Chapter 10: Acids and Bases and Equilibrium

10.1: Acids and Bases

Acids

• Arrhenius Acids: Substances that produce hydrogen ions when they dissolve in water.
  • Because acids produce ions in water, they are also electrolytes.
• Svante Arrhenius — A Swedish chemist to first describe acids.

Naming Acids

• When an acid dissolves in water to produce a hydrogen ion and a simple nonmetal anion, the prefix hydro– is used before the name of the nonmetal, and its –ide ending is changed to –ic acid.

  • Ex.: Hydrochloric acid, hydrobromic acid and hydrocyanic acid

• The most common form of an oxygen-containing acid has a name that ends with –ic acid.

  • The name of its polyatomic anion ends in –ate.
  • Ex.: Sulfate, carbonate, acetate and phosphate

• An acid that contains one less oxygen atom than the common form is named as an –ous acid.

  • Ex.: Phosphorous acid, sulfurous acid, and chlorous acid.

• The name of its polyatomic anion ends with –ite.

  • Ex.: Sulfite, nitrite, phosphite, and chlorite.

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Bases

• Arrhenius Bases: These are ionic compounds that dissociate into metal ions and hydroxide ions when they dissolve in water. These bases are considered electrolytes as well.
• Most Arrhhenius bases are formed from Group 1A and Group 2A metals.

Naming Bases

• Typical Arrhenius bases are named as hydroxides.
  • Lithium Hydroxide
  • Sodium Hydroxide
  • Potassium Hydroxide
  • Calcium Hydroxide
  • Aluminum Hydroxide

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Brønsted–Lowry Acids and Bases

• J. N. Brønsted and T. M. Lowry — They expanded the definition of acids and bases in 1923.

• Brønsted–Lowry acid: It can donate a hydrogen ion to another substance.

• Brønsted–Lowry base: It can accept hydrogen ion.

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Conjugate Acid–Base Pairs

• According to the Brønsted–Lowry theory, a conjugate acid–base pair consists of molecules or ions related by the loss of one H+ by an acid, and the gain of one H+ by a base.
• Amphoteric: Substances that can act as acids and bases.

10.2: Strengths of Acids and Bases

• The strength of an acid is determined by the moles of H3O+ that are produced for each mole of acid that dissolves.
• The strength of a base is determined by the moles of OH- that are produced for each mole of base that dissolves.
• Strong acids: These are examples of strong electrolytes because they donate H+ so easily that their ionization in water is virtually complete.
• Weak Acids: These are weak electrolytes because they ionize slightly in water, which produces only a few ions.
• Diprotic Acid: Carbonic acid that has two H+, which ionize one at a time
• Strong Bases: These are ionic compounds that dissociate in water to give an aqueous solution of metal ions and hydroxide ions.
• Weak Bases: These are weak electrolytes that are poor acceptors of hydrogen ions and produce very few ions in solution.

10.3: Acid-Base Equilibrium

• Equilibrium: The rates of the forward and reverse reactions become equal.
  • The reactants form products at the same rate that the products form reactants.
  • It has been reached when no further change takes place in the concentrations of the reactants and products.
• Le Châtelier’s principle: It states that when equilibrium is disturbed, the rates of the forward and reverse reactions change to relieve that stress and reestablish equilibrium.

10.4: Ionization of Water

• In pure water, a few water molecules transfer H+ to other water molecules, producing small, but equal, amounts of H3O+ and OH-.

• In pure water, the molar concentrations of H3O+ and OH- are each 1.0 x 10^-7 M.

• The ion product constant for water:

  

• In acidic solutions, the [H3O+] is greater than the [OH-].

• In basic solutions, the [OH-] is greater than the [H3O+].

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10.5: The pH Scale

• pH Scale: A range of numbers typically from 0 to 14, which represents the [H3O+] of the solution.
• A neutral solution has a pH of 7.0.
• In acidic solutions, the pH is below 7.0.
• In basic solutions, the pH is above 7.0.
• Mathematically, pH is the negative logarithm of the hydronium ion concentration, pH = –log [H3O+].

10.6: Reactions of Acid and Bases

• Salt: An ionic compound that does not have H+ as the cation or OH- as the anion.
• Acids react with certain metals to produce hydrogen gas and salt.
  • Metals that react with acids include potassium, sodium, calcium, magnesium, aluminum, zinc, iron, and tin.
• When an acid is added to a carbonate or bicarbonate, the products are carbon dioxide gas, water, and an ionic compound.
• Neutralization: A reaction between an acid and a base to produce water and salt.
• Titration: A laboratory procedure in which we neutralize an acid sample with a known amount of base.
  • In the titration, we neutralize the acid by adding a volume of base that contains a matching number of moles of OH-.
  • We know that neutralization has taken place when the phenolphthalein in the solution changes from colorless to pink. This is called the neutralization endpoint.
  • From the volume added and molarity of the NaOH solution, we can calculate the number of moles of NaOH, the moles of acid, and then the concentration of the acid.
• Antacids: These are substances used to neutralize excess stomach acid.

10.7: Buffers

• Buffer Solution: It maintains pH by neutralizing small amounts of added acid or base.
• A buffer contains either a weak acid and its salt or a weak base and its salt.
• When an acid or base is added to a buffer solution, there is little change in pH.
• Acidosis: A condition that occurs when there’s an increase in the CO2 level that leads to a low blood pH concentration.
• Alkalosis: A condition that occurs when there’s a decrease in the CO2 level that leads to a high blood pH concentration.

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