Chapter 16 Acids and Bases

Acidsand

Bases

Chapter 16Acids and Bases

John D. Bookstaver

St. Charles Community College

St. Peters, MO

2006, Prentice Hall, Inc.

Chemistry, The Central Science, 10th editionTheodore L. Brown; H. Eugene LeMay, Jr.;

and Bruce E. Bursten

Acidsand

Bases

HW • CHAPTER 16 –ACID-BASE EQUILIBRIUM• Bronsted Lowry 15,16,17,19,23,25• Kw 29• pH scale 35, 37• Strong acids and bases 41, 43 (a to c)• 45, 47• WEAK ACIDS 53, 55, 57, 61-a , 65• 73, 75, 81, 85, 87• 91 to 110 red problems ONLY

Acidsand

Bases

Electrolytes

• Substances that dissociate into ions when dissolved in water.

• A nonelectrolyte may dissolve in water, but it does not dissociate into ions when it does so.

Acidsand

Bases

Electrolytes and Nonelectrolytes

Soluble ionic compounds tend to be electrolytes.

Acidsand

Bases

Electrolytes and Nonelectrolytes

Molecular compounds tend to be nonelectrolytes, except for acids and bases.

Acidsand

Bases

Electrolytes

• A strong electrolyte dissociates completely when dissolved in water.

• A weak electrolyte only dissociates partially when dissolved in water.

Acidsand

Bases

Strong Electrolytes Are…

• Strong acids

Acidsand

Bases

Strong Electrolytes Are…

• Strong acids• Strong bases

Acidsand

Bases

Some Definitions

• ArrheniusAcid: Substance that, when dissolved

in water, increases the concentration of hydrogen ions.

Base: Substance that, when dissolved in water, increases the concentration of hydroxide ions.

Acidsand

Bases

Some Definitions

• Brønsted–LowryAcid: Proton donorBase: Proton acceptor

Acidsand

Bases

A Brønsted–Lowry acid…

…must have a removable (acidic) proton.

A Brønsted–Lowry base…

…must have a pair of nonbonding electrons.

Acidsand

Bases

If it can be either…

...it is amphiprotic.

HCO3−

HSO4−

H2O

Acidsand

Bases

What Happens When an Acid Dissolves in Water?

• Water acts as a Brønsted–Lowry base and abstracts a proton (H+) from the acid.

• As a result, the conjugate base of the acid and a hydronium ion are formed.

Acidsand

Bases

Conjugate Acids and Bases:

• From the Latin word conjugare, meaning “to join together.”

• Reactions between acids and bases always yield their conjugate bases and acids.

Acidsand

Bases

Conjugate Acid-Base Pairs• Whatever is left of the acid after the proton is donated is

called its conjugate base.• Similarly, whatever remains of the base after it accepts a

proton is called a conjugate acid.• Consider

– After HA (acid) loses its proton it is converted into A- (base). Therefore HA and A- are a conjugate acid-base pair.

– After H2O (base) gains a proton it is converted into H3O+ (acid). Therefore, H2O and H3O+ are a conjugate acid-base pair.

• Conjugate acid-base pairs differ by only one proton.

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

Acidsand

Bases

Example – Identify the acid and the base in each equation, and identify each acid-base pair.

1. HNO3 + NH3 NO3- + NH4

+

2. CH3COOH + OH- H2O + CH3COO-

Identify the acid and base for the reverse reaction in each example.

Acidsand

Bases

Acid and Base Strength

• Strong acids are completely dissociated in water. Their conjugate bases are quite

weak.

• Weak acids only dissociate partially in water. Their conjugate bases are weak

bases.The weaker the acid the stronger its conjugate base.

Acidsand

Bases


• Substances with negligible acidity do not dissociate in water.Their conjugate bases are

exceedingly strong.

Acidsand

Bases

Relative Strengths of Acids and Bases• The stronger the acid, the weaker the conjugate base. • H+ is the strongest acid that can exist in equilibrium in

aqueous solution.• OH- is the strongest base that can exist in equilibrium in

aqueous solution.

Acidsand

Bases


In any acid-base reaction, the equilibrium will favor the reaction that moves the proton to the stronger base.

HCl(aq) + H2O(l) H3O+(aq) + Cl−(aq)

H2O is a much stronger base than Cl−, so the equilibrium lies so far to the right K is not measured (K>>1).

Acidsand

Bases


Acetate is a stronger base than H2O, so the equilibrium favors the left side (K<1).

HC2H3O2(aq) + H2O H3O+(aq) + C2H3O2−(aq)

Acidsand

Bases

Autoionization of Water

• As we have seen, water is amphoteric.

• In pure water, a few molecules act as bases and a few act as acids.

• This is referred to as autoionization.

H2O(l) + H2O(l) H3O+(aq) + OH−(aq)

Acidsand

Bases

Ion-Product Constant

• The equilibrium expression for this process is

Kc = [H3O+] [OH−]

• This special equilibrium constant is referred to as the ion-product constant for water, Kw.

• At 25°C, Kw = 1.0 10−14

Acidsand

Bases

• In most solutions [H+(aq)] is quite small.• We define

• In neutral water at 25 C, pH = pOH = 7.00.• In acidic solutions, [H+] > 1.0 10-7, so pH < 7.00.• In basic solutions, [H+] < 1.0 10-7, so pH > 7.00.• The higher the pH, the lower the pOH, the more basic the

solution.

The pH ScaleThe pH Scale

]OHlog[pOH ]Hlog[]OHlog[pH -3

Acidsand

Bases

pH

• Therefore, in pure water,pH = −log (1.0 10−7) = 7.00

• An acid has a higher [H3O+] than pure water, so its pH is <7

• A base has a lower [H3O+] than pure water, so its pH is >7.

Acidsand

Bases

pH

These are the pH values for several common substances.

Acidsand

Bases

Other “p” Scales

• The “p” in pH tells us to take the negative log of the quantity (in this case, hydrogen ions).

• Some similar examples arepOH −log [OH−]pKw −log Kw

Acidsand

Bases

Watch This!

Because

[H3O+] [OH−] = Kw = 1.0 10−14,

we know that

−log [H3O+] + −log [OH−] = −log Kw = 14.00

or, in other words,

pH + pOH = pKw = 14.00

Acidsand

Bases

February 28• Measuring pH: a/b indicators

• pH meter

• Section 16.5 Strong a/b – key points

• Section 16.6 Weak acidsKa Problems a) Calculating Ka and %

Ionization from measured pH and initial concentration

Acidsand

Bases

HW

• WEAK ACIDS 53, 55, 57, 61-a , 65

Acidsand

Bases

Daily Quiz

• 1. Write the formation of nitrous acid from its anhydride

• 2. Calculate the pH of an aqueous solution of LiOH that has a pH of 12.5

Acidsand

Bases

How Do We Measure pH?

For more accurate measurements, one uses a pH meter, which measures the voltage in the solution.

Acidsand

Bases

• Most pH and pOH values fall between 0 and 14.• There are no theoretical limits on the values of pH or

pOH. (e.g. pH of 2.0 M HCl is -0.301.)

Examples:

1. Consider a solution with [H+] = 6.2 x 10-3.• Calculate the pH, pOH, and [OH-]• Is this solution acidic or basic?

2. Consider a solution with pOH = 13.65• Calculate the pH, [H+], and [OH-]• Is this solution acidic or basic?

Acidsand

Bases

Measuring pH• Most accurate method to measure pH is to use a pH

meter.• However, certain dyes change color as pH changes.

These are indicators.• Indicators are less precise than pH meters.• Many indicators do not have a sharp color change as a

function of pH.

Acidsand

Bases

How Do We Measure pH?

• For less accurate measurements, one can useLitmus paper

• “Red” paper turns blue above ~pH = 8

• “Blue” paper turns red below ~pH = 5

An indicator

Acidsand

Bases

Strong Acids

• You will recall that the seven strong acids are HCl, HBr, HI, HNO3, H2SO4, HClO3, and HClO4.

• These are, by definition, strong electrolytes and exist totally as ions in aqueous solution.

• For the monoprotic strong acids,

[H3O+] = [acid].

Acidsand

Bases

Strong Bases

• Strong bases are the soluble hydroxides, which are the alkali metal and heavier alkaline earth metal hydroxides (Ca2+, Sr2+, and Ba2+).

Acidsand

Bases

• Strong bases are strong electrolytes and dissociate completely in solution.

• The pOH of a strong base is given by the initial molarity of the hydroxide ion. Be careful of stoichiometry.

• In order for a hydroxide to be a base, it must be soluble. • Bases do not have to contain the OH- ion:

O2-(aq) + H2O(l) 2OH-(aq)

H-(aq) + H2O(l) H2(g) + OH-(aq)

N3-(aq) + 3H2O(l) NH3(aq) + 3OH-(aq)

Acidsand

Bases

Strong basic solutions

• Ionic metal oxides Na2O and CaO are used in industry to produce strong basic solutions.

• Find the pH of a solution formed by dissolving 0.01 mol of Na2O in enough water to produce a liter of solution.

• 12.3

Acidsand

Bases

Examples: Calculate the pH and pOH of each:

1. 0.25 M NaOH

2. 12.00 M HCl

3. 6.00 M KOH

4. 0.0050 M HNO3

5. 6.5 x 10-12 M RbOH

6. 0.10 M HClO4

Acidsand

Bases

Examples: Calculate the pH and pOH of each:

pH pOH

1. 0.25 M NaOH 13.40 0.60

2. 12.00 M HCl -1.08 15.08

3. 6.00 M KOH 14.78 -0.78

4. 0.0050 M HNO3 2.30 11.70

5. 6.5 x 10-12 M RbOH 7.00 7.00

6. 0.10 M HClO4 1.00 13.00

Acidsand

Bases

• Weak acids are only partially ionized (dissociated) in solution.

• There is a mixture of ions and unionized acid in solution.• Therefore, weak acids are in equilibrium:

Weak AcidsWeak Acids


HA(aq) H+(aq) + A-(aq)

]HA[]A][OH[ -

3

aK

]HA[]A][H[ -

aK

Acidsand

Bases

Dissociation Constants

• For a generalized acid dissociation,

the equilibrium expression would be

• This equilibrium constant is called the acid-dissociation constant, Ka.

[H3O+] [A−][HA]

Kc =

HA(aq) + H2O(l) A−(aq) + H3O+(aq)

Acidsand

Bases

Dissociation Constants• The larger the Ka the stronger the acid (i.e. the

more ions are present at equilibrium relative to unionized molecules).

• If Ka >> 1, then the acid is completely ionized and the acid is a strong acid

Acidsand

Bases

Percent Ionization• Percent ionization is another method to assess acid

strength.• For the reaction


100]HA[

]OH[ionization %

0

3

eq

Acidsand

Bases

Calculating Ka from pH

• Weak acids are simply equilibrium calculations.• The pH is used to calculate the equilibrium concentration

of H+.– Write the balanced chemical equation clearly showing the

equilibrium.

– Write the equilibrium expression.

– Use an ICE Chart / find the [H+] from the pH

– Find the value for Ka.

Acidsand

Bases

Calculating Ka from the pH

• The pH of a 0.10 M solution of formic acid, HCOOH, at 25°C is 2.38. Calculate Ka for formic acid at this temperature.

• We know that

[H3O+] [COO−][HCOOH]

Ka =

Acidsand

Bases


• The pH of a 0.10 M solution of formic acid, HCOOH, at 25°C is 2.38. Calculate Ka for formic acid at this temperature.

• To calculate Ka, we need the equilibrium concentrations of all three things.

• We can find [H3O+], which is the same as [HCOO−], from the pH.

Acidsand

Bases


pH = −log [H3O+]

2.38 = −log [H3O+]

−2.38 = log [H3O+]

10−2.38 = 10log [H3O+] = [H3O+]

4.2 10−3 = [H3O+] = [HCOO−]

Acidsand

Bases


Now we can set up a table…

[HCOOH], M [H3O+], M [HCOO−], M

Initially 0.10 0 0

Change −4.2 10-3 +4.2 10-

3

+4.2 10−3

At Equilibrium

0.10 − 4.2 10−3

= 0.0958 = 0.10

4.2 10−3 4.2 10−3

Acidsand

Bases


[4.2 10−3] [4.2 10−3][0.10]

Ka =

= 1.8 10−4

Acidsand

Bases

Calculating Percent Ionization

• Percent Ionization = 100

• In this example

[H3O+]eq = 4.2 10−3 M

[HCOOH]initial = 0.10 M

[H3O+]eq

[HA]initial

Acidsand

Bases

Calculating Percent Ionization

Percent Ionization = 1004.2 10−3

0.10

= 4.2%

Acidsand

Bases

Examples: Calculate Ka.

1. A 0.25 M solution of a hypothetical acid, HA, has a pH of 4.06.

3.0 x 10-8

2. A 0.10 M solution of acetic acid (CH3COOH) has a pH of 2.87.

1.8 x 10-5

Acidsand

Bases

Using Ka to calculate pH from initial concentration

5% Rule. When to make approximations and when to use quadratic equation.

Percent ionizationPolyprotic acids

Acidsand

Bases

Using Ka to Calculate pH

• Percent ionization relates the equilibrium H+ concentration, [H+]eqm, to the initial HA concentration, [HA]0.

• The higher percent ionization, the stronger the acid.• Percent ionization of a weak acid decreases as the molarity

of the solution increases.• For acetic acid, 0.05 M solution is 2.0 % ionized whereas a

0.15 M solution is 1.0 % ionized.• Use ICE Charts to determine [H+]• -log [H+] = pHSubstitute into the equilibrium constant expression and solve.

Remember to turn x into pH if necessary.

Acidsand

Bases

Calculating pH from Ka

Calculate the pH of a 0.30 M solution of acetic acid, HC2H3O2, at 25°C.

HC2H3O2(aq) + H2O(l) H3O+(aq) + C2H3O2−(aq)

Ka for acetic acid at 25°C is 1.8 10−5.

Acidsand

Bases


The equilibrium constant expression is

[H3O+] [C2H3O2−]

[HC2H3O2]Ka =

Acidsand

Bases


We next set up a table…

[C2H3O2], M [H3O+], M [C2H3O2−], M

Initially 0.30 0 0

Change −x +x +x

At Equilibrium 0.30 − x 0.30 x x

We are assuming that x will be very small compared to 0.30 and can, therefore, be ignored.

Acidsand

Bases


Now,

(x)2

(0.30)1.8 10−5 =

(1.8 10−5) (0.30) = x2

5.4 10−6 = x2

2.3 10−3 = x

Acidsand

Bases


pH = −log [H3O+]pH = −log (2.3 10−3)pH = 2.64

Acidsand

Bases

5% RULE

• As a general rule if x is greater than 5% than the initial value we can not considered negligible and a quadratic equation has to be solved.

• If x is less than 5% it is ok to considered negligible.

• SOLVE THE PROBLEM CONSIDERING X NEGLIGIBLE AND CHECK!

Acidsand

Bases

Examples – Calculate pH, pOH and percent dissociation for each of the following:

1. 0.35 M HF (Ka = 7.2 x 10-4)

pH = 1.80, pOH = 12.20, 4.5% ionized

2. 0.75 M HF

pH = 1.63, pOH = 12.37, 3.1% ionized

3. 0.62 M NH4+ (from NH4Cl) (Ka = 5.6 x 10-10)

pH = 4.73, pOH = 9.27, .0030% ionized

4. 1.08 M HNO2 (Ka = 4.0 x 10-4)

pH = 1.68, pOH = 12.32, 1.9% ionized

5. 6.00 M HNO3

pH = -0.78, pOH = 14.78, 100% ionized

Acidsand

Bases

Percent Ionization• Percent ionization is another method to assess acid

strength.• For the reaction


100]HA[

]OH[ionization %

0

3

eq

Acidsand

Bases

Acidsand

Bases

Polyprotic Acids

• Have more than one acidic proton.

• If the difference between the Ka for the first dissociation and subsequent Ka values is 103 or more, the pH generally depends only on the first dissociation.

Acidsand

Bases

Dilute Sulfuric Acid Solutions

• First Dissociation is Strong (large Ka)

• Second Dissociation is weak (Ka2 = 1.2 x 10-2)Make ICE chart for second dissociation – initial [H+]

is equal to initial [HSO4-]

Determine x (using quadratic if necessary)Solve for [H+]eq

Example – Calculate the pH of 0.010 M H2SO4

Acidsand

Bases

Polyprotic Acids

Acidsand

Bases

• Weak bases

• Relationship between Ka and Kb

• Hydrolysis of salts

Acidsand

Bases

Weak Bases

Bases react with water to produce hydroxide ion.

Acidsand

Bases

Weak Bases

The equilibrium constant expression for this reaction is

[HB] [OH−][B−]

Kb =

where Kb is the base-dissociation constant.

Acidsand

Bases

Kb can be used to find [OH−] and, through it, pH.

Acidsand

Bases

Types of Weak Bases• Bases generally have lone pairs or negative charges in

order to attack protons.• Most neutral weak bases contain nitrogen.• Amines are related to ammonia and have one or more N-

H bonds replaced with N-C bonds (e.g., CH3NH2 is methylamine).

• Anions of weak acids are also weak bases. Example: ClO- is the conjugate base of HClO (weak acid):

ClO-(aq) + H2O(l) HClO(aq) + OH-(aq) Kb = 3.3 x 10-7

Acidsand

Bases

Calculating pH and pOH of Weak Base Solutions

• Use ICE Chart to determine [OH-]eq

• Use [OH-] to find pOH, and thus pH

Examples – Calculate pOH and pH of each of the following solutions:

1. 0.15 M NH3 (Kb = 1.8 x 10-5)

pH = 11.22 pOH = 2.78

2. 0.38 M HONH2 (Kb = 1.1 x 10-8)

pH = 9.81 pOH = 4.19

Acidsand

Bases

pH of Basic Solutions

What is the pH of a 0.15 M solution of NH3?

[NH4+] [OH−]

[NH3]Kb = = 1.8 10−5

NH3(aq) + H2O(l) NH4+(aq) + OH−(aq)

Acidsand

Bases


Tabulate the data.

[NH3], M [NH4+], M [OH−], M

Initially 0.15 0 0

At Equilibrium 0.15 - x 0.15

x x

Acidsand

Bases


(1.8 10−5) (0.15) = x2

2.7 10−6 = x2

1.6 10−3 = x2

(x)2

(0.15)1.8 10−5 =

Acidsand

Bases


Therefore,

[OH−] = 1.6 10−3 M

pOH = −log (1.6 10−3)

pOH = 2.80

pH = 14.00 − 2.80

pH = 11.20

Acidsand

Bases

• Relationship between strength of acid and conjugate base:

• When two reactions are added to give a third, the equilibrium constant for the third reaction is the product of the equilibrium constants for the first two:

Reaction 1 + reaction 2 = reaction 3

has

Relationship Between KRelationship Between Kaa

and Kand Kbb

213 KKK

Acidsand

Bases

• For a conjugate acid-base pair

• Therefore, the larger the Ka, the smaller the Kb. That is, the stronger the acid, the weaker the conjugate base.

• Taking negative logarithms:

baw KKK

baw pKpKpK

Acidsand

Bases

Acidsand

Bases

• Nearly all salts are strong electrolytes.• Salts exist entirely of ions in solution.• Acid-base properties of salts are a consequence of the

reaction of their ions in solution.

• The reaction in which ions produce H+ or OH- in water is

called hydrolysis.

Acid-Base Properties of Acid-Base Properties of Salt SolutionsSalt Solutions

Acidsand

Bases

HYDROLYSIS OF SALTS

• STRONG PARENT ACID-PARENT BASE NEUTRAL SALT

• STRONG PARENT ACID – WEAK PARENT BASE ACIDIC SALT

• WEAK PARENT ACID – STRONG PARENT BASE BASIC SALT

• WEAK AND WEAK DEPENDS ON WHICH IS GREATER Ka OR Kb

Acidsand

Bases

An Anion’s Ability to React with Water• Anions, X-, can be considered conjugate bases from

acids, HX.• For X- comes from a strong acid, then it is neutral.• If X- comes from a weak acid, then

The pH of the solution can be calculated using equilibrium!

Anions from strong acids are neutral.Anions from weak acids are basic.

X-(aq) + H2O(l) HX(aq) + OH-(aq)

Acidsand

Bases

Reactions of Cations with Water

• Cations with acidic protons (like NH4

+) will lower the pH of a solution.

• Most metal cations that are hydrated in solution also lower the pH of the solution.

Acidsand

Bases

Reactions of Cations with Water

• Attraction between nonbonding electrons on oxygen and the metal causes a shift of the electron density in water.

• This makes the O-H bond more polar and the water more acidic.

• Greater charge and smaller size make a cation more acidic.

Acidsand

Bases

Hydrolysis of Metal Ions• Metal ions are positively charged and attract water

molecules (via the lone pairs on O).• The higher the charge, the smaller the metal ion and the

stronger the M-OH2 interaction.

• Hydrated metal ions act as acids:

• The pH increases as the size of the ion increases (e.g. Ca2+ vs. Zn2+) and as the charge increases (Na+ vs. Ca2+ and Zn2+ vs. Al3+).

Fe(H2O)63+(aq) Fe(H2O)5(OH)2+(aq) + H+(aq)

Ka = 2 x 10-3

Acidsand

Bases

Hydrolysis of Metal Ions

Acidsand

Bases

A Cation’s Ability to React with Water• Polyatomic cations with ionizable protons can be

considered conjugate acids of weak bases.

• Some metal ions react in solution to lower pH.

Combined Effect of Cation and Anion in Solution• An anion from a strong acid has no acid-base properties.• An anion that is the conjugate base of a weak acid will

cause an increase in pH (basic ion).

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

Acidsand

Bases

• A cation that is the conjugate acid of a weak base will cause a decrease in the pH of the solution (acidic ion).

• Metal ions will cause a decrease in pH except for the alkali metals and alkaline earth metals.

• When a solution contains both cations and anions from weak acids and bases, use Ka and Kb to determine the final pH of the solution.

• Use ICE charts to calculate pH and pOH (as for weak acid and weak base solutions)

Acidsand

Bases

Effect of Cations and Anions

1. An anion that is the conjugate base of a strong acid will not affect the pH.

2. An anion that is the conjugate base of a weak acid will increase the pH.

3. A cation that is the conjugate acid of a weak base will decrease the pH.

Acidsand

Bases

Effect of Cations and Anions

4. Cations of the strong Arrhenius bases will not affect the pH.

5. Other metal ions will cause a decrease in pH.

6. When a solution contains both the conjugate base of a weak acid and the conjugate acid of a weak base, the affect on pH depends on the Ka and Kb values.

Acidsand

Bases

Examples – State whether each will be acidic, basic, or neutral in aqueous solution:

1. NaClO

2. Fe(NO3)3

3. KBr

4. LiClO4

5. NaHCO3

6. CaCl27. NH4NO3

8. NiI2

Acidsand

Bases

Examples – State whether each will be acidic, basic, or neutral in aqueous solution:

1. NaClO basic

2. Fe(NO3)3 acidic

3. KBr neutral

4. LiClO4 neutral

5. NaHCO3 basic

6. CaCl2 neutral

7. NH4NO3 acidic

8. NiI2 acidic

Acidsand

Bases

Examples – Calculate pH and pOH for each solution:

1. 0.25 M NH4Cl

2. 0.60 M KBr

3. 0.18 M KC2H3O2

4. 0.50 M NaHCO3

Acidsand

Bases

Factors that Affect Acid Strength

Consider H-X. For this substance to be an acid we need:• H-X bond to be polar with H+ and X- (if X is a metal

then the bond polarity is H-, X+ and the substance is a base),

• the H-X bond must be weak enough to be broken,• the conjugate base, X-, must be stable.• The 3 factors are then 1)Polarity of H—X bond

2)Bond strenght

3)Stability of conjugate base

Acidsand

Bases

Binary Acids

H—X strength most important factor

The bigger the atom X the weaker the strength of the bond the more acidic the acid.

HF is a weak acid because the bond energy is high, HCl is more acidic because the bond strength is weaker.

• Acid strength increases down a group.• In a period the acidity strength increases with the

increase in the polarization of the bond. That depends on the electronegativity of the element.

• In a period the acidity increases across a period

Acidsand

Bases

Base Strength

• Conversely, base strength decreases across a period and down a group.

• The electronegativity difference between C and H is so small that the C-H bond is non-polar and CH4 is neither an acid nor a base.

Acidsand

Bases

Binary Acids

Acidsand

Bases

Oxyacids

In oxyacids, in which an OH is bonded to another atom, Y, the more electronegative Y is, the more acidic the acid.

Acidsand

Bases

For a series of oxyacids, acidity increases with the number of oxygens.

Acidsand

Bases

Carboxilic Acids

Resonance in the conjugate bases of carboxylic acids stabilizes the base and makes the conjugate acid more acidic.

Acidsand

Bases

The strength of carboxilic acid increase as the number of

electronegative atoms in the acid increases.

Trifluoroacetic Ka=5 x 10 -1

Acidsand

Bases

• Brønsted-Lowry acid is a proton donor.• Focusing on electrons: a Brønsted-Lowry acid can be

considered as an electron pair acceptor.• Lewis acid: electron pair acceptor (the acceptor in a

coordinate bond).• Lewis base: electron pair donor ( electron rich – the

donor in a coordinate bond).• Note: Lewis acids and bases do not need to contain

protons.• Therefore, the Lewis definition is the most general

definition of acids and bases.

Lewis Acids and BasesLewis Acids and Bases

Acidsand

Bases

• Lewis acids generally have an incomplete octet (e.g. BF3).

• Transition metal ions are generally Lewis acids.• Lewis acids must have a vacant orbital (into which the

electron pairs can be donated).• Compounds with -bonds can act as Lewis acids:

H2O(l) + CO2(g) H2CO3(aq)

Acidsand

Bases

Lewis Acids

• Lewis acids are defined as electron-pair acceptors.

• Atoms with an empty valence orbital can be Lewis acids.

Acidsand

Bases

Lewis Bases

• Lewis bases are defined as electron-pair donors.• Anything that could be a Brønsted–Lowry base is

a Lewis base.• Lewis bases can interact with things other than

protons, however.

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Chapter 16 Acids and Bases