1. TOPIC 12-ACIDS AND BASES

2. THE ARRHENIUS THEORY In an aqueous solution a strong electrolyte exists only in the form of ions, whereas a weak electrolyte exists partly as ions and partly as molecules. When the acid HCl dissolves in water, the HCl molecules ionize completeley, yielding H + ions. When the base NaOH dissolves in water, the Na + and OH - ions become dissociated

3. BRONSTED-LOWRY THEORY OF ACIDS AND BASES In their theory an acid is a proton donor and a base is a proton acceptor

4. BRONSTED-LOWRY THEORY OF ACIDS AND BASES

5. We identify acids and bases in some typical acid-base reactions. In working through this example, notice the following additional features: 1.Any species that is an acid by the Arrhenius theory remains an acid in the Bronsted-Lowry theory, the same is true of bases. 2.Certain species, even though they do not contain the OH - group, produce OH - in aqueous solution, e.g OCl -. As such they are Bronsted-Lowry bases. amphiprotic. The Arrhenius theory does not account for amphiprotic behaviour. The Bronsted-Lowry theory accounts for substances that can either react as an acid or a base; they are said to be amphiprotic. The Arrhenius theory does not account for amphiprotic behaviour. Bronsted-Lowry acid-base reaction is favored in the direction from the stronger to the weaker acid-base combination.

6. i ncreasing acid strength increasing base strength RELATIVE STRENGTHS OF SOME COMMON ACIDS AND BASES

7. THE COMPARISON OF THE STRENGTH OF ACIDS Both HCl and HClO 4 are strong acids. To determine whether HCl or HClO 4 is stronger acid, we need to use a solvent that is a weaker base than water, a solvent that will accept protons from the stronger of the two acids more readily than water from the weaker one. In the solvent diethyl ether, (C 2 H 5 ) 2 O, HClO 4 is completely ionized but HCl is only partially ionized. HClO 4 is stronger acid than HCl.

8. THE SELF-IONIZATION OF WATER AND THE PH SCALE For each H 2 O molecule that acts as an acid another acts as a base, and hydronium, H 3 O +, and hydroxide, OH - ions are formed. In the reverse reaction H 3 O + donates a proton to OH -. Equilibrium is displaced far to the left since H 3 O + and OH - are much stronger than the other acid-base conjugate The equilibrium constant expression for pure water The significance of the equation is that it applies to all aqueous solutions

9. pH and pOH In 1909 the Danish biochemist Soren Sorensen proposed the term pH to refer to the « potential of hydrogen ion». He defined pH as the negative logarithm of [H + ]. Restated in termS of [H 3 O + ]. pH = - log [H 3 O + ] Thus, in a solution that is 0,0025 M HCl [H 3 O + ] = 2,5 x 10 -3 M and pH= -(log 2,5x 10 -3 )= 2,60 To determine [H 3 O + ] corresponding to pH value, we do an inverse calculation Log [H 3 O + ] = - 4,50 [H 3 O + ]= 10 -4,50 = 3,2 x 10 -5 pOH = -log [OH - ] pK w = pH + pOH =14,00 We say that pure water and all aqueous solutions with pH=7 are pH neutral. If the pH is less than 7,00, the solution is acidic, if the pH is greater than 7, the solution is basic or alkaline

10. STRONG ACIDS AND STRONG BASES

12. WEAK ACIDS AND WEAK BASES

13. SOME COMMON WEAK ACIDS

14. SOME COMMON WEAK BASES

15. DETERMINATION OF K a Example: Butyric acid, HC 4 H 7 O 2 is used to make compounds employed in artificial flavorings and syrups. A 0,25 M aqueous solution of HC 4 H 7 O 2 is found to have a pH of 2,72. Determine K a for butyric acid initial conc changes equil. conc.

16. Calculation of pH of a weak base without neglecting the value of X What is the pH of a methylamine(CH 3 NH 2 ) solution that is 0,0025 M? K b = 4,2 x 10 -4

17. Rate of Ionization Acid Concentration, M Percent ionization gives the proportion of ionized molecules.The percent ionization of a weak acid or a weak base increases as the solution becomes more dilute. PERCENT IONIZATION Percent ioniz.= [H 3 O] + [HA] X 100 pH of the solution

18. PERCENT IONIZATION Example: What is the percent ionization of acetic acid in 1,0 M, 0,1 M, 0,010 M HC 2 H 3 O 2 ? K a =1,8x10 -5

19. POLYPROTIC ACIDS Some acids have more than one ionizable H atom per molecule. These are called polyprotic acids. For example : Phosphoric acid, H 3 PO 4, it has three ionizable H atoms; it ionizes in three steps and for each step we can write an ionization equation and an ionization constant expression with a value of K a : 1.K a1 is so much larger than K a2 and K a3 that essentially all the H 3 O + is produced in the first ionization step alone. 2.So little of the H 2 PO 4 - formed in the first ionization step ionizes any further that we can assume [H 2 PO 4 - ]=[H 3 O + ] 3.[HPO 4 2- ] ̰͂ K a2 regardless of the molarity of the acid

20. POLYPROTIC ACIDS Calculating Ion Concentrations in a Polyprotic Acid Solution: For a 3,0 M H 3 PO 4 solution, calculate a) [H 3 O + ] b) [H 2 PO 4 - ] c) [HPO 4 2- ] and d) [PO 4 3- ] Solution:

21. IONIZATION CONSTANTS OF SOME COMMON POLYPROTIC ACIDS

22. BUFFER SOLUTIONS There are some water solutions, called buffer solutions, whose pH values change only very slightly upon the addition of small amounts of either an acid or a base. What buffer solutions require are two components, one of which is able to neutralize acids, and the other bases. But, of course, the two components must not neutralize each other. This rules out mixtures of a strong acid and a strong base. Instead, common buffer solutions are described either as a mixture of 1.A weak acid and its conjugate base 2.A weak base and its conjugate acid

23. An important example of a buffered system is that found in blood, which is maintained at a pH=7,4 and called HCO 3 /H 2 CO 3 buffer solution. It helps that the pH of the blood remains constant at 7,4 ( A change of the value 0,2 in pH in blood may cause lethal disinfections or diseases, therefore it has a very important role in human health.) Other applications with buffer solutions are protein studies which aim to measure enzymatic activities. These studies are performed under buffer media since the magnitude and kind of electric charges of the protein molecules depend on the pH. BUFFER SOLUTIONS

24. ACID-BASE INDICATORS An acid-base indicator is a substance whose color depends on the pH of the solution to which it is added. When a small amount of indicator is added to the solution, the indicator does not affect the pH of the solution because so little of it is present. Acid-base indicators find their greatest use where only an approximate pH determination is needed. E.g in soil testing kits to establish the pH of soil. In swimming pools, the optimal pH=7,4 and phenol red is common indicator used in testing swimming pool water.

25. IndicatorLow pH colorTransition pH rangeHigh pH color Gentian violetGentian violet (Methyl violet 10B)Methyl violet 10B yellow0.0–2.0blue-violet Leucomalachite greenLeucomalachite green (first transition) yellow0.0–2.0green Leucomalachite greenLeucomalachite green (second transition) green11.6–14colorless Thymol blueThymol blue (first transition) red1.2–2.8yellow Thymol blue (second transition) yellow8.0–9.6blue Methyl yellowred2.9–4.0yellow Bromophenol blueyellow3.0–4.6purple Congo redblue-violet3.0–5.0red Methyl orangered3.1–4.4orange Bromocresol greenyellow3.8–5.4blue Methyl redred4.4–6.2yellow Methyl redred4.5–5.2green Azolitminred4.5–8.3blue Bromocresol purpleyellow5.2–6.8purple Bromothymol blueyellow6.0–7.6blue Phenol redyellow6.4–8.0red Neutral redred6.8–8.0yellow Naphtholphthaleincolorless to reddish7.3–8.7greenish to blue Cresol Redyellow7.2–8.8reddish-purple Phenolphthaleincolorless8.3–10.0fuchsia Thymolphthaleincolorless9.3–10.5blue Alizarine Yellow Ryellow10.2–12.0red