Chapter 10 Acids and Bases. Acids produce H + ions in water H 2 O HCl(g) H+(aq) + Cl (aq) they are...
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Transcript of Chapter 10 Acids and Bases. Acids produce H + ions in water H 2 O HCl(g) H+(aq) + Cl (aq) they are...
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
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