ACIDS & BASES. ACID/BASE THEORY Acids and bases are solutions which can be described differently by multiple theories. So far, we have treated everything.

ACIDS & BASES

ACID/BASE THEORY

• Acids and bases are solutions which can be described differently by multiple theories.

• So far, we have treated everything that utilizes • “H” as a cation = acid • H3PO4

• “OH” as an anion = base• NaOH

ARRHENIUS THEORY

• The theory that anything capable of producing H+ ions are acids and those producing OH- anions are bases is Arrhenius’ Theory of acids/bases.

• Based on the notion that these substances dissociate in solution into their ions.

• H2SO4 2H+ + SO4-2

H20

ACID OR BASE?

•HNO3 KOH• Acid Base

• Ca(OH)2 HBr• Base Acid

•HC2H3O2 Al(OH)3

• Acid Base

BRONSTAD LOWRY ACIDS/BASES

•More general description of acids and bases, incorporates all Arrhenius acid/bases• Acids = H+ donors• Bases = H+ acceptors

Example:NH3 (aq) + H2O (l) NH4

+ (aq) + OH-(aq)

• Would solutions of NH3 be acidic or basic?

CONJUGATE ACIDS/BASES

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

Conjugate Acid: The compound remaining after H+ acceptance by a base.

NH3 NH4+

Base Conjugate Acid

Conjugate Base: The compound remaining after H+ donation by an acid.

H2O OH-

Acid Conjugate Base

LEWIS ACID/BASES

•Describes Acid/Bases based upon electron donation.• Lewis Acid: A substance that can accept a

pair of electrons to form a covalent bond.• Lewis Base: A substance that can donate a

pair of electrons to form a covalent bond.

H + + OH - H2O

STRENGTH OF ACIDS/BASES

• For this class, we will mostly discuss Arrhenius acid/bases (H+ cation/OH- anion)

• The strength of these solutions is determined by the concentration (Molarity) of H+ or OH- ions in the solution.• Dependent upon…

1) Amount of dissolved solute donating the ion2) Likelihood of ion dissociation

CONCENTRATION OF SOLUTE

• Just like calculating the molarity of non acid/base solutions.• M = mol solute / liter solution

• Example: What is the concentration of a H2SO4

solution if 13.2 g of solute are dissolved in 100mL of di-water?

CALCULATING [H+ ] & [OH-]

•However, we cannot assume the molarity (concentration) of H+ /OH- ions is the same as the molarity of the solution for 2 reasons:

1. Not all acids/bases completely dissociate into their ions.

2. Not all acidic/basic compounds donate 1 ion per molecule

• Complete dissociation: all solute dissociates into ions, meaning the concentration of ions can be calculated with the molar ratio.

HCl H+ + Cl-

1 mol HCl produces 1 mol H+

• So a 0.2M HCl solution would have a H+ concentration of 0.2M as well

DISSOCIATION

COMPLETE DISSOCIATION

• Few acids/bases completely dissociate• If they do, they are termed “Strong Acids/Bases” because they produce high concentrations of their ions upon dissociation.

• Strong Acids (7): HCl, HBr, HI, H2SO4, HNO3, HClO3, HClO4

• Strong Bases (8): LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

CALCULATING ION DONATION

•Not all acids/bases donate 1 H+ or OH- into solution when they dissociate!

•How many ions would the following acids/bases donate if they completely dissociate?

• HBr• Ca(OH)2

• H2SO4

[H+] EXAMPLE

Using the previous H2SO4 example (13.2 g of solute are dissolved in 100mL of di-water), calculate the [H3O+] of the solution.

PARTIAL DISSOCIATION

•Weak acids/bases do not completely dissociate in water. They are called weak, because a lower concentration of H+ or OH- ions results.

• Example: Only about 1% of Phosphoric Acid (H3PO4) dissociates into H+ + H2PO4.

DISSOCIATION OF STRONG & WEAK ACIDS/BASES

• In other words, you would need to know what percentage of an acidic/basic compound is likely to donate its ions into solution in order to calculate the [H+ ] & [OH-].

• For this class, we will be working with strong acids/bases. So complete dissociation (ion donation) can be assumed.

POTENTIAL OF HYDROGEN

• pH = potential of hydrogen

• pH is an expression of the concentration of H+ ions in solution. Since H+ readily reacts with water to form H3O+, we calculate…

pH = −log [H3O+]

[H3O+] = 10-pH

PH & POH

• Just as pH is an expression of the concentration of H+ ions in solution, pOH expresses the concentration of OH- ions in solution…

pOH = −log [OH-][OH-] = 10-pOH

• pH & pOH always add up to 14 and thus are related through the equation:

pH + pOH = 14

PH SCALE

• pH < 7 acidic• pH = 7 neutral• pH > 7 basic

•High [H3O+] acid

• Low [H3O+] base

• [H3O+] = ~1x10-7

neutral

ACID/BASE REACTIONS

• Also known as neutralization reactions, because the pH of the final solution becomes neutralized towards pH=7

• Either [H3O+] or [OH-] is decreased by forming more neutral products.

Acid + Base Salt + Water

ACID/BASE REACTIONS

• 2 Types you will need to know…

1) Acid + Base Salt + H2O

2) Acid + Carbonate Salt + H2O + CO2

ACID/BASE REACTION #1

• Example: Hydrochloric acid reacts with sodium hydroxide.

HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

ACID/BASE REACTION #2

• Example: Potassium carbonate reacts with sulfuric acid.

H2SO4(aq)+ K2CO3(aq) K2SO4(aq) + H2O(l) + CO2(g)

Acid Carbonate Salt Water Carbon

Dioxide