

(aq)

2) Bronsted-Lowry TheoryAdded to Arrhenius definition of bases

Bronsted-Lowry acid: molecule or ion that is a proton donor

Bronsted-Lowry base: molecule or ion that is a proton acceptor.

Bases don't necessarily have to supply OH-

NH3, ammonia, can now be recognized as a base

Bronsted-Lowry Theory: General acid/base reaction

Acid + Base Conjugate base + Conjugate Acid

Conjugate base: particle that remains after a proton that is released by the acid.

Conjugate acid: particle that remains after a base has acquired a proton from the acid.

ExamplesIdentify the conjugate pairs.

a) H2SO4(aq) + H2O(l) HSO4-(aq) + H3O+(aq)

b) H2O(l) + F-(aq) OH-(aq) + HF(aq)

3) Lewis TheoryThe broadest definition of acids and bases

Definitions:Acid: electron-pair acceptor • Doesn't necessarily have to supply (H+) or be a proton

donorBase: electron-pair donor

• Doesn't necessarily have to supply (OH-) or be a proton acceptor

May 09, 2013

ExampleBF3(g) + NH3(g) F3BNH3(g)

Type Acid Base

Arrhenius H+ or H3O+ producer

OH- producer

Bronsted-Lowry proton (H+) donor proton (H+) acceptor

Lewis electron-pair acceptor electron-pair donor

Amphoteric CompoundsAny substance that can behave as an acid or as a base.

The most common example is water.

Monoprotic and Polyprotic AcidsMonoprotic: an acid that can only donate one proton per molecule.

Polyprotic: is an acid that can donate more than one proton per molecule.

(Diprotic can donate two protons.)

(Triprotic can donate three protons.)

Ionization of polyprotic acidsProtons can only be donated one at a time to water. As a result each hydrogen ion will leave the acid in its own reaction step. The first proton leaves in the first ionization, the second proton is removed by the second ionization, etc.)

• One proton=one reaction step• Two protons=two reaction steps• Three protons=three reaction steps

Monoprotic, Diprotic and Triprotic

HCl

H2SO4

An acid that contains "x" number of hydrogens/molecule able to dissociate in water.

May 09, 2013

Strong Acids and Bases• Strength in acids and bases does NOT refer to how

corrosive or dangerous they are• Strong acids and bases dissociate (break apart)

completely in water.• At equilibrium, there is no acid or base left that has not

ionized (separated into its ions).• In contrast, one of the most corrosive and dangerous

acids, HF, is actually a weak acid.

Acid Base Strength• Strong acids and bases will dissociate completely in

aqueous solution.• Reaction equations will show a one way arrow and there will

be no equilibrium• Good conductors of electricity

• Weak acids and bases do no ionize completely in aqeous solution.

• Reaction will show a two way arrow and equilibrium will be established.

• Not good conductors of electricity.

Acid Base Concentration• Concentration refers to the amount of solute in a given

amount of solution

• Concentrated Acid-more solute in a given amount of solution than another solution of the same volume

• Dilute Acid-less solute in a given amount of solution than another solution of the same volume.

Equilibrium Constants• Law of Chemical Equilibrium-at a given temperature, at

equilibrium, a ratio of reactant to product concentrations has a constant value.

• Equilibrium constants (Keq)- the numeric value that compares reactant and product concentrations

• Equilibrium constant expressions can only be written for reactions at equilibrium and solids and liquids are always excluded from the equilibrium constant expression.

May 09, 2013

Equilibrium Constants• If the value of Keq is <1, there is more reactant at

equilibrium• If the value of Keq is >1, there is more product at

equilibrium.• Equilibrium Constant Expression: aA + bB cC + dD

> A, B - reactants> C, D - products> a, b, c, d- coefficients of the balanced equation

Example2H2S (g) 2H2(g) + S2(g)

• Write the equilibrium constant expression.

• If [H2S]= 0.184M, [H2]=0.0377M and [S2]=0.0540M, calculate the value of Keq.

• Does the reaction favor the production of reactants or products?

Acid Ionization Constant (Ka)• Ka is the value of the equilibrium constant expression for

the ionization of a weak acid.

• The base ionization constant (Kb) is the value of the equilibrium constant expression for a weak base.

• Large value of Ka or Kb=stronger acid or base• Small value of Ka or Kb=weaker acid or base

Ionization ConstantsKa: Quantitative expression for the strength of an acid

Kb: Quantitative expression for the strength of a base

Kw: Ionization constant for water

Kw=[H3O+][OH-]=1.0Χ10-14

Ion Product Constant for Water (Kw)We can relate the three values to find either the Ka or Kb value.

Kw=KaKb

Kw=[H+][OH-]

May 09, 2013

Self Ionization of Water

Hydronium Ion• The H+ is so small it immediately

attaches to one of the unshared pair of electrons on a water molecule forming H3O+

• Studies have shown that hydrogen "bonds" form with other water molecules to form even larger complexes.

• As far as we're concerned, we'll symbolize the hydronium ion

> H+ or H3O+

> We will use these interchangeably

Example: at 298K an aqueous solution has a [H+] of 1.0 x 10-5 M. What is the [OH-] concentration?

• pH and pOH calculations

May 09, 2013

pH values into perspective• Because the pH scale is

a log scale, changes by one pH unit are changes by a factor of 10.

• Assume that a pH=1 solution, which [H+]=1x10-1 is similar in size as the length of a sports field.

• As pH increases, the [H+] decreases by the same factor as the comparative scale.

pH [H+] Comparison1 1 x 10-1 100 yards2 1 x 10-2 10 yards3 1 x 10-3 1 yard4 1 x 10-4 3.6 inches5 1 x 10-5 0.36 inches6 1 x 10-6 0.036 inches

7 1 x 10-7 0.0036 inches

pH• pH is a number scale from 0-14 that represents the

acidity or basicity of a solution• pH is a measure of the hydrogen/hydronium ion

concentration of a solution

• pH=-log[H+]• the lowercase p stands for power

pH + pOH=14[H+][OH-]=1x10-14

pH and pOH table with H+ and OH- values

pH OH-H+ pOH

1x10-7

1x10-5

1x10-5

6

Example pH1. Find the pH of a solution with a 7.01Χ10-6 M H3O+ concentration.

• [H3O+]= 7.01Χ10-6 M• pH= -log[H3O+]• pH= -log[7.01Χ10-6 M]• pH= 5.154

2. Find the pH of a solution with 5.025 Χ10-4 M H3O+ concentration.

Honors: See weak acid/base ice table!

May 09, 2013

Measuring pH: Indicators1. Indicators are organic bases and acids whose colors differ from their conjugate acids or bases.

2. Many organic substances can be made into solutions that can be used to identify acids and bases.

3. Example: Red cabbage will be red if it is acidic and green if it is basic.

Transition IntervalA transition interval is the pH range at which the indicator is changing color.

Not every indicator will work in the same pH range.

pH meter: Makes a rapid, accurate pH measurement. This instrument is usually easier to read than liquid or pH indicator strips.

Neutralization/Titration

Titration: a way to determine the amount of acid/base needed to completely neutralize a substance.(Honors: More notes to follow)

Buffers• A buffer system is a solution that can absorb moderate amounts of acid or base without a significant change in its pH• Example: Enzymes

Blood: buffer of carbonic acid (H2CO3) and bicarbonate (HCO3-) to maintain a pH between 7.35 and 7.45.

May 09, 2013