Chapter 10 Acids and Bases

Acids produce H+ ions in water H2O HCl(g) H+(aq) + Cl(aq)

they are electrolyteshave a sour taste turn litmus redneutralize 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 waterare electrolytesfeel soapy and slipperyneutralize acids

NaOH sodium hydroxideKOH potassium hydroxide

sodium and potassium hydroxide release 1 OH- /molecule

other bases such as ammonium hydroxide (NH4OH) release far fewer OH-

fewer

Milk of Magnesia (Mg(OH)2

Tums Ca(OH)2

vinegar C2H4O2

Cola H3PO4

Strong acids completely ionize (100%) in aqueous solutions. HCl(g) + H2O(l) H3O+(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.

H2CO3(aq) + H2O(l) H3O+(aq) + HCO3

−(aq)

NH3(g) + H2O(l) NH4+(aq) + OH−(aq) Windex weak base

H2CO3 + OH- HCO3- + H2O CO3

= + H3O+ Baking Soda weak base

NaOH Na+ + OH- Drano strong base

Water reacts with itself in the following manner:

H+ is transferred from one H2O molecule to another ;one water molecule acts as an acid, while another acts as a base

H2O + H2O H3O+ + OH− .. .. .. .. H:O: + H:O: H:O:H+ + :O:H−

.. .. .. .. H H H water water hydronium hydroxide

ion(+) ion(-)

The concentration of

H3O+ = OH- = 10-7 mols/L

pHThe 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[H3O+ concentration]; pOH = -log [OH- concentration]

when the H3O+ concentration is

expressed in mols/L

pH + pOH = 14

Reactions of acids and bases

Acid + Base = Salt + Water

Mg(OH)2 + HCl (gastric juice) = MgCl2 + H2O

Mg(OH)2 + 2HCl (gastric juice) = MgCl2 + 2 H2O

CaCO3 + HCl = CaCl2 + H2CO3 = CaCl2 + H2O + CO2

CaCO3 + 2HCl = CaCl2 + 2H2CO3 = CaCl2 + 2H2O + 2CO2

+ burp

OH- (equivalents)

0.00 0.02 0.04 0.06 0.08 0.10 0.12

pH

0

1

2

3

4

5

6

7

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 = H2O + NaCl

buffered in this region

0.1

0.01

0.001

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)?

OH- (equivalents)

0.00 0.02 0.04 0.06 0.08 0.10

pH

2

3

4

5

6

7

8

9

HOAc + NaOH = H2O + 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 10-4 moles of HCl or 10-4 moles of NaOH;

What would be the resulting pH assuming no volume change with HCl addition?

H3O+ = 10-4 mol/L; pH= 4

What would be the resulting pH assuming no volume change with NaOH addition?

OH- = 10-4 mol/L; pOH = 4

pH = 14- pOH = 10

alternatively

[H+][OH-] = 1 *10-14; [H+] = 10-14/10-4;

[H+] = 10-10; 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 = MgCl2 + H2O

Mg(OH)2 + 2HCl = MgCl2 + 2 H2O Balanced 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 MgCl2 + H2O

0.01 moles of HCl requires 0.005 moles of Mg(OH)2

What is the pH of a vinegar solution that is 0.1 M?

HOAc + H2O = H3O+ + OAc-

What is the equilibrium expression?

[H3O+][OAc-]/[HOAc] = K

K = 18*10-6

if we let x = [H3O+]; the X also = [OAc-]

x2/[0.1-x] = 18*10-6

lets assume that x is very small in comparison to 0.1

x2 = 1.8 * 10-6; x ≈ 1.3*10-3 pH = 2.74

• H2CO3 is a very weak acid; however both hydrogens can be removed in the presence of strong base; the pH of a solution of NaHCO3 is very close to physiological pH; in the presence of an acid the HCO3- ion tends to pick up the proton, thus buffering the solution and preventing the solution to become too acidic.

• H+ + HCO3- H2CO3 CO2 + H2O

In the presence of a base, the HCO3- ion can

lose its proton as H+ and thus neutralize the strong base; thus the HCO3

- ion can buffer the solution in both directions

HCO3- + OH- CO3

-2 + H2O

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 CO2 is in the form of HCO3-

HCO3- can react with either acid or base

HCO3- + H3O+ H2CO3 CO2 + H2O

HCO3- + OH- H2O + CO3-2

In this manner, HCO3-

stabilizes the pH and does not allow it to become too acidic or to basic; it acts as a buffer

CO2 + H2O = H2CO3 = H+ + HCO3-