Acids, Bases, &

SaltsChapter 20Lesson 1

What is an ACID?• pH less than 7• Neutralizes bases• Forms H

+ ions in solution

• Corrosive-reacts with most metals to form hydrogen gas

• Good conductors of electricity

Acids Generate Ions

HNO3 + H2O H3O+ + NO3

Weak vs. Strong Acids

• Weak Acids do not ionize completely: Acetic, Boric, Nitrous, Phosphoric, Sulfurous

• Strong Acids ionize completely: Hydrochloric, Nitric; Sulfuric, Hydriodic

Common Acids • HCl- hydrochloric- stomach acid

• H2SO4- sulfuric acid - car batteries

• HNO3 – nitric acid - explosives

• HC2H3O2- acetic acid - vinegar

• H2CO3-carbonic acid – sodas

• H3PO4- phosphoric acid -flavorings

What is a BASE?

• pH greater than 7• Feels slippery• Dissolves fats and oils• Usually forms OH- ions in

solution• Neutralizes acids

Weak vs. Strong Bases

• Weak Bases: ammonia; potassium carbonate, sodium carbonate

• Strong Bases: sodium hydroxide; sodium phosphate; barium hydroxide; calcium hydroxide

Common Bases

• NaOH- sodium hydroxide (LYE) soaps, drain cleaner

• Mg (OH)2 - magnesium hydroxide-antacids

• Al(OH)3-aluminum hydroxide-antacids, deodorants

• NH4OH-ammonium hydroxide- “ammonia”

Types of Acids and Bases

• In the 1800’s chemical concepts were based on the reactions of aqueous solutions.

• Svante Arrhenius developed a concept of acids and bases relevant to reactions in H2O.

• Arrhenius acid – produces hydrogen ions in water.• Arrhenius base – produce hydroxide ions in water.

A broader ,more modern concept of acids and bases was developed later.

Bronsted-Lowry acid- donates a hydrogen ion in a reaction.

Bronsted – Lowry base – accepts a hydrogen in a reaction.

• Conjugate acid- compound formed when an base gains a hydrogen ion.

• Conjugate base – compound formed when an acid loses a hydrogen ion.

12

BrØnsted-Lowry Acids and Bases

• acids donate a proton (H+).

• bases accept a proton (H+).

13

Conjugate Acid-Base Pairs

In the reaction of HF and H2O,

• one conjugate acid-base pair is HF/F−.

• the other conjugate acid-base pair is H2O/H3O+.

• each pair is related by a loss and gain of H+.

14

Learning Check

Write the conjugate base of the following.

1. HBr

2. H2S

3. H2CO3

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Solution

Remove H+ to write the conjugate base.

1. HBr Br-

2. H2S HS-

3. H2CO3 HCO3-

pH Scale

pH of Common Substances

Timberlake, Chemistry 7th Edition, page 335

Acid – Base Reactions

• A reaction between an acid and a base is called neutralization. An acid-base mixture is not as acidic or basic as the individual starting solutions.

Reactions with indicators

Indicator Acid color

Neutral color

Base color

Phenolphthalein Colorless Faint pink Dark pink

Bromthymol blue

Yellow Green Blue

Litmus Red ----- Blue

Kw of Water

Kw is the ion-product

• Kw is aka dissociation constant of water

• It has been found experimentally that at 25 °C,

• [H+ ]=[OH-] and they both equal 1X10-7M

• Since Kw= [H+ ][OH-] = [1X10-7M]2

• Then Kw = 1X 10-14.

What does this mean?

• This means that any aqueous solution at 25 ° C, no matter what it (the water) contains, the product of [H+] and [OH-] MUST always equal 1.0 X 10-14.

• There are 3 possibilities:

• A neutral solution where [H+]=[OH-].

• An acid solution where [H+]>[OH-].

• A basic solution where [H+]<[OH-].

In all 3 of the situations:

• Kw = [H+][OH-] = 1.0 X 10-14.

• So in any given aqueous situation, one may calculate the [H+] or [OH-] as required for any solution at 25°C.

• State if Acidic, Basic or Neutral.

• A. 1.0X10-5 M OH- 1.0 X10-9 M H+

• B. 1.0X10-7 M OH- 1.0 X10-7 M H+

• C. 1.0X10-15 M OH- 10.0 M H+

Answers

• A. Basic

• B. Neutral

• C. Acid

How to Solve for the [ions]

• Kw = [H+][OH-]

• 1.0 X 10-14= [H+][OH-]

• [H+] = [OH-]/(1.0 X 10-14)

pH scale is an easy way to represent acidity.

• pH = -log[H+]

• At a neutral solution at 25 °C

• [H+] = [OH-] = (1.0 X 10-14)½ = 1.0 X 10-7

• What is the pH of this?

• pH = -log[H+]

• = -log(1.0 X 10-7)

• Take out you calculator and what do you get?

• 7.00

Sig Figs in Log problems• The number of sig figs in an original number

equals the number of decimal places in the pH.

• Example:

• If sample is Kw= [H+] = 1.0 X 10-7

• How many sig fig?

• 2

• This pH is …

7.00 2 decimal places for the two sig figs.

pH vs. pOH

• If pH is = -log [H+]

• Then pOH = -log [OH-]

• And pK = -log K

• Note that pH changes by 1 for every power of 10 in the change of concentration.

Examples

• Calculate the pH and pOH of each• 1.0 X 10-3 M OH-• pOH = 3.00 pH = 11.00• H+ = Kw/[OH-] = (1.0 X 10-14)/(1.0 X 10-3)• = 1.0 x 10-11

• 1.0 M H+ • pH = 0.00 [OH-]=Kw/[H+]= 1.0x10-14/1• pOH = 14

Remember

• Kw =[H+][OH-]

• And therefore,

• -log K = -log [H+] + -log [OH-]

• log K = log[H+] + log[OH-]

• pKw = pH = pOH therefore

• pH + pOH = 14

• At 25 °C pKw = 14.00 (1.0 X 10-14)

• Thus pH = pOH = 14 at 25 °C

Example Problem

• If the [H+] in a solution is 1.0 x 10-5M, is the solution acidic, basic, or neutral? What is the [OH-] of this solution?