Chapter 16. Some Definitions. Some Definitions. What Happens When an Acid Dissolves in Water? 27/07/2014. If it can be either HCO 3 HSO 4 H 2 O

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1 Chemistry, The Central Science, 11th edition Theodore L. Brown, H. Eugene LeMay, Jr., Bruce E. Bursten Some Definitions Chapter 16 AP Chemistry North Nova Education Centre Mr. Gauthier Arrhenius An acid is a substance that, when dissolved in water, increases the concentration of hydrogen ions. A base is a substance that, when dissolved in water, increases the concentration of hydroxide ions. Some Definitions Brønsted-Lowry An acid is a proton donor. A base is a proton acceptor. A Brønsted-Lowry acid must have a removable (acidic) proton. A Brønsted-Lowry base must have a pair of nonbonding electrons. If it can be either it is amphiprotic. HCO 3 - HSO 4 - H 2 O What Happens When an Acid Dissolves in Water? Water acts as a Brønsted-Lowry base abstracts a proton (H + ) from the acid. As a result, the conjugate base of the acid a hydronium ion are formed. 1

2 Conjugate Acid Base Strength The term conjugate comes from the Latin word conjugare, meaning to join together. Reactions between acids bases always yield their conjugate bases acids. Strong acids are completely dissociated in water. Their conjugate bases are quite weak. Weak acids only dissociate partially in water. Their conjugate bases are weak bases. Acid Base Strength Acid Base Strength Substances with negligible acidity do not dissociate in water. Their conjugate bases are exceedingly strong. In any acid-base reaction, the equilibrium will favor the reaction that moves the proton to the stronger base. HCl (aq) + H 2 O (l) H 3 O + (aq) + Cl - (aq) H 2 O is a much stronger base than Cl -, so the equilibrium lies so far to the right that K is not measured (K>>1). Acid Base Strength Autoionization of Water In any acid-base reaction, the equilibrium will favor the reaction that moves the proton to the stronger base. CH 3 CO 2 H (aq) + H 2 O (l) H 3 O + (aq) + CH 3 CO 2 - (aq) Acetate is a stronger base than H 2 O, so the equilibrium favors the left side (K<1). As we have seen, water is amphoteric. In pure water, a few molecules act as bases a few act as acids. H 2 O (l) + H 2 O (l) H 3 O + (aq) + OH - (aq) This is referred to as autoionization. 2

3 Ion-Product Constant ph The equilibrium expression for this process is K c = [H 3 O + ] [OH - ] This special equilibrium constant is referred to as the ion-product constant for water, K w. At 25 C, K w = ph is defined as the negative base-10 logarithm of the concentration of hydronium ion. ph = -log [H 3 O + ] ph In pure water, K w = [H 3 O + ] [OH - ] = Since in pure water [H 3 O + ] = [OH - ], ph Therefore, in pure water, ph = -log ( ) = 7.00 An acid has a higher [H 3 O + ] than pure water, so its ph is <7. A base has a lower [H 3 O + ] than pure water, so its ph is >7. [H 3 O + ] = = ph Other p Scales These are the ph values for several common substances. The p in ph tells us to take the negative base-10 logarithm of the quantity (in this case, hydronium ions). Some similar examples are poh: -log [OH - ] pk w : -log K w 3

4 Watch This! Because [H 3 O + ] [OH - ] = K w = , we know that -log [H 3 O + ] + -log [OH - ] = -log K w = or, in other words, ph + poh = pk w = How Do We Measure ph? For less accurate measurements, one can use Litmus paper Red paper turns blue above ~ph = 8 Blue paper turns red below ~ph = 5 Or an indicator. How Do We Measure ph? For more accurate measurements, one uses a ph meter, which measures the voltage in the solution. Strong You will recall that the seven strong acids are HCl, HBr, HI, HNO 3, H 2 SO 4, HClO 3, HClO 4. These are, by definition, strong electrolytes exist totally as ions in aqueous solution. For the monoprotic strong acids, [H 3 O + ] = [acid]. Strong Dissociation Constants Strong bases are the soluble hydroxides, which are the alkali metal heavier alkaline earth metal hydroxides (Ca 2+, Sr 2+, Ba 2+ ). Again, these substances dissociate completely in aqueous solution. For a generalized acid dissociation, HA (aq) + H 2 O (l) the equilibrium expression would be K c = [H 3 O + ] [A - ] [HA] A - (aq) + H 3 O + (aq) This equilibrium constant is called the acid-dissociation constant, K a. 4

5 Dissociation Constants Calculating K a from the ph The greater the value of K a, the stronger is the acid. The ph of a 0.10 M solution of formic acid, HCOOH, at 25 C is Calculate K a for formic acid at this temperature. We know that K a = [H 3 O + ] [COO - ] [HCOOH] Calculating K a from the ph Calculating K a from the ph The ph of a 0.10 M solution of formic acid, HCOOH, at 25 C is Calculate K a for formic acid at this temperature. To calculate K a, we need the equilibrium concentrations of all three things. We can find [H 3 O + ], which is the same as [HCOO - ], from the ph. ph = -log [H 3 O + ] 2.38 = -log [H 3 O + ] = log [H 3 O + ] = 10 log [H 3O+] = [H 3 O + ] = [H 3 O + ] = [HCOO - ] Calculating K a from ph Calculating K a from ph Now we can set up a table [HCOOH], M [H 3 O + ], M [HCOO - ], M Initially Change K a = [ ] [ ] [0.10] = At Equilibrium = =

6 Calculating Percent Ionization Calculating ph from K a [H 3 O + ] eq Percent Ionization = [HA] 100 initial In this example [H 3 O + ] eq = M [HCOOH] initial = 0.10 M Percent Ionization = Calculate the ph of a 0.30 M solution of acetic acid, HC 2 H 3 O 2, at 25 C. HC 2 H 3 O 2 (aq) + H 2 O (l) H 3 O + (aq) + C 2 H 3 O 2 - (aq) K a for acetic acid at 25 C is = 4.2% Calculating ph from K a The equilibrium constant expression is Calculating ph from K a We next set up a table [C 2 H 3 O 2 ], M [H 3 O + ], M [C 2 H 3 O 2- ], M K a = [H 3 O + ] [C 2 H 3 O 2- ] [HC 2 H 3 O 2 ] Initially Change -x +x +x At Equilibrium x 0.30 x x We are assuming that x will be very small compared to 0.30 can, therefore, be ignored. Calculating ph from K a Calculating ph from K a Now, = ( ) (0.30) = x 2 (x) 2 (0.30) ph = -log [H 3 O + ] ph = -log ( ) ph = = x = x 6

7 Polyprotic have more than one acidic proton If the difference between the K a for the first dissociation subsequent K a values is 10 3 or more, the ph generally depends only on the first dissociation. Weak react with water to produce hydroxide ion. Weak The equilibrium constant expression for this reaction is Weak K b can be used to find [OH - ], through it, ph. K b = [HB] [OH - ] [B - ] where K b is the base-dissociation constant. ph of Basic Solutions ph of Basic Solutions What is the ph of a 0.15 M solution of NH 3? NH 3 (aq) + H 2 O (l) NH 4 + (aq) + OH - (aq) [NH 4+ ] [OH - ] K b = = [NH 3 ] -5 Tabulate the data. [NH 3 ], M [NH 4+ ], M [OH - ], M Initially At Equilibrium x 0.15 x x 7

8 ph of Basic Solutions ph of Basic Solutions = ( ) (0.15) = x = x = x 2 (x) 2 (0.15) Therefore, [OH - ] = M poh = -log ( ) poh = 2.80 ph = ph = K a K b Reactions of Anions with Water K a K b are related in this way: K a K b = K w Therefore, if you know one of them, you can calculate the other. Anions are bases. As such, they can react with water in a hydrolysis reaction to form OH - the conjugate acid: X - (aq) + H 2 O (l) HX (aq) + OH - (aq) Reactions of Cations with Water Reactions of Cations with Water Cations with acidic protons (like NH 4+ ) will lower the ph of a solution. Most metal cations that are hydrated in solution also lower the ph of the solution. Attraction between nonbonding electrons on oxygen the metal causes a shift of the electron density in water. This makes the O-H bond more polar the water more acidic. Greater charge smaller size make a cation more acidic. 8

9 Effect of Cations Anions Effect of Cations Anions 1. An anion that is the conjugate base of a strong acid will not affect the ph. 2. An anion that is the conjugate base of a weak acid will increase the ph. 3. A cation that is the conjugate acid of a weak base will decrease the ph. 4. Cations of the strong Arrhenius bases will not affect the ph. 5. Other metal ions will cause a decrease in ph. 6. When a solution contains both the conjugate base of a weak acid the conjugate acid of a weak base, the affect on ph depends on the K a K b values. Factors Affecting Acid Strength Factors Affecting Acid Strength The more polar the H-X bond /or the weaker the H-X bond, the more acidic the compound. So acidity increases from left to right across a row from top to bottom down a group. In oxyacids, in which an -OH is bonded to another atom, Y, the more electronegative Y is, the more acidic the acid. Factors Affecting Acid Strength Factors Affecting Acid Strength Resonance in the conjugate bases of carboxylic acids stabilizes the base makes the conjugate acid more acidic. For a series of oxyacids, acidity increases with the number of oxygens. 9

10 Lewis Lewis Lewis acids are defined as electron-pair acceptors. Atoms with an empty valence orbital can be Lewis acids. Lewis bases are defined as electron-pair donors. Anything that could be a Brønsted-Lowry base is a Lewis base. Lewis bases can interact with things other than protons, however. 10

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