Chapter 14 Acids and Bases

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1 Properties of Acids and Bases Chapter 14 Acids and Bases Svante Arrhenius ( ) First to develop a theory for acids and bases in aqueous solution Arrhenius Acids Compounds which dissolve (dissociate) in water to produce H + ions Examples: HCl(aq) H + (aq) + Cl - (aq) HNO 3 (aq) H + (aq) + NO 3- (aq) Arrhenius Bases Ionic compounds which dissolve (dissociate) in water to produce OH - ions Examples: NaOH(aq) Na + (aq) + OH - (aq) Ca(OH) 2 (aq) Ca 2+ (aq) + 2OH - (aq) Arrhenius acids react with Arrhenius bases to form neutral salt solutions Ammonia can also react with an Arrhenius acid to form a neutral salt solution! It s behaving like a base but contains no OH -! 1

2 The Brønsted-Lowry Model A more sophisticated definition of acids and bases is needed Brønsted-Lowry Acids Compounds which donate protons (H + ) in aqueous solution Protons cannot exist on their own in aqueous solution. Instead, they combine with water molecules forming the hydronium ion, H 3 O + : HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq) Brønsted-Lowry Bases Compounds which accept protons (H + ) in aqueous solution. These protons are typically taken from water molecules forming the hydroxide ion, OH - : B(aq) + H 2 O(l) BH + (aq) + OH - (aq) 2

3 Conjugate Acid-Base Pairs Two substances related by the loss or gain of a single proton, H + Acid-Base Equilibria When an acid (HA) donates H + to a base (B), the products are A - and BH + : HA(aq) + B(aq) A - (aq) + BH + (aq) Since the products are also an acid and a base, the reverse reaction can occur in which the acid BH + donates H + to the base A - : A - (aq) + BH + (aq) HA(aq) + B(aq) We therefore have an equilibrium reaction: HA(aq) + B(aq) A - (aq) + BH + (aq) Write the formula and name of the conjugate base for each of the following acids: a. HCN(aq) b. NH 4+ (aq) c. H 2 CO 3 (aq) d. HSO 4- (aq) e. HClO 2 (aq) Write the formula and name of the conjugate acid for each of the following bases: a. CO 2-3 (aq) b. SH - (aq) c. H 2 PO 3- (aq) d. Br - (aq) e. NO 2- (aq) Acid Dissociation Constant, K a For: HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq) K a = [H 3 O + ][A - ]/[HA] = [H + ][A - ]/[HA] H + is normally used to represent the hydrated proton in acid-base reactions rather than H 3 O + The size of K a depends on the strength of the conjugate base, A -, compared to H 2 O If A - is a much weaker base than H 2 O, the equilibrium position will be far to the right (K a large) with most of the dissolved acid in ionized form If A - is a much stronger base than H 2 O the equilibrium position will be far to the left (K a small) with most of the dissolved acid in molecular form 3

4 Acid Strength Determined by the size of the acid dissociation constant, K a, for the dissociation (ionization) equilibrium: HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq) Strong Acids The equilibrium position of its dissociation reaction lies far to the right: HCl(g) + H 2 O(l) H 3 O + (aq) + Cl - (aq) K a large Strong acids are excellent conductors of electricity since they almost completely dissociate into ions [H 3 O + ] ~ [HA] 0 Conjugate base much weaker base than H 2 O Weak Acids The equilibrium position of its dissociation reaction lies far to the left: CH 3 CO 2 H(aq) + H 2 O(l) H 3 O + (aq) + CH 3 CO 2- (aq) K a small Weak acids are poor conductors of electricity since they hardly dissociate and remain mostly in molecular form [H 3 O + ] «[HA] 0 Conjugate base much stronger base than H 2 O 4

5 Summary The Six Common Strong Acids HCl(aq): hydrochloric acid HBr(aq): hydrobromic acid HI(aq): hydroiodic acid HNO 3 (aq): nitric acid H 2 SO 4 (aq): sulfuric acid HClO 4 (aq): perchloric acid You can assume that all the others are weak! Types of Acids Strong acids have weak conjugate bases! The conjugate base of a strong acid has a low attraction for protons It is a much weaker base than water so the water molecule wins the competition for H + ions: HCl(g) + H 2 O(l) H 3 O + (aq) + Cl - (aq) HCl is a strong acid Cl - is a weak conjugate base Weak acids have strong conjugate bases! The conjugate base of a weak acid has a high attraction for protons It is a much stronger base than water so it wins the competition for H + ions over water: CH 3 CO 2 H(aq) + H 2 O(l) H 3 O + (aq) + CH 3 CO 2- (aq) CH 3 CO 2 H is a weak acid CH 3 CO 2- is a strong conjugate base Strengths of Acids and Conjugate Bases 5

6 Amphoteric Substances A substance is referred to as amphoteric if it can as act as both an acid and a base Water is most common example! Water accepts H + in the presence of a stronger acid: HF(g) + H 2 O(l) H 3 O + (aq) + F - (aq) Water donates H + in the presence of a stronger base: NH 3 (g) + H 2 O(l) NH 4+ (aq) + OH - (aq) Autoionization of Water Since water is amphoteric it is possible for it to react with itself forming a conjugate acid-base pair: H 2 O(l) + H 2 O (l) H 3 O + (aq) + OH - (aq) The equilibrium position is far over to the left so only a very tiny amount of H 3 O + and OH - exist in pure water at room temperature Ion-product Constant for Water, K w H 2 O(l) + H 2 O (l) H 3 O + (aq) + OH - (aq) The autoionization of water leads to the equilibrium expression: K w = [H 3 O + ][OH - ] = [H + ][OH - ] where K w is called the ion-product constant for water In pure water at 25 ºC, experiments show that: [H + ] = [OH - ] = 1.0 x 10-7 M so K w = [H + ][OH - ] = (1.0 x 10-7 )(1.0 x 10-7 ) = 1.0 x K w is a measure of the strength of water as an acid 6

7 Applications of K w Since K w is an equilibrium constant, in any aqueous solution at 25 C, regardless of what it contains, the product of [H + ] and [OH - ] must always be equal to 1.0 x Therefore: If [H + ] then [OH - ] and If [H + ] then [OH - ] so that [H + ][OH - ] remains equal to 1.0 x Neutral, Acidic and Basic Solutions 1. [H + ] = [OH - ] solution is neutral 2. [H + ] > [OH - ] solution is acidic 3. [OH - ] > [H + ] solution is basic Compositions of Acidic, Basic and Neutral Solutions For all solutions at 25 C, regardless of whether they are acid, basic or neutral: K w = [H + ][OH - ] = 1.0 x = constant 7

8 Calculating Ion Concentrations in Water If [H + ] or [OH - ] is given, K w can be used to calculate the corresponding [OH - ] or [H + ] concentration at 25 C: K w = [H + ][OH - ] = 1.0 x Since K w is an equilibrium constant, it varies with temperature and so a different value of Kw has to be used if the temperature is not 25 C The ph Scale ph is a number representing the [H + ] concentration in a solution: Acidic Solution ph < 7 Neutral Solution ph = 7 Basic Solution ph > 7 [H + ] > 1.0 x 10-7 M [H + ] = 1.0 x 10-7 M [H + ] < 1.0 x 10-7 M Examples of ph values Measuring ph ph meter Litmus paper Indicators 8

9 Calculating ph The ph of a solution can be calculated if the [H + ] concentration is known from the following formula: ph = -log[h + ] It is a logarithmic scale in which a ph change of 1 corresponds to a factor of 10 change in [H + ] concentration When calculating ph values we need to remember that the number of decimal places in the ph is equal to the number of significant figures in the concentration Example: [H + ] = 1.00 x 10-8 M (3 sf) ph = -log[h + ] = -log(1.00 x 10-8 ) = (3 dp) Other logarithmic scales: poh= -log[oh - ] pk = -log K Relationship between ph and poh Scales Calculating [H + ] from ph ph = -log[h + ] [H + ] = antilog(-ph) = 10 -ph When calculating [H + ] from ph values we need to remember that the number of decimal places in the ph value becomes the number of significant figures in the [H + ] concentration. Its just the opposite of calculating ph from [H + ] Example: ph = 3.45 (2 dp) [H + ] = 10 -ph = = 3.5 x 10-4 M (2 sf) 9

10 Solving Acid-Base Problems Although acid-base problems vary greatly in complexity, the first things to do for any problem are always to: 1. Write the major species in solution 2. Write balanced equations for reactions forming H + 3. Find the species that forms the most H + Major Species in Solution Solution components which are present in the largest amounts Used to determine which components are important and which can be ignored The first and most important step in solving acid-base (or any aqueous solution) problems! Composition Pure Water Strong Acid Weak Acid Major Species H 2 O H +, A -, H 2 O HA, H 2 O What are the major species present in 1.0 M solutions of the following: a. hydrosulfuric acid b. perchloric acid c. permanganic acid d. hydroiodic acid e. sodium chloride Calculating the ph of Strong Acid Solutions The major species in solution are: H +, A - and H 2 O (HA and OH - are present in negligible amounts) The acid-base equilibria present in a strong acid solution are: HA(aq) H + (aq) + A - (aq) K a very large H 2 O(l) H + (aq) + OH - (aq) K w very small The large amount of H + from the strong acid will shift the second equilibrium even further to the left so the tiny amount of H + coming from the autoionization of water can be safely ignored The only important source of H + therefore comes from the strong acid, and assuming [H + ] ~ [HA] 0, we can calculate the ph of the solution 10

11 Calculating the ph of Weak Acid Solutions Procedure Calculating the ph of Weak Acid Solutions The major species in solution are: HA and H 2 O The acid-base equilibria present in a weak acid solution are: HA(aq) H + (aq) + A - (aq) K a small H 2 O(l) H + (aq) + OH - (aq) K w very small The most important source of H + normally comes from the weak acid, allowing us to calculate [H + ] and hence the ph of the solution using the same method as used in Chapter 13 for systems with small equilibrium constants 11

12 Calculating the ph of a Mixture of Weak Acids If one of the weak acids is much stronger than the other (K a1» K a2 ) then we can assume that the dominant source of H + will be from the strongest of the two and follow the usual method For solutions of any weak acid HA, H + decreases as [HA] 0 decreases, but the percent dissociation increases as [HA] 0 decreases: Percent Dissociation (Ionization) Used the specify the amount of weak acid dissociated (ionized) at equilibrium: Calculating % dissociation involves the same calculation as used in checking the 5% rule! Percent dissociation can be used to calculate K a! 12

13 Base Dissociation Constant, K b For: B(aq) + H 2 O(l) BH + (aq) + OH - (aq) K b = [BH + ][OH - ]/[B] The position of the equilibrium depends on the competition between the two bases B and OH - for the proton If B is a much stronger base than OH - the equilibrium position will be far to the right with most of the dissolved base in ionized form If OH - is a much stronger base than B the equilibrium position will be far to the left with most of the dissolved base in molecular form 13

14 Strong Bases Hydroxides of Group 1A metals (Li, Na, K, Rb and Cs) are strong bases which completely dissociate in aqueous solution: LiOH(aq) Li + (aq) + OH - (aq) Hydroxides of Group 2A metals (Ca, Sr and Ba) are also strong bases forming 2 moles of OH - per mole base: Ca(OH) 2 (aq) Ca 2+ (aq) + 2OH - (aq) However, Group 2A hydroxides have low solubilities Calculating the ph of Strong Base Solutions The major species in solution are: M +, OH - and H 2 O (MOH and H + are present in very small amounts) The equilibria present in a strong base solution are: MOH(aq) M + (aq) + OH - (aq) K b very large H 2 O(l) H + (aq) + OH - (aq) K w very small The large amount of OH - from the strong base will shift the second equilibrium even further to the left so the tiny amount of OH - coming from the autoionization of water can be safely ignored The only important source of OH - therefore comes from the strong base, and assuming [OH - ] ~ [MOH] 0, we can then calculate [H + ] followed by the ph of the solution 14

15 Weak Bases Most weak bases contain nitrogen atoms with lone pairs onto which protons can attach: Calculating the ph of Weak Base Solutions The major species in solution are: B and H 2 O The acid-base equilibria present in a weak acid solution are: B(aq) +H 2 O(aq) BH + (aq) + OH - (aq) K b small H 2 O(l) H + (aq) + OH - (aq) K w very small The most important source of OH - normally comes from the reaction of the weak base with water which allows us to calculate [OH - ], [H + ] and hence the ph of the solution using a similar method to that used for weak acids Polyprotic Acids Acids that can produce more than one proton Polyprotic acids always dissociate in a stepwise manner with the conjugate base of one step becoming the acid in the next step Example for a generic triprotic acid: H 3 A(aq) H + (aq) + H 2 A - (aq) K a1 = [H + ][H 2 A - ]/[H 3 A] H 2 A - (aq) H + (aq) + HA 2- (aq) K a2 = [H + ][HA 2- ]/[H 2 A - ] HA 2- (aq) H + (aq) + A 3- (aq) K a3 = [H + ][A 3- ]/[HA 2- ] Typically K a1 > K a2 > K a3 with each acid becoming successively weaker since the increasing negative charge on the acid makes it more difficult to remove the proton 15

16 Since K a1 is usually much larger than the others, only the first dissociation step makes a significant contribution to [H + ] and hence the ph This simplifies the calculations required for most polyprotic acids Sulfuric Acid (H 2 SO 4 ) Sulfuric acid is unique among the polyprotic acids in that it is a strong acid in its first dissociation step and a weak acid in its second step: All other polyprotic acids are weak acids in each step 16

17 Summary Acid-Base Properties of Salts When an ionic compound dissolves in water, it dissociates into ions. These solutions can be acid, basic or neutral depending on the acidic and basic properties of the cations and anions present Will a solution of NaNO 3 be neutral, acidic or basic? Na + is a cation of a strong base and has no affinity for H +, nor can it produce H + so it have no effect on [H + ] and therefore the ph NO 3- is weak conjugate base of a strong acid so it WILL NOT remove H + from water since it prefers to remain dissociated as NO 3- (aq) The solution forms neither H + or OH - and will therefore be neutral with a ph of 7 Salts that produce Neutral Solutions A solution of a salt containing the cation from a strong base and the anion from a strong acid will be neutral Examples: Alkali metal chlorides, bromides, iodides, nitrates, sulfates and perchlorates Will a solution of NaC 2 H 3 O 2 be neutral, acidic or basic? Na + is a cation of a strong base and has no affinity for H +, nor can it produce H + so it have no effect on [H + ] and therefore the ph C 2 H 3 O 2- is a strong conjugate base of a weak acid and so it will remove H + from water forming OH- since it prefers to remain in its molecular form: C 2 H 3 O 2- (aq) + H 2 O(l) HC 2 H 3 O 2 (aq) + OH - (aq) The solution forms OH - and will therefore be basic with a ph of >7 17

18 Salts that produce Basic Solutions A solution of a salt containing the cation from a strong base and the anion from a weak acid will be basic Examples: Alkali metal fluorides, cyanides, carbonates, nitrites, acetates, sulfides, phosphates Will a solution of NH 4 Cl be neutral, acidic or basic? Since Cl - is a weak conjugate base of a strong acid it will not remove H + from water since it prefers to remain dissociated as Cl - (aq) Since NH 4+ is a strong conjugate acid of a weak base it will donate H + to water since it prefers to remain in the form of NH 3 molecules: NH 4+ (aq) + H 2 O(aq) NH 3 (aq) + H 3 O + (aq) Salts that produce Acidic Solutions A solution of a salt containing the cation from a weak base and the anion from a strong acid will be acidic Examples: Ammonium compounds The solution forms H + and will therefore be acidic with a ph of <7 Relationship between K a and K b for Weak Acids For a weak acid: HA(aq) H + (aq) + A - (aq) K a = [H + ][A - ]/[HA] For the conjugate base of this weak acid: A - (aq) + H 2 O(l) HA(aq) + OH - (aq) K b = [HA][OH - ]/[A - ] Multiplying K a and K b : K a x K b = ([H + ][A - ]/[HA]) x ([HA][OH - ]/[A - ]) = [H + ][OH - ] = K w = 1.0 x For any weak base and its conjugate acid: K a x K b = K w 18

19 Hydrated Cations Salts that contain small, highly charged metal ions (Al 3+, Fe 3+, Cr 3+ ) in combination with anions from strong acids will also be acidic since the hydrated cations formed in water are weak acids: Al(H 2 O) 6 3+ (aq) Al(OH)(H 2 O) H + (aq) Will a solution of NH 4 F be neutral, acidic or basic? Since F - is a strong conjugate base of a weak acid it will try to remove H + from water since it prefers to remain undissociated as HF(aq) Since NH 4+ is a strong conjugate acid of a weak base it wants to donate H + to water since it prefers to remain in the form of NH 3 (aq) molecules It is impossible to tell from this information alone since it will depend on the relative strengths of the acid (NH 4+ ) and the base (F - ) Solutions of salts containing the cations of weak bases and the anions of weak acids can be acidic, basic or depending on the relative size of K a for the acidic ion compared to K b for the basic ion Their ph s can be predicted qualitatively as follows: Predict whether an aqueous solutions of the following salts will be acidic, basic or neutral: NH 4 C 2 H 3 O 2 K a for NH 4+ = 5.6 x 10-10, K b for C 2 H 3 O 2- = 5.6 x NH 4 CN K a for NH 4+ = 5.6 x 10-10, K a for HCN = 6.3 x Al 2 (SO 4 ) 3 K a for Al(H 2 O) 6 3+ = 1.4 x 10-5, K a2 for H 2 SO 4 = 1.2 x 10-2 Quantitative calculations on these salts are complicated! 19

20 This method can in fact be used to determine the qualitative ph of any salt solution! Classification of Cations and Anions in Solution Summary 20

21 C-H bonds are both strong and non-polar so molecules containing these bonds do not show acidic properties Effect of Structure on Acid-Base Properties Any molecule containing a hydrogen atom can potentially act as an acid There are two factors which determine whether a molecule containing an X-H will behave as an acid: Hydrogen halides on the other hand have polar bonds: H-F > H-Cl > H-Br > H-I most polar least polar Since the H-F bond is the most polar we might expect that HF would be the strongest acid, but it is in fact a weak acid due the fact that the H-F bond is very strong 1. Bond strength 2. Bond polarity The stronger the bond and the lower the polarity the lower the probability that the proton will be released Oxyacids are very common and contain the H-O-X functional group For these acids, their K a values increase as the number of electronegative oxygen atoms attached to the central X atoms increase since they are able to draw electrons away from the O-H bond and thereby weaken it: This is also true for electronegative atoms other than oxygen: 21

22 Acid-Base Properties of Oxides When covalent oxides dissolve in water they produce acidic solutions since the strong covalent bonds within the molecule remain intact while the O-H bonds in water break to produce protons: These compounds are called acidic oxides When ionic oxides dissolve in water they produce basic solutions: This occurs because the oxide ion reacts with water to form the hydroxide ion: These compounds are called basic oxides The Lewis Acid-Base Model The Lewis model is even more general than the Brønsted-Lowry Model: A Lewis Acid is an electron-pair acceptor since it has an empty atomic orbital that it can use to accept (share) from another molecule which has a lone pair A Lewis Base is an electron-pair donor since it has a lone pair of electrons which can be donated (shared) with another molecule which has an empty atomic orbital G. N. Lewis The Lewis Model has the advantage that it covers many reactions which do not involve Brønsted-Lowry acids: 22

23 Acid-Base Theories Compared 23

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