Chapter 10 Acids and Bases

Arrhenius Acids and Bases In 1884, Svante Arrhenius proposed these definitions acid: a substance that produces H3O+ ions aqueous solution base: a substance that produces OH- ions in aqueous solution

Arrhenius Acids and Bases when HCl, for example, dissolves in water, its reacts with water to give hydronium ion and chloride ion we use curved arrows to show the change in position of electron pairs during this reaction

Arrhenius Acids and Bases With bases, the situation is slightly different many bases are metal hydroxides such as KOH, NaOH, Mg(OH)2, and Ca(OH)2 these compounds are ionic solids and when they dissolve in water, their ions merely separate other bases are not hydroxides; these bases produce OH- by reacting with water molecules

Arrhenius Acids and Bases we use curved arrows to show the transfer of a proton from water to ammonia

Acid and Base Strength Strong acid: one that reacts completely or almost completely with water to form H3O+ ions Strong base: one that reacts completely or almost completely with water to form OH- ions here are the six most common strong acids and the four most common strong bases

Acid and Base Strength Weak acid: a substance that dissociates only partially in water to produce H3O+ ions acetic acid, for example, is a weak acid; in water, only 4 out every 1000 molecules are converted to acetate ions Weak base: a substance that dissociates only partially in water to produce OH- ions ammonia, for example, is a weak base

Brønsted-Lowry Acids & Bases Acid: a proton donor Base: a proton acceptor Acid-base reaction: a proton transfer reaction Conjugate acid-base pair: any pair of molecules or ions that can be interconverted by transfer of a proton

Brønsted-Lowry Acids & Bases Brønsted-Lowry definitions do not require water as a reactant

Brønsted-Lowry Acids & Bases we can use curved arrows to show the transfer of a proton from acetic acid to ammonia

Brønsted-Lowry Acids & Bases Note the following about the conjugate acid-base pairs in the table 1. an acid can be positively charged, neutral, or negatively charged; examples of each type are H3O+, H2CO3, and H2PO4- 2. a base can be negatively charged or neutral; examples are OH-, Cl-, and NH3 3. acids are classified a monoprotic, diprotic, or triprotic depending on the number of protons each may give up; examples are HCl, H2CO3, and H3PO4

Brønsted-Lowry Acids & Bases carbonic acid, for example can give up one proton to become bicarbonate ion, and then the second proton to become carbonate ion 4. several molecules and ions appear in both the acid and conjugate base columns; that is, each can function as either an acid or a base

Brønsted-Lowry Acids & Bases the HCO3- ion, for example, can give up a proton to become CO32-, or it can accept a proton to become H2CO3 a substance that can act as either an acid or a base is said to be amphiprotic the most important amphiprotic substance in Table 8.2 is H2O; it can accept a proton to become H3O+, or lose a proton to become OH- 5. a substance cannot be a Brønsted-Lowry acid unless it contains a hydrogen atom, but not all hydrogen atoms in most compounds can be given up acetic acid, for example, gives up only one proton

Brønsted-Lowry Acids & Bases 6. there is an inverse relationship between the strength of an acid and the strength of its conjugate base the stronger the acid, the weaker its conjugate base HI, for example, is the strongest acid in Table 8.2, and its conjugate base, I-, is the weakest base in the table CH3COOH (acetic acid) is a stronger acid that H2CO3 (carbonic acid); conversely, CH3COO- (acetate ion) is a weaker base that HCO3- (bicarbonate ion)

Acid-Base Equilibria we know that HCl is a strong acid, which means that the position of this equilibrium lies very far to the right in contrast, acetic acid is a weak acid, and the position of its equilibrium lies very far to the left but what if the base is not water? How can we determine which are the major species present?

Acid-Base Equilibria To predict the position of an acid-base equilibrium such as this, we do the following identify the two acids in the equilibrium; one on the left and one on the right using the information in Table 10.1, determine which is the stronger acid and which is the weaker acid also determine which is the stronger base and which is the weaker base; remember that the stronger acid gives the weaker conjugate base, and the weaker acid gives the stronger conjugate base the stronger acid reacts with the stronger base to give the weaker acid and weaker base; equilibrium lies on the side of the weaker acid and weaker base

Acid-Base Equilibria identify the two acids and bases, and their relative strengths the position of this equilibrium lies to the right

Acid-Base Equilibria Example: predict the position of equilibrium in this acid-base reaction

Acid-Base Equilibria Example: predict the position of equilibrium in this acid-base reaction Solution: the position of this equilibrium lies to the right

Acid Ionization Constants when a weak acid, HA, dissolves in water the equilibrium constant, Keq, for this ionization is because water is the solvent and its concentration changes very little when we add HA to it, we treat [H2O] as a constant equal to 1000 g/L or 55.5 mol/L we combine the two constants to give a new constant, which we call an acid ionization constant, Ka

Acid Ionization Constants Ka for acetic acid, for example is 1.8 x 10-5 because the acid ionization constants for weak acids are numbers with negative exponents, we commonly express acid strengths as pKa where the value of pKa for acetic acid is 4.75 values of Ka and pKa for some weak acids are given in Table 10.2 as you study the entries in this table, note the inverse relationship between values of Ka and pKa the weaker the acid, the smaller its Ka, but the larger its pKa

Properties of Acids & Bases Neutralization acids and bases react with each other in a process called neutralization. Reaction of acids with metals strong acids react with certain metals (called active metals) to produce a salt and hydrogen gas, H2

Properties of Acids & Bases Reaction with metal hydroxides reaction of an acid with a metal hydroxide gives a salt plus water the reaction is more accurately written as omitting spectator ions gives this net ionic equation

Properties of Acids & Bases Reaction with metal oxides strong acids react with metal oxides to give water plus a salt

Properties of Acids & Bases Reaction with carbonates and bicarbonates strong acids react with carbonates to give carbonic acid, which rapidly decomposes to CO2 and H2O strong acids react similarly with bicarbonates

Properties of Acids & Bases Reaction with ammonia and amines any acid stronger than NH4+ is strong enough to react with NH3 to give a salt

Self-Ionization of Water pure water contains a very small number of H3O+ ions and OH- ions formed by proton transfer from one water molecule to another the equilibrium expression for this reaction is we can treat [H2O] as a constant = 55.5 mol/L

Self-Ionization of Water combining these constants gives a new constant called the ion product of water, Kw in pure water, the value of Kw is 1.0 x 10-14 this means that in pure water

Self-Ionization of Water the product of [H3O+] and [OH-] in any aqueous solution is equal to 1.0 x 10-14 for solutions as well. for example, if we add 0.010 mole of HCl to 1 liter of pure water, it reacts completely with water to give 0.010 mole of H3O+ in this solution, [H3O+] is 0.010 or 1.0 x 10-2 this means that the concentration of hydroxide ion is

pH and pOH we commonly express these concentrations as pH, where pH = -log [H3O+] we can now state the definitions of acidic and basic solutions in terms of pH acidic solution: one whose pH is less than 7.0 basic solution: one whose pH is greater than 7.0 neutral solution: one whose pH is equal to 7.0

pH and pOH just as pH is a convenient way to designate the concentration of H3O+, pOH is a convenient way to designate the concentration of OH- pOH = -log[OH-] the ion product of water, Kw, is 1.0 x 10-14 taking the logarithm of this equation gives pH + pOH = 14 thus, if we know the pH of an aqueous solution, we can easily calculate its pOH

pH and pOH pH of some common materials

pH of Salt Solutions When some salts dissolve in pure water, there is no change in pH from that of pure water Many salts, however, are acidic or basic and cause a change the pH when they dissolve We are concerned in this section with basic salts and acidic salts

pH of Salt Solutions Basic salt: raises the pH as an example of a basic salt is sodium acetate when this salt dissolves in water, it ionizes; Na+ ions do not react with water, but CH3COO- ions do the position of equilibrium lies to the left nevertheless, there are enough OH- ions present in 0.10 M sodium acetate to raise the pH to 8.88

pH of Salt Solutions Acidic salt: lowers the pH an example of an acidic salt is ammonium chloride chloride ion does not react with water, but the ammonium ion does although the position of this equilibrium lies to the left, there are enough H3O+ ions present to make the solution acidic

Acid-Base Titrations Titration: an analytical procedure in which a solute in a solution of known concentration reacts with a known stoichiometry with a substance whose concentration is to be determined

Acid-Base Titrations An acid-base titration must meet these requirement 1. we must know the equation for the reaction so that we can determine the stoichiometric ratio of reactants to use in our calculations 2. the reaction must be rapid and complete 3. there must be a clear-cut change in a measurable property at the end point (when the reagents have combined exactly) 4. we must have precise measurements of the amount of each reactant

Acid-Base Titrations As an example, let us use 0.108 M H2SO4 to determine the concentration of a NaOH solution requirement 1: we know the balanced equation requirement 2: the reaction between H3O+ and OH- is rapid and complete requirement 3: we can use either an acid-base indicator or a pH meter to observe the sudden change in pH that occurs at the end point of the titration requirement 4: we use volumetric glassware

Acid-Base Titrations experimental measurements doing the calculations

pH Buffers pH buffer: a solution that resists change in pH when limited amounts of acid or base are added to it a pH buffer as an acid or base “shock absorber” a pH buffer is common called simply a buffer the most common buffers consist of approximately equal molar amounts of a weak acid and a salt of the conjugate base of the weak acid for example, if we dissolve 1.0 mole of acetic acid and 1.0 mole of its conjugate base (in the form of sodium acetate) in water, we have an acetate buffer

pH Buffers How an acetate buffer resists changes in pH if we add a strong acid, such as HCl, added H3O+ ions react with acetate ions and are removed from solution if we add a strong base, such as NaOH, added OH- ions react with acetic acid and are removed from solution

pH Buffers The effect of a buffer can be quite dramatic consider a phosphate buffer prepared by dissolving 0.10 mole of NaH2PO4 (a weak acid) and 0.10 mole of Na2HPO4 (the salt of its conjugate base) in enough water to make 1 liter of solution

pH Buffers Buffer pH if we mix equal molar amounts of a weak acid and a salt of its conjugate base, the pH of the solution will be equal to the pKa of the weak acid if we want a buffer of pH 9.14, for example, we can mix equal molar amounts of boric acid (H3BO3), pKa 9.14, and sodium dihydrogen borate (NaH2BO3), the salt of its conjugate base

pH Buffers Buffer capacity depends both its pH and its concentration

Blood Buffers The average pH of human blood is 7.4 any change larger than 0.10 pH unit in either direction can cause illness To maintain this pH, the body uses three buffer systems carbonate buffer: H2CO3 and its conjugate base, HCO3- phosphate buffer: H2PO4- and its conjugate base, HPO42- proteins: discussed in Chapter 21

Henderson-Hasselbalch Eg. Henderson-Hasselbalch equation: a mathematical relationship between pH, pKa of the weak acid, HA concentrations HA, and its conjugate base, A- It is derived in the following way taking the logarithm of this equation gives

Henderson-Hasselbalch Eg. multiplying through by -1 gives -log Ka is by definition pKa, and -log [H3O+] is by definition pH; making these substitutions gives rearranging terms gives

Henderson-Hasselbalch Eg. Example: what is the pH of a phosphate buffer solution containing 1.0 mole of NaH2PO4 and 0.50 mole of Na2HPO4 dissolved in enough water to make 1.0 liter of solution

Henderson-Hasselbalch Eg. Example: what is the pH of a phosphate buffer solution containing 1.0 mole of NaH2PO4 and 0.50 mole of Na2HPO4 in enough water to make one liter of solution Solution the equilibrium we are dealing with and its pKa are substituting these values in the H-H equation gives