Chapter 10 Acids and Bases

Acids produce H + ions in water H 2 O HCl(g) H+(aq) + Cl  (aq) they are electrolytes have a sour taste turn litmus red neutralize bases

Some acids like sulfuric and phosphoric release more than 1 H + in water; other like acetic acid (vinegar) release far less than 1 H + per molecule

Bases produce OH − ions in water are electrolytes feel soapy and slippery neutralize acids NaOHsodium hydroxide KOHpotassium hydroxide sodium and potassium hydroxide release 1 OH - /molecule other bases such as ammonium hydroxide (NH 4 OH) release far fewer OH - fewer

Milk of Magnesia (Mg(OH) 2 Tums Ca(OH) 2 vinegar C 2 H 4 O 2 Cola H 3 PO 4

Strong acids completely ionize (100%) in aqueous solutions. HCl(g) + H 2 O(l) H 3 O + (aq) + Cl−(aq) Amount of acid added Weak acids dissociate only slightly in water to form a solution of mostly molecules and a few ions. H 2 CO 3 (aq) + H 2 O(l) H 3 O + (aq) + HCO 3 − (aq)

NH 3 (g) + H 2 O(l) NH 4 + (aq) + OH − (aq) Windex weak base H 2 CO 3 + OH - HCO H 2 O CO 3 = + H 3 O + Baking Soda weak base NaOH Na + + OH - Drano strong base

Water reacts with itself in the following manner: H + is transferred from one H 2 O molecule to another ; one water molecule acts as an acid, while another acts as a base H 2 O + H 2 O H 3 O + + OH − H:O: + H:O: H:O:H + + :O:H − H H H water water hydronium hydroxide ion(+) ion(-) The concentration of H 3 O + = OH - = mols/L

pH The pH of a solution is used to indicate the acidity of a solution; it has values that usually range from 0 to 14; the solution is acidic when the values are less than 7; the solution is neutral with a pH of 7; the solution is basic when the values are greater than 7

How is the numerical value of pH determined? pH = - log[H 3 O + concentration]; pOH = -log [OH- concentration] when the H 3 O + concentration is expressed in mols/L pH + pOH = 14

Reactions of acids and bases Acid + Base = Salt + Water Mg(OH) 2 + HCl (gastric juice) = MgCl 2 + H 2 O Mg(OH) 2 + 2HCl (gastric juice) = MgCl H 2 O CaCO 3 + HCl = CaCl 2 + H 2 CO 3 = CaCl 2 + H 2 O + CO 2 CaCO 3 + 2HCl = CaCl 2 + 2H 2 CO 3 = CaCl2 + 2H 2 O + 2CO 2 + burp

How does the pH vary if we add NaOH (0.1 mol/L) dropwise to a solution of HCl (0.1 mol/L)? pH of resulting solution HCl + NaOH = H 2 O + NaCl buffered in this region

How does the pH vary is we add NaOH (0.1 mol/L) dropwise to a solution of the weak acid acetic acid (0.1 mol/L)? HOAc + NaOH = H 2 O + NaOAC pH of resulting solution buffered in this entire region

Suppose we have a liter of water and we either add a drop of water containing moles of HCl or moles of NaOH; What would be the resulting pH assuming no volume change with HCl addition? H 3 O + = mol/L; pH= 4 What would be the resulting pH assuming no volume change with NaOH addition? OH - = mol/L; pOH = 4 pH = 14- pOH = 10 alternatively [H + ][OH - ] = 1 * ;[H + ] = /10 -4 ; [H + ] = ; pH = 10

How much Mg(OH) 2 would be required to neutralize 100 mL of HCl that is 0.1 M? Mg(OH) 2 + HCl = MgCl 2 + H 2 O Mg(OH) 2 + 2HCl = MgCl H 2 OBalanced equation How many moles of HCl are their in 100 mL of 0.1 M HCl ? 0.1 M HCl = 0.1 mol/L; 100 mL = 0.1 L 0.1 mol/L *0.1 L = 0.01 moles of HCl 0.5 Mg(OH) 2 + HCl = 0.5 MgCl 2 + H 2 O 0.01 moles of HCl requires moles of Mg(OH) 2

What is the pH of a vinegar solution that is 0.1 M? HOAc + H 2 O = H 3 O+ + OAc - What is the equilibrium expression? [H 3 O + ][OAc - ]/[HOAc] = K K = 18*10 -6 if we let x = [H 3 O + ]; the X also = [OAc - ] x 2 /[0.1-x] = 18*10 -6 lets assume that x is very small in comparison to 0.1 x 2 = 1.8 * 10 -6; x ≈ 1.3*10 -3 pH = 2.74

H 2 CO 3 is a very weak acid; however both hydrogens can be removed in the presence of strong base; the pH of a solution of NaHCO 3 is very close to physiological pH; in the presence of an acid the HCO 3 - ion tends to pick up the proton, thus buffering the solution and preventing the solution to become too acidic. H + + HCO 3 - H 2 CO 3 CO 2 + H 2 O In the presence of a base, the HCO 3 - ion can lose its proton as H+ and thus neutralize the strong base; thus the HCO 3 - ion can buffer the solution in both directions HCO OH - CO H 2 O

The pH in living systems is very important. For example the pH of blood is kept at 7.4 and must be maintained within ±0.5 pH units. How is this done? At a pH of 7.4, most CO 2 is in the form of HCO 3 - HCO 3 - can react with either acid or base HCO H 3 O + H 2 CO 3 CO 2 + H 2 O HCO OH - H 2 O + CO 3 -2 In this manner, HCO 3 - stabilizes the pH and does not allow it to become too acidic or to basic; it acts as a buffer

CO 2 + H 2 O = H 2 CO 3 = H + + HCO 3 -