CHEMISTRY - SILBERBERG 8E CH.18 - ACID-BASE...

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CHEMISTRY - SILBERBERG 8E

CH.18 - ACID-BASE EQUILIBRIA

CONCEPT: ACID IDENTIFICATION

The most common feature of an acid is that many possess an H+ ion called the _______________________________ .

When it comes to acids there are 2 MAJOR TYPES that exist:

_______________________ are acids where the H+ ion is attached to an electronegative element.

• These types of acids lack the element __________________ and usually possess no __________________ .

• The most common type of these particular acids are the haloacids: _______ , _______ , _______ & _______ .

_______________________ are acids that contain the ________________ , ________________ & ________________.

• They are created by the hydration of nonmetal oxides.

PRACTICE: Which of the following compound(s) cannot be classified as an acid?

a) H2S b) HCN c) H2 d) C6H6 e) All are acids.



CONCEPT: BINARY ACID STRENGTH

STRONG ACIDS are considered _________________ Electrolytes so they ionize completely in water.

HCl (aq) H2O

H+ (aq) + Cl – (aq)

WEAK ACIDS are considered __________________ Electrolytes so they don’t completely ionize in water.

HF + H2O F – (aq) + H3O+ (aq)

The strength of a BINARY ACID is based on the _________________________ or ________________ of the nonmetal.

• For elements in the same period then look at their __________________ . The ________, the ________ acidic.

• For elements in the same group then look at their __________________ . The ________, the ________ acidic.

BINARY ACID STRENGTH

PRACTICE 1: Which is the weakest acid from the following?

a) H2S b) H2Se c) H2Te d) All would have the same acid strength.

PRACTICE 2: Which of the following acids would be classified as the strongest?

a) CH4 b) NH3 c) H2O d) HF e) PH3



CONCEPT: OXYACID STRENGTH

The strength of OXYACIDS is based on the number of _____________ or the _____________________ of the nonmetal.

• RULE: If my oxyacid has 2 or More ___________ than ____________ then my oxyacid is a __________ ACID.

HNO3 ___ Oxygens – ___ Hydrogens

C6H5OH ___ Oxygens – ___ Hydrogens

HBrO4 ___ Oxygens – ___ Hydrogens

When comparing the strengths of different oxyacids remember:

• If they have different number of oxygens then the _________ oxygen the ___________ acidic

• If they have the same number of oxygens then the _________ electronegative the nonmetal the ________ acidic.

Electronegativity

H2C2O4 ___ Oxygens – ___ Hydrogens

HSO4 –

___ Oxygens – ___ Hydrogens

3 Exceptions

HIO3 ___ Oxygens – ___ Hydrogens

PRACTICE: Rank the following oxyacids in terms of increasing acidity.

a) HClO3 b) HBrO4 c) HBrO3 d) HClO4



CONCEPT: BASE STRENGTHS

STRONG BASES are considered _________________ Electrolytes so they ionize completely in water.

NaOH (aq) H2O

Na+ (aq) + OH – (aq)

WEAK BASES are considered __________________ Electrolytes so they don’t completely ionize in water.

NH3 + H2O NH4+ (aq) + OH – (aq)

Bases possess THREE major features: __________________ or __________________ or __________________ .

Group ________:

• Any Group ______ metal when combined with OH –, H –, O2– or NH2 – makes a STRONG BASE.

Group ________:

• Any Group ______ metal, from _____ to _____ , when combined with OH –, H –, O2– or NH2 – makes a STRONG

BASE.

_____________:

• ____________________________________ are considered WEAK BASES.

Ex:

• ____________________________________ are considered WEAK ACIDS.

Ex:



PRACTICE: ACID & BASE IDENTIFICATION

EXAMPLE: Classify each of the following as a strong acid, weak acid, strong base or weak base.

a) HCHO2 c) H2NNH2

b) (CH3CH2)3NH+ d) HBrO3

PRACTICE 1: Classify each of the following as a strong acid, weak acid, strong base or weak base.

a) KOCH3 b) CH3OH


a) HOCN b) H5IO6


a) NaN3 b) SrH2



CONCEPT: ARRHENIUS ACIDS & BASES

The most general definition for acids and bases was developed by Svante Arrhenius near the end of the 19th century.

• According to him, the _______ cation and the _________ anion are fundamental to the concept of acids and bases.

• His definition however failed to describe acidic and basic behavior in nonaqueous media.

The Arrhenius definition states an acid is a compound that increases _______________ when dissolved in a solvent.

The Arrhenius definition states a base is a compound that increases _______________ when dissolved in a solvent.

PRACTICE 1: Which ions are formed from the dissociation of the following compound?

a) Sr(OH)2 (s) Dissolves in H2O


a) H2SO4 (l) Dissolves in H2O


a) HBO32-

Dissolves in H2O



CONCEPT: BRONSTED LOWRY ACIDS & BASES

In 1923, Johannes Brønsted and Thomas Lowry developed a new set of definitions for acids and bases.

According to the Bronsted-Lowry definition, acids are considered _____________________________ and bases are

considered _____________________________.

• Unlike Arrhenius acids and bases, they are not limited to aqueous solutions.

• Every Arrhenius acid is a Brønsted-Lowry acid (and likewise for the bases).

• Brønsted-Lowry acids and bases always occur in pairs called _____________________________________ .

EXAMPLE 1: Write the formula of the conjugate base for the following compound:

HSO4 –

EXAMPLE 2: Write the formula of the conjugate acid for the following compound:

V2O52-

PRACTICE 1: Write the formula of the conjugate base for the following compound:

H2Se

PRACTICE 2: Write the formula of the conjugate for the following compound:

NH2NH2



PRACTICE: BRONSTED LOWRY ACIDS & BASES (CALCULATIONS)

EXAMPLE 1: Identify the acid, base, conjugate acid and conjugate base in the following reactions:

a) HF (aq) + H2O (aq) F – (aq) + H3O+ (aq)

EXAMPLE 2: Identify the acid, base, conjugate acid and conjugate base in the following reactions:

a) CN – (aq) + H2O (aq) HCN (aq) + OH – (aq)

PRACTICE 1: Which of the following is a Bronsted-Lowry acid? a) CH4 b) HCN c) NH3 d) Br2

PRACTICE 2: Determine the chemical equation that would result when carbonate, CO32-, reacts with water.



CONCEPT: AMPHOTERIC SPECIES

An amphoteric, or _________________________, is a species that can act as a(n) ACID or BASE.

• Water is prime example of an amphoteric species.

Partially dissociated conjugate bases of polyprotic acids are also amphoteric.

• These compounds possess _________________ and a __________________________.

Ex:

PRACTICE: Which of the following species is/are amphoteric?

a) CO32– b) HF c) NH4

+ d) HPO32- e) CH3O –



CONCEPT: LEWIS…THE FINAL TYPE OF ACID & BASE

In the 1920s, Gilbert Lewis proposed a new set of definitions for acids and bases.

A Lewis acid is a(n) ______________________________________.

• ________ acts as a Lewis acid when connected to an electronegative element: ___ , ___ , ___ , ___ , or ________

• _____________________ charged hydrogen or metals.

• If your central element has _________________ 8 valence electrons.

A Lewis base is a(n) ______________________________________.

• Compounds with _________________________ .

NH3 H2O CH3OH CH3OCH3

• Compounds with a _________________________ .

CN – OH – CH3O – N3–



PRACTICE: LEWIS….THE FINAL TYPE OF ACID & BASE (CALCULATIONS)

EXAMPLE: Identify each of the compounds in the following chemical equation.

H3CH2C O

CH2H3C

Al

Br

Br Br

H3CH2C O

CH2H3C

Al

Br

Br

Br

PRACTICE 1: Identify the Lewis acids and bases in the following reactions.

a) H+ + OH – H2O

b) Cl – + BCl3 BCl4–

c) SO3 + H2O H2SO4

PRACTICE 2: Identify each of the following compounds as either a Lewis acid, a Lewis base or neither.

a) ZnCl2 b) CN –

c) NH4+ d) Co3+



CONCEPT: pH and pOH

To deal with incredibly small concentration values of [H+] and [OH-] we can use the pH scale.

• Under normal conditions, the pH scale operates within the range of ______ to ______ .

By taking the – log of [H+] and [OH-] we can find pH and pOH.

pH = − log[H+ ] pOH = − log[OH− ] p = − log

By recognizing the relationship between [H+] and [OH-] with pH and pOH we can create new formula relationships.

pH = − log[H+ ] pOH = − log[OH− ]

A species with a pH greater than 7 is classified as _____________ and the [H+] is ___________________ than the [OH-].

• The ______________ the base then the ______________ the pH.

A species with a pH less than 7 is classified as _______________ and the [H+] is ____________________ than the [OH-].

• The ______________ the acid then the ______________ the pH.

A species with a pH equal to 7 is classified as ______________ and the [H+] is _____________________ than the [OH-].

By using – log with the equilibrium expression for water a relationship between pH and pOH can be created.

pH+ pOH =14



PRACTICE: pH and pOH (CALCULATIONS 1)

EXAMPLE: What is the hydroxide ion and hydronium ion concentration of an aqueous solution that has a pH equal to 6.12?

PRACTICE 1: Which of the following solutions will have the lowest concentration of hydronium ions?

a) 0.100 moles C6H5NH2

b) 0.100 moles Be(OH)2

c) 0.100 moles SrH2

d) 0.100 moles (CH3)2NH

PRACTICE 2: Which of the following statements about aqueous solutions is/are true?

a) For an basic solution the concentration of H3O+ is greater than the concentration of OH –.

b) The pH of a neutral aqueous solution is 7.00 at all temperatures.

c) An acidic solution under normal conditions has a pH value less than 7.00.

d) If the concentration of H3O+ decreases then the concentration of OH – will also decrease.

e) The pH of aqueous solutions is less than 7.



PRACTICE: pH and pOH (CALCULATIONS 2)

EXAMPLE: A solution is prepared by dissolving 0.235 mol Sr(OH)2 in water to produce a solution with a volume of 750 mL.

a) What is the [OH-]?

b) What is the [H+]?

PRACTICE: What is the Kw of pure water at 20.0°C, if the pH is 7.083?

a) 8.26 × 10-8 b) 6.82 × 10-15 c) 7.23 × 10-14 d) 1.00 × 10-14



CONCEPT: AUTO IONIZATION OF WATER

Water can react with itself in a reaction called self–ionization where ______________ and ______________ are produced.

H2O (l) + H2O (l)

This reaction is usually written more simply as:

H2O (l) + H2O (l)

The equilibrium equation for water is called the ________________________ (KW) for water and is given by the following:

KW = [H+ ][OH− ]

At 25oC, KW = ___________________, but remember KW, like all other constants K, is temperature dependent.

• Increasing the temperature will ______________ KW.

Constant

0oC

10 oC

50 oC

100 oC

KW

1.14 x 10-14

2.93 x 10-14

5.476 x 10-14

5.13 x 10-13

EXAMPLE: Determine the concentration of hydronium ions for a neutral solution at 25oC and at 50oC.



CONCEPT: CALCULATING pH and pOH OF STRONG SPECIES

STRONG ACIDS & BASES are considered _________________ Electrolytes so they ionize completely in water.

HCl (aq) H2O

H+ (aq) + Cl – (aq)

NaOH (aq) H2O

Na+ (aq) + OH – (aq)

EXAMPLE 1: Calculate the pH of a 0.0782 M solution of CaH2.

EXAMPLE 2: Calculate the pH of a 0.000550 M HBr solution to the correct number of significant figures.

a) 3.3

b) 3.26

c) 3.260

d) 3.2596

e) All are correct

PRACTICE: Calculate the pH of 50.00 mL of 4.3 x 10-7 M H2SO4.



CONCEPT: CALCULATING pH and pOH OF WEAK SPECIES

WEAK ACIDS & BASES are considered __________________ Electrolytes so they don’t completely ionize in water.

HF + H2O F – (aq) + H3O+ (aq)

NH3 + H2O NH4+ (aq) + OH – (aq)

EXAMPLE 1: Pryridine, an organic molecule, is a very common weak base.

C5H5N (aq) + H2O (l) C5H5NH+ (aq) + OH- (g)

Assume you have a 0.0225 M aqueous solution of pyridine, C5H5N, determine its pH. The Kb value for the

compound is 1.5 x 10-9.



PRACTICE: CALCULATING pH and pOH OF WEAK SPECIES (CALCULATIONS 1)

EXAMPLE: An unknown weak base has an initial concentration of 0.750 M with a pH of 8.03. Calculate its equilibrium base constant.

PRACTICE: Determine the pH of a solution made by dissolving 6.1 g of sodium cyanide, NaCN, in enough water to make a

500.0 mL of solution. (MW of NaCN = 49.01 gmol

). The Ka value of HCN is 4.9 x 10-10.



CONCEPT: ACID & BASE CONSTANTS As you might already realize, there are relatively few strong acids. The great majority of acids are weak acids.

Consider a weak monoprotic acid, HA, and its ionization in water:

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

The equilibrium expression for this ionization would be:

Ka =

ProductsReac tan ts

=

Where Ka represents the _________________________________________ and it measures the strength of weak acids.

When looking at weak bases we don’t use Ka, but instead _______, which represents the __________________________.

• The relationship between Ka and Kb can be expressed with the following equation:

In general, the __________________ the Ka the stronger the acid and the __________________ the concentration of H+.

In general, the __________________ the pKa the stronger the acid and the __________________ the concentration of H+.

PRACTICE: If the Kb of NH3 is 1.76 x 10-5, determine the acid dissociation constant of its conjugate acid.

KW =Ka ⋅Kb



PRACTICE: ACID & BASE CONSTANTS (CALCULATIONS 1)

EXAMPLE 1: Knowing that HF has a higher Ka value than CH3COOH, determine, if possible, in which direction the following equilibrium lies.

HF (aq) + CH3COO – (aq) F – (aq) + CH3COOH (aq)

a) Equilibrium lies to the left.

b) Equilibrium lies to the right.

c) Equilibrium is equal and balanced.

d) Not enough information given.

EXAMPLE 2: What is the equilibrium constant for the following reaction and determine if reactants or products are favored.

HCN (aq) + ClO2 – (aq) CN – (aq) + HClO2 (aq)

The acid dissociation constant of HCN is 4.9 x 10-10 and the acid dissociation of HClO2 is 1.1 x 10-2.

HCN (aq) + H2O (aq) CN – (aq) + H3O+ (aq)

HClO2 (aq) + H2O (aq) ClO2 – (aq) + H3O+ (aq)



PRACTICE: ACID & BASE CONSTANTS (CALCULATIONS 2)

EXAMPLE: Which of the following solutions will have the lowest pH?

a) 0.25 M HC2F3O2

b) 0.25 M HIO4

c) 0.25 M HC3H5O3

d) 0.25 M H2CO3

e) 0.25 M HSeO4–

PRACTICE 1: Which Bronsted-Lowry base has the greatest concentration of hydroxide ions?

a) C2H8N2 (Kb = 8.3 x 10-5)

b) C5H5N (Kb = 1.7 x 10-9)

c) (CH3)3N (Kb = 1.0 x 10-6)

d) C3H7NH2 (Kb = 3.5 x 10-4)

e) C6H5NH2 (Kb = 3.9 x 10-10)

PRACTICE 2: Which Bronsted-Lowry acid has the weakest conjugate base?

a) HCNO (Ka = 2.0 x 10-4)

b) HF (Ka = 3.5 x 10-4)

c) HN3 (Ka = 2.5 x 10-5)

d) H2CO3 (Ka = 4.3 x 10-7)



CONCEPT: ACID & BASE NEUTRALIZATION

When an acid neutralizes a base an ionic compound called a _______________ is formed.

• These solutions can be neutral, acidic or basic, based on acid-base properties of the cations and anions formed.

RULES FOR IDENTIFYING YOUR IONS

CATIONS (POSITIVE IONS)

1) Transition Metals: If your transition metal has a charge of +2 or higher it is acidic. If the charge is less than +2 then it is

neutral.

EX:

2) Main-Group Metals: If your main-group metal has a charge of +3 or higher it is acidic. If the charge is less than +3 then

it is neutral.

EX:

3) Positive Amines are acidic.

EX:

ANIONS (NEGATIVE IONS)

1) NEGATIVE ION: If you have a negative ion then add an H+ to it. If you create a weak acid then your negative ion is

basic.

EX:



PRACTICE: ACID & BASE NEUTRALIZATION 1

EXAMPLE: Determine if each of the following compounds will create an acidic, basic or neutral solution.

a) NaOCl b) PbCl4

PRACTICE 1: Determine if each of the following compounds will create an acidic, basic or neutral solution.

a) LiC2H3O2 b) C6H5NH3Br


a) Co(HSO4)2 b) Sr(HSO3)2


a) C3H7NH3F



PRACTICE: ACID & BASE NEUTRALIZATION 2

EXAMPLE 1: Determine whether each compound will become more soluble in an acidic solution.

a) NaBr b) LiCl c) KIO

EXAMPLE 2: Determine the pH of a 0.50 M NH4Cl solution. The Kb of NH3 is 1.75 x 10-5.

PRACTICE: Determine the pH of a 0.55 M NaCN solution. The Ka of hydrocyanic acid, HCN, is 4.9 x 10-10.



CONCEPT: DIPROTIC ACIDS

Diprotic acids and bases are compounds that can donate or accept _______ H+ ion.

For diprotic acids ________ their equations can be illustrated by:

For diprotic bases ________ their equations can be illustrated by:

Based on these equations the relationship between the different forms of diprotic species are:

As a result of these equations for diprotic acids and bases the relationship between Ka and Kb will be:

Ka1 ⋅Kb2 =Kw

Ka2 ⋅Kb1 =Kw

When dealing with diprotic acids:

1) H2A can be treated as a monoprotic acid and we use ________ can be used to find pH.

2) HA – represents the intermediate form and we use ________ can be used to find pH.

3) A2– represents the basic form and we use ________ can be used to find pH.

H2A (aq) + H2O (l) HA– (aq) + H3O+ (aq) Ka1 =ProductsReac tan ts

=

HA– (aq) + H2O (l) A2– (aq) + H3O+ (aq) Ka2 =ProductsReac tan ts

=

A2– (aq) + H2O (aq) HA – (aq) + OH – (aq) Kb1 =ProductsReac tan ts

=

HA – (aq) + H2O (aq) H2A (aq) + OH – (aq) Kb2 =ProductsReac tan ts

=



PRACTICE: DIPROTIC ACIDS CALCULATIONS 1

EXAMPLE 1: Sulfurous acid, H2SO3, represents a diprotic acid with a Ka1 = 1.6 x 10-2 and Ka2 = 4.6 x 10-5. Calculate the pH and concentrations of H2SO3, HSO3– and SO32– when given 0.250 M H2SO3.

EXAMPLE 2: Determine the pH of 0.115 M Na2S. Hydrosulfuric acid, H2S, contains Ka1 = 1.0 x 10-7 and Ka2 = 9.1 x 10-8.



CONCEPT: POLYPROTIC ACIDS

Our understanding of diprotic acids and bases can be used to understand polyprotic acids and bases.

For polyprotic acids ________ their equations can be illustrated by:

For polyprotic bases ________ their equations can be illustrated by:

As a result of these equations for polyprotic acids and bases the relationship between Ka and Kb will be:

Ka1 ⋅Kb3 =Kw

Ka2 ⋅Kb2 =Kw

Ka3 ⋅Kb1 =Kw

When dealing with polyprotic acids:

• H3A can be treated as a monoprotic acid and we use ________ can be used to find pH.

• A3– represents the basic form and we use ________ can be used to find pH.

H2A−

[H+ ] ≈ Ka1Ka2[ ]0 +Ka1Kw

Ka1 + [ ]0 HA2−

[H+ ] ≈ Ka2Ka3[ ]0 +Ka2Kw

Ka2 + [ ]0

H3A (aq) + H2O (l) H2A– (aq) + H3O+ (aq) Ka1 =ProductsReac tan ts

=

H2A– (aq) + H2O (l) HA2– (aq) + H3O+ (aq) Ka2 =ProductsReac tan ts

=

HA2– (aq) + H2O (l) A3– (aq) + H3O+ (aq) Ka3 =ProductsReac tan ts

=

A3– (aq) + H2O (l) HA2– (aq) + OH – (aq) Kb1 =ProductsReac tan ts

=

HA2– (aq) + H2O (l) H2A – (aq) + OH – (aq) Kb2 =

ProductsReac tan ts

=

H2A – (aq) + H2O (l) H3A (aq) + OH – (aq) Kb3 =

ProductsReac tan ts

=



PRACTICE: POLYPROTIC ACIDS CALCULATIONS

EXAMPLE 1: Determine the pH of 0.300 M sodium hydrogen phosphate, Na2HPO4. Phosphoric acid, H3PO4, contains Ka1 = 7.5 x 10-3, Ka2 = 6.2 x 10-8 and Ka3 = 4.2 x 10-13.

EXAMPLE 2: Determine the pH of 0.300 M citric acid, H3C6H5O7 it possesses Ka1 = 7.4 x 10-4, Ka2 = 1.7 x 10-5 and Ka3 =

4.0 x 10-7.



CHEMISTRY - SILBERBERG 8E CH.18 - ACID-BASE...

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