Acid and Base Chemistry

Arrhenius Theory of Acids

• An acid is a substance that gives off H+ in water.

This is a hydronium ion or proton.

• H2SO4 is sulfuric acid.

H2SO4 2H+ + SO4 2-

• has one proton and one electron. H+ has lost the electron and so we call it a proton.

H+ + H2O H3O+ Hydronium Ion

H2SO4 + 2H2O 2H3O+ + SO4 2-

Arrhenius Bases

• According to the Arrhenius Theory, a base

is a substance that dissolves in water to

make OH-.

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

Ca(OH)2(s) Ca+2(aq) + 2OH-(aq)

Overall Arrhenius Theory

• Acid + Base → Salt + Water + Heat • Salt is an ionic compound that results from the reaction between an acid and a base. • A solution is acidic if there are more moles of H+ than OH-. • A solution is basic if there are more moles of OH - than H+. • A solution is neutral if the number of moles of H+ and OH- are the same. They then neutralize, or cancel out, each other.

Examples of Neutralization Equations

• HCl + NaOH → NaCl + H2O

• 2HBr + Ca(OH)2 → CaBr2 + 2H2O

• HNO3 + KOH → KNO3 + H2O

• 3H2SO4 + 2Al(OH)3 → Al2(SO4)3 + 6H2O

Acid Nomenclature

• Binary Acids = H and one nonmetal element. • Name always begins hydro-, root is based on the nonmetal name with the suffix -ic, followed by the word acid. HCl = Hydrochloric Acid HBr = Hydrobromic Acid H2Se = Hydroselenic Acid

Oxyacids

These are “ternary” acids, and contain 3 elements: H, O and a nonmetal. • Example: H2SO4

• Ternary or oxyacids do not have hydro as a prefix. • Name starts with root. The suffix depends on the anion from which the acid is derived. • -ate anion = -ic acid. • -ite anion = -ous acid.

Examples

• HClO3 = H+ + ClO3 -

(Chlorate ion) Chloric Acid

• HNO2 = Nitrous Acid

• H2CO3 = Carbonic Acid

Titration

Titration is a process in which the amount of acid

and base are carefully measured to exactly

produce a neutral solution.

Technically speaking, the substance being added

(either acid or base) is slowly added until the

solution is neutral, and the volume that is added is

carefully measured and recorded.

Indicator

Substance that changes color as the relative

concentration of acid and base changes.

The best indicator will change color from

acid to neutral or from base to neutral.

• Example--phenolthalein. It is clear in an

acidic solution or neutral solution, and pink

in a basic solution.

Math on Titration Problems

In a titration problem, the moles of H+ and the moles of OH- must be equal. • Molarity • Volume = moles, so MaVa = MbVb Ma = molarity of H+ Va = volume of acid Mb = molarity of OH- Vb = volume of base

Practice Problem

23.1 mL of a .100 M HCl solution is titrated with

15.7 mL of a NaOH base solution. Find the

concentration of the base.

• MaVa = MbVb

• (.100 M)•(23.1mL)=Mb•(15.7mL)

• Mb= .147 M

Logarithms

Logarithm Defined

• X is the logarithm (or log) of A if A can be

written as 10X.

• A = 10X X = log A

• What is the log of 1,000,000?

• Can 1,000,000 be written as a power of 10?

• Yes, 1,000,000 = 10e6.

• So log 1,000,000 = 6.

pH

• Definition: pH is the - log of the [H+] • pH = - log [H+] • pH is a measure of the acidity or basicity of a solution. • For a neutral solution, pH = 7. • For an acidic solution, pH < 7. • For a basic solution, pH > 7.

pH and Logarithms

104 = 10,000 4 is the logarithm of 10,000

• 10e3 = 1,000 3 is the logarithm of 1,000

• 10e2 = 100 2 is the logarithm of 100

• 10e1 = 10 1 is the logarithm of 10

• 10e0 = 1 0 is the logarithm of 1

• 10e-1 = .1 -1 is the logarithm of .1

• 10e-2 = .01

• 10e-3 = .001

What is the log of 4,258?

• 4,258 = 10X

• What is X?

Use calc.---- log(4258)=3.63

• 10e3.63 = 4,258

Other important equations

pH = - log [H+] • pOH = - log [OH-] • Kw = [H+][OH-] = 1 x 10-14 • pH + pOH = 14 • [H+] = 10e-pH • [OH-] = 10e-pOH

Dissociation Equilibrium for Water

• H2O(l) ↔ H+(aq) + OH-(aq)

• Kw = [H+][OH-] = 1 x 10-14

So as water dissociates, [H+]=[OH-]. • Then, [H+]2 = 1 x 10-14, and

• [H+] = [OH-] = 1 x 10-7.

• This is true for a neutral solution

Practice Problem A solution is .20 M in HCl. Find the pH, pOH and the [OH-]. • HCl → H+ + Cl- • .20 M HCl = .20 M H+ = .20 M Cl- HCl → H+ + Cl- • pH = -log [.20 M] = .70 • [OH-] = 1x10e-14/.20 = 5.0x10e-14 • pOH = -log [5.0x10e-14] = 13.30 or pH + pOH = 14 pOH= 14-.70= 13.30

A solution has a pH of 3.20. Find the concentration of H+ and OH-. Also find the pOH. • [H+] = 10e-pH = 10e-3.20 or antilog of -3.20 • [H+] = 6.31 x 10e-4 • pH + pOH = 14 pOH = 10.80. • [OH-] = 10e-pOH or [H+][OH-]=1x10e-14 • [OH-] = 1.58 x 10-11

Write an equation for dissolving and

dissociating the following acids in water:

HCl, HNO3, H2SO4, H3PO4

HCl → H+(aq) + Cl-(aq) or HCl + H2O → H3O+(aq) + Cl-(aq)

HNO3 → H+(aq) + NO3-(aq) or HNO3 + H2O → H3O+(aq) + NO3-(aq)

H2SO4 → 2H+(aq) + SO4 2-(aq) orH3PO4 ↔ 3H+(aq) + PO4 3-(aq)

Mono-, Di- and Triprotic Acids

A monoprotic acid gives off one H+ for each acid molecule. HCl and HNO3 are examples of monoprotic acids. • A diprotic acid gives off two H+ ions for each acid molecule. H2SO4 is an example of a diprotic acid. • A triprotic acid gives off three H+ ions for each acid molecule. H3PO4 is an example of a triprotic acid.

Concentration Calculations with Di- and Triprotic Acids

When you have di- and triprotic acids, the concentration of H+ is greater than the concentration of the original acid. • Find the molarity H+ and the pH of a .200 M H2SO4 solution. H2SO4 → 2H+(aq) + SO4 2-(aq) (.200 moles H2SO4) (2 moles H+) = .400 M H+ ( 1 liter soln.) (1 mole H2SO4) pH = - log (.400) = .398

Dibasic and Tribasic Substances

Bases that have more than one OH- per

molecule follow a similar pattern to the

acids.

Ca(OH)2 → Ca 2+ + 2OH-

Al(OH)3 → Al3+ + 3OH-

Bronsted-Lowry Theory of Acids and Bases

A Bronsted-Lowry Acid is a substance the donates a proton or H+ ion to another molecule or ion. • A Bronsted-Lowry Base is a substance that accepts a proton or H+ ion from another molecule or ion. • An Acid-Base reaction must include both a proton donor and a proton acceptor.

Examples of Bronsted-Lowry Acid Base Reactions

Acetic Acid reacts with Ammonia to

produce Ammonium Acetate.

CH3CO2H + NH3 ↔ NH4CH3COo

CH3CO2H + NH3 ↔ NH4+ + CH3COo -

Lewis Acid and Base

• Lewis Acid-an atom, ion, or molecule that accepts an electron pair to form a covalent bond

H+ + NH3 -> NH4+

• Lewis Base-an atom, ion, or molecule that donates an electron pair to form a covalent bond

BF3 + F- -> BF4+ • Lewis Acid-Base Reaction-the formation of one or

more covalent bonds between an electron-pair donor and an electron-pair acceptor

• Do page 482

Conjugate Acids and Bases

A conjugate acid is an acid (as a product) that results from a Bronsted acid-base reaction. • A conjugate base is a base (as a product) that results from a Bronsted acid-base reaction. • A Bronsted Base accepts a proton to produce its conjugate acid. • A Bronsted Acid donates a proton to produce its conjugate base.

Example

CH3CO2H + NH3 NH4+ + CH3COo-

The Bronsted acid is CH3COoH.

The Bronsted base is NH3. The conjugate acid is NH4+.

The conjugate base is CH3COo-.

Weak Acids and Bases

Weak Acids do not fully dissociate when dissolved in water. This means that the concentration of the H+ ion is less than expected due to the stoichiometry. • Examples of Weak Acids: Acetic Acid (CH3CO2H) Phosphoric Acid (H3PO4) • Weak Bases do not fully accept one H+ per molecule as expected.

Strong and Weak Acids and Bases

Strong Acids fully dissociate in water, and donate all possible H+ ions readily. • Strong Bases fully accept all donated H+ ions readily. • Weak Acids do not readily donate all possible H+ ions. • Weak Bases do not readily accept all possible H+ ions.

Buffers

• A buffer is a solution that resists changes in the pH. • A buffer is made from a weak acid and its conjugate base or a weak base and its conjugate acid. • Since it is made of both a weak acid and a weak base, there is both acid and base in solution to react with any strong acid or base that is added to it. This minimizes the effect of adding acid or base.

A .001 M acetic acid solution is 5% dissociated. What is the H+ concentration? % Dis. = M of dissociated acid(100) M of acid total CH3CO2H ↔ H+ + CH3CO2- % Dis. = [CH3CO2-](100) [CH3CO2H]

Ionization Constant

HA ↔ H+ + A-

Ka= [H+]•[A-]

[HA]

Ka = acid dissociation constant

A- + H2O ↔ HA + OH

Kb= [HA]•[OH-]

[A-]

Kb = base dissociation constant