Chapter Ten Acids, Bases, and Salts. Chapter 10 | Slide 2 of 66 Acid-Base Theories Arrhenius...

Chapter TenChapter Ten

Acids, Bases, and Salts

Chapter 10 | Slide 2 of 66

Acid-Base Theories

• Arrhenius Acid-Base Theory

• Bronsted-Lowry Acid-Base Theory


Arrhenius Acids

• Arrhenius Acids:– substances that ionize to form H+ in solution (e.g. HCl, HNO3,

CH3CO2H, lemon, lime, vitamin C).

• Ionization: the process in which individual positive and negative ions are produced from a compound that is dissolved in solution

• HCl (g) H+ (aq) + Cl- (aq)

• Characteristics– Sour taste

– Change blue litmus to red

– Quite corrosive


Arrhenius Bases

• Arrhenius Bases– Hydroxide-containing compounds that, in water,

produces hydroxide ions– NaOH (s) Na+ (aq) + OH- (aq)

• Characteristics– Bitter taste– Change red litmus blue– Slippery to the touch


The difference between the aqueous solution processes of ionization and dissociation.

Acids, Bases, and Salts cont’d


Litmus is a vegetable dye obtained from certain lichens found principally in the Netherlands.



Bronsted-Lowry Acid-Base Theory

• Bronsted-Lowry Acids– A substance that can donate a proton (H+ ion) to some other

substance

• Bronsted-Lowry Base– A substance that can accept a proton (H+ ion) from another

substance

Which is the acid? The base?


Conjugate Acid-Base Pairs

• Conjugate acid– The species formed when a Bronsted-Lowry base

accepts a proton

• Conjugate base– The species formed when a Bronsted-Lowry acid loses

a proton

HF (aq) + H2O (l) H3O+ (aq) + F- (aq)


Practice: Conjugate Acids and Bases

• Write the chemical formula of each of the following:– The conjugate acid of ClO3

-

– The conjugate base of NH4+

– The conjugate acid of PO43-

– The conjugate base of HS-


Amphoteric Substances

• Molecules that are able to function as either Bronsted-Lowry acids or bases

– HNO3 (aq) + H2O (l) H3O+ (aq) + NO3- (aq)

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


Mono-, Di-, and Polyprotic Acids

– Acids with one acidic proton are called monoprotic (e.g., HCl).

– Acids with two acidic protons are called diprotic (e.g., H2SO4).

– Acids with many acidic protons are called polyprotic.

(e.g., citric acid)


Strong v. Weak Acids and Bases

• Strong acids and bases completely ionize or dissociate (break into their ions) in solution.– 100% of strong acid molecules break up into their ions

• Weak acids and bases partially ionizes in solution.– Usually, less than 5% of weak acid molecules break up into their

ions. The majority of the molecules remain in the molecular form


Fig. 10.5 A comparison of the number of acidic species present in strong acid and weak acid solutions of the same concentration.



Strong Acids and Bases

Memorize these!


Ionization Constants

• Acid ionization constant– The equilibrium constant for the reaction of a weak acid with water

– HA (aq) + H2O (l) H3O+ (aq) + A- (aq)

][

]][[ 3

HA

AOHKa

The higher the Ka, the stronger the acid.


Ionization Constants

• Base ionization constant– The equilibrium constant for the reaction of a weak base with water

– B (aq) + H2O (l) BH+ (aq) + OH- (aq)

][

]][[

B

OHBHKb

The higher the Kb, the stronger the base.


→ Fig. 10.6 The acid-base reaction between sulfuric acid and barium hydroxide produces the insoluble salt barium sulfate.



← Fig. 10.3 A white cloud of finely divided solid NH4Cl is produced by the acid-base reaction that results when the colorless gases HCl and NH3 mix.


Ken O’Donoghue © Houghton Mifflin Company


Formation of water by the transfer of protons from H3O+ ion to OH- ions. A Neutralization or double replacement reaction.



Self-Ionization of Water

• A small percentage of water molecules interact with each other to form ions

[H+] = [OH-] = 1 x 10-7 M


Ion Product Constant for Water

1427 101)101(]][[ xxOHHKw

][

101][

14

3

OH

xOH

We can use the ion product constant for water to calculate the concentration of hydroxide or hydronium ions in any solution

][

101][

3

14

OH

xOH


The relationship between (H3O+) and (OH-) in aqueous solution is an inverse proportion; when (H3O+) is increased,

(OH-) decreases, and vice versa.


Relationship Between [H3O+] and [OH-] in Neutral, Acidic, and Basic Solutions.


CC 10.1



pH

• Hydronium ion concentrations can range from very small numbers to very large numbers.

It is very inconvenient to work with numbers extended over a large range

pH scale was developed to give a practical way to deal with a large range of numbers

• pH = -log[H3O+]

“p” = “take the negative logarithm of”

• [H3O+] = 10-pH


As the hydronium ion concentration increases, the pH AND the hydroxide ion concentration decrease.


Practice Calculating pH

• Calculate the pH for the following solutions:[H3O+] = 1.00 x 10-11

[H3O+] = 1.00 x 10-5

[H3O+] = 1.00 x 10-6

• Calculate the hydronium ion concentration of solutions with the following pH’s

pH = 3

pH = 9


Interpreting pH Values

• Acids– Have pH less than 7

• Bases– Have pH greater than 7

• Neutral solutions– Have pH = 7


Interpreting pH Values

• A change of 1 unit in pH always corresponds to a tenfold change in hydronium ion concentration

0.1 M HCl 10-1 M pH = 1

0.01 M HCl 10-2 M pH = 2

10 fold difference difference of 1


→ Fig. 10.10 Most fruits and vegetable are acidic.



→ Fig. 10.12 pH values of selected common liquids.



Fig. 10.13 A pH meter gives an accurate measurement of pH values.



pKa and Acid Strength

• Ka is the acid ionization constant

– The higher the Ka, the stronger the acid

• pKa = -log(Ka)

– The lower the pKa, the stronger the acid

• Determine the pKa for hydrocyanic acid, HCN, given that the Ka for this acid is 4.4 x 10-10.

-9.4


Salts

• An ionic compound containing a metal or polyatomic ion as the positive ion and a nonmetal or polyatomic ion (except hydroxide) as the negative ion– Ionic compounds with hydroxides as the negative ions

are bases

• All common soluble salts completely dissociate into ions in solution


“Neutralization” Reactions• When solutions of an acid and a base are mixed, the products of the reaction have none of the characteristic properties of the acid or the base.– Typically, water and a salt are formed (a neutral

solution) when an acid reacts with a metal hydroxide base.

• Neutralization reactions are double replacement reactions

AX + BY AY + BX


“Neutralization” Reaction Examples

• Write a balanced complete chemical equation for the reaction between aqueous solutions of HCl and KOH.

• Write a balanced complete chemical equation for the reaction between aqueous solutions of acetic acid (HC2H3O2) and barium hydroxide [Ba(OH)2].

• Write a balanced complete chemical equation for the reaction between aqueous solutions of H3PO4 and NaOH.



CAG 10.1


Buffers

• An aqueous solution consisting of a weak acid and its conjugate base in approximately equal amounts that prevent major changes in solution pH when small amounts of acid or base are added to it– The weak acid can react with added base to neutralize it– The conjugate base can react with added acid to

neutralize it


A Comparison of pH Changes in Buffered and Unbuffered Solutions.


← CC 10.2Acid Rain



Table 10.8



Fig. 10.14 (a) The buffered and unbuffered solutions have the same pH level.


Fig. 10.14 (b) After adding 1mL of a 0.01 M HCl solution, the pH of the buffered solution has not perceptibly changed, but the unbuffered solution has become acidic.

Ken O’Donoghue © Houghton Mifflin Company Ken O’Donoghue © Houghton Mifflin Company



CAG 10.2



CC 10.4


Buffers and Le Chatelier’s Principle

• HC2H3O2 (aq) + H2O (l) H3O+1 (aq) + C2H3O2-1 (aq)

• What happens if we add acid to the equilibrium?

• What happens if we add base to the equilibrium?


Buffers Practice

• Predict whether each of the following pairs of substances could function as a buffer system in aqueous solution– HCl and NaOH

– HC2H3O2 and KC2H3O2

– NaCl and NaCN

– HCN and HC2H3O2


Strong v. Weak Electrolytes

• Electrolyte: A substance whose aqueous solution conducts electricity

• Strong electrolytes: completely dissociate in solution.– For example:

• Weak electrolytes: produce a small concentration of ions when they dissolve.– These ions exist in equilibrium with the unionized substance.– For example:

• Nonelectrolytes: do not ionize in solutionHC2H3O2(aq) H+(aq) + C2H3O2

-(aq)

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


This simple device can be used to distinguish among strong electrolytes, weak electrolytes, and nonelectrolytes.



CC 10.5Electrolyte and Body Fluids



Titration

• A laboratory procedure to determine the concentration of a compound in solution. If required, the total amount of the compound present can be determined from the experimentally determined concentration.– Must have a balanced chemical equation to describe the reaction.– Must know the volume of the solution for which the unknown concentration

of the compound is to be determined.– Must have a means to indicate the endpoint of the reaction or when all of

the compound of unknown concentration has completely reacted.– The volume of a solution of known concentration is determined that will

completely react with the compound for which the concentration or amount is unknown using a titration equipment setup.

– Use the concept of Volume in L X concentration in moles / L = moles


Titration

This now becomes a stoichiometry calculation.

grams X mole / gram = moles

L X moles / L = moles

Use stoichiometric mole ratios to determine the number of moles of the compound of unknown amount or concentration.


← Fig. 10.16Diagram showing setup for titration procedures.



Titrations: Reactions Used to Determine an Unknown Concentration


Fig. 10.17 An acid-base titration using an indicator that is yellow in acidic solution and red in basic solution.



Titration Calculations: An Example

• Suppose you are titrating an unknown sample of sulfuric acid (H2SO4) with sodium hydroxide. It takes 25 mL of 0.24 M sodium hydroxide to titrate a 15 mL sample of the acid. What is the molarity of the acid?

– 2NaOH (aq) + H2SO4 (aq) 2H2O (l) + Na2SO4 (aq)


Titration Calculations: Another Example

• Determine the molarity of a NaOH solution when the following amount of acid neutralizes 25.0 mL of the NaOH solution: 23.76 mL of 1.00 M HCl.


Group Work: Titration Problems

Suppose you are titrating an unknown sample of hydrochloric acid with magnesium hydroxide. It takes 14.28 mL of 1.35 M magnesium hydroxide to titrate a 23 mL sample of the acid. What is the molarity of the acid?

2HCl (aq) + Mg(OH)2 (aq) MgCl2 (aq) + 2 H2O (l)

Chapter Ten Acids, Bases, and Salts. Chapter 10 | Slide 2 of 66 Acid-Base Theories Arrhenius...

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Transcript of Chapter Ten Acids, Bases, and Salts. Chapter 10 | Slide 2 of 66 Acid-Base Theories Arrhenius...