Chapter 16 Acids and Bases. Chapter 16 Acids and Bases

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1 . Chapter 16 Acids and Bases 1

2 Some Definitions Arrhenius Acid: Substance that, when dissolved in water, increases the concentration of hydrogen ions. Base: Substance that, when dissolved in water, increases the concentration of hydroxide ions. But not all acid base reactions involve water, and many bases (NH3, carbonates) do not contain any OH. Some Definitions Brønsted Lowry theory defines acids and bases in terms of proton (H+) transfer. Acid: Proton donor must have a removable (acidic) proton. Base: Proton acceptor must have a pair of nonbonding electrons. If it can do both... it is AMPHIPROTIC. HCO3, HSO4, H2O 2

3 Conjugate Acids and Bases The conjugate acid of a base is the base plus the attached proton and the conjugate base of an acid is the acid minus the proton Acid conjugate base Base conjugate acid Ionization of HCl Ionization of NH3 3

4 Water is AMPHIPROTIC When an Acid Dissolves in water, water acts as a Brønsted Lowry base and abstracts a proton (H+) from the acid. As a result, the conjugate base of the acid and a hydronium ion are formed. HW: 16.1 PE HW: 16.2 PE 4

5 Strength of Conjugate Acid Base Pairs A stronger acid can donate H+ more readily than a weaker acid. The stronger an acid, the weaker is its conjugate base. The stronger a base, the weaker is its conjugate acid. 5

6 Based on their behavior in water: 1. Strong acids are completely dissociated in water. Their conjugate bases are quite weak having a negligible tendency to be protonated in aqueous solution. 2.Weak acids only dissociate partially in water. Their conjugate bases are weak bases having a slight tendency to remove protons from water. 3. Substances with negligible acidity do not demonstrate any acidic behavior. Their conjugate bases are strong bases, reacting completely with water, removing protons to form OH ions Acid Base Equilibria For equilibrium constant expressions, Ka is used to represent the acid ionization constant and Kb is used to represent the base ionization constant Ka values are used to compare the strengths of weak acids; K, strength 6

7 In any acid base reaction, the equilibrium will favor the reaction that moves the proton from the stonger acid to the stronger base. H2O is a much stronger base than Cl, so the equilibrium lies so far to the right K is not measured (K>>1). H+ is the stongest acid that can exist in equilibrium in an aqueous solution. Acetate is a stronger base than H2O, so the equilibrium favors the left side (K<1). The strong acids HCl, HBr, HI, HNO3, H2SO4, HClO4 are considered strong because they ionize completely in water. The strong acids all appear above H3O+ in the table. The strong acids are leveled to the same strength to that of H3O+ when they are placed in water. 7

8 8

9 Autoionization of Water As we have seen, water is amphoteric. In pure water, a few molecules act as bases and a few act as acids. This is referred to as autoionization Ion Product Constant 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. 14 At 25 C, Kw = 1.0 x 10 = [H3O+][OH ] This equilibrium constant is very important because it applies to all aqueous solutions acids, bases, salts, and nonelectrolytes not just to pure water. 9

10 HW: 16.4 PE HW: 16.5 PE ph ph is defined as the negative base 10 logarithm of the hydronium ion concentration. 10

11 11

12 EX and 16.7 with PE. How Do We Measure ph? For more accurate measurements, one uses a ph meter, which measures the voltage in the solution. 12

13 Strong Acids In solution, the strong acid is usually the only source of H+. Therefore, the ph of a monoprotic acid may be calculated directly from the initial molarity of the acid. Strong Bases 13

14 A student determines the ph of milk of magnesia, a suspension of solid magnesium hydroxide in its saturated aqueous solution, and obtains a value of What is the molarity of Mg(OH)2 in its saturated aqueous solution? The suspended, undissolved Mg(OH)2(s) does not affect the measurement. 14

15 Weak Acids For a generalized acid dissociation, the equilibrium expression would be This equilibrium constant is called the aciddissociation constant, Ka. Weak Acids Weak acids are only partially ionized in aqueous solution. There is a mixture of ions and un ionized acid in solution. Therefore, weak acids are in equilibrium: We can write an equilibrium constant expression for this dissociation: HA(aq) + H2O(l) HA(aq) H3O+(aq) + A (aq) Or: H+(aq) + A (aq) 15

16 Equilibrium Constant Expresasion for Weak Acids Ka is called the acid dissociation constant. The larger the Ka, the stronger the acid. Ka is larger since there are more ions present at equilibrium relative to un ionized molecules. If Ka >> 1, then the acid is completely ionized and the acid is a strong acid. These calculations are similar to the equilibrium calculations performed in Chapter 14. An equation is written for the reversible reaction. Data are organized, often in an I/C/E format. Changes that occur in establishing equilibrium are assessed. Simplifying assumptions are examined (the 5% rule ). Equilibrium concentrations, equilibrium constant, etc. are calculated. 16

17 Calculating Ka from ph In order to find the value of Ka, we need to know all of the equilibrium concentrations. The ph gives the equilibrium concentration of H+. Thus, to find Ka we use the ph to find the equilibrium concentration of H+ and then use the stoichiometric coefficients of the balanced equation to help us determine the equilibrium concentration of the other species. We then substitute these equilibrium concentrations into the equilibrium constant expression and solve for Ka. Calculating Ka and % Ionization from the ph The ph of a 0.10 M solution of formic acid, HCOOH, at 25 C is Calculate Ka for formic acid at this temperature. b) what % of the acid is ionized in this 0.10 M solution? 17

18 To calculate Ka, we need the equilibrium concentrations of all three things. We can find [H3O+], which is the same as [HCOO ], from the ph. Now we can set up a table 18

19 HW: PE Using Ka, we can calculate the concentratio n of H+ (and hence the ph). 1. Write the balanced chemical equation clearly showing the equilibrium. 2.Write the equilibrium expression. Look up the value for Ka (in a table). 3 Write down the initial and equilibrium concentrations for everything except pure water. 4. We usually assume that the equilibrium concentration of H+ is x. 5. Substitute into the equilibrium constant expression and solve. 6. Remember to convert x to ph if necessary. 7. What do we do if we are faced with having to solve a quadratic equation in order to determine the value of x? 8. Often this cannot be avoided. 9. However, if the Ka value is quite small, we find that we can make a simplifying assumption. Assume that x is negligible compared with the initial concentration of the acid. Once we have the value of x, check to see how large it is compared with the initial concentration. If x is <5% of the initial concentration, the assumption is probably a good one. If x>5% of the initial concentration, then it may be best to solve the quadratic equation or use successive approximations. 19

20 Calculating ph from Ka set up a table We are assuming that x will be very small compared to 0.30 and can, therefore, be ignored. 20

21 Calculate the ph of a 0.20M solution of HCN. Ka for ionization of HCN is 4.9 x

22 22

23 Weak Bases Bases react with water to produce hydroxide ion. The equilibrium constant expression for this reaction is where Kb is the base dissociation constant. Weak Bases Kb can be used to find [OH ] and, through it, ph. What is the ph of a 0.15 M solution of NH3? Write the ionization reaction and the corresponding Eq. const. expression. 23

24 HW: PE Types of Weak Bases Weak bases generally fall into one of two categories. 1. Neutral substances with a lone pair of electrons that can accept protons. Most neutral weak bases contain nitrogen. Amines are related to ammonia and have one or more N H bonds replaced with N C bonds (e.g., CH3NH2 is methylamine). 2. Anions of weak acids are also weak bases. Example: ClO is the conjugate base of HClO (weak acid): 24

25 Using ph to Determine the Concentration of a Salt A solution made by adding solid sodium hypochlorite (NaClO) to enough water to make 2.00 L of solution has a ph of Using the information in Equation 16.37, calculate the number of moles of NaClO that were added to the water We can quantify the relationship between the strength of an acid and the strength of its conjugate base. Consider the following equilibria: NH4+(aq) NH3(aq) + H+(aq) NH3(aq) + H2O(l) NH4+(aq) + OH (aq) We can write equilibrium expressions for these reactions: If we add these equations together: NH4+(aq) NH3(aq) + H+(aq) NH3(aq) + H2O(l) NH4+(aq) + OH (aq) The net reaction is the autoionization of water. H2O(l) H+(aq) + OH (aq) Recall that: Kw = [H+][OH ] We can use this information to write an expression that relates the values of Ka, Kb, and Kw for a conjugate acid base pair: Ka x Kb = Kw Alternatively, we can express this as: pka + pkb = pkw = (at 25 C) Thus, the larger Ka (and the smaller pka), the smaller Kb (and the larger pkb). The stronger the acid, the weaker its conjugate base and vice versa 25

26 Calculating Ka or Kb for a Conjugate Acid Base Pair Calculate (a) the base dissociation constant, Kb, for the fluoride ion (F ); (b) the acid dissociation constant, Ka, for the ammonium ion (NH4+) Acid Base Properties of Salt Solutions Acid base properties of salts are a consequence of the reactions of their ions in solution. Many salt ions can react with water to form OH or H+ by a process called hydrolysis. An Anion s Ability to React with Water Consider an anion, X, as the conjugate base of an acid. Anions from weak acids are basic. They will cause an increase in ph. Anions from strong acids are neutral.they do not cause a change in ph. Anions with ionizable protons (e.g., HSO4 ) are amphiprotic. They are capable of acting as an acid or a base. If Ka > Kb the anion will tend to decrease the ph. If Kb > Ka the anion will tend to increase the ph. 26

27 Reactions of Cations with Water Cations with acidic protons (like NH4+) 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 and the metal causes a shift of the electron density in water. This makes the O H bond more polar and the water more acidic. Greater charge and smaller size make a cation more acidic. Effect of Cations and Anions An anion that is the conjugate base of a strong acid will not affect the ph. An anion that is the conjugate base of a weak acid will increase the ph. A cation that is the conjugate acid of a weak base will decrease the ph. Cations of the strong Arrhenius bases will not affect the ph. Other metal ions will cause a decrease in ph. When a solution contains both the conjugate base of a weak acid and the conjugate acid of a weak base, the affect on ph depends on the Ka and Kb values. 27

28 Predicting the Relative Acidity of Salt Solutions List the following solutions in order of increasing ph: (i) 0.1 M Ba(C2H3O2)2, (ii) 0.1 M NH4Cl, (iii) 0.1 M NH3CH3Br, (iv) 0.1 M KNO

29 29

30 Resonance in the conjugate bases of carboxylic acids stabilizes the base and makes the conjugate acid more acidic. 30

31 Combined Effect of Cation and Anion in Solution The ph of a solution may be qualitatively predicted using the following guidelines: 1. Salts derived from a strong acid and a strong base are neutral. Examples are NaCl and Ca(NO3)2. 2. Salts derived from a strong base and a weak acid are basic. Examples are NaClO and Ba(C2H3O2)2. 3. Salts derived from a weak base and a strong acid are acidic. An example is NH4Cl. 4. Salts derived from a weak acid and a weak base can be either acidic or basic. Equilibrium rules apply! We need to compare Ka and Kb for hydrolysis of the anion and the cation. For example, consider NH4CN. Both ions undergo significant hydrolysis. Is the salt solution acidic or basic? The Ka of NH4+ is smaller than the Kb of CN, so the solution should be basic. 31

32 What types of compounds can act as Lewis acids? 1. Lewis acids generally have an incomplete octet (e.g., BF3). 2. Transition metal ions are generally Lewis acids. 3. Lewis acids must have a vacant orbital (into which the electron pairs can be donated). 4. Compounds with multiple bonds can act as Lewis acids. For example, consider the reaction: H2O(l) + CO2(g) H2CO3(aq) Water acts as the electron pair donor and carbon dioxide as the electron pair acceptor in this reaction. Overall, the water (Lewis base) has donated a pair of electrons to the CO2 (Lewis acid). Hydrolysis of Metal Ions The Lewis concept may be used to explain the acidic properties of many metal ions. Metal ions are positively charged and attract water molecules (via the lone pairs on the oxygen atom of water). This interaction is called hydration. Hydrated metal ions act as acids. For example: Fe(H2O)63+(aq) Fe(H2O)5(OH)2+(aq) + H+(aq) Ka = 2 x In general: the higher the charge is, the stronger the M OH2 interaction. the smaller the metal ion is, the more acidic the ion. Thus, the ph of an aqueous solution increases as the size of the ion increases (e.g., Ca2+ vs. Zn2+) and as the charge increases (e.g., Na+ vs. Ca2+ and Zn2+ vs. Al3+). 32

33 The Amphiprotic Behavior of Amino Acids Amino acids are the building blocks of proteins. Each contains a carboxyl group AND an amine group. Thus, amino acids have both acidic and basic groups. They undergo a proton transfer in which the proton of the carboxyl is transferred to the basic nitrogen atom of the amine group. A zwitterion or dipolar ion results. 33

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