1 Acids and Bases Original: L. Scheffler Modified: Swiftney.

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Acids and Bases

Original: L. Scheffler Modified: Swiftney

8.1 Theories of acids and bases 2 hours

Assessment statement Obj Teacher’s notes

8.1.1 Define acids and bases according to the Brønsted–Lowry and Lewis theories.

1 TOK: Discuss the value of using different theories to explain the same phenomenon. What is the relationship between depth and simplicity?

8.1.2 Deduce whether or not a species could act as a Brønsted–Lowry and/or a Lewis acid or base.

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8.1.3 Deduce the formula of the conjugate acid (or base) of any Brønsted–Lowry base (or acid).

3 Students should make clear the location of the proton transferred, for example, CH3COOH/CH3COO– rather than C2H4O2/C2H3O2

–.

2Original: L. Scheffler Modified: Swiftney

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Arrhenius DefinitionArrhenius

Acid - Substances in water that increase the concentration of hydrogen ions (H+).

Base - Substances in water that increase concentration of hydroxide ions (OH-).

Categorical definition – easy to sort substances into acids and bases

Problem – many bases do not actually contain hydroxides


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Bronsted-Lowry DefinitionAcid - neutral molecule, anion, or cation that donates

a proton.

Base - neutral molecule, anion, or cation that accepts a proton.

HA + :B HB+ + :A-

Ex. HCl + H2O H3O+ + Cl-

Acid Base Conj Acid Conj Base

Operational definition - The classification depends on how the substance behaves in a chemical reaction.


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Conjugate Base - The species remaining after an acid has transferred its proton.

Conjugate Acid - The species produced after base has accepted a proton.

* HA & :A- - conjugate base/acid pair

* :A- - conjugate base of acid HA

* :B & HB+ - conjugate acid/base pair

* HB+ - conjugate acid of base :B

Conjugate Acid Base Pairs


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* Note: Water can act as acid or base (Amphoteric)

Acid Base Conjugate Acid Conjugate Base

HCl + H2O H3O+ + Cl-

H2PO4- + H2O H3O+ + HPO4

2-

NH4+ + H2O H3O+ + NH3

Base Acid Conjugate Acid Conjugate Base :NH3 + H2O NH4

+ + OH-

PO43- + H2O HPO4

2- + OH-

Examples of Bronsted-Lowry Acid Base Systems


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Lewis

Acid - an electron pair acceptor

Base - an electron pair donor

*Note: form dative bonding

G.N. Lewis Definition


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Acid Base

Arrhenius: H+ OH- (based on H2O)

B-L: H+ donor H+ acceptor (e.g. NH3)

Lewis: Electron Pair Electron Pair

Acceptor Donor

Acid Base Definition Summary


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8.2 Properties of acids and bases 1 hour


8.2.1 Outline the characteristic properties of acids and bases in aqueous solution.

2 Bases that are not hydroxides, such as ammonia, soluble carbonates and hydrogen carbonates, should be included.Alkalis are bases that dissolve in water.Students should consider the effects on indicators and the reactions of acids with bases, metals and carbonates.


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Acids and Bases The concepts acids and bases were loosely defined as

substances that change some properties of water.

One of the criteria that was often used was taste.

Substances were classified salty-tasting sour-tasting sweet-tasting bitter-tasting

Sour-tasting substances would give rise to the word 'acid', which is derived from the Greek word oxein, which mutated into the Latin verb acere, which means 'to make sour'

• Vinegar is a solution of acetic acid. Citrus fruits contain citric acid.


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Acids• React with certain metals to produce hydrogen

gas.• 2HCl + 2Na 2NaCl + H2

• React with carbonates and bicarbonates to produce carbon dioxide gas• MgCO3 + 2HCl MgCl2 + H2CO3 H2O + CO2

• Have a bitter taste• Feel slippery. • Many soaps contain bases.

Bases


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Properties of Acids

þ Produce H+ (as H3O+) ions in water (the hydronium ion

is a hydrogen ion attached to a water molecule)

þ Taste sour

þ Corrode metals

þ Good Electrolytes

þ React with bases to form a salt and water

þ pH is less than 7þTurns blue litmus paper to red “Blue to Red A-CID”

(BRA)


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Properties of Bases Generally produce OH- ions in water

Taste bitter, chalky

Are electrolytes

Feel soapy, slippery

React with acids to form salts and water

pH greater than 7

Turns red litmus paper to blue “Basic Blue”


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8.3 Strong and weak acids and bases 2 hours


8.3.1 Distinguish between strong and weak acids and bases in terms of the extent of dissociation, reaction with water and electrical conductivity.

2 Aim 8: Although weakly acidic solutions are relatively safe, they still cause damage over long periods of time. Students could consider the effects of acid deposition on limestone buildings and living things.

8.3.2 State whether a given acid or base is strong or weak.

1 Students should consider hydrochloric acid, nitric acid and sulfuric acid as examples of strong acids, and carboxylic acids and carbonic acid (aqueous carbon dioxide) as weak acids.

Students should consider all group 1 hydroxides and barium hydroxide as strong bases, and ammonia and amines as weak bases.

8.3.3 Distinguish between strong and weak acids and bases, and determine the relative strengths of acids and bases, using experimental data.

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Acid Base DissociationAcid-base reactions are equilibrium processes.

The ratio of the concentrations of the reactants and products is constant for a given temperature at equilibrium

It is known as the Acid or Base Dissociation Constant.

The stronger the acid or base, the larger the value of the dissociation constant.

]B[:

][OH [HB] K O][H K

[HA]

][H ]A[: K O][H K

constant. is solutions dilute in O][H

][H OH

:Note

O][H ]B:[

]OH][HB[ K

O][H HA][

]OH][ A[: K

waterin base a For waterin acid an For

-

-

b2eq

-

a2eq

2

3

2-

-

eq2

3-

eq


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Acid Strength Strong Acid - Transfers all of its protons to water;

- Completely ionized (dissociated); - Strong electrolyte; - The conjugate base is weaker and has a negligible tendency to be protonated.

Weak Acid - Transfers only a fraction of its protons to water;

- Partially ionized (dissociated); - Weak electrolyte; - The conjugate base is stronger, readily accepting protons from water

As acid strength decreases, base strength increases. The stronger the acid, the weaker its conjugate base The weaker the acid, the stronger its conjugate baseOriginal: L. Scheffler Modified: Swiftney

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Acid Dissociation ConstantsDissociation constants for some weak acids


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Base Strength Strong Base - all molecules accept a proton; - Completely ionized (dissociated); - strong electrolyte; - conjugate acid is very weak, negligible tendency to donate protons.

Weak Base - fraction of molecules accept proton; - Partially ionized (dissociated); - weak electrolyte; - the conjugate acid is stronger. It more

readily donates protons. As base strength decreases, acid strength increases. The stronger the base, the weaker its conjugate acid. The weaker the base the stronger its conjugate acid.Original: L. Scheffler Modified: Swiftney

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Common Strong Acids/Bases

Strong BasesSodium Hydroxide

Potassium Hydroxide

*Barium Hydroxide

*Calcium Hydroxide

*While strong bases they are not very soluble

Strong AcidsHydrochloric Acid

Nitric Acid

Sulfuric Acid

Perchloric Acid


8.4 The pH scale 1 hour


8.4.1 Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale.

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8.4.2 Identify which of two or more aqueous solutions is more acidic or alkaline using pH values.

2 Students should be familiar with the use of a pH meter and universal indicator.

8.4.3 State that each change of one pH unit represents a 10-fold change in the hydrogen ion concentration [H+

(aq)].

1 Relate integral values of pH to [H+(aq)] expressed as powers of 10.Calculation of pH from [H+(aq)] is not required.TOK: The distinction between artificial and natural scales could be discussed.

8.4.4 Deduce changes in [H+(aq)] when the pH of a solution changes by more than one pH unit.

3 Aim 8: A study of the effects of small pH changes in natural environments could be included.


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Acid Base Equilibrium

Strong acid

Weak Acid

HA(aq) H3O+(aq) + A-(aq)

0.1M 0.1 M 0.1 M

pH = - log [H3O+]

e.g. HCl, HNO3

HA(aq) H3O+((aq) + A-(aq)

0.1 M <<0.1 M <<0.1 Me.g. CH3COOH

assume 100% dissociation

***The pH scale is used to describe the concentration of acid present in a solution pH is used to make an “ugly” number into something simple (between 0-14).


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The pH Scale * pH = - log [H3O+]

pH [H3O+ ] [OH- ] pOH


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pH and aciditypH = - log [H3O+] or pH = - log [H+]

The pH values of several common substances are shown at the right.

Many common foods are weak acids

Some medicines and many household cleaners are bases.


pH of Common Substances


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pH and acidity1. Acidity or Acid Strength depends on Hydronium Ion

Concentration [H3O+]

2. The pH system is a logarithmic representation of the Hydrogen Ion concentration [H+] (or [OH-]) as a means of avoiding using large numbers and powers.

* pH = - log [H3O+] = log(1 / [H3O+])

* pOH = - log [OH-] = log(1 / OH-])

3. In pure water, [H3O+] = 1 x 10-7 mol / L (at 20oC)

pH = - log(1 x 10-7) = - (0 - 7) = 7

4. pH range of solutions: 0 - 14

pH < 7 (Acidic) [H3O+]; pH > 7 (Basic) [H3O+]Original: L. Scheffler Modified: Swiftney

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Calculating the pHpH = - log [H3O+]

Example 1: If [H3O+] = 1 X 10-10

pH = - log 1 X 10-10

pH = - (- 10)

pH = 10

Example 2: If [H3O+] = 1.8 X 10-5

pH = - log 1.8 X 10-5

pH = - (- 4.74)

pH = 4.74


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pH Practice1) pH of a solution containing 25 g HCl dissolved in 1.5 L of H2O?

2) pH of a solution containing 1.32 g of HNO3 dissolved in 750 mL?

3) pH of a solution containing 1.2 moles of Nitric Acid and 1.7 moles of Hydrochloric acid dissolved in 1000 mL?

4) If a solution has a [H+] of 4.5 x 10-7 M, is this acidic or basic? Explain


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pH and acidity

* Kw = [H3O+] [OH-] = 1.0 x10-14

In pure water

* [H3O+] = [OH-] = 1.0 x10-7

* pH + pOH = 14


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Neutralization An acid will neutralize a base, giving a salt and water

as products (THIS IS STOICHIOMETRY) Examples Acid Base Salt water

HCl + NaOH NaCl + H2O

H2SO4 + 2 NaOH Na2SO4 + 2 H2O

H3PO4 + 3 KOH K3PO4 + 3 H2O

2 HCl + Ca(OH) 2 CaCl2 + 2 H2O

A salt is an ionic compound that is formed from the positive ion (cation) of the base and the negative ion (anion) of the acid


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Neutralization Calculations

If the concentration of acid or base is expressed in Molarity or mol dm-3

then:--The volume in dm3 multiplied by the concentration

yields moles (mol) . -- If the volume is expressed in cm3

the same product yields millimoles (mmol)

mol dm-3 x dm3 = mole

* mol dm-3 x cm3 = (0.001) x mole = mmol


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Neutralization Problems

The volume of solution in dm3 multiplied by concentration in moles dm-3 will yield moles.

If an acid and a base combine in a 1 to 1 ratio, the moles of acid will equal the moles of base.

*** This is a DIFFERENT way of doing Acid/Base Stoichiometry Therefore the volume of the acid multiplied by the

concentration of the acid is equal to the volume of the base multiplied by the concentration of the base.

Vacid C acid = V base C base

If any three of the variables are known, it is possible to determine the fourth.Original: L. Scheffler Modified: Swiftney

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Neutralization Problems Example 1: Hydrochloric acid reacts with potassium hydroxide according to the following reaction:

HCl + KOH KCl + H2O If 15.00 cm3 of 0.500 M HCl exactly neutralizes 24.00 cm3 of

KOH solution, what is the concentration of the KOH solution?

Solution:Vacid Cacid = Vbase Cbase

(15.00 cm3 )(0.500 M) = (24.00 cm3 ) Cbase

Cbase = (15.00 cm3 )(0.500 M) (24.00 cm3 ) Cbase = 0.313 M


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Neutralization Problems

Whenever an acid and a base do not combine in a 1 to 1 ratio, a mole factor must be added to the neutralization equation

n Vacid C acid = V base C base

The mole factor (n) is the number of times the moles the acid side of the above equation must be multiplied so as to equal the base side. (or vice versa)

Example

H2SO4 + 2 NaOH Na2SO4 + 2 H2O

The mole factor is 2 and goes on the acid side of the equation. The number of moles of H2SO4 is one half that of NaOH. Therefore the moles of H2SO4 are multiplied by 2 to equal the moles of NaOH.


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Neutralization Problems Example 2: Sulfuric acid reacts with sodium hydroxide according to the following reaction:

H2SO4 + 2 NaOH Na2SO4 + 2 H2O If 20.00 cm3 of 0.400 M H2SO4 exactly neutralizes 32.00 cm3 of

NaOH solution, what is the concentration of the NaOH solution?

Solution:In this case the mole factor is 2 and it goes on the acid side, since the mole ratio of acid to base is 1 to 2. Therefore

2 Vacid Cacid = Vbase Cbase

2 (20.00 cm3 )(0.400 M) = (32.00 cm3 ) Cbase

Cbase = (2) (20.00 cm3 )(0.400 M) (32.00 cm3 )

Cbase = 0.500 MOriginal: L. Scheffler Modified: Swiftney

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Neutralization Problems Example 3: Phosphoric acid reacts with potassium hydroxide according to the following reaction:

H3PO4 + 3 KOH K3PO4 + 3 H2O If 30.00 cm3 of 0.300 M KOH exactly neutralizes 15.00 cm3 of

H3PO4 solution, what is the concentration of the H3PO4 solution?

Solution:In this case the mole factor is 3 and it goes on the acid side, since the mole ratio of acid to base is 1 to 2. Therefore

3 Vacid Cacid = Vbase Cbase

3 (15.00 cm3 )(Cacid) = (30.00 cm3 ) (0.300 M) Cacid = (30.00 cm3 )(0.300 M) (3) (15.00 cm3 ) Cacid = 0.200 M


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Neutralization Problems Example 4: Hydrochloric acid reacts with calcium hydroxide according to the following reaction:

2 HCl + Ca(OH)2 CaCl2 + 2 H2O If 25.00 cm3 of 0.400 M HCl exactly neutralizes 20.00 cm3 of

Ca(OH)2 solution, what is the concentration of the Ca(OH)2 solution?

Solution:In this case the mole factor is 2 and it goes on the base side, since the mole ratio of acid to base is 2 to 1. Therefore

Vacid Cacid = 2 Vbase Cbase

(25.00 cm3) (0.400) = (2) (20.00 cm3) (Cbase)

Cbase = (25.00 cm3 ) (0.400 M) (2) (20.00 cm3 )

Cbase = 0.250 MOriginal: L. Scheffler Modified: Swiftney

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Weak Acid EquilibriaA weak acid is only partially ionized.Both the ion form and the unionized form exist at equilibrium HA + H2O H3O+ + A-

The acid equilibrium constant is

Ka = [H3O+ ] [A-] [HA]

Ka values are relatively small for most weak acids. The greatest part of the weak acid is in the unionized form


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Water has the ability to act as either a Bronsted- Lowry acid or base.

Autoionization – spontaneous formation of low concentrations of [H+] and OH-] ions by proton transfer from one molecule to another.

Equilibrium Constant for Water

7--

o14--w

o14--3w

-3

22c

22

-3

c

10 x 0.1 ][OH ][H

: WaterPure In

C)25(at 10 x 0.1 ][OH ][H K

C)25(at 10 x 0.1 ][OH ]O[H K

][OH ]O[H O]H[K

O]H[

][OH ]O[H K

Water as an Equilibrium System


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Amphoteric Solutions A chemical compound able to react with both an

acid or a base is amphoteric. Water is amphoteric. The two acid-base couples

of water are H3O+/H2O and H2O/OH-

It behaves sometimes like an acid, for example

And sometimes like a base :

Hydrogen carbonate ion HCO3- is also amphoteric,

it belongs to the two acid-base couples H2CO3/HCO3

- and HCO3-/CO3

2-Original: L. Scheffler Modified: Swiftney

1 Acids and Bases Original: L. Scheffler Modified: Swiftney.

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Transcript of 1 Acids and Bases Original: L. Scheffler Modified: Swiftney.