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