chapter 19 acids, bases, and salts. acids and bases acids vinegar citrus fruits carbonated drinks...

Chapter 19Acids, Bases, and Salts

Acids and Bases

AcidsAcids vinegar citrus fruits carbonated drinks car battery lemon juice tea

BasesBases calcium hydroxide in mortar antacids household cleaning agents

Properties of Acids

1.Give foods a tart or sour taste

a)lemon & vinegar for example

2.Aqueous solutions of acids are electrolytes (conduct electricity)

3.Acids cause certain chemical indicators to change color.

4.Acid + Base Salt + water

Properties of Bases

1.Bases have a bitter taste

a)soap

2.Bases have a slippery feel

3.Aqueous solutions of bases are electrolytes (conduct electricity)

4.Bases cause certain chemical indicators to change color.

5.Acid + Base Salt + water

Arrhenius Acids & 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

In 1887, Swedish chemist Svante Arrhenius proposed a revolutionary way of defining and

thinking about acids and bases

Chemists recognized the properties of acids and bases, but were unable to propose a theory to explain their behavior.

Arrhenius Acids

Monoprotic acidsMonoprotic acids – acids that contain one ionizable hydrogen

HNO3 – nitric acid

Diprotic acidsDiprotic acids – acids that contain two ionizable hydrogens

H2SO4 – sulfuric acid

Triprotic acidsTriprotic acids – acids that contain three ionizable hydrogens

H3PO4 – phosphoric acid

Arrhenius Acids

• Not all compounds that contain hydrogen are acidsEx. CH4 – methane has weak polar C – H bonds and no ionizable hydrogens. Not an acid.

• Not all hydrogens in an acid may be released as hydrogen ions.

• Only hydrogens in very polar bonds are ionizable. In the case where hydrogen is joined to a very electronegative element.

Ex. HCl hydrogen chloride very polar covalent molecule

Arrhenius Acids

• When HCL dissolves in water, it releases hydrogen ions because the hydrogen ions are stabilized by solvation.

H2O

H – Cl (g) H+ (aq) + Cl- (aq)Hydrogen Hydrogen Chloride chloride ion ion

Ionizes to form an aqueous solution of hydronium ions and chloride ions

HCl + H2O H3O+ + Cl-

Arrhenius Acids • Ethanoic acid CH3COOH is a monoprotic acid due

to its structure

H O

H C C O H

H

The three H attached to the carbon are in weak polar bonds. They do not ionize.

Only the H bonded to the highly electronegative O can be ionized

Arrhenius Bases Sodium hydroxide dissociates into sodium ions and hydroxide ions in aqueous solution.

H2O

NaOH (s) Na+ (aq) + OH- (aq)

Sodium Sodium HydroxideHydroxide Ion ion

Potassium hydroxide dissociates into sodium ions and hydroxide ions in aqueous solution.

H2O

KOH (s) K+ (aq) + OH- (aq)

Potassium Potassium HydroxideHydroxide Ion ion

Arrhenius Bases

Group IA, the alkali metals, react with water to produce solutions that are basic.

Group IA metals are very soluble in water and can produce concentrated solutions.

Group 2A metals are not very soluble in water. Their solutions are always very dilute.

Bronsted-Lowry Acids and Bases

Arrhenius’ definition of acids and bases is not a very comprehensive one.

If defines acids and bases narrowly and does not include certain substances that have acidic or basic properties.

Na2CO3 (aq) is basic


The Bronste-Lowry theory defines

acid – a hydrogen-ion donor

base – a hydrogen-ion acceptor

All acids and bases included in the Arrhenius theory are also acids and bases according to the Bronsted-Lowry theory.

Ammonia as a BaseBronsted-Lowry Theory

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

•ammonia is the hydrogen-ion acceptor and therefore a BL base

•water is the hydrogen-ion donor and therefore a BL acid.

•Hydrogen ions are transferred from water to ammonia, which causes the hydroxide-ion concentration to be greater than it is in pure water.

Conjugate Acids and Bases

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

base acid conjugate acid conjugate base

•When ammonia dissolves and reacts with water, NH4

+ is the conjugate acid of the base NH3.

•OH- is the conjugate base of acid H2O

Conjugate Acids and Bases

HCl (g) + H2O (l) Ý H3O+ (aq) + Cl- (aq)

acid base conjugate acid conjugate base

•HCl is the hydrogen-ion donor – thus a BL acid.

•Water is the hydrogen-ion acceptor – thus BL base

Conjugate Acid-Base Pair

Conjugate acid – the particle formed when a base gains a hydrogen ion

Conjugate base – the particle that remains when an acid has donated a hydrogen ion..

Conjugate acids and bases are always paired with a base or an acid, respectively.

Conjugate acid-base pairs consists of two substances related by the loss or gain of a single hydrogen ion.

Common Conjugate Acid-Base Pairs

Acid BaseHCl Cl-

H2SO4 HSO4-

H3O+ H2O

HSO4- SO4

2-

CH3COOH CH3COO-

H2CO3 HCO3-

HCO3- CO3

2-

NH4+ NH3

H2O OH-


A water molecule that gains a hydrogen ion becomes a positively charged hydronium ionhydronium ion (H3O+)

AmphotericAmphoteric – a substance that can act as both an acid and a base

Ex: water

H2SO4 + H2O H3O+ + HSO4-

NH3 + H2O NH4+ + OH-

Lewis Acids and Bases

Gilbert Lewis proposed a third Acid Base theory

AcidAcid – accepts a pair of electrons during a reaction

Base Base – donates a pair of electrons during a reaction

Concept is more general than either the Arrhenius theory or the Bronsted-Lowry theory.

Lewis Acids and Bases

Lewis AcidLewis Acid – a substance that can accept a pair of electrons to form a covalent bond.

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

..H+ + -:O – H :O:

.. H H

Lewis LewisAcid Base

Acid Base Definitions

Type Acid Base

Arrhenius H+ producer OH- producer

Bronsted Lowry H+ H+ acceptor

LewisElectron-pair

acceptorElectron-pair

donor

End of Section 19.1

Hydrogen Ions From Water

Water molecules are highly polar and are in continuous motion.

Occasionally, the collisions between water molecules are energetic enough to transfer a hydrogen ion from one water molecule to another.

Self ionization of waterSelf ionization of water – the reaction in which water molecules produce ions

Hydrogen Ions From Water

A water molecule that loses a hydrogen ion becomes a negatively charged hydroxide ion

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

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

Hydroxide ion Hydroxide ion

Self ionization of waterSelf ionization of water – the reaction in which water molecules produce ions

Self Ionization of Water

Hydrogen ions in aqueous solution have several names. • Some chemists call them protons •Some chemists call them hydrogen ions or hydronium ions.

For our purposes, either H+ or H3O+ will represent hydrogen ions in aqueous solution.

H2O + H2O H3O+ + OH-

Self Ionization of Water

The self-ionization of water occurs to a very small extent.

• In pure water at 25˚C, the equilibrium concentration of hydrogen ions and hydroxide ions are each only 1 x 10-7.

• In other words the concentration of OH- and H+ are equal in pure water

Neutral Solutions

Any aqueous solution in which H+ and OH- are equal is a neutral solution.neutral solution.

Ion Product Constant for Water

When [H+] increases [OH-] decreases

When [H+] decreases [OH-] increases

LeChatelier’s principle LeChatelier’s principle – when a stress is applied to a system in dynamic equilibrium, the system changes in a way that relieves the stress

If additional ions (either H+ or OH-) are added to a solution, the equilibrium shifts. The concentration of the other type of ion decreases. More water molecules are formed in the process.

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


For aqueous solutions, the product of the hydrogen ion concentration and the hydroxide ion concentration equals 1.0 x 10-14

[H+] x [OH-] = 1.0 x 10-14

This equation is true for all dilute aqueous solutions at 25˚C.

Ion-Product Constant for Water Ion-Product Constant for Water (K(Kw) ) – the product of the concentrations of the hydrogen ions and hydroxide ions in water

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


Not all solutions are neutral

When some substances dissolve in water, they release hydrogen ions.

When hydrogen chloride dissolves in water, it forms hydrochloric acid.

H2O

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


In the previous HCl solution, the hydrogen-ion concentration is greater than the hydroxide-ion concentration.

Acidic SolutionAcidic Solution – one in which [H+] is greater than [OH-].

The [H+] of an acidic solution is greater than 1 x 10-7


When sodium hydroxide dissolves in water, it forms hydroxide ions in solution.

H20

NaOH(s) Na+(aq) + OH-(aq)

In the above solution, the hydrogen-ion concentration is less than the hydroxide-ion concentration.

Basic SolutionBasic Solution – one in which [H+] is less than [OH-]

The [H+] of a basic solution is less than 1 x 10-7

Basic solutions are also known as alkaline solutionsalkaline solutions.

The pH Concept

The pH scale was proposed by Danish Scientist Soren Sorensen in 1909.

The pH scale is used to express [H+]

1 2 3 4 5 6 7 8 9 10 11 12 13 14

Strongly Neutral Strongly Acidic Basic

Calculating pH

The pH of a solution is the negative logarithm of the hydrogen-ion concentration.

pH = -log[H+]

Calculating pH

In neutral solution, the [H+] = 1 x 10-7M. The pH is 7

pH = -log[H+] pH = -log(1 x 10-7)

pH = -(log 1 + log 10-7)pH = -(0.0 + -7.0)

pH = 7.0

Classifying Solutions

A solution in which [H+] is greater than 1 x 10-7 has a pH less than 7.0 and is acidic.

A solution in which [H+] is less than 1 x 10-7 has a pH greater than 7.0 and is basic.

The pH of pure water or a neutral aqueous solution is 7.0

Acidic solution: pH < 7.0 [H+] > 1 x 10-7MNeutral solution: pH = 7.0 [H+] equals 1 x 10-7MBasic solution: pH > 7.0 [H+] < 1 x 10-7

Calculating pH

pH can be read from the value of [H+] if it is written in scientific notation and has a coefficient of 1.

Then the pH of the solution equals the exponent, with the sign changed from minus to plus

[H+] = 1 x 10-2 has a pH of 2.0

[H+] = 1 x 10-13 has a pH of 13.0

Calculating pH

If the pH is an integer, it is also possible to directly write the value of [H+].

pH = 9.0 then [H+] of 1 x 10-9M

pH = 4 then [H+] = 1 x 10-4M

Calculating pOH

Acidic solution: pOH > 7.0 [OH-] < 1 x 10-7MNeutral solution: pOH = 7.0 [OH-] equals 1 x 10-7MBasic solution: pOH < 7.0 [OH-] > 1 x 10-7

The pOH of a solution equals the negative logarithm of the hydroxide-ion concentration

pOH = -log [OH-]

A neutral solution has a pOH of 7

pH and pOH Relationship

pOH + pH – 14

pH= 14 – pOH

pOH = 14 - pH

pH Significant Figures

For pH calculation, you should express the hydrogen-ion concentration in scientific notation

[H+] = 0.0010M should be written 1.0 x 10-3

0.0010M has two sig figs

Write pH = 3.00 with 2 zeros to the right of the decimal place representing the 2 sig figs

Problem Example

Colas are slightly acidic. If the [H+] in a solution is 1.0 X 10-5 M , is the solution acidic, basic or neutral. What is the [OH-] of this solution?

[H+] = 1.0 X 10-5 M which is greater than 1.0 X 10-7 M so solution is acidic

Kw = [OH-] x [H+] = 1.0 X 10-14

[OH-] = 1.0 X 10-14 / [H+]

[OH-] = 1.0 X 10-14 / 1.0 X 10-5 [OH-] = 1.0 X 10-9

Problem Example

What is the pH of a solution with a hydrogen-ion concentration of 4.2 x 10-10 M?

pH = -log [H+]pH = -log (4.2 x 10-10)

pH = -(9.3765)pH = 9.38

Problem Example

pH of an unknown solution is 6.35. What is its hydrogen-ion concentration?

pH = -log [H+]6.35 = -log [H+]-6.35 = log [H+]

Using calculator find the antilog of -6.354.5 x 10-7 M = [H+]

Problem Example

What is the pH of a solution if the [OH-] =4.0X10-11M?

Kw = [H+] x [OH-]= 1 x 10-14

[H+] = 1 x 10-14 / [OH-]

[H+] = 1 x 10-14 / 4.0 x 10-11

[H+] =0.25 x 10-3 M

[H+] = 2.5 x 10-4 M

Problem Example (con’t)

What is the pH of a solution if the [OH-] =4.0X10-11M?

pH = -log [H+]

pH = -log (2.5 x 10-4)

pH = - (-3.60205)

pH = 3.60

Acid-Base IndicatorsIndicatorIndicator - (HIn) is an acid or a base tht undergoes dissociation in a know pH range

An indicator is a valuable tool for measuring pH because its acid form and base form have different color in solution.

OH-

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

acid H+ base form form

The acid form dominates the dissociation equilibrium at low pH (high [H+]), and the base form dominates the equilibrium at high pH (high [OH-])

Acid-Base Indicators

For each indicator, the change from dominating acid from to For each indicator, the change from dominating acid from to dominating base form occurs in a narrow range of dominating base form occurs in a narrow range of approximately two pH units. approximately two pH units.

Within this range, the color of the solution is a mixture of the colors of the acid and the base forms.

Knowing the pH range over which this color change occurs, can give you a rough estimate of the pH of the solution.

Acid-Base Indicators

Many different indicators are needed to span the entire pH Many different indicators are needed to span the entire pH spectrum. spectrum.

Indicator characteristics that limit their usefulness. • Listed pH values of indicators are usually given for 25ºC.

At other temperatures, an indicator may change color at a different pH.

• If the solution being tested is not colorless, the color of the indicator may be distorted.

• Dissolved salts in a solution may also affect the indicator’s dissociation.

Using indicator strips can help overcome these problems.

pH Meters

A pH meter makes rapid, accurate pH A pH meter makes rapid, accurate pH measurements. measurements.

• often easier to use than liquid indicators or indicator

strips.

• Measurements of pH obtained with a pH meter are typically accurate to within 0.01 pH unit of the true pH.

• Color and cloudiness of the unknown solution do not affect the accuracy of the pH value

• If the solution being tested is not colorless, the color of the indicator may be distorted.

End of section 19.2

Strong Acids

Acids are classified as strong or weak depending Acids are classified as strong or weak depending on the degree to which they ionize in water.on the degree to which they ionize in water.

• In general, strong acids are completely ionized in

aqueous solution.

HNO3 - nitric acid HCl - hydrochloric acidH2SO4 - sulfuric acid HClO4 - perchloric acid HBr - hydrobromic acidHI - hydroiodic acid

HCl(g) + H2O(l) H3O+(aq) + Cl-(aq)

Weak Acids

Weak acids ionize only slightly in aqueous solution.Weak acids ionize only slightly in aqueous solution. • Some Weak Acids

Acetic Acid H3COOH

Boric Acid H3BO3 (all three are weak)

Phosphoric Acid H3PO4 (all three are weak)

Sulfuric Acid HSO4- (first ionization is strong)

CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-

(aq)

ethanoic acid water hydronium ethanoate

ion ion

Acid Strength

A strong acid completely dissociates in water A strong acid completely dissociates in water ([H([H33OO++] is high).] is high).

A weak acid remains largely undissociated.A weak acid remains largely undissociated. ([H([H33OO++] is low).] is low).

Equilibrium Constant (Keq)

Write the equilibrium-constant expression from the Write the equilibrium-constant expression from the balanced chemical equation.balanced chemical equation.

CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq)

Keq = [H3O+] x [ CH3COO- ]

[H3COOH] x [H2O] [H2O] constant in dilute solutions

Acid Dissociation Constant (Ka)

KKaa = Ratio of the concentration of the dissociated = Ratio of the concentration of the dissociated

form of an acid to the concentration of the form of an acid to the concentration of the undissociated form. undissociated form.

H3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq)

Acid Dissociation ConstantAcid Dissociation ConstantKa = [H3O+] x [ CH3COO- ]

[CH3COOH]


Acid dissociation constant reflects the fraction of an Acid dissociation constant reflects the fraction of an acid in the ionized form. (Kacid in the ionized form. (Kaa sometimes called sometimes called

ionization constant) ionization constant) If the value of the Ka is small, then the degree of

dissociation or ionization of the acid in the solution is small.

Weak acids – small Ka values

Stronger the acid – larger the Ka


Nitrous acid (HNONitrous acid (HNO22) has a K) has a Kaa of 4.4 x 10 of 4.4 x 10-4-4

Acetic acid (CHAcetic acid (CH33COOH) has a KCOOH) has a Kaa of 1.8 x 10 of 1.8 x 10-5-5

Nitrous acid is more ionized in solution and a

stronger acid

Acids

Strong AcidsStrong Acids• Have high [HHave high [H33OO++] ] • Large dissociation constantLarge dissociation constant

Weak AcidsWeak Acids• Have low [HHave low [H33OO++] ] • Small dissociation constantSmall dissociation constant

Acids

Diprotic and triprotic acids lose their hydrogens one Diprotic and triprotic acids lose their hydrogens one at a time. at a time.

Each ionization reaction has a separate dissociation constant.

H3PO4 – 3 separate dissociation constants.

Base Dissociation Constant (Kb)

Strong bases dissociate completely into metal ions Strong bases dissociate completely into metal ions and hydroxide ions in aqueous solution. and hydroxide ions in aqueous solution.

• Some strong bases are not very soluble in water Some strong bases are not very soluble in water (calcium hydroxide and magnesium hydroxide)(calcium hydroxide and magnesium hydroxide)

• Small amounts that do not dissolve dissociate Small amounts that do not dissolve dissociate completelycompletely

Weak bases react with water to form the hydroxide

ion and the conjugate acid of the base.

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

Ammonia Water Ammonium Ion Hydroxide ion

Base Dissociation Constant (Kb)

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

Ammonia Water Ammonium Ion Hydroxide ion

Only about 1% of ammonia is present as Only about 1% of ammonia is present as NH4+

Equilibrium ConstantEquilibrium Constant Keq = [NH4

+] x [OH- ]

[NH3] x [H2O] [H2O] constant in dilute solutions

Base Dissociation ConstantBase Dissociation ConstantKb = [NH4

+] x [OH- ]

[NH3]

Concentration and Strength

The words concentrated and dilute indicate how much of an acid or base is dissolved in solution.

• Number of moles of the acid or base in a given volume

The words strong and weak refer to the extent of ionization or dissociation of an acid or base

• How many of the particles ionize or dissociate into ions

A sample of HCl added to a large volume of water A sample of HCl added to a large volume of water becomes more dilute, but it is still a strong acid. becomes more dilute, but it is still a strong acid.

Vinegar is a dilute solution of a weak acid.Vinegar is a dilute solution of a weak acid.

End of section 19.3

Acid-Base Reactions

If you mix a solution of a strong acid containing hydronium ions with a solution of a strong base that has an equal number of hydroxide ions, a neutral solution results.

• Final solution has properties that are characteristic of neither an acidic nor a basic solution.

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

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

Neutralization Reactions

Reactions of weak acids and weak bases do not usually produce a neutral solution.

In general, reactions with which an acid and a base react in an aqueous solution to produce a salt and water are called neutralization reactionsneutralization reactions.

Making Salts

Prepare potassium chloride by mixing equal molar quantities of hydrochloric acid and potassium hydroxide.

HCl + KOH KCl + H20

Heating the solution to evaporate the water will leave the salt potassium chloride.

In general, the reaction of an acid with a base produced water and salt

TitrationThe number of moles of hydrogen ions provided by

the acid are equivalent to the number of hydroxide ions provided by the base.

HCl(aq) + NaOH(aq) NaCl (aq) + H20 (l) 1 mole 1 mole 1 mole 1 mole

H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H20 (l) 1 mole 2 mole 1 mole 2 mole

When and acid & base are mixed, the Equivalence Equivalence pointpoint is when the number of moles of hydrogen ions equals the number of moles of hydroxide ions.

Sample ProblemHow many moles of sulfuric acid are required to

neutralize 0.50 mol of sodium hydroxide?

H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H20 (l) • Mole ratio of H2SO4 to NaOH is 1:2

0.50 mol NaOH 1 mol H2SO4 = 0.25 mol H2SO4

2 mol NaOH

Practice ProblemHow many moles of potassium hydroxide are needed

to completely neutralize 1.56 mol of phosphoric acid?

H3PO4(aq) + 3KOH(aq) K3PO4(aq) + 3H2O(l)

1.56 mol H3PO4 3 mol KOH = 4.68 mol H3PO4

1 mol H3PO4

TitrationYou can determine the concentration of acid or base

in a solution by performing a neutralization reaction.

You must use an appropriate acid-base indicator to show when neutralization has occurred.

In the lab, typically phenolphthalein for acid base

neutralization reactions.

Solutions that contain phenolphthalein turn from colorless to deep pink as the pH of the solution changes from acidic to basic.

Titration

a.Measured volume of an acid solution of unknown concentration is added to a flask

Titration

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

Titration

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

Titration

TitrationTitration – the process of adding a known amount of solution of known concentration to determine the concentration of another solution.

Standard solutionStandard solution – the solution of known concentration

End pointEnd point – the point at which the indicator changes color

You can also use titration to find the concentration of a base using a standard acid.

Titration

TitrationTitration – the process of adding a known amount of solution of known concentration to determine the concentration of another solution.

Standard solutionStandard solution – the solution of known concentration

End pointEnd point – the point at which the indicator changes color. The point of neutralization

Equivalence pointEquivalence point – the point in a titration where the number of moles of hydrogen ions = number of moles of hydroxide ions..

Sample Problem

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

H2SO4(aq) + 2NaOH(aq) Na2SO4(aq) + 2H20 (l) 0.018 L NaOH 1.0 mol NaOH 1 mol H2SO4 =

1L NaOH 2 mol NaOH 0.025L

0.36M H2SO4

Practice Problem

How many milliliters of 0.45M HCl will neutralize 25.0ml of 1.00M KOH?

HCl + KOH H2O + KCl 0.025 L KOH 1.0 mol KOH 1 mol HCl 1L KOH 1 mol KOH

1 L HCl 1000 ml HCl = 56 ml HCl

0.45 mol HCl 1 L HCl

Practice Problem

What is the molarity of H3PO4 if 15.0 ml is completely neutralized by 38.5 ml of 0.150 M?

H3PO4 + 3NaOH 3H2O + Na3PO4

0.0385 L NaOH 0.150 mol NaOH 1 mol H3PO4

1L NaOH 3 mol NaOH

= 0.129M H3PO4

0.015L H3PO4

End of section 19.4

Salt Hydrolysis

A salt consists of an anion from an acid and a cation from a base.

The salt forms as a result of a neutralization reaction

Although solutions of many salts are neutral, some are acidic and others are basic.

..

Salt Hydrolysis

Salt HydrolysisSalt Hydrolysis – the cations or anions of a dissociated salt remove hydrogen ions from or donate hydrogen ion to water.

Hydrolyzing salts are usually derived from a strong acid and weak base or from a weak acid and a strong base.

In general, salts that produce acidic solutions contain positive ions that release protons to water.

Salts that produce basic solutions contain negative ions that attract protons from water.

Salt Hydrolysis

CH3COONa (aq) CH3COO- (aq) + Na+ (aq)

Sodium ethanoate ethanoate ion sodium ion

CH3COONa is the salt from a weak acid CH3COOH and a strong base NaOH

In solution the salt is completely ionized.

Salt Hydrolysis

Salt HydrolysisSalt Hydrolysis – the cations or anions of a dissociated salt remove hydrogen ions from or donate hydrogen ion to water.

CH3COO-(aq) + H2O(l) CH3COOH (aq) + OH- (aq)

BL base BL acid makes hydrogen-ion hydrogen-ion solution acceptor donor basic

This process is called hydrolysis because it splits a hydrogen ion off a water molecule.

Resulting solution contains a hydroxide-ion concentration greater than the hydrogen-ion concentration. Thus the solution is basic

Salt Hydrolysis

NH4Cl (aq) NH4+ (aq) + Cl- (aq)

Ammonium Ammonium ion Chloride ion chloride

NH4Cl is the salt from a strong acid (hydrochloric acid, HCl) and a weak base (ammonia, NH3)

In solution the salt is completely ionized.

Salt Hydrolysis

NH4+(aq) + H2O(l) NH3(aq) + H3O+(aq)

BL acid BL base makes hydrogen-ion hydrogen-ion solution donor acceptor acidic

This process is also called hydrolysis because it splits a hydrogen ion off a water molecule.

Resulting solution contains a hydrogen-ion concentration greater than the hydroxide-ion concentration. Thus the solution is acidic

Salt Hydrolysis

Equivalence Point

Strong Acid Strong Base

pH= 7 neutral

Weak Acid

Strong BasepH > 7 basic

Strong Acid

Weak BasepH < 7 acidic

Equivalence pointEquivalence point – the point in a titration where the number of moles of hydrogen ions = number of moles of hydroxide ions

Buffers

BufferBuffer – a solution in which the pH remains relatively constant when small amounts of acid or base are added.

A buffer is a solution of a weak acid and one of its salts, or a solution of a weak base and one of its salts.

A buffer solution is better able to resist drastic changes in pH than is pure water.

Buffers

A solution of ethanoic acid (CH3COOH) and sodium ethanoate (CH3COONa) is an example of a typical buffer.

CH3COOH and CH3COO- (source is the completely ionized CH3COONa) act as reservoirs of neutralizing power.

Buffers

CH3COO-(aq) + H+(aq) CH3COOH (aq)

ethanoate ion hydrogen ion ethanoic acid

When an acid is added to the solution, the ethanoate ions act as a hydrogen-ion sponge.

CH3COOH (aq) + OH-(aq) CH3COO-(aq) + H2O (l) Ethanoic acid hydroxide ion ethanoate ion water

When a base is added to the solution, the ethanoic acid and the hydroxide ions react to produce water and the ethanoate ion.

Buffers

The ethanoate ion is not strong enough base to accept hydrogen ions from water extensively.

The buffer solution cannot control the pH when too much acid is added, because no more ethanoate ions are present to accept hydrogen ions.

Buffer also become ineffective when too much base is added. No more ethanoic acid molecules are present to donate hydrogen ions.

Buffers

When too much acid or base is added, the buffer capacity is exceeded.

Buffer capacityBuffer capacity – the amount of acid or base than can be added to a buffer solution before a significant change in pH occurs.

Buffers

When a base is added to a buffered solution, the acidic form removes hydroxide ions from the solution.

When an acid is added to a buffered solution, the basic form removes hydrogen ions from the solution.

Buffers & Your Blood

Your body function properly only when the pH of your blood lies between 7.35 and 7.45.

Your blood contains buffers (hydrogen carbonate ions and carbonic acid)

HCO3- (aq) + H+ (aq) H2CO3 (aq)

hydrogen hydrogen ions carbonic acid carbonate ion

As long as there are hydrogen carbonate ions available, the excess hydrogen ions are removed, and the pH of the blood changes very little.

End of Section Chapter 19

chapter 19 acids, bases, and salts. acids and bases acids vinegar citrus fruits carbonated drinks...

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