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Acids and Bases
Chemical Opposites
Questions
What are the chemical properties of acids and bases? What does it mean to say that an acid or base is weak or strong? What happens when acids react with bases? What does pH mean and how is it used? What is acid rain? What are buffers? How do you determine the concentration of acids and bases?
Acids
"Acid" comes from the Latin acidus, "sour" or "tart" In water acids have a sour or tart taste.
The acid of vinegar is acetic acid.
Behavior of Acids
Aqueous solutions of acids react with some metals such as zinc and magnesium to liberate hydrogen gas.
Zn(s) + H 2 SO 4 (aq) → ZnSO 4 (aq) + H 2 (g)
Aqueous solutions of acids react with some bases such as sodium hydroxide to produce salt and water.
HCl(aq) + NaOH(aq) → NaCl(aq) + H 2 O(l)
The earliest explanations for the behavior of acids was from Svanté Arrhenius in 1887.
Arrhenius Acids
"An acid ionizes in aqueous solution to produce hydrogen ions, H+, and anions."
In water, an acid will release a hydrogen ion that will combine with water molecules and form the hydronium ion:
H+^ + H 2 O → H 3 O+
Compounds like HCl gas will react with water to form this ion.
HCl(g ) + H 2 O(l) → H 3 O+(aq) + Cl-(aq)
Acidic Oxides
Nonmetal oxides are acidic oxides because they react with water to produce acids.
CO 2 reacts with water to form carbonic acid making rainwater slightly acidic.
CO 2 (g) + H 2 O(l) → H 2 CO 3 (aq)
Sulfuric Acid
SO 2 reacts with water to form sulfuric acid, the number one chemical of commerce.
The Frasch process is used to bring sulfur in underground deposits to the surface by injecting superheated steam and compressed air into the deposit and forcing the molten sulfur out.
Sulfuric Acid
The sulfur is then oxidized and water is added to make sulfuric acid, called the contact process.
S(s) + O 2 (g) → SO 2 (g)
2 SO 2 (g) + O 2 (g) → 2 SO 3 (g) using Pt catalyst
SO 3 (g) + H 2 SO 4 (aq) → H 2 S 2 O 7 (l)
H 2 S 2 O 7 (l) + H 2 O(l) → 2 H 2 SO 4 (aq)
Phosphoric Acid
P 4 O 10 reacts with water to form phosphoric acid.
P 4 (s) + 5 O 2 (g) → P 4 O 10 (s)
P 4 O 10 (s) + H 2 O(l) → 2 H 2 SO 4 (aq)
90% of phosphoric acid is used for fertilizers.
Nitric Acid
NO 2 reacts with water to form nitric acid used in fertilizers, explosives, and plastics.
Ammonia is used to form NO gas, the NO gas is oxidized to NO 2 and then mixed with water to make nitric acid.
4 NH 3 (g) + 5 O 2 (g) → 4 NO(g) + 6 H 2 O(g)
2 NO(g) + O 2 (g) → 2 NO 2 (g)
3 NO 2 (g) + H 2 O(l) → 2 HNO 3 (aq) + NO(g)
Bases
Water solutions of bases feel slippery or soapy to the touch. They react with acids to form salt and water.
Arrhenius definition - a base is a substance that ionizes in aqueous solution to produce hydroxide ions, OH-, and cations.
The most common are the hydroxides of the alkali and alkaline earth metals, NaOH, KOH, Ca(OH) 2 , Mg(OH) 2
"Alkali" comes from the Arabic word al-qali, plant ashes. Ashes from wood fires contain potassium carbonate and it dissolves in water to form a slippery, bitter solution.
Basic Oxides
Oxides of metals are called basic oxides because they react with water to form bases.
Calcium oxide (lime) reacts with water to give calcium hydroxide, called slaked lime. This forms a strongly basic solution.
CaO(s) + H 2 O(l) → Ca(OH) 2 (aq)
Lime is fifth in the top chemicals produced in the US. Lime is one of the oldest construction materials. Mortar in Great Wall of China, Roman ruins, Crete.
Weak Acids and Bases
Weak acids and bases are only about 1% ionized and will be weak electrical conductors.
HF H 2 O H 2 S
HCN CH 3 COOH
H 2 CO 3
NH 3 H 2 O
Strengths of Conjugate Acids
and Bases
Strong acids have weak conjugate bases. HCl Cl-
Weak acids have strong conjugate bases. CH 3 COOH CH 3 COO-
Strong bases have weak conjugate acids. OH-^ H 2 O
Weak bases have strong conjugate acids. H 2 O H 3 O+
Polyprotic Acids
Some acids, like sulfuric and phosphoric, have more than one ionizable hydrogen ion.
H 2 SO 4 H 3 PO 4 H 2 SO 3 H 2 S H 2 CO 3
Concentration of Acid and
Base Solutions
The concentration of a solution is the quantity of a solute dissolved in a specific quantity of a solvent.
Molarity is often used to express concentration. Moles per liter or M
A 1 M solution contains 1 mole of solute dissolved in 1 liter of solution.
A 0.5 M solution contains 1/2 mole of solute dissolved in 1 liter of solution.
Examples
Consider a 0.15 M solution of HCl
How many moles of solute are dissolved in 200 mL of the solution?
How many grams of solute are dissolved in 200 mL of the solution?
Examples
How many grams of solute are needed to prepare 0.350 L of 1.1 M Ca(OH) 2?
The pH Scale
Pure water is neutral; it contains equal numbers of hydronium ions and hydroxide ions. About 1 in 550,000,000 water molecules is ionized at a time in pure water.
The concentration of [H+] in pure water is 1.0 x 10-7^ M. The concentration of [OH-] in pure waer is 1.0 x 10-7^ M.
The pH Scale
When we discuss a system at equilibrium, we can write an equilibrium constant that divides the product of the product concentrations by the product of the reactant concentrations.
For the reaction of the ionization of water,
H 2 O + H 2 O ↔ H 3 O+^ + OH-
Kw = [H 3 O+][OH-^ ]/[H 2 O]^2 , since water is 1,
Kw = [H 3 O+][OH-^ ] = 1 x 10-
This value is a constant.
Solutions of Acids or Bases
If acid is added, [H 3 O+] wil be greater than 1.0 x 10-7M and [OH-^ ] will be less than 1.0 x 10-7M. The product of the two must always be 1 x 10-14.
The pH is defined as the negative log of the hydronium ion concentration.
pH = - log [H 3 O+]
The pH scale ranges from 0 - 14 because of the concentration of hydronium ion in pure water.
Solutions of Acids or Bases
The pH of pure water is 7 because the concentration of [H+] in pure water is 1.0 x 10-7M.
7 = -log [1.0 x 10-7]
Because pH is a log scale, each decrease in the pH scale represents a 10-fold increase in [H+].
pH and A cidity/Basicity
if pH < 7.0 acid if pH > 7.0 base if pH = 7.0 neutral
Acid-Base Buffers
Buffer solutions contain a base and an acid that can react with added acid or base, respectively, and maintain a pH very close to the original value. In this way is acts as a pH shock absorber.
The most important buffer is the carbonate system. It controls the pH of blood and the pH of lakes and streams.
HCO 3 - (aq) + H 3 O+(aq) ↔ H 2 CO 3 (aq) + H 2 O(l)
HCO 3 - (aq) + OH-(aq) ↔ CO 3 =(aq) + H 2 O(l)
