Acids, Bases, and Salts

Chapter 14, 15

Some Properties of Acids1. The word acid comes from the Latin

word acere, which means "sour." All acids taste sour.

2. In 1663, Robert Boyle wrote that acids would make a blue vegetable dye called "litmus" turn red.

3. Acids react with bases (they destroy the chemical properties of bases).

4. Acids conduct an electric current. 5. Upon chemically reacting with an

active metal, acids will evolve hydrogen gas (H2).

Some Properties of Bases

1. The word "base" has a more complex history and its name is not related to taste. All bases taste bitter.

2. Bases are substances which will restore the original blue color of litmus after having been reddened by an acid.

3. Bases destroy the chemical properties of acids (will react with acids)

4. Bases will conduct an electric current.5. Bases feel “slippery” (soap, bleach) on

your skin.

This is because they dissolve the fatty acids and oils from your skin and this

cuts down on the friction between your fingers as you rub them together.

Some Properties of Salts

1. A salt is the combination of an anion (- ion) and a cation (+ ion).

2. Salts are products of the reaction between acids and bases.

3. Solid salts are usually crystalline.4. If a salt dissolves in water solution, it

usually dissociates into the anions and cations that make up the salt (depends on Ksp)

The Acid Base Theory

The three main theories regarding acids and bases are:

1. Arrhenius 3. Lewis

2. Brønsted-Lowry

Arrhenius Theory – late 1890s

DEFINITIONS:Acid - any substance which donates

hydrogen ions (H+) to water (produces hydronium ions, H3O+):

HA → H+ + A¯ Base - any substance which produces

hydroxide ions (OH¯) in water.XOH → X+ + OH¯

When acids and bases react, they neutralize each other, forming water and a salt:

HA + XOH → H2O + XA

Problems with Arrhenius Theory

The theory did not explain why ammonia (NH3) was a base.

The theory only considers water as a solvent. We know that an acid added to benzene will not dissociate. Solvents are crucial to acid definition.

The end result of mixing certain acids and bases can be a slightly acidic or basic solution. Arrhenius had no explanation for this phenomenon (degrees of acidity).

Brønsted – Lowry Theory – Early 1920s

Two chemists, independent of one another, proposed a new definition of an acid and a base:

An acid is a substance from which a proton can be removed (donates protons).

A base is a substance that can remove a proton from an acid (proton acceptor).

*This definition does not require acids and bases to be in aqueous solutions.

Reactions Based on Bronsted - Lowry

Which are the acids and bases?: HCl + H2O → H3O+ + Cl¯

HCl - this is an acid, because it has a proton available to be transferred (it can give a proton).

H2O - this is a base, since it gets the proton that the acid lost (it has the capacity to accept a proton).

Conjugate acid-base pairs

Example: HCl + H2O → H3O+ + Cl¯

Notice that each pair (HCl and Cl¯ as well as H2O and H3O+ differ by one proton (H+). These pairs are called conjugate pairs.

Example: HNO3 + H2O → H3O+ + NO3¯

The acids are HNO3 and H3O+ and the bases are H2O and NO3¯. What are the pairs?

Bases and Conjugate Acid

Base NameConjugate acid Name

C2H3O2- Acetate ion CH3COOH Acetic acid

NH3 Ammonia NH4+ Ammonium

H2PO4- Dihydrogen

phosphate ionH3PO4 Phosphoric

acid

HSO4- Hydrogen sulfate

ionH2SO4 Sulfuric acid

OH- Hydroxide ion H20 water

NO3- Nitrate ion HNO3 Nitric acid

H2O water H30+ Hydronium ion

Lewis Theory –Early 1920s

Remember drawing Lewis Dot Structures for ionic and covalent compounds?

Lewis Theory focuses on the nature of electrons rather than proton transfer.

DEFINITIONS: An acid is an electron pair acceptor and

a base as an electron pair donor. Lewis Theory is much more general and

apply to reactions that do not involve hydrogen or hydrogen ions.

Lewis acid-base reaction:BF3 accepts an electron pair from

ammonia:

A Lewis acid must have an empty orbital to accept an electron pair.

A Lewis base must have a pair of unshared electrons that can be donated. Typical Lewis bases are OH-, H2O, NH3, Cl-, CN-… due to lone pair electrons.

Lewis AB reactions and Formation of Coordinate

ComplexesThe metal ion is a Lewis acid and the

ligands coordinated to the ion are Lewis bases.

Autoionization of Water

Strong Acids and Bases

Strong acids are those that ionized completely in water.

The dissociation of a strong base also looks like the diagram at the right in that it dissociates into positive and negative ions.

7 Strong Acids HNO3 - nitric acid

HCl - hydrochloric acidHBr - hydrobromic acidHI - hydroiodic acid

H2SO4- sulfuric acid

HClO4 - perchloric acid HClO3 - chloric acid (wanna be)

•Strong acids are assumed to ionize completely (100%) in water. They exist as H3O+ ions in water. This is known as “the leveling effect”. Water has a greater affinity for H+

than the conjugate bases do.

ANIMATION LINKS

Acid ionization equilibrium demo

Weak Acids and Bases

Some acids and bases ionize only slightly in water.

These are considered weak.

The most important weak base is ammonia.

Balance of ions in solutions

Acidic NeutralSolution Solution

Strong BasesLiOH - lithium hydroxideNaOH - sodium hydroxideKOH - potassium hydroxideRbOH - rubidium hydroxideCsOH - cesium hydroxideBa(OH)2 - barium hydroxideSr(OH)2 - strontium hydroxideCa(OH)2 - calcium hydroxide

GROUP 1 hydroxides

SomeGROUP 2 hydroxides

•Strong bases also ionize completely in water, except for Sr(OH)2 and Ca(OH)2 which are only slightly soluble (remember Mg(OH)2 is insoluble).

Uncommonin labs becausetoo expensive

Polyprotic Acids Polyprotic acids dissociate in a stepwise

fashion with different Ka values for each step… In the second and subsequent ionizations the acids are always weak, whether or not the original is a strong or weak acid.

For most of these acids (ex. H3PO4), the first dissociation contributes the significant amount of H+ for pH calculations, and the rest are negligible (except for H2SO4 where second ionization is significant).

Naming Acids -REVIEW

-ide ending (elements): “hydro____ic acid”

ex. chloride (HCl): hydrochloric acid

-ate ending (polyatomics): “______ic acid”ex. chlorate (HClO3): chloric acid

-ite ending(polyatomics): “______ous acid”ex. chlorite (HClO2): chlorous acid

Net Ionic Equations -REVIEW

For aqueous acid-base reactions reactions, it is common to write equations in the net ionic form.

Standard form: NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l)

Ionic form: Na+(aq) + OH-(aq) + H+(aq) + Cl-(aq) Na+

(aq) + Cl-(aq) + H2O(l)

Net ionic form: OH-(aq) + H+ (aq) H2O(l)

(No spectator ions are included)

Things to remember when writing Net Ionic Equations

Binary Acids: HCl, HBr, and HI are strong: all other binary acids and HCN are weak. Strong acids are written in ionic form; weak acids are written in molecular form.

Ternary Acids: If the number of oxygen atoms in an inorganic acid molecule exceeds the number of hydrogen atoms by two or more, the acid is strong (complete dissociation). We will consider all organic acids as weak.

Strong: HClO3, HClO4, H2SO4, HNO3

Weak: HClO, H3AsO4, H2CO3, H4SiO4, HNO2

Polyprotic Acids: (acids that contain more than one ionizable hydrogen atom. Ex: H2SO4, H3PO4, H2CO3).

Bases: Hydroxides of Group 1 and 2 elements (except Be(OH)2 and Mg(OH)2) are strong bases. All others including ammonia, hydroxlamine, and organic bases are weak.

Salts: Salts are written in ionic form if soluble, and in undissociated form if insoluble. *Know the solubility rules.

Oxides: Oxides are always written in molecular or undissociated form (ex: MgO).

Gases: Gases are always written in molecular form (ex: SO2).

Practice Net Ionic Equations

1. AgNO3 (aq) + H2SO4 (aq)

2. H4SiO4 (aq) + NaOH (aq)

3. HBr (aq) + KOH (aq)

1. Ag+ + HSO4- → AgHSO4(s)

2. H4SiO4 + OH- → H3SiO4- + H2O

3. H+ + OH- → H2O

Weak acids and bases will have Ka or Kb values less than one, but greater than water dissociation, Kw

Relationship between Ka and Kb

Ka x Kb = Kw

For any acid and it’s conjugate base, this relationship can be used to determine Ka or Kb.

Ex: NH3 + H2O ↔ NH4+ + OH-

NH4+ + H2O ↔ NH3 + H3O+

Kb(NH3)=[NH4+][OH-] Ka(NH4

+)=[NH3][H3O+]

[NH3] [NH4+]

Therefore, Ka x Kb = [OH-][H3O+]= Kw

Strength of Acid-base pairs

Strong acids yield WEAK conjugate bases… they have a low affinity for H+

Weak acids yield STRONG conjugate bases

Strong bases yield WEAK conjugate acids

Weak bases yield STRONG conjugate acids

pH Scale

The pH scale (potential hydrogen scale) is a measure of hydronium ion (H3O+) concentration.

Hydronium ion concentration indicates acidity. Each increase in pH # means a 10-fold decrease in [H+].

The higher the [H3O+], the higher the acidity.

Soren Sorenson (1868-1939) invented the pH scale while creating a way to test the acidity of beer. Beer has a pH of about 4.5.

pH scale and [H+]

Calculating pH The concentration (M or mol/L) of H3O+

is expressed in powers of 10, from 10-14 to 100.

Scientists use pH which is the negative log of [H3O+].

pH = -log[H3O+]

Note: The significant figures for logarithmic numbers are given after the decimal, and the numbers preceding the decimal give the exponent.

Calculating pH of a strong acid:

Ex: Given a solution of 0.50M HCl, what is the pH?

Step 1: Find [H3O+] in mol/L

0.50mol/L = 5.0 x 10-1 mol/LStep 2: Place value in equation and

solve. pH = -log[5.0 x 10-1] = 0.30

Practice pH Calculations Find pH of the following solutions if

[H3O+] is:

1. 1.00 x 10-3

2. 6.59 x 10-6

3. 9.47 x 10-10

Find [H3O+] if the pH is:

1. 6.678 3. 10.02. 2.533 4. 2.56

pOH

You can calculate the pH of a solution if you know the concentration of hydronium ion. [OH-]

If we use the ion product constant of water we can derive this equation:

[pH][pOH] = 1.00 x 10-14

Working with this equation leads to:pH + pOH = 14

Calculating pH of a strong base:

Ex: Find the pH of a solution with an [NaOH] of 1.0 x 10-8.

Step 1: Solve for [H3O+] in equation:

[H3O+] = 10-14

[OH-]Step 2: Place values in: [H3O+] = 10-14 = 10-6 M

[1.0 x 10-8]

Step 3: Solve for pH by placing [H3O+] in pH = –log[H3O+]

pH = -log(1.0 x 10-6)pH = 6.0

Practice pH Calculations Using pOH

Find the pH of the following solutions with [OH] of:

1. 1.00 x 10-4

2. 2.64 x 10-13

3. 5.67 x 10-2

4. 3.45 x 10-11

Calculating pH, pOH, [H+] and [OH-]

If one of the these values is known, all others can be found using the following relationships:

[OH-] [H+]

pH pOH pH + pOH = 14

[H+] [OH-] = Kw

pH=-log[H+] [H+]=10-pH pOH=-log[OH-] [OH-]=10-pOH

Calculating pH of a weak acid:

Ex: Find the pH of a 1.00 M HF solution (Ka=7.2 x10-4):

HF(aq) ↔ H+(aq) + F-(aq)

Ka= [H+][F-]

[HF]Use ICE box method: HF(aq) ↔ H+(aq) + F-(aq)

I 1.00M 0 0C -x +x +xE 1.00 -x x x

Ka= [H+][F-] = (x)(x) ≈ x2

[HF] (1.00-x) (1.00)

x2 = Ka(1.00) = (7.2 x10-4)(1.00)

x ≈ 2.7 x10-2

CHECK: 5% rule is valid since 2.7 x10-2/1.00 = 2.7%pH = -log[H+] = 1.57

Try AP 2005 practice problem!!

Acid-base properties of salts

Salt hydrolysis the reaction in which a salt dissolved in water produces an acid or basic solution (opposite of a neutralization reaction):

Ex1: AlCl3 + H2O → Al(OH)3 + 3HCl

A salt that contains the conjugate base

of a strong acid will produce a slightly acidic solution when dissolved in water.

weak base strong acid

Acid-base properties of salts

Ex2: NaC2H3O2 + H2O → NaOH + HC2H3O2

A salt that contains the conjugate

base of a weak acid will produce a slightly basic solution when dissolved in water.

What will happen when NaCl dissolves in water?

strong base weak acid

Percent dissociation

Percent dissociation = amount dissociated x 100%

initial concentration

Practice Problem: find Ka given % dissociation of a weak acid:

Find [H+] first using % dissociation Use formula Ka=x2/[acid]0 - x to find Ka

The effect of structure on acid-base properties

Relative acidity of oxyacids (hydrogen is attached to oxygen, and acid contains one other element):

Acid strength increases as the O-H bond is weakened (or with increasing # of oxygen atoms on central atom). HClO4>HClO3>HClO2>HClO

Bond strength of H-F is too strong (F is so small) to make it a strong acid (will not dissociate easily). HI>HBr>HCl>>HF

BUFFERS

A buffer solution is one which resists changes in pH when small quantities of an acid or a base are added to it.

How do buffer solutions work? A buffer solution has to contain

things which will remove any hydrogen ions or hydroxide ions that you might add to it - otherwise the pH will change.

buffer demo

ex: HF and NaFex: NH3 and NH4Cl

Calculating pH of a buffer Taking the log of the Ka equation

(rearranged to solve for H+)…

…. yields the Henderson-Hasselbalch equation:

This is used to find the pH of a buffer solution

H-H Special case… What happens with equimolar

amounts of acid and conjugate base ion?

This situation simplifies to pH = pKa, since log 1=0.

Preparing buffers: A buffer can be prepared by adding a

common ion (the conjugate base) to a weak acid.

ex: CH3COOH + NaCH3COO A buffer can also be prepared by

partially neutralizing a weak acid with a strong base to produce the conjugate base anion.

ex: CH3COOH + OH- → CH3COO- + H2O

Practice Problem A buffer problem is simply a weak acid

equilibrium problem that involves a common ion..

Ex: Calculate the [H+], pH and percent dissociation of HF in a buffer solution that contains 1.0 M HF (Ka= 7.2 x10-4) and 1.0 M NaF.

ANS: HF ↔ H+ + F-

Ka= 7.2 x10-4 = [H+][F-] = x(1.0 + x) ≈ x(1.0)

[HF] 1.0 – x 1.0

Solving for x = 7.2 x10-4 [H+] = 7.2 x10-4 and pH = 3.14

% dissociation is 7.2 x10-4 x100% = 0.072%

1.0 M

Buffer Capacity

The buffering capacity of a solution represents the amount of H+ or OH- the buffer can absorb without significant changes in pH.

The pH of a buffered solution is determined by the ratio [A-]/[HA]. The capacity of a buffered solution is determined by the magnitudes of [HA] and [A-].

Therefore, a good buffer solution will have relatively high concentrations of BOTH components.

Practice Problem: Buffer Preparation

A chemist needs a solution buffered at pH 4.30 and can choose from the following acids and their sodium salts:a) CH2ClCOOH, cloroacetic acid (Ka = 1.35x10-3)

b) CH3CH2COOH, propanoic acid (Ka = 1.3x10-5)

c) C6H5COOH, benzoic acid (Ka = 6.4x10-5)

d) HOCl, hypochlorous acid (Ka = 3.5x10-8)Calculate the ratio [HA]/[A-] required for each system to yield a pH of 4.30. Which will work best?

A pH of 4.30 corresponds to [H+] = 5.0x10-5 M, and [H+] = Ka[HA]/[A-]

Using Ka values above, the [HA]/[A-] ratio for each of the salts are:a) cloroacetic acid: 3.7x10-2

b) propanoic acid: 3.8c) benzoic acid: 0.78d) hypochlorous acid: 1.4 x103

best since [HA]/[A-] ratio is closest to 1

Indicators

pH and titration curves

Adding strong acid to strong base Adding strong base to strong acid

Adding strong acid to weak base Adding weak acid to strong base

Adding weak base to strong acid Adding strong base to weak acid

More complicated titration curves…

Adding weak acid to weak base

Neutralization Reaction

Acid-Base Titration