Chapter 9 - Acids and Bases

Chapter Nine: Acids and Bases Introduction to Acids and Bases the earliest definition was given by arrhenius: • an acid contains a hydrogen atom and dissolves in water to form a hydrogen ion, H+ • a base contains hydroxide and dissolves in water to form OH - the arrhenius definition correctly predicts the behavior of many acids and bases. however, this definition is limited and sometimes inaccurate. for example, H + does not exist in water. instead, it reacts with water to + form the hydronium ion, H3O . the bronsted-lowry definition is more widely used: • a bronsted-lowry acid is a proton (H+ ) donor • a bronsted-lowry base is a proton (H+ ) acceptor HCl is a bronsted acid H2O is a bronsted base Bronsted-Lowry Acids a bronsted-lowry acid must contain a hydrogen atom common bronsted-lowry acids (HA): • HCl (hydrochloric acid) • HBr (hydrobromic acid) • HNO3 (nitric acid) • H2SO4 (sulfuric acid) the names of common acids are derived from the anions formed when they dissolve in water. anions whose names end in “-ide,” add the prefix “hydro” and change the “-ide” ending to “-ic acid.” • ex: chloride (Cl - ) → hydrochloric acid (HCl) polyatomic anions whose names end in “-ate,” change the “-ate” ending to “-ic acid.” • ex: sulfate (SO4 2- ) → Sulfuric acid (H2SO4) polyatomic anions whose names end in “-ite,” change the “-ite” ending to “-ous acid.” • ex: sulfite (SO3 2- ) → Sulfurous acid (H2SO3) monoprotic acid: contains one acidic proton (HCl) diprotic acid: contains two acidic protons (H2SO4) triprotic acid: contains three acidic protons (H3PO4) a bronsted-lowry acid may be neutral, or it may carry a net positive or + negative charge (Hcl, H3O , HSO4 - ) Bronsted-Lowry Bases a bronsted-lowry base is a proton acceptor, so it must be able to form a bond to a proton. a base must contain a lone pair of electrons that can be used to form a new bond to the proton. common bronsted-lowry bases (B:): lone pairs make these neutral compounds bases: .. • ammonia (NH3) .. . • water (H2O.) the OH- is the base in each metal salt: • sodium hydroxide (NaOH) The Reaction of a • potassium hydroxide (KOH) Bronsted-Lowry Acid with • magnesium hydroxide (Mg(OH)2) a Bronsted-Lowry Base • calcium hydroxide (Ca(OH)2) conjugate base: the product formed by loss of a proton from an acid conjugate acid: the product formed by gain of a proton by a base HBr and Br - are a conjugate acid-base pair. + H2O and H3O are a conjugate acid-base pair. the net charge must be the same on both sides of the equation. when a species gains a proton (H+ ), it gains a +1 charge. + when a species loses a proton (H ), it gains a -1 charge. amphoteric compound: a compound that contains both a hydrogen atom and a lone pair of electrons; it can be either an acid or base Relating Acid and Base Strength dissociation: the splitting apart of a covalent molecule (or ionic compound) into its ions when a strong acid dissolves in water, all of the acid dissociates into ions. a single reaction arrow is used because the product is greatly favored at equilibrium. common strong acids are HI, HBr, HCl, H2SO4, and HNO3. when a weak acid dissolves in water, only a small fraction of the acid dissociates into ions. unequal reaction arrows are used because the reactants are usually favored at equilibrium. common weak acids are H3PO4, HF, H2CO3, and HCN. when a strong base dissolves in water, all of the base dissociates into ions. common strong bases are NaOH and KOH. when a weak base dissolves in water, only a small fraction of the base dissociates into ions. a strong acid readily donates a proton, forming a weak conjugate base. • strong acid: HCl - • weak conjugate -base: Cl a strong base readily accepts a proton, forming a weak conjugate acid. • strong base: OH - • weak conjugate acid: H2O Using Acid Strength to Predict the Direction of Equilibrium a bronsted-lowry acid-base reaction represents an equilibrium the position of the equilibrium depends upon the strengths of the acids and base. the stronger acid reacts with the stronger base to form the weaker acid and the weaker base. when the stronger acid and base are the reactants on the left side, the reaction readily occurs and the reaction proceeds to the right. a larger forward arrow means that products are favored. compare acid and conjugate acid goes to the weak. if an acid-base reaction would form the stronger acid and base, equilibrium favors the reactants and little product forms. a larger reverse arrow means that reactants are favored. example: are the reactants or products favored in the following acid-base reaction? step one: identify the acid in the reactants and the conjugate acid in the products. step two: determine the relative strength of the acid and the conjugate acid. from table 9.2, HCN is a stronger acid than H2O. step three: equilibrium favors the formation of the weaker acid. Equilibrium and Acid Dissociation Constants multiplying both sides by [H2O] forms a new constant, called the acid dissociation constant, Ka. the stronger the acid, the larger the Ka value. equilibrium favors formation of the weaker acid, the acid with the smaller Ka value. Dissociation of Water water can behave as both a bronsted-lowry acid and a bronsted-lowry base. thus, the two water molecules can react together in an acid-base reaction: from the reaction of two water molecules, the following equilibrium constant expression can be written: 2 multiplying both sides by [H2O] yields Kw, the ionproduct constant for water. experimentally, it can be shown that: Kw is a constant, 1.0 x 10-14 , for all aqueous solutions at 25 °C. to calculate [OH - ] when [H3O+ ] to calculate [H3O + ] when [OH - ] is known: is known: if the [H3O+ ] in a cup of coffee is 1.0 x 10-5 M, then the [OH - ] can be calculated as follows: in a cup of coffee, therefore [H3O+ ] > [OH - ], and the solution is acidic overall. Calculating pH the lower the pH, the higher the concentration of H3O + : acidic solution: pH < 7 → [H3O + ] > 1 x 10-7 neutral solution: pH = 7 → [H3O + ] = 1 x 10 -7 basic solution: pH > 7 → [H3O + ] < 1 x 10 -7 Calculating pH Using a Calculator a logarithm has the same number of places after the decimal as there are digits in the original number example: if [H3O+ ] = 1.2 x 10 -5 M for a solution, what is its pH? pH = -log [H3O+ ] pH = -log (1.2 x 10 -5 ) pH = - (-4.92) 2 digits 2 decimal places pH = 4.92 the solution is acidic because pH < 7 if the pH of a solution is 8.50, what is the [H3O+ ]? pH = -log [H3O+ ] 8.50 = -log [H3O+ ] + -8.50 = log [H3O ] + anti log (-8.50) = [H3O ] 2 decimal places 10 -8.5 = [H3O+ ] [H3O+ ] = 3.2 x 10 -9 M 2 digits the solution is basic because [H3O+ ] < 1 x 10 -7 M Reaction of Acids with Hydroxide Bases neutralization reaction: an acid-base reaction that produces a salt and water as products the acid HA donates a proton (H + ) to the OH - base to form H2O. the anion A - from the acid combines with the cation M + from the base to form the salt MA. How to Draw a Balanced Equation for a Neutralization Reaction Between HA and MOH example: write a balanced equation for the reaction of Mg(OH)2 with HCl. step one: identify the acid and base in the reactants and draw H2O as one product step two: determine the structure of the salt the salt is formed from the parts of the acid and base that are not used to form H2O. Mg2+ and Cl - combine to form MgCl2 step three: balance the equation Net Ionic Equations a net ionic equation contains only the species involved in a reaction written as individual ions: omit the spectator ions: Na+ and Cl - . what remains is the net ionic equation: Reaction of Acids with Bicarbonate Bases a bicarbonate base, HCO3 -, reacts with one H +to form carbonic acid, H2CO3. carbonic acid then decomposes into H2O and CO2. for example: a carbonate base, CO32- , reacts with two H + to form carbonic acid, H2CO3. for example: The Acidity and Basicity of Salt Solutions a salt can form an acidic, basic, or neutral solution depending on whether its cation and anion are derived from a strong or weak acid and base. for the salt M + A - : • the cation M + comes from the base, such as NaOH or KOH. • the anion A - comes from the acid, such as HCl or H2SO4. a salt derived from a strong acid and strong base forms a neutral solution (pH = 7). a salt derived from a strong base and a weak acid forms a basic solution (pH > 7). a salt derived from a weak base and a strong acid forms an acidic solution (pH < 7). thus, the ion derived from the stronger acid or base determines whether the solution is acidic or basic. Buffers buffer: a solution whose pH changes very little when acid or base is added most buffers are solutions composed of roughly equal amounts of: • a weak acid • the salt of its conjugate base the buffer resists change in pH because • added base, OH - , reacts with the weak acid • added acid, H3O , reacts with the conjugate base + General Characteristics of a Buffer if an acid is added to the following buffer equilibrium, then the excess acid reacts with the conjugate base, so the overall pH does not change much. if a base is added to the following buffer equilibrium, then the excess base reacts with the conjugate acid, so the overall pH does not change much. Calculating the pH of a Buffer the effective pH range of a buffer depends on its Ka. rearranging this expression to solve for [H3O + ]: hydronium ion concentration determines the buffer pH example: calculate the pH of a buffer that contains 0.20 M CH3COOH and 0.20 CH3COO - Na + . Buffers in the Blood normal blood pH is between 7.35 and 7.45. the principle buffer in the blood is carbonic acid/bicarbonate (H2CO3/HCO3 - ). CO2 is constantly produced by metabolic processes in the body. the amount of CO2 is related to the pH of the blood. respiratory acidosis results when the body fails to eliminate enough CO2 due to lung disease or failure. respiratory alkalosis is caused by hyperventilating; very little CO2 is produced by the body.