Chapter 15: Acids and Bases

Transcription

1 Chapter 15: Acids and Bases Essentials of General Chemistry Ebbing Gammon Ragsdale 2nd Edition Dr. Azra Ghumman Memorial University of Newfoundland 1 Acids and Bases 15.1 Arrhenius Concept of Acids and Bases 15.2 Bronsted Lowery Concept of Acids and bses 15.3 Lewis Concept of Acids and Bases 15.4 Relative Strengths of Acids and Bases 15.6 Self-Ionization of Water 15.7 Solutions of a Strong Acid or Base 15.8 The ph of a Solution Section 15.5 is not covered 2 1

2 Acid Base concepts In 1884 Svante Arrhenius explained the actual cause of acidity and basicity in terms of the effect these compounds have on water. 1. Arrhenius Acid-Base concept. 2. Bronsted-Lowery Acid-Base concept. 3. Lewis Acid-Base concept Arrhenius Concept of Acids and Bases According to Arrhenius concept, An acid is a substance that, when dissolved in water, increases the concentration of hydronium ion (H 3 O ). HA(aq) H 2 O(l)Æ H 3 O (aq) A - (aq) Hydronium Ion H 3 O,chemists often use the notation H (aq) for the H 3 O (aq) ion, and call it the hydrogen ion. The aqueous hydrogen ion is too reactive and actually chemically bonds to water, that is H 3 O. Hydronium ion in turn bonds to other H 2 O molecules through hydrogen bonding (hydrate of H ion H (H 2 O)). 4 2

3 Arrhenius Concept of Acids and Bases The H 3 O is shown here hydrogen bonded to three water molecules Forming hydrate of hydrogen ion with general formula[h (H 2 O) n ]. where n=1,2,3,4 5 Arrhenius Concept of Acids In the Arrhenius concept, A strong acid is a substance that ionizes completely (almost 100%) in aqueous solution to give H 3 O (aq) and an anion. An example is perchloric acid, HClO 4. HClO4 (aq) H2O(l) H3O (aq) ClO4 (aq) Other strong acids include HCl, HBr, HI, HNO 3, and H 2 SO 4. All other acids are classified as weak acids. A weak acid is a substance that does not ionizes completely in aqueous solution 6 3

4 Arrhenius Base Arrhenius base, is a substance that, when dissolved in water, increases the concentration of hydroxide ion, OH - (aq). H ( ) 2 O BOH s B ( aq) OH ( aq) A strong base is a substance that ionizes completely (almost 100%) in aqueous solution to give OH - (aq) and a cation. An example is sodium hydroxide, NaOH. H 2 O NaOH(s) Na (aq) OH (aq) Other strong bases include LiOH, KOH, Ca(OH) 2, Sr(OH) 2, and Ba(OH) 2. Mostly strong bases are hydroxide of Group 1A and Group IIA elements(except Be). 7 Common Strong Acids and Bases 8 4

5 Neutralization Reaction The neutralization Reaction is the reaction between a strong acid and a strong base forming water and a salt. is essentially the reaction of H 3 O (aq) and OH - (aq). For example the neutralization of HCl(aq) with NaOH(aq) in ionic form, The net ionic equation is as follows H 3 O (aq) OH - (aq) 2H 2 O(l) 9 Weak Acids and Bases Most other acids and bases that we encounter are weak. Weak Acids are not completely ionized and exist in equilibrium with the corresponding ions. Aqueous solution mainly contains un-dissociated acid in equilibrium with some ions. An example is acetic acid, HC 2 H 3 O 2 or HF(aq). HC 2 H 3 O 2 (aq) H 2 O(l) H 3 O (aq) C 2 H 3 O 2- (aq) Ionize less than 1% Weak bases are not completely ionized and exist in equilibrium with the some ions in solution. 10 5

6 Limitation of Arrhenius concept The Arrhenius concept is limited in that; it looks at acid and base reactions in aqueous solutions only. It singles out the OH - ion as the source of base character, when other species can play a similar role e.g. NH 3 Broader definitions of acids and bases are required 11 Reaction of HCl(g) and NH 3 (g) to form NH 4 Cl(s) Gas phase reaction of HCl and NH 3. Figure

7 15.2 Brønsted-Lowry Concept of Acids and Bases In 1923 J. N. Bronsted, a Danish chemist and Thomas M. Lowery, a British chemist pointed out that acid base reactions are proton-transfer reactions. According to the Brønsted-Lowry concept, B-L acid is the species donating a proton in a protontransfer reaction (proton donor). HA(aq) H 2 O(l) H 3 O (aq) A - (aq) B-L base is the species accepting a proton in a protontransfer reaction (proton acceptor). In any reversible acid-base reaction, both forward and reverse reactions involve proton transfer. 13 Brønsted-Lowry Concept of Acids and Bases B-L base, NH 3 dissolves in H 2 O; Base acid acid base NH3(aq) H2O(l) NH4 (aq) OH (aq) H H In the forward reaction, NH 3 accepts a proton from H 2 O. Thus, NH 3 is a base and H 2 O is an acid In the reverse reaction, NH 4 donates a proton to OH -. The NH 4 ion is an acid and OH - is a base 14 7

8 Brønsted-Lowry Concept of Acids and Bases Representation of the Reaction H 3 O (aq) NH 3 (aq) H 2 O (l) NH 4 (aq) This reaction can occurs in aq. solution, in benzene, and in gas phase. The charges associated for ions are overall charges 15 A Conjugate Acid Base Pair A conjugate acid-base pair consists of two species in an acid-base reaction, one acid and one base, that differ by the loss or gain of a proton. base acid acid base NH3 (aq) H2O(l) NH4 (aq) OH (aq) The species NH 4 and NH 3 are a conjugate acid-base pair. Here NH 4 is the conjugate acid of NH 3 and NH 3 is the conjugate base of NH 4. The Brønsted-Lowry concept defines a species as an acid or a base according to its function in the proton-transfer reaction. 16 8

9 Conjugate Acid-base pairs Examples:Bronsted acids (molecular and ionic species) Conjugate acid Conjugate base HF F- HSO - 4 SO 2-4 NH 4 NH 3 HCN CN - 17 Amphiprotic Species Amphiprotic species is a species that can act either as an acid or a base (it can gain or lose a proton). For example, HCO 3- acts as a proton donor (an acid) in the presence of OH - and can act as a proton acceptor (a base) in the presence of HF. 2 HCO 3 (aq) OH (aq) CO3 (aq) H2O(l) H HCO 3 (aq ) HF (aq ) H 2CO 3 (aq ) F (aq ) H 18 9

10 Amphiprotic Species The amphiprotic characteristic of water is important in the acid-base properties of aqueous solutions. H [ O H ] H H O H H H O H hydroxide water hydroxide ion ion Water can react as a base with the acid HF and also as an acid with base NH 3. Amphoteric A species that can act as an acid or a base. Need not be amphiprotic (contain a proton) e.g. Al 2 O 3 [ ] 19 Importance of Bronsted Lowery Cocept 1. A base is a species that accepts proton; OH - is only one example of a base. 2. Acids and bases can be ions as well as molecular substances. 3. Acid-base reactions are not restricted to aqueous solution. 4. Some species can act as either acids or bases depending on what the other reactant is

11 Identifying Acid and Base Species Give the conjugate acid of each of the following species regarded as bases. a. BrO - b. PH 3 c. HPO 4 2- d. PH 2 - e. ClO 2-21 Lewis Concept of Acids and Bases According to Lewis concept (G.N. Lewis) Lewis acid is a species that can form a covalent bond by accepting an electron pair from another species (an electron pair acceptor). Lewis base is a species that can form a covalent bond by donating an electron pair to another species (an electron pair donor). This concept broadened the scope of acid-base theory to include reactions that did not involve H. The Lewis concept explains many reactions that we might not think of as acid-base reactions e.g. Na 2 O(s) SO 3 (g) Na 2 SO 4 (s) 22 11

12 Lewis Concept of Acids and Bases The reaction of boron trifluoride with ammonia is an example : : F B : : : : : F F : H : N H H Boron trifluoride accepts the electron pair, so it is a Lewis acid. Ammonia donates the electron pair, so it is the Lewis base. The formation of complex ions can also be looked as Lewis acid-base reactions. : : : F B : : : : : F F : : H N H H 23 Reaction of Boron Trifluoride with Ammonia BF 3 NH 3 (Lewis acid) (Lewis base) 24 12

13 Hydrated metal ion or complex ions The formation of complex ions can also be looked as Lewis acidbase reactions. Metal ion bonds to e - pairs from molecules or ions like H 2 O, NH 3 or CN - e.g. Al(H 2 O) Identifying Lewis Acid and Base Species Example: In the following reactions, identify the Lewis acid and Lewis Base. Write the chemical equation using electron dot formula. a. BF 3 and CH 3 OH b. O 2- and CO 2 Solution: Write the equations using Lewis electron-dot formula. Then identify the e - pair acceptors or Lewis acids and e - pair donor, Lewis base

14 15.4 Relative strengths of Acids and Bases The Brønsted-Lowry concept introduced the idea of conjugate acid-base pairs and proton-transfer reactions. Such acid-base reactions are consider to be a competition between species for hydrogen ions. From this point of view, we can order acids by their relative strength as hydrogen ion donors. 27 Relative strengths of Acids and Bases Stronger acids are those that lose their hydrogen ions more easily than other acids. Stronger bases are those that hold onto hydrogen ions more strongly than other bases. If an acid loses its H, the resulting anion is now in a position to re-accept a proton, making it a Brønsted- Lowry base. if an acid is considered strong, its conjugate base (that is, its anion) would be weak, since it is unlikely to accept a hydrogen ion. Levelling Effect 28 14

15 Relative strengths of Acids and Bases Consider the equilibrium below: HC2 H3O 2(aq) H2O(l) acid base H 3O (aq) C2H 3O 2 acid base (aq) In this example, H 3 O is the stronger of the two acids. Consequently, the equilibrium is skewed toward reactants. Strongest bases have the weakest conjugate acid and the reaction will go towards weak acid 29 Table 15.2 The relative strength of some common acids and their conjugate bases. This concept of conjugate pairs is fundamental to understanding why certain salts can act as acids or base

16 Deciding whether reactants or products are favoured in an Acid-Base reaction You can decide whether a forward reaction (right side) is favoured or a reverse reaction is favoured by using table (Examples will be solved in class) Self-Ionization of Water and ph Self-ionization is a reaction in which two like molecules react to give ions. In the case of water, the following equilibrium is established: H 2O (l) H 2O (l) H 3O (aq ) OH (aq ) The equilibrium-constant expression for this system is: [H 3O ][OH ] K c = 2 [H 2O] 32 16

17 Self-Ionization of Water The concentration of ions is extremely small, so the concentration of H 2 O remains essentially constant. This gives: 2 H O] K = [H O ][OH ] [ 2 c 3 constant The equilibrium value for the ion product [H 3 O ][OH - ] the ion-product constant for water, which is written K w. K w = [H3O ][OH ] At 25 C, the value of Kw is 1.0 x Like any equilibrium constant, K w varies with temperature 33 Ionization constant of Water Using K w you can calculate the concentrations of H and OH - ions in pure water. These ions are produced in equal numbers in pure water, so if we let x = [H ] = [OH - ] x = 14 = (x)(x) 14 at 25 = If you add acid or base to water then [H ]and [OH - ] are no longer equal but the K w expression still holds. o 7 C 34 17

18 15.7 Solutions of Strong Acid or Base In a solution of a strong acid you can normally ignore the self-ionization of water as a source of H (aq). The H (aq) concentration is usually determined by the strong acid concentration. However, the self-ionization still exists and is responsible for a small concentration of OH - ion. You can use K w to calculate this concentration ( concentration of OH - ion) 35 Calculating the concentrations of H 3 O and OH - ions in solutions of a strong Acid or Base. Example :Calculate the concentrations of H 3 O and OHions in a mol L -1 solution of sodium hydroxide(a strong base)

19 Sample Problem Example: Calculate the concentration of OH - ion in 0.10 M HCl. Because you started with 0.10 M HCl (a strong acid) the reaction will produce 0.10 M H (aq). HCl ( aq ) H ( aq ) Cl = ( 0.10 )[ OH ( aq Substituting [H ] = 0.10 into the ion-product expression, we get: [OH ] = = mol L -1 ] ) 37 Solutions of Strong Acid or Base By dissolving substances in water, you can alter the concentrations of H (aq) and OH - (aq). In a neutral solution, the concentrations of H 3 O (aq) and OH - (aq) are equal, as they are in pure water. At 25 C [H 3 O ] = [OH - ] = 1.0 x 10-7 mol L -1 In an acidic solution, the concentration of H 3 O (aq) is greater than that of OH - (aq). [H 3 O ] > 1.0 x 10-7 mol L -1 In a basic solution, the concentration of OH - (aq) is greater than that of H 3 O (aq). [H 3 O ] < 1.0 x 10-7 mol L

20 15.8 The ph of a Solution Although you can quantitatively describe the acidity of a solution by its [H ], it is often more convenient to give acidity in terms of ph. The ph of a solution is defined as the negative logarithm of the molar hydrogen-ion concentration. ph = log[h ] S.P.L.Sorensen devised the ph scale, while working on brewing of beer 39 The ph of a Solution For a solution in which the hydrogen-ion concentration is 1.0 x 10-3, the ph is: ph = log( ) = 3.00 Note that the number of decimal places in the ph equals the number of significant figures in the hydrogen-ion concentration

21 The ph of a Solution In a neutral solution, whose hydrogen-ion concentration is 1.0 x 10-7, the ph = neutral solution ph = 7.00 For acidic solutions, the hydrogen-ion concentration is greater than 1.0 x 10-7 acidic solutions ph < 7.00 (less than) In basic solution, the hydrogen-ion concentration is less than 1.0 x10-7, basic solution ph > Figure 15.5 shows a diagram of the ph scale and the ph values of some common solutions. 41 Figure 15.5: The ph Scale 42 21

22 Calculating the ph from H 3 O ion concentration A sample of orange juice has a hydrogen-ion concentration of 2.9 x 10-4 mol L -1. What is the ph? 43 Calculating the concentration H ion from ph The ph of human arterial blood is What is the hydrogen-ion concentration? [ H ] = anti log( ph ) 44 22

23 poh or Concentration OH - ions A measurement of the hydroxide ion concentration, similar to ph, is the poh. The poh of a solution is defined as the negative logarithm of the molar hydroxide-ion concentration. poh = log[ OH ] Then because K w = [H ][OH - ] = 1.0 x at 25 o C, you can show that ph poh = Sample Problem An ammonia solution has a hydroxide-ion concentration of 1.9 x 10-3 M. What is the ph of the solution? 46 23

24 Measuring ph The ph of a solution can accurately be measured using a ph meter. 1. ph meter translate H ion concentrate of a solution into an electrical signal that is converted into either a digital display or a deflection on a meter that reads ph directly. 2. Acid-base indicators are dyes whose acid form has one coloue and base form has another colour. 3. ph paper strips or litmus paper are coated with a mixture of ph sensitive dyes used to measure ph of food, ground water and biological stuff over a wide range of ph. 47 ph Meter 48 24

25 Measuring ph 2. Acid-base indicators. Although less precise, acid-base indicators are often used to measure ph because they usually change color within a narrow ph range. The color change of an indicator involves an equilibrium between an acid form and base form that have different colours HIn(aq) H 2 O(l) H 3 O(l) In - (aq) Acid form of Phenolphthalien is colourless and basic form is pink. Universal indicator is a mixture of acid base indicators. It is deep red, in acidic solution and it changes to yellow and green at ph Colour changes of various acid-base indicators. Figure

26 51 Operational Skills Identifying acid and base species Identifying Lewis acid and base species Deciding whether reactants or products are favored in an acid-base reaction Calculating the concentration of H 3 O and OH - in solutions of strong acid or base Calculating the ph from the hydronium-ion concentration, and vice versa 52 26