What are Acids and Bases?

What are some common acids you

know?

What are some common bases you

know?

Where is it common to hear about pH

balanced materials?

Historically, classified by their observable

properties

› Acids:

Have a sour taste – like lemons or sour candy

Corrode metals – learned not to store vinegar

or fruit juices in metal containers

Changed blue litmus dye to red

› Bases:

Bitter in taste

Slippery in feel

Changed red litmus dye to blue

1. Arrhenius Definition:

› Experimented with electrolytes

› Aqueous solutions of acids and bases

conduct electricity

› Therefore, the compounds were forming

positive and negative ions in solution

Arrhenius Model of Acids

An aqueous solution that produced hydrogen

ions, H+

Example: HCl (g) H+ (aq) + Cl- (aq)

Arrhenius Model of Bases

An aqueous solution that produced hydroxide

ions, OH-

Example: NaOH (s) Na+ (aq) + OH- (aq)

The Arrhenius model explains how acids

and bases neutralize each other

› H+ (aq) + OH- (aq) H2O (l)

He did earn the 1903 Nobel Prize in

Chemistry

› Insisting that the H+ (aq) and OH- (aq) were

important in acid and base behavior

Fundamental Problems:

› H+ ion: essentially a proton with a small radius & positive charge

› Therefore, H+ are unlikely to exist as free ions

in aqueous solutions

› Instead they exist with surrounding water

molecules resulting in: Hydronium ion, H3O+

(aq) as we know them today

Fundamental Problems:

› Assumes that all bases contain OH- ions

› Many ionic compounds (salts) have basic

properties such as the ability to neutralize acids

› Examples: metal oxides, carbonates,

fluorides, ammonia (NH3)

Binary Acids – contain hydrogen and

one other element

Use: hydro_______ic acid

› Ex: HCl = hydrochloric acid

› Ex: H2S = hydrosulfuric acid

Oxyacids – contain hydrogen and a poly atomic ion › Use ending of ion for naming:

-ite -ous acid “I bite a delicous apple”

-ate -ic acid “I ate something icky”

› Ex: HNO3 = nitric acid

› Ex: HNO2 = nitrous acid

› Ex: HC2H3O2 = acetic acid

› Ex: H3PO3 = phosphorous acid

Strong Acid or Base:

› A strong electrolyte and completely ionizes

or dissociates in water

Weak Acid or Base:

› A weak electrolyte and only partially ionizes

in water

Examples:

›HCl – hydrochloric acid: stomach

acid, pools

›HBr – hydrobromic acid

›H2SO4 – sulfuric acid: car battery

acid, acid rain

Strong Bases: completely ionize in water

Most of the common strong bases are

the ionic hydroxides from group 1 and 2

metals.

Dissociate completely win water to form

OH- and the cation it was bonded to

Example: H2O

NaOH (s) Na+ (aq) + OH- (aq)

Examples:

›NaOH – sodium hydroxide: drain

cleaners

› KOH – potassium hydroxide

›Mg(OH)2 – magnesium hydroxide:

used in antacids

Examples:

›Acetic Acid (CH3CO2H) –

vinegar, sour wine

›Carbonic acid (H2CO3) – soda,

blood

›Citric acid (H3C6H5O7) – fruit,

soda

Examples:

›Ammonia (NH3) – glass cleaners

›Calcium carbonate (CaCO3) – antacids, minerals

›Calcium hypochlorite (Ca(OCl)2) – chlorine source for swimming pools

Brǿnsted-Lowry Acid:

› Any substance that can donate an H+ ion to

another substance

Brǿnsted-Lowry Base:

› Any substance that can accept an H+ ion

from another substance

Polyprotic Acid:

› An acid containing more then one acidic hydrogen

› Examples: Phosphoric acid: H3PO4 – 3 acidic

hydrogens

Carbonic acid: H2CO3 – 2 acidic hydrogens

Sulfuric acid: H2SO4 – 2 acidic hydrogens

Polyprotic acids do not lose all their

acidic hydrogen atoms in water to

the same extent

Example: Sulfuric Acid complete ionization….

H2SO4 (aq) + H2O (l) H3O+ (aq) + HSO4

- (aq)

Once HSO4- (aq) forms, it also acts as an acid, but as a weak

acid:

HSO4- (aq) + H2O (l) H3O

+ (aq) + SO42- (aq)

Lewis Acid: substance that is an

electron–pair acceptor

› To avoid bonding based on hydrogen

› BF3 (aq) + F- (aq) BF4- (aq)

Lewis Base: substance that is an

electron-pair donor

Acid Base

Arrhenius H+ donor OH- donor

Bronsted-Lowry

p+ donor p+ acceptor

Lewis e- pair acceptor e- pair donor

Conjugate Acid:

› The product that forms as a result of gaining

an p+

Conjugate Base:

› The product that forms as a result of losing

an p+

Example:

HCl (g) + H2O (l) H3O+ (aq) + Cl- (aq)

acid base conjugate acid conjugate

base

Conjugate acid-base pairs always differ

by one H+ ion

› Conjugate acids has one more H+

Has one more H atom in its formula

Increase in charge by 1

› Conjugate base has one less H+

Has one less H atom in its formula

Decrease in charge of 1

Amphoteric Substances:

› A substance that can act as either an acid

or a base

› Examples:

Water (most common)

Acid: donates H+ forming OH-

Base: accepts H+ forming H3O+

Examples › Bicarbonate ion, HCO3

-

Found in sodium bicarbonate, used to neutralize both acids and bases

› When mixed with a basic solution, it acts as an acid

HCO3- (aq) + OH- (aq) CO3

2- (aq) + H2O (l) acid conjugate base

› When mixed with an acidic solution, it acts as a base

HCO3- (aq) + H3O

+ (aq) H2CO3 (aq) + H2O (l) base conjugate acid

Strong acids completely dissociate to form H3O

+ and strong bases completely dissociate to form OH-

HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)

Neutralization reactions: a reaction where hyronium ions and hydroxide ions form water molecules

Salt: ionic compound composed of a cation from a base and an anion from an acid

Self-Ionization of Water

› Two water molecules interact to produce a

hydronium ion and a hydroxide ion by

proton transfer

2 H2O H3O+ (aq) + OH- (aq)

› At 25oC, 1 mole of hydronium and hydroxide

ions exist in 107 liters of water

› Therefore: 1 mole ions = 1 x 10-7 M

107 L water

Water is neutral when the [H3O+] = [OH-]

Water dissociation constant (Kw) – constant rate at which water dissociates › Different at each temperature

Kw = [H3O+][OH-]

= [1 x 10-7 M][1 x 10-7 M]

= 1 x 10-14 M2

pH: The negative logarithm (base 10) of

the [H3O+]

› Equation: pH = - log [H3O+]

› Example: pure water

pH = - log [H3O+]

pH = - log (1.0 x 10-7)

pH = 7

Pure water has [H3O+] = [OH-] => pH = 7

Acidic solutions have pH < 7

Basic solutions have pH > 7

The pH and pOH total is equal to 14.00

pH + pOH = 14.00

This relationship allows us to determine the pH if the pOH is known

pH = 14.00 – pOH = 14.00 – 2.00

= 12.00

Some common

substances, their

pH and their

[H3O+]

French: “pouvior

hyrogene,”

meaning

“hydrogen

power”

Consider the pH values of

solutions that range in [H3O+] from

1.0x10-1 M to 1.0x10-14

M

Notice that the pH value = the

exponent in the [H3O+] but with a

positive value

› Only allows for calculation if the

[H3O+] is a power of ten

Example: What is the pH of each of the

following solutions? Once calculated,

check to make sure you answer makes

sense.

› A) 0.0010 M HBr

› B) 0.035 M HNO3

› C) 0.035 M KOH

Answer: A) 0.0010 M HBr

› HBr is a strong acid, which ionizes

completely, so [H3O+] = [HBr]

› pH = - log [H3O+]

› pH = - log (0.0010)

› pH = 3.00

› Acid!

Answer: B) 0.035 M HNO3

› Nitric acid is also a strong acid, so the [H3O+]

= [HNO3]

› pH = - log [H3O+]

› pH = - log (0.035)

› pH = 1.46

› Acid!

Answer: C) 0.035 M KOH

› Potassium hydroxide is a strong base, which

dissociates completely to form K+ (aq) and

OH- (aq)

› The [KOH] = [OH-] = 0.035 M

› In order to calculate pH, we need [H3O+]

› [H3O+] = Kw_ = 1.0 x 10-14 = 2.9 x 10-14 M

[OH-] 0.0035 M

Answer: C) 0.035 M KOH

› [H3O+] = 2.9 x 10-14 M

› pH = - log [H3O+]

› pH = - log (2.9 x 10-14 )

› pH = 12.54

› Base!

If you can calculate the pH from your

[H3O+], can you do the reverse? Yes!

How? Rearrange your parent equation!

pH = - log [H3O+]

-pH = log [H3O+]

Inverse log (-pH) = [H3O+]

10-pH = [H3O+]

Example: Olivia measures the pH of a

soil sample solution to have a pH = 6.20,

what is the [H3O+] ?

Answer:

› [H3O+] = 10-pH

› [H3O+] = 10-6.20

› [H3O+] = 6.3 x 10-7 M

If you can calculate the pOH from your

[OH-], can you do the reverse? Yes!

How? Rearrange your parent equation!

pOH = - log [OH-]

-pOH = log [OH-]

Inverse log (-pOH) = [OH-]

10-pOH = [OH-]

Example: Jake measured the pH of water in a swimming pool as 8.10. What is the OH-

concentration in the pool water?

Answer:

Determine the pOH from the pH

pH + pOH = 14.00

8.10 + pOH = 14.00

pOH = 5.90

Answer:

pOH = 5.90

Then calculate the [OH-]

› [OH-] = 10-pOH

› [OH-] = 10-5.90

› [OH-] = 1.3 x 10-6 M

Several Methods:

1. pH Meter:

Very accurate to within hundredths of a pH unit

Measures the voltage that develops when electrodes are dipped into the solution

2. pH Indicators or Litmus Strips:

Less accurate but more convenient and cost friendly

Brightly colored organic dyes that are weak acids or bases › In solution they form an equilibrium with their

conjugate bases

› Color of the indicator depends on whether the dye is in its acidic or basic form

2. pH Indicators or Litmus Strips:

Ex) Phenolphthalein

HIn (aq) + H2O (l) H3O+ (aq) + In- (aq)

colorless pink

In the acidic form: HIn (aq) = colorless

In the basic form: In- (aq) = pink

Changes from colorless to pink between

pH 8.2 and 10

An indicator reveals if the pH of a

solution is above or below a certain

value

Also disclose a specific pH within the

indicators color-change range

Subtle differences in hues are discernible

at slightly different pH values

A mixture of indicators having a variety of colors and color-change ranges can be used to measure the pH of any solution

Broad-range pH paper is treated with several indicators › The user reads the pH by comparing the

color the paper turns to a chart of reference colors and pH values

The formation of water is quite common

from the addition of hydrogen ion (H+)

and the hydroxide ion (OH-)

Titration: › The process of determining the

concentration of one substance in a solution by reacting it with a solution of another substance that has a known concentration.

› Add the known substance until the reaction between the two substances is complete: equivalence point

› Shown by an indicator: changes color due to sensitivities of acids and bases

› End point: the point at which the indicator changes color

Phenolphthalein

Molarity = moles / Liter

Macid x Vacid = Mbase x Vbase

**Only true for a 1 to 1 mole

ratio between the acid and

base

Normality (N) – number of equivalents of

solute per liter of solution

Equation: N = n * M

normality = number of equiv * Molarity

What is the molarity of a 0.090 N Ca(OH)2

soln?

› N = n * M

› 0.090 N = 2 equiv * M

› M = 0.045 M Ca(OH)2

Chapter 14 p. 491-493

› #15, 23, 24

Chapter 15 p. 523-525

› #6, 9, 12, 15, 26, 31, 36