Acids And Bases(a)

A Rong Rong time ago…… ……there were chemists, people were aware of groups of compounds that later were called acids and bases. Acids were sour to the taste (probably first observed as the taste of wine that has oxidized to give vinegar), dissolved rocks such as limestone and marble, and dissolved metals. Bases were also known as compounds that were bitter to the taste, slimy to the touch and neutralized acids. ACS(I) Pre-IB-1 Chemistry

April 2005

And not so long ago….

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Acids and Bases Acids and Bases are generally classified in 3 ways 1

Arrhenius Theory

2

Brönsted-Lowry Theory

3

Lewis Theory

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Acids and Bases 1 2 3

Arrhenius Theory Brönsted-Lowry Theory Lewis Theory

Arrhenius Acids and Bases Acids produce H+ in aqueous solutions HCl ACS(I) Pre-IB-1 Chemistry

water

H+(aq)+ Cl -(aq)

Acids and Bases 1 2 3

Arrhenius Theory Brönsted-Lowry Theory Lewis Theory

Arrhenius Acids and Bases Acids produce H+ in aqueous solutions HCl

water

H+(aq)+ Cl

-

(aq)

Bases produce OH- in aqueous solutions water

NaOH ACS(I) Pre-IB-1 Chemistry

Na+(aq)+ OH-(aq)

Acids and Bases 1 2 3

Arrhenius Theory Brönsted-Lowry Theory Lewis Theory

Arrhenius Acids and Bases Acids produce H+ in aqueous solutions water

HCl

H+(aq)+ Cl

-

(aq)

Bases produce OH- in aqueous solutions water

NaOH

Na+(aq)+ OH-(aq)

Acids neutralize Bases producing salts and water

H+ (aq) + OH- (aq) → H2O (aq)

NaOH + HCl

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NaCl

+ HO

Refinement to Arrhenius Theory It is highly improbable that a proton could exist independently in aqueous solution. So, HCl + H2O → H3O+ + Cl¯ (hydroxonium ion) CH3COOH + H2O ⇄ H3O+ + CH3COO¯  

Recognises the role of “solvent” water molecules in the dissociation of acids and, Water is necessary for an acid to display its properties.

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Limitations to the Arrhenius Theory Aqueous:  NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)  NH3(aq) + HCl(aq) → NH4Cl(aq)   

OH- ions are reacting with the H+ ions in 1st eqn. However, no OH- ions are present in 2nd eqn. However, can be argued because of NH4OH.

Gaseous:  NH3(g) + HCl(g) → NH4Cl(s)  Arrhenius theory does not count this as an acid-base ACS(I) Pre-IB-1 Chemistry

Acids and Bases 1 2 3

Arrhenius Theory Brönsted-Lowry Theory Lewis Theory

Brönsted-Lowry Acids and Bases Acid : a proton-donor (or hydrogen ion donor)

Base : a proton-acceptor (or hydrogen ion acceptor) ACS(I) Pre-IB-1 Chemistry

Brönsted-Lowry Acids and Bases Acid : a proton-donor (any substance that can donate a hydrogen ion)

An acid has a conjugate base. ( The particle left when the acid has given away its proton)

HCl + H2O

H3O+ + Cl-

acid ACS(I) Pre-IB-1 Chemistry

conjugate base They differ by a H+ ion

Brönsted-Lowry Acids and Bases is really a chloride ion with an extra proton   As the H2O gets close to the HCl, the HCl donates a proton to the water molecule

So, HCl can donate one proton to a water molecule. We end up with a "hydroxonium ion“and a chloride ion.

Co-ordinate Bond or Dative Covalent Bond ACS(I) Pre-IB-1 Chemistry

HCl + H2O → H3O+ + Cl-

This proton is transferred… ACS(I) Pre-IB-1 Chemistry

… to this lone pair.

H3O+ + OH- → 2H2O

How about the HCl and NH3 problem?

NH3 accepts a proton and HCl donates its proton. Now bases cover more substances like NH3, H2O and anions like Cl-. ACS(I) Pre-IB-1 Chemistry

Brönsted-Lowry Acids and Bases Acid : a proton-donor Base : a proton-acceptor

An acid has a conjugate base. ( The particle left when the acid has given away its proton) A base has a conjugate acid. ( The particle formed after bonding with the proton )

HCl + H2O acid ACS(I) Pre-IB-1 Chemistry

base

H3O+ + Clconjugate acid

conjugate base

Brönsted-Lowry Acids and Bases Acid : a proton-donor Base : a proton-acceptor

An acid has a conjugate base. ( The particle left when the acid has given away its proton) A base has a conjugate acid. ( The particle formed after bonding with the proton )

H2O + NH3 acid ACS(I) Pre-IB-1 Chemistry

base

NH4+ + OHconjugate acid

conjugate base





A compound can be both a Brønsted acid and a Brønsted base Water is the perfect example of this behavior because it simultaneously acts as an acid and a base when it forms the H3O+ and OH- ions.

self-ionization of water ACS(I) Pre-IB-1 Chemistry

Self-Ionization of Water H2O(l) + H2O(l)

H3O+(aq) + OH-(aq)

When this reaction takes place, we get two ions. H3O+ ---- the hydroxonium ion OH- ---- the hydroxide ion •for each hydroxonium ion that is made there is also a hydroxide ion made, there are equal numbers of each one present in pure water. •Of course there are very few of either ion present in pure water so water doesn't act as an acid or a base in most cases. ACS(I) Pre-IB-1 Chemistry

Self-Ionization of Water

H2O(l) + H2O(l)

H3O+(aq) + OH-(aq)

Water as the perfect example of a Brønsted acid and a Brønsted base

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Acids and Bases 1 2 3

Arrhenius Theory Brönsted-Lowry Theory Lewis Theory

Lewis Acids and Bases Acid: an electron pair acceptor Base: an electron pair donor A "Lewis acid" is any atom, ion, or molecule which can accept electrons (usually electron pair) and a "Lewis base" is any atom, ion, or molecule capable of donating electrons. ACS(I) Pre-IB-1 Chemistry

Octet deficient molecules like BF3 are often strong Lewis acids because they can achieve an octet configuration by accepting a lone pair from a Lewis base ACS(I) Pre-IB-1 Chemistry

Compounds involving elements in periods lower than the second period can act as Lewis acids as well by expanding their valence shells. Thus, SnCl4 acts as a Lewis acid ACS(I) Pre-IB-1 Chemistry

What controls the Strength of an Acid?

•The bond strength of the hydrogen to be dissociated/ protonated. •The stability of the anion formed. HA ACS(I) Pre-IB-1 Chemistry

H + + A-

Weak Acid versus Strong Acid

What about dilute acids? ACS(I) Pre-IB-1 Chemistry

What about concentrate d acids?

The strength of an acid is measured by the extent it dissociates/protonates in water. The strength of an acid is directly proportionate to the concentration of hydrogen ions present in the solution, [H+].

pH = - log10[H+] or more correctly, ACS(I) Pre-IB-1 Chemistry

pH = - lg[H+]

Question 1 





The concentration of H+ ions in a bottle of table wine was 3.2 X 10-4 M right after the cork was removed. Only half of the wine was consumed. The other half, after it had been standing open to the air for a month, was found to have a [H] ion concentration equal to 1.0 X 10-3 M. Calculate the pH of the wine on these two occasions. pH = 3.49 and 3.00 Explain?

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Question 2 The pH of rainwater collected in a northern region of Singapore on a particular day was 4.82. Calculate the H+ ion concentration of the rainwater.  [H+] = 1.5 X 10-5 M. 

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Question 3 

In a NaOH solution [OH-] is 2.9 X 10-4 M. Calculate the pH of the solution.



pH = 10.46

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

When we add a strong acid to water, the concentration of the H3O+ ion increases. HCl(aq) + H2O(l)



H3O+(aq) + Cl-(aq)

At the same time, the OH- ion concentration decreases because the H3O+ ions produced in this reaction neutralize some of the OH- ions in water.

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Indicators Acid-base indicators are weak acids or bases themselves which establish an equilibrium between their molecular and ionic forms. The molecular form has a different colour than the ionic form. Changes in the pH cause a shift in the equilibrium which favours one species over the other . ACS(I) Pre-IB-1 Chemistry

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[H+]

Acids

Neutral

Bases

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pH

Example

1 X 100

0

HCl

1 x 10-1

1

Stomach acid

1 x 10-2

2

Lemon juice

1 x 10-3

3

Vinegar

1 x 10-4

4

Soda

1 x 10-5

5

Rainwater

1 x 10-6

6

Milk

1 x 10-7

7

Pure water

1 x 10-8

8

Egg whites

1 x 10-9

9

Baking Soda

1 x 10-10

10

Tums® antacid

1 x 10-11

11

1 x 10-12

12

Ammonia Mineral Lime - Ca(OH)2

1 x 10-13

13

Drano®

1 x 10-14

14

NaOH

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BASICITY of an ACID The basicity is the number of hydrogen atoms in an acid that are replaceable by a metal. Hydrochloric acid ( HCl ) has one replaceable hydrogen, hence it is monobasic. Sulphuric acid ( H2SO4 ) is dibasic Give an example of a tribasic acid? ACS(I) Pre-IB-1 Chemistry

Phosphoric acid Citric acid