Chapter 10 Acids and Bases Arrhenius Acids and

Chapter 10 Acids and Bases

Arrhenius Acids and Bases l In 1884, Svante Arrhenius proposed these definitions l l acid: a substance that produces H 3 O+ ions aqueous solution base: a substance that produces OH- ions in aqueous solution

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

Arrhenius Acids and Bases l With bases, the situation is slightly different l l l many bases are metal hydroxides such as KOH, Na. OH, 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 l we use curved arrows to show the transfer of a proton from water to ammonia

Acid and Base Strength l l l Strong acid: one that reacts completely or almost completely with water to form H 3 O+ 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 l Weak acid: a substance that dissociates only partially in water to produce H 3 O+ ions l l 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 l ammonia, for example, is a weak base

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

Brønsted-Lowry Acids & Bases l 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 H 3 O+, H 2 CO 3, and H 2 PO 42. a base can be negatively charged or neutral; examples are OH-, Cl-, and NH 3 3. acids are classified a monoprotic, diprotic, or triprotic depending on the number of protons each may give up; examples are HCl, H 2 CO 3, and H 3 PO 4 l

Brønsted-Lowry Acids & Bases l 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 l l l the HCO 3 - ion, for example, can give up a proton to become CO 32 -, or it can accept a proton to become H 2 CO 3 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 H 2 O; it can accept a proton to become H 3 O+, 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 l 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 l l l 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 CH 3 COOH (acetic acid) is a stronger acid that H 2 CO 3 (carbonic acid); conversely, CH 3 COO- (acetate ion) is a weaker base that HCO 3 - (bicarbonate ion)

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

Acid-Base Equilibria l To predict the position of an acid-base equilibrium such as this, we do the following l l 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 l identify the two acids and bases, and their relative strengths l the position of this equilibrium lies to the right

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

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

Acid Ionization Constants l when a weak acid, HA, dissolves in water l the equilibrium constant, Keq, for this ionization is l because water is the solvent and its concentration changes very little when we add HA to it, we treat [H 2 O] 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 l

Acid Ionization Constants l l l 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 p. Ka where the value of p. Ka for acetic acid is 4. 75 values of Ka and p. Ka 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 p. Ka the weaker the acid, the smaller its Ka, but the larger its p. Ka

Properties of Acids & Bases l Neutralization l l acids and bases react with each other in a process called neutralization. Reaction of acids with metals l strong acids react with certain metals (called active metals) to produce a salt and hydrogen gas, H 2

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

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

Properties of Acids & Bases l Reaction with carbonates and bicarbonates l strong acids react with carbonates to give carbonic acid, which rapidly decomposes to CO 2 and H 2 O l strong acids react similarly with bicarbonates

Properties of Acids & Bases l Reaction with ammonia and amines l any acid stronger than NH 4+ is strong enough to react with NH 3 to give a salt

Self-Ionization of Water l pure water contains a very small number of H 3 O+ ions and OH- ions formed by proton transfer from one water molecule to another l the equilibrium expression for this reaction is l we can treat [H 2 O] as a constant = 55. 5 mol/L

Self-Ionization of Water l 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 l this means that in pure water l

Self-Ionization of Water l l the product of [H 3 O+] 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 H 3 O+ in this solution, [H 3 O+] is 0. 010 or 1. 0 x 10 -2 this means that the concentration of hydroxide ion is

p. H and p. OH l l l we commonly express these concentrations as p. H, where p. H = -log [H 3 O+] we can now state the definitions of acidic and basic solutions in terms of p. H acidic solution: one whose p. H is less than 7. 0 basic solution: one whose p. H is greater than 7. 0 neutral solution: one whose p. H is equal to 7. 0

p. H and p. OH l l just as p. H is a convenient way to designate the concentration of H 3 O+, p. OH is a convenient way to designate the concentration of OHp. OH = -log[OH-] the ion product of water, Kw, is 1. 0 x 10 -14 taking the logarithm of this equation gives p. H + p. OH = 14 thus, if we know the p. H of an aqueous solution, we can easily calculate its p. OH

p. H and p. OH l p. H of some common materials

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

p. H of Salt Solutions l Basic salt: raises the p. H l l 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 CH 3 COO- 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 p. H to 8. 88

p. H of Salt Solutions l Acidic salt: lowers the p. H l l l an example of an acidic salt is ammonium 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 H 3 O+ ions present to make the solution acidic

Acid-Base Titrations l 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 l 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 l As an example, let us use 0. 108 M H 2 SO 4 to determine the concentration of a Na. OH solution l requirement 1: we know the balanced equation l requirement 2: the reaction between H 3 O+ and OH- is rapid and complete requirement 3: we can use either an acid-base indicator or a p. H meter to observe the sudden change in p. H that occurs at the end point of the titration requirement 4: we use volumetric glassware l l

Acid-Base Titrations l experimental measurements l doing the calculations

p. H Buffers l p. H buffer: a solution that resists change in p. H when limited amounts of acid or base are added to it l l a p. H buffer as an acid or base “shock absorber” a p. H 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

p. H Buffers l How an acetate buffer resists changes in p. H l if we add a strong acid, such as HCl, added H 3 O+ ions react with acetate ions and are removed from solution l if we add a strong base, such as Na. OH, added OH- ions react with acetic acid and are removed from solution

p. H Buffers l The effect of a buffer can be quite dramatic l consider a phosphate buffer prepared by dissolving 0. 10 mole of Na. H 2 PO 4 (a weak acid) and 0. 10 mole of Na 2 HPO 4 (the salt of its conjugate base) in enough water to make 1 liter of solution

p. H Buffers l Buffer p. H l l if we mix equal molar amounts of a weak acid and a salt of its conjugate base, the p. H of the solution will be equal to the p. Ka of the weak acid if we want a buffer of p. H 9. 14, for example, we can mix equal molar amounts of boric acid (H 3 BO 3), p. Ka 9. 14, and sodium dihydrogen borate (Na. H 2 BO 3), the salt of its conjugate base

p. H Buffers l Buffer capacity depends both its p. H and its concentration

Blood Buffers l The average p. H of human blood is 7. 4 l l any change larger than 0. 10 p. H unit in either direction cause illness To maintain this p. H, the body uses three buffer systems l l l carbonate buffer: H 2 CO 3 and its conjugate base, HCO 3 phosphate buffer: H 2 PO 4 - and its conjugate base, HPO 42 proteins: discussed in Chapter 21

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

Henderson-Hasselbalch Eg. l multiplying through by -1 gives l -log Ka is by definition p. Ka, and -log [H 3 O+] is by definition p. H; making these substitutions gives l rearranging terms gives

Henderson-Hasselbalch Eg. l Example: what is the p. H of a phosphate buffer solution containing 1. 0 mole of Na. H 2 PO 4 and 0. 50 mole of Na 2 HPO 4 dissolved in enough water to make 1. 0 liter of solution

Henderson-Hasselbalch Eg. l l Example: what is the p. H of a phosphate buffer solution containing 1. 0 mole of Na. H 2 PO 4 and 0. 50 mole of Na 2 HPO 4 in enough water to make one liter of solution Solution l the equilibrium we are dealing with and its p. Ka are l substituting these values in the H-H equation gives