Chapter 15Acid-Base Theories

Properties of Acids and Bases• Acids

– Give foods a tart or sour taste• What acidic foods might you eat?

– Aqueous solutions of acids are electrolytes• Conduct Electricity

– Some are strong electrolytes (strong acids)– Some are weak electrolytes (weak acids)

• Cause indicator dyes to change colors• Many metals react with acids producing hydrogen gas• React with compounds containing hydroxide ions to form

water and a salt

Properties of Acids and Bases

• Bases– Have bitter taste, and slippery feel– Aqueous solutions of bases are also

electrolytes• Conduct Electricity

– Some are strong electrolytes (strong bases)– Some are weak electrolytes (weak bases)

• Cause indicator dyes to change colors• Water and salt are formed when a base that

contains hydroxide ions react with an acid

Arrhenious Acids and Bases

• Acids are hydrogen-containing compounds that ionize to yield hydrogen ions (H+) in aqueous solution

• Bases are compounds that ionize to yield hydroxide ions (OH-) in aqueous solution

Arrhenious Acids• Can be monoprotic, diprotic, or triprotic

– Monoportic: HNO3 → H+ + NO3-

• Ionization yields one hydrogen ion– Diprotic: H2SO4 → 2H+ + SO4

2-

• Complete ionization yields 2 hydrogen ions– Triprotic: H3PO4 → 3H+ + PO4

3-

• Complete ionization yields 3 hydrogen ions• Not all the hydrogens in an acid may be released as hydrogen ions• Not all hydrogen containing compounds are acids

– Only hydrogens joined to very electronegative elements, and thus have very polar bonds, are ionizable in water

H

H

H

H+C C O-

O

Ethanoic Acid

Nonionizable HydrogenIonizable Hydrogen

Arrhenious Bases• NaOH → Na+(aq) + OH-(aq)• KOH → K+(aq) + OH-(aq)

– Bases formed with group one metals are very soluble and caustic

– Can be made by reacting group one metals with water• Na + H2O → Na+(aq) + OH-(aq) H2 (g)

• Bases of group 2 metals are very weak resulting low solubility– Examples are Ca(OH)2 and Mg(OH)2

Bronsted-Lowry Acids and Bases

• Arrhenious definition of acids and bases is not very comprehensive and does not explain why certain substances have basic or acidic properties – Ammonia (NH3) is a base, but there is no hydroxide (OH-) in the

compound to ionize

• The Bronsted-Lowry theory defines an acid as a hydrogen-ion donor, and a base as a hydrogen-ion acceptor– Why ammonia is a base

NH3(aq) + H2O(l) → NH4+(aq) + OH-(aq)

Hydrogen ion aceptor, Bronsted-

Lowry Base

Hydrogen ion donar, Bronsted-

Lowry Acid

Makes the solution basic

Conjugate Acids and Bases

• A conjugant acid is the particle formed when a base gains a hydrogen ion• A conjugant base is the particle that remains when an acid has donated a

hydrogen ion• A conjugate acid-base pair consists of two substances related by the loss or

gain of a single hydrogen ion

NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

Base Acid Conjugate Acid

Conjugate Base

conjugate acid-base pair

conjugate acid-base pair

HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq)Acid Base Conjugate

AcidConjugate

Base

conjugate acid-base pair

conjugate acid-base pair

Conjugate Acids and Bases

• A water molecule that gains a hydrogen ion becomes a positively charged hydronium ion H3O+

– In above equation, what is the hydrogen ion donor (acid) and which is the hydrogen ion acceptor (base)

• Notice, water can both accept and donate a hydrogen ion and thus act as an acid and a base– A substance that can act as both an acid and a base is said to be

amphoteric • Amino Acids as an example

HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq)Acid Base Conjugate

AcidConjugate

Base

conjugate acid-base pair

conjugate acid-base pair

Lewis Acids and Bases

• Acids accept a pair of electrons during a reaction while a base donates a pair of electrons– Lewis acid – a substance that can accept a

pair of electrons to form a covalent bond– Lewis base – a substance that can donate a

pair of electrons to form a covalent bond

NH3 + BF3 → NH3BF3

Identify the Lewis Acid and the Lewis Base in the above equation

Acid-Base Definitions Review

Type Acid Base

Arrhenius H+ producer OH- producer

Bronsted-Lowry H+ donor H+ acceptor

Lewis electron-pair acceptor electron-pair donor

Hydrogen Ions and Acidity• Occasionally collusions between water molecules

cause them to react forming hydroxide ions and hydronium ions– The reaction in which water molecules produce ions is called

the self ionization of water

• In aqueous solution, hydrogen ions H+ are always joined to a water molecule as hydronium ions

• In pure (neutral) water, the self-ionization of water results in 1 x 10-7 M of H+ ions and 1 x 10-7 M of OH- ions– Any aqueous solution in which H+ and OH- ions are equal is

described as a neutral solution

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

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

Ion Product Constant for Water• For aqueous solutions, the product of the hydrogen-ion concentration and

the hydroxide-ion concentration equals 1.0x10-14

Kw = [H+] x [OH-] = 1.0x10-14 M

• The product of the concentrations of the hydrogen ions and hydroxide ions in water is called the ion-product constant

• Aqueous acids and bases sift the ratio of hydrogen ions to hydroxide ions in solution causing it to become either acidic or basic– In a basic solution aka alkaline solution, the hydroxide ion (OH-) is greater than

1x10-7 M and the hydrogen ion (H+) is less 1x10-7 M– In a acidic solution, the hydrogen ion (H+) is greater than 1x10-7 M and the

hydroxide ion (OH-) is less 1x10-7 M• Regardless of the acidity or alkalinity of the solution, the product of the

Molarity (M) concentration of H+ and OH- always equals 1x10-14 at 25ºC

– If the hydrogen ion H+ concentration in a soft drink is 1 x 10-5 M, what is the concentration of the hydroxide ion OH-?

– Is the solution basic, neutral, or acidic?

The pH Concept• Expressing hydrogen-ion concentrations in molarity is cumbersome

– Soren Sorensen suggested that the hydrogen ion concentration be expressed as the negative log of the hydrogen-ion concentration giving us much smaller numbers to work with

pH = -log[H+] • The pH of a solution is the negative logarithm of the hydrogen ion

concentration• A neutral solution H+ = 1x10-7 has a pH = -log[1x10-7]= 7• A solution in which the [H+] is greater than 1x10-7 M and has a pH less than

7.0 is acidic• The pH of pure water or a neutral solution has a pH of 7• A solution with a pH greater than 7 is basic and has a [H+] concentration of

less than 1x10-7 M– You can also calculate pOH which is the negative logarithm of hydroxide

ion concentration

pOH = -log[OH-]

Work some pH problems (pH of 1 x10-5 M H+?)

pH + pOH = 14

pH = 14 – pOH

pOH = 14 - pH

Relationship between pH and pOH

pH and significant figures

• Hydrogen ion concentrations should always be reported to two significant figures

• pH and pOH calculations should always be reported to two decimal places

– Rules are due to the sensitivity of pH meters

Acid-Base IndicatorsDyes

• An indicator (HIn) is an acid or base that undergoes dissociation in a known pH range– An indicator is a valuable tool for measuring pH because its

acid form and base form have different colors in solution

– The acid form of the indicator dominates the disassociation equilibrium at low pH

– The basic form of the indicator dominates the disassociation equilibrium at high pH

– Color change of an indicator occurs in a narrow pH range ≈ 2 pH units

– Thus it takes many indicators to span the entire pH spectrum• Indicator dyes have limitations

HIn (aq) H+ (aq) + In- (aq)OH-

H+

Acid Form Base Form

• Makes rapid, accurate pH measurements

• Can record pH continuously over time when performing a reactions

• Measures pH to two decimal places• Color and cloudiness of solution does

not interfere with reading• Are many different types specialized

for different jobs were pH measurements are required

Acid-Base IndicatorspH Meter

http://personals.galaxyinternet.net/tunga/Meter.jpg

http://www.vittbi.com/photogallery/biotech/PH-Meter.jpg

Strengths of Acids and BasesStrong and Weak Acids and Bases• Acids are classified as strong or weak

depending on the degree to which they ionize in water– A strong acid completely ionizes in water– Weak acids ionize only slightly in aqueous solution– What are some example of strong acids and weak

acids

HCl(q) + H2O(l) H3O+(aq) + Cl-(aq) 100% ionized

CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq) <1% ionized

pH of 0.10 M Solutions of Common Acids and Bases• Compound  pH• HCl (hydrochloric acid)  1.1• H2SO4 (sulfuric acid)  1.2• NaHSO4 (sodium hydrogen sulfate)  1.4• H2SO3 (sulfurous acid)  1.5• H3PO4 (phosphoric acid)  1.5• HF (hydrofluoric acid)  2.1• CH3CO2H (acetic acid) 2.9• H2CO3 (carbonic acid) 3.8 (saturated solution)• H2S (hydrogen sulfide)  4.1• NaH2PO4 (sodium dihydrogen phosphate)  4.4• NH4Cl (ammonium chloride)  4.6• HCN (hydrocyanic acid)  5.1• Na2SO4 (sodium sulfate)  6.1• NaCl (sodium chloride)  6.4• NaCH3CO2 (sodium acetate) 8.4• NaHCO3 (sodium bicarbonate)  8.4• Na2HPO4 (sodium hydrogen phosphate)  9.3• Na2SO3 (sodium sulfite)  9.8• NaCN (sodium cyanide)  11.0• NH3 (aqueous ammonia)  11.1• Na2CO3 (sodium carbonate)  11.6• Na3PO4 (sodium phosphate)  12.0• NaOH (sodium hydroxide, lye)  13.0http://www.cartage.org.lb/en/themes/sciences/chemistry/Inorganicchemistry/AcidsBases/Common/Common.htm

Acid Disassociation Constant• The equilibrium constant for weak acids (HA) can be written as:

• For dilute solutions, the concentration of water is a constant, and can be combined with Keq to give the acid dissociation constant (Ka)

Keq[H3O+] X [A-][HA] X [H2O]

=

Keq X H2O = Ka

[H3O+] X [A-][HA]

=

• Ka reflects the fraction of an acid in the ionized form and thus is sometimes referred to as the ionization constant

• Weak acids have small Ka values, while stronger acids have larger Ka values; why?

H3O+(aq) + A-(aq)HA(aq) + H2O(l)Acid Conjugate base

Base Disassociation Constant• The equilibrium constant for weak Bases (B) can be written as:

• For dilute solutions, the concentration of water is a constant, and can be combined with Keq to give the base dissociation constant (Kb)

Keq[BH+] X [HO-]

[B] X [H2O]=

Keq X H2O = Kb

[BH+] X [OH-][B]

=• Kb is the ratio of the concentration of the conjugate acid times the

concentration of the hydroxide ion to the concentration of the base

• The magnitue of Kb indicates the ability of a weak base to compete with the very strong base OH- for hydrogen ions

• The smaller the Kb the weaker the base

BH+(aq) + HO-(aq)B(aq) + H2O(l)base Conjugate acid

Concentration and Strength• Remember, the word strong and weak acids and bases

refers to the number particles of the acid or base that completely dissociate into their respective ions in solution

• Concentration and dilute refer to how many moles of an acid or base is diluted in a constant volume of solution

• Even though an acid may be “weak”, if it is highly concentrated, it will result in much lower pH of the solution it is dissolved in than a dilute solution of the same weak acid

Calculating Dissociation Constants• Disassociation constants are calculated from

experimental data

• To find the Ka of weak acid or the Kb of a weak base, substitute the measured concentrations of all the substances present at equilibrium into the expression for Ka or Kb

• A 0.1000M solution of ethanoic acid is only partially ionized and has a pH of 2.87. What is the acid dissociation constant (Ka) or ethanoic acid

Neutralization ReactionsAcid Base Reactions

• Reactions in which an acid and a base react in an aqueous solution to produce a salt and water are called neutralization reactions

HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

H2SO4(aq) + 2KOH(aq) → K2SO4(aq) + 2H2O(l)

Titrations• Titration - the process used to determine the concentration of

solution (often an acid or base) in which a solution of known concentration (the standard) is added to a measured amount of the solution of unknown concentration until an indicator signals the end point – In titrations, it is important to know the mole ratios that the acid and

base in question react– When an acid and base mix, the equivalence point is when the number

of moles of hydrogen ions equals the number of moles hydroxide ions giving a pH of 7

HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

H2SO4(aq) + 2KOH(aq) → K2SO4(aq) + 2H2O(l)

2HCl(aq) + CaOH(aq) → CaCl2(aq) + 2H2O(l)

Mole ratio 1:1

Mole ratio 1:2

Mole ratio 2:1

Steps in Titrating a Neutralization Reaction

1. A measured volume of an acid solution of unknown concentration is added to a flask

2. Several drops of the indicator are added to the solution while the flask is gently swirled

3. Measured volumes of a base of known concentration are mixed into the acid until the indicator just barely changes color

• The solution of known concentration is called the standard solution

• The point at which the indicator changes color is the end point

• The point of neutralization is the end point aka equivalence point of the titration

Titration of Strong Acid with a Strong Base

http://www.bwsd.k12.wi.us/astitt/Chemistry/AcidsandBases/TitrationCurve-StrongAcid.gif

Problem

• A 25 ml solution of H2SO4 is completely neutralized by 18 ml of 1.0M NaOH. What is the concentration of the H2SO4 solution?

Salts in SolutionSalt hydrolysis

• Salt consist of anion from an acid and a cation from a base– Solutions of many salts are neutral while other

salt solutions are not