acids and bases. what are acids and bases? lemons, grapefruit, vinegar, etc. taste sour because they...
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Acids and Bases
What are acids and bases?
• Lemons, grapefruit, vinegar, etc. taste sour because they contain acids.
• Acid in our stomach helps food digestion• Acid from bacteria turn milk sour and are used to make
yogurt and cheese. o Bases neutralize acids (pH = 7)o Antacids are taken to offset the effect of too much acid
in the stomach. Other examples are drain cleaners and oven cleaners.
Naming Acids
• Acids are substances that dissolve in water to produce hydrogen ions (H+) and a simple nonmetal anion.
HCl (g) H+ (aq) + Cl- (aq)
• Hydro…… is used before the name of the nonmetal, and its …..ide ending is changed to ………ic acid.
When the acid contains a polyatomic ion the name comes from the polyatomic ion…
If it ends in …..ate, replace with …….ic acid If it ends in …..ite, replace with …..ous acid
Naming Bases
• Bases are ionic compounds that dissociate into a metal ion and hydroxide ions (OH-) when they dissolve in water.
NaOH (s) Na+ (aq) + OH- (aq)
• Bases are named as hydroxides.
Bronsted-Lowry acids and bases
• Bronsted-Lowry Acid- donates a proton (hydrogen ion, H+) to another substance
• Bronsted-Lowry Base• - donates accepts a proton. • The proton does NOT actually exist in water………a
hydronium ion (H3O+) is formed when the proton is attracted
to the polar water molecule.• • H - O + H+ H - O - H +
l l H HWater proton hydronium ion
Bronsted-Lowry acids and bases
• Examples:
HCl + H2O H30+ + Cl-
• When ammonia (NH3) reacts with water, the nitrogen has a stronger attraction for the proton than water.
NH3 + H2O NH4
+ + OH-
Water can be an acid or a base!
• Identify the reactant that is an acid ( H+ donor) and the reactant that is a base (H+ acceptor).
HBr + H20 H30
+ + Br-
H2O + CN- HCN + OH-
Thus, H2O can be an acid or a base!
Conjugate acid-base pairs
• When molecules are related by the loss or gain of one H+ (proton) they make a conjugate acid-base pair. The protons are transferred both forward and reverse!
Conjugate acid-base pair
HA + B A- + BH+
Acid 1 Base 2 Base 1 Acid 2
H+ donor H+ acceptor H+ acceptor H+ donor
Problem Time!
• Identify the conjugate acid-base pairs in the following reactions;
HF + H2O F- + H30+
NH3 + H20 NH4+ + OH-
Bronsted-Lowry acids and bases
Bronsted-Lowry acids and bases
Classifying acids and bases
• Acids- classified according to their ability to donate protons
• Strong acids………give up protons easily• Weak acids………..give up only a few protons and most
molecules keep their protons. • Bases
- classified in terms of their ability to accept protons.
• Strong bases……..have a strong attraction for protons• Weak bases……….have little attraction for protons
Strong and Weak Acids
• Strong acids dissociate almost completely to give H30+
ions and anions.
HCl + H20 H30+ + Cl-
• Weak acids dissociate only slightly and produce small
concentrations of H30+ ions.
CH3COOH + H2O H30+ + CH3COO-
Strong and Weak Bases
• Strong bases are ionic compounds that dissociate in water to give an aqueous solution of a metal ion and hydroxide ion.
KOH (s) K+(aq) + OH-
(aq)
• Weak bases are poor acceptors of protons.
NH3 (g) + H2O (l) NH4+
(aq) + OH-(aq)
Ionization of Water
• Remember water can act as BOTH an acid AND a base…………….how?
• One water molecule can donate a proton to another to
produce H3O+ and OH-
Conjugate acid-base pair
H2O(l) + H20(l) H3O
+(aq) + OH-
(aq)
Conjugate acid-base pair
Ion-product constant
• In pure water the transfer of proton between two water molecules produces EQUAL numbers of hydronium and hydroxide ions.
[H30+] = [OH-] = 1.0 x 10-7M
Ion-product constant for water (Kw) = [H3O+] x [OH-] = 1.0 x 10-14
(no units) • The Kw values applies to any aqueous solution. ALL aqueous
solutions have have H3O+ and OH- ions.
• [H3O+] = [OH-] ……solution is neutral
• [H3O+] > [OH-]……………………. ……..acid
• [H3O+] < [OH-] ……………………..…..basic
• Remember the sum of both ions is always = 1.0 x 10-14 at 25 oC.
pH Scale
• On the pH scale a number between 0 and 14 represents the H3O
+ concentration.
• Acidic solution pH < 7 [H3O+] > 1.0 x 10-7M
• Neutral solution pH = 7 [H3O+] = 1.0 x 10-7M
• Basic solution pH > 7 [H3O+] < 1.0 x 10-7M
Calculating the pH of solutions
• pH scale is a log scale. ………………… pH = -log[H3O+]
Example…………. lemon juice solution has [H3O+] = 1.0 x 10-2M
pH = - log[ 1.0 x 10-2]= - (- 2.00)= 2.00
• Remember the number of significant figures……….answer should
have the same number of sig. figs. as the [H3O+]
[H3O+] = 1.0 x 10-3 pH = 3.00
2 significant figs. 2 decimal places
Measuring pH
Reactions of Acids and Bases
• Acids will react with • - metals• - carbonates and bicarbonates• - hydroxides
Acids + Metals
• Acid + “active metals” “salt” + hydrogen gas
These are single replacement reactions. • Active metals are ; potassium, sodium, calcium,
magnesium, aluminium, zinc, iron, tin. Example:
Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g)
Metal acid salt hydrogen
Acids + carbonates & bicarbonates
Strong acid + carbonate/bicarbonate CO2 + H2O + salt
Example:
• HCl(aq) + NaHCO3(aq) CO2(g) + H2O + NaCl(aq)
Acids + Hydroxides
• Neutralization occurs between an acid and a base to form a salt and water.
HCl + NaOH NaCl + H2O
( H+ Cl- Na+ OH- Na+ Cl- H2O )
• One can write the net ionic equation ( omitting the metal ions and chloride ions that are not reacting) as follows:
H+ + OH- H2O
Buffers
• BUFFER SOLUTIONS resist changes in pH when small amounts of acids or base are added.
• Buffers contain an acid to react with any OH- that is
added and a base to react with any H3O+ added………
BUT that acid and base must not neutralize each other. • So, a combination of an acid-base pair conjugate is used
in buffers ie. a weak acid and its salt and a weak base and its salt.
CH3COOH (aq) + H2O (l) H3O+ (aq) + CH3COO- (aq)
Large amount Large amount
Buffers
1. If a small amount of acid is added then:
CH3COOH (aq) + H2O (l) H3O+ (aq) + CH3COO- (aq)
• Acid will combine with the acetate ion and shift the equilibrium…
[CH3COO-] [CH3COOH] no change in H3O+
2. If a small amount of base is added then:
CH3COOH (aq) + H2O (l) H3O+ (aq) + CH3COO- (aq)
• Base is neutralized by the acetic acid to produce water and shift the equilibrium.
[CH3COO-] [CH3COOH] no change in H3O+
Buffers in the Blood
• Cells can only function properly when the pH is between 6.8 – 8.0.
• The normal pH of arterial blood is 7.35 – 7.45.
• The bicarbonate/carbonic acid system is an important buffer system in the blood.
H2O
CO2 + H2O H2CO3 H3O+ + HCO3
-
• Excess H3O+ entering the body fluids reacts with the
HCO3- and excess OH- reacts with carbonic acid.
Buffers in the Blood
H2O
CO2 + H2O H2CO3 H3O+ + HCO3
-
CO2 H3O+ pH……………………. Acidosis
…….emphysema, difficulty with breathing, medulla affected by depressive drugs or accident trauma.
CO2 H3O+ pH……………………. Alkalosis
…….hyperventilation causes expiration of a lot of carbon dioxide eg. excitement, trauma, high temperatures