chemistry ms. agostine

CHAPTER 20

Chemistry

Ms. Agostine

ACIDS AND BASES AND THE CONCEPT OF pH

Properties of Acids & Bases

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?

What are Acids and Bases?

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

Some Acidic Foods

Review: Naming Acids

Binary Acids – contain hydrogen and one

other element

Use: hydro_______ic acid

HCl = hydrochloric acid

H2S = hydrosulfuric acid

Anion Ending Acid Name

-ide hydro-(stem)-ic acid

-ate (stem)-ic acid

-ite (stem)-ous acid

Acid Nomenclature Review

Binary

Ternary

An easy way to remember which goes with which…

“In the cafeteria, you ATE something ICky”

“You bITE a deliciOUS apple”

Acid Nomenclature Flowchart

Acids

Start with “H”

2 elements

Hydro- prefix

-ic ending

-ate ending Becomes

-ic ending

-ite ending Becomes

-ous ending

3 elements

No hydro- prefix

Acid Nomenclature Review

HBr → hydrobromic acid

H2CO3 → carbonic acid

H2SO3 → sulfurous acid

Part 1

Acid and Base Definitions

Acid and Base Definitions

1. Arrhenius Definition:

He 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 & Bases

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)

1. Arrhenius Model of Acids & Bases

The Arrhenius model explains how acids and

bases neutralize each other

H+ (aq) + OH- (aq) H2O (l)

He earned the 1903 Nobel Prize in Chemistry

Insisted 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+ ions 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

Hydronium Ion


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)

Strong and Weak Acids and Bases

Strong Acid or Base:

Is a strong electrolyte and completely ionizes or

dissociates in water

Weak Acid or Base:

Is a weak electrolyte and only partially ionizes in

water

Strong Acids

Strong Acids: 100% Dissociated

Examples:

HCl – hydrochloric acid: stomach acid,

pools

HBr – hydrobromic acid

H2SO4 – sulfuric acid: car battery acid,

acid rain

Strong Bases: 100% Dissociated

Strong Bases: completely ionize in water

Most of the common strong bases are the

ionic hydroxides from group 1 and 2

metals.

Dissociate completely in water to form OH-

and the cation it was bonded to

Strong Base: 100% Dissociated

Example:

H2O

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

Strong Bases: 100% Dissociated

Examples:

NaOH – sodium hydroxide: drain

cleaners

KOH – potassium hydroxide

Mg(OH)2 – magnesium hydroxide: used in

antacids

Weak Acids: < 100% Dissociated

Examples:

Acetic Acid (CH3CO2H) – vinegar, sour

wine

Carbonic acid (H2CO3) – soda, blood

Citric acid (H3C6H5O7) – fruit, soda

Weak Bases: <100% Dissociated

Examples:

Ammonia (NH3) – glass cleaners

Calcium carbonate (CaCO3) – antacids, minerals

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

Part 2

Acids and Base Definitions

2. Brǿnsted-Lowry Acids & Bases

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 Acids

Polyprotic Acid:

An acid containing more than one acidic

hydrogen

Examples:

Phosphoric acid: H3PO4 – 3 acidic hydrogens

Carbonic acid: H2CO3 – 2 acidic hydrogens

Sulfuric acid: H2SO4 – 2 acidic hydrogens

Polyprotic Acids

Polyprotic acids do not lose all their acidic

hydrogen atoms in water to the same extent

Example: Sulfuric Acid (strong acid)

Has 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)

Conjugate Acid-Base Pairs

Acid-Base Reactions

Conjugate Acid:

The product that forms as a result of gaining a p+

Conjugate Base:

The product that forms as a result of losing a p+

Example:

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

acid base conjugate acid conjugate base

Conjugate Acid-Base Pairs

Conjugate acid-base pairs always differ by one H+ ion

Conjugate acid 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

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+

Amphoteric Water (Acts as Acid)

Amphoteric Water (Acts as Base)

Amphoteric Substances

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 Acid-Strong Base Neutralization

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 hydronium ions and hydroxide ions form water molecules

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

Acid + Base Salt + Water

3. Lewis Acid & Base Definitions

Lewis Acid: is an atom, ion, or molecule that accepts an electron pair to form a covalent bond (electron pair acceptor)

Includes as acids substances that do not contain hydrogen at all.

BF3 (aq) + F- (aq) BF4- (aq)

Lewis Base: is an atom, ion, or molecule that is an electron-pair donor

Definitions of Acids & Bases

Acid Base

Arrhenius Releases H+ Releases OH-

Bronsted-Lowry

p+ donor p+ acceptor

Lewis e- pair

acceptor e- pair donor

The pH Scale and pH Calculations

The Concept of pH

Aqueous Solutions and the Concept of pH

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 (10 million) liters of water

Therefore: [H+] = [OH-] = 1 𝑚𝑜𝑙𝑒 𝑖𝑜𝑛𝑠

107𝐿 𝑊𝑎𝑡𝑒𝑟 = 1 𝑥 10−7 𝑀

So at 25 0C: [H+] = 1 x 10-7 M and [OH-] = 1 x 10-7 M

Self-Ionization

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

Water dissociation constant (Kw) is the rate at which water dissociates

At 25oC:

Kw = [H3O+][OH-] = [1 x 10-7 M][1 x 10-7 M]

Kw = 1 x 10-14 M2

Self-Ionization of Water

Water dissociation constant (Kw) is the rate at which water dissociates, and it varies with temperature.

The value of Kw is different at different temperatures

Since:

Kw = [H3O+][OH-]

Calculating pH

Definition of pH:

The negative logarithm (base 10) of the [H3O+]

Equation: pH = - log10 [H3O+]

Example: pure water at 25oC

pH = - log10 [H3O+]

pH = - log10 (1.0 x 10-7)

pH = 7

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

Logarithms

𝑙𝑜𝑔10 1 = 0 100 = 1

pH Values

Acidic solutions have pH < 7

Basic solutions have pH > 7

pH and pOH

Since at 25oC, Kw = 1 x 1014 M2

The total of pH and pOH is equal to 14.00

pH + pOH = 14.00

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

If the pOH is 2.00, what is the pH?

pH = 14.00 – pOH = 14.00 – 2.00

= 12.00

The pH Scale

Some common substances, their pH

and their [H3O+]

pH: “pouvior hyrogene” from French

meaning “hydrogen power”

Understanding the Log Function

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 of the [H3O

+] but with a positive value

Only allows for calculation if the [H3O+] is a

power of ten

If you are given a value like 0.03 M, use your calculator log function!

Calculating Concentrations from pH

Since 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+]

Because: log10[H30+] = -pH

So 10-pH = [H3O+] AND 10-pOH = [OH-]

Calculating Concentrations from pOH

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-]

Equations for Calculations

pH = - log [H3O+]

pOH = - log [OH-]

[OH-] = 10-pOH

[H3O+] = 10-pH

pH + pOH = 14

Worksheet Tutorial

Brǿnsted-Lowry Acids & Bases


H2O + H2O ↔ H3O+ + OH-

H+ H+

Base Acid Conjugate

Acid

Conjugate

Base


H3PO4 + H2O ⟷ H2PO4- + H3O

+

H+ H+

Base Acid Conjugate

Acid

Conjugate

Base

Determining pH and Titrations

Chapter 20

15.2 – Determining pH and Titrations

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

Measuring pH

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

Measuring pH

2. pH Indicators or Litmus Strips:

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

Acid-Base Indicators

Measuring pH

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

Universal Indicators

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

Universal Indicator

pH of Households Lab Image

Acid-Base Titration and pH

Acid-Base Titrations

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

Acid-Base Titrations

Acid-Base Indicator

Phenolphthalein

Equations

Molarity = moles / Liter

Macid x Vacid = Mbase x Vbase

Note: Only true for a 1:1 mole

ratio between the acid and base

We will use this equation for lab!

Normality (N)

Normality (N) – number of equivalents of solute per liter

of solution

Equation: N = n * M

normality = number of equiv * Molarity

What is the Normality of a 0.050 M Ca(OH)2 soln?

N = n * M

N = 2 equiv * M

N = 0.10 N Ca(OH)2

chemistry ms. agostine

Documents