Past Zoom Webinar

Why Carboxylic Acids undergo Nucleophilic Substitution

In this JC2 webinar we want to understand why carboxylic acids and derivatives undergo nucleophilic substitution instead of addition.

We can deduce the mechanism of a functional group based on the charge and degree of saturation of carbon.

1. Charge of Functional Group Carbon

Positive or partial positive carbon will react with electron rich nucleophiles.

Examples include halogenoalkanes, alcohols, carbonyl compounds, acids and derivatives, and nitrogen compounds where carbon is attached to more electronegative atoms such as halogen, oxygen or nitrogen.

Negative or electron rich carbon will react with electron deficient electrophiles.

Examples include alkenes and arenes.

Neutral or non-polar carbon will only react with free radicals hence they are generally unreactive.

Example will be alkanes.

2. Degree of Saturation of Functional Group Carbon

Saturated or sp3 hybridised carbon will tend to undergo substitution.

Examples are alkanes, halogenoalkanes, alcohols and amines.

Unsaturated or sp2 hybridised carbon will tend to undergo addition.

Examples are alkenes and carbonyl compounds.

We do have exceptions for this in benzene and acids and derivatives, where the carbon is unsaturated but prefers substitution.

Benzene prefers substitution to addition as it wants to retain its resonance stability.

Addition will mean that all carbons in benzene will be saturated, and there will be no more delocalised pi electrons to stabilise benzene.

Therefore substitution is preferred for benzene.

Let's consider nucleophilic substitution of acids and derivatives in detail.

acid nucleophilic substitution reaction

The reason why acids prefer substitution lies in the stability of the intermediate formed.

Let's take a look at the mechanism.

acid nucleophilic substitution mechanism

In the first step the nucleophile attacks acyl carbon.

The pi bond between carbon and oxygen opens up and both electrons go to oxygen.

This causes the C=O to become C-O, and oxygen to be negatively charged.

Notice this first step is exactly the same as Step 2 in the nucleophilic addition mechanism of carbonyl compounds.

acid nucleophilic substitution compare nucleophilic addition of carbonyl cpd

Check out the following video lesson for a more detailed discussion of nucleophilic addition of carbonyl compounds.

So it is still true that since both reactive carbons for acids and carbonyl compounds are partial positive and unsaturated, they both undergo nucleophilic addition.

The difference lies in the stability of the intermediates formed.

For carbonyl compound, intermediate formed has a saturated carbon which is bonded to only 2 electronegative species (O and CN).

Carbon is relatively stable hence will remain saturated, ie addition reaction.

For acid, intermediate formed has a saturated carbon bonded to 3 electronegative species (O, Y and Nu).

Carbon is very positively charged and will eliminate one electronegative Y group to increase stability, ie substitution reaction.

Drawing nucleophilic substitution mechanism for acids and derivatives is not in A Level syllabus but it's useful for understanding organic chem mechanism in general.

Topic: Acids and Derivatives, Organic Chemistry, A Level Chemistry, Singapore

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3 Rules to write out Electronic Configuration using Quantum Model

In this JC1 webinar we want to learn how to write out electronic configuration using Quantum Model.

There are 3 rules that we can use.

1. Aufbau Principle

Electrons are first placed in the orbital of the lowest energy, then the orbital of the next lowest energy and so on.

This idea is exactly the same as Bohr's Model, where we fill electrons to the first shell first which has the lowest energy, then the second shell, and so on.

The distance from the nucleus and energy level of subshells from principle quantum shells n=1 to n=4 are shown below:

3 rules electronic config quantum model energy level of subshells

The subshells in increasing energy and order of filling electrons will therefore be:

1s 2s 2p 3s 3p 4s 3d 4p ...

Notice 4s subshell is more stable or closer to the nucleus than 3d subshell when empty.

Hence we add electrons to 4s subshell first, then to 3d subshell.

We can also determine the maximum number of electrons each subshell can hold.

3 rules electronic config quantum model max no of electrons in each subshell

- s subshell with 1 orbital can hold 2 electrons
- p subshell with 3 orbitals can hold 6 electrons
- d subshell with 5 orbitals can hold 10 electrons

2. Pauli Exclusion Principle

Each orbital can only hold a maximum of 2 electrons.

If there are 2 electrons in the same orbital, they must have opposite spins.

So these are the allowed and not allowed configurations we can have for an orbital.

3 rules electronic config quantum model pauli exclusion principle allowed and not allowed

3. Hund's Rule

Electrons are added to orbitals in a subshell singly and with parallel spins before orbitals are occupied in pairs.

Using p subshell as an example, the first electron can be placed in the first orbital with spin up.

3 rules electronic config quantum model hunds rule 2p1

Second electron should be placed in another empty orbital with the same spin as the first one to minimise repulsion.

3 rules electronic config quantum model hunds rule 2p2

Third electron should be placed in the last empty orbital with same spin.

3 rules electronic config quantum model hunds rule 2p3

Fourth electron now has to be paired with another electron in the first orbital so it has to be spin down or of opposite spin.

3 rules electronic config quantum model hunds rule 2p4

Fifth electron will go to the second orbital with spin down.

3 rules electronic config quantum model hunds rule 2p5

Finally the sixth electron will go to the third orbital with spin down.

3 rules electronic config quantum model hunds rule 2p6

A combination of these 3 rules will allow us to write out electronic configurations of atoms or ions using the Quantum Model.

Topic: Atomic Structure, Physical Chemistry, A Level Chemistry, Singapore

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What is Oxidation State?

In this JC1 webinar we want to understand what oxidation state is.

Oxidation state of an atom is the charge that it carries as a monatomic ion for ionic compounds, or the hypothetical charge it would have if it were an ion for covalent molecules.

Ionic Compounds

The oxidation state of the element will be its charge as an ion.

For instance in NaCl, charge of sodium ion is +1 hence oxidation state is +1; charge of chloride ion is -1 hence oxidation state is -1.

oxidation state ionic compound

Covalent Compounds

The oxidation state of the element will be the hypothetical charge if it were an ion.

In this case we have to consider the breaking of all covalent bonds and assign the shared electrons to the more electronegative atom for each covalent bond broken.

We can use this Periodic Table to help compare electronegativity.

oxidation state periodic table

Fluorine is the most electronegative element and electronegativity decreases as we move away from fluorine.

Let's have a few examples to deduce oxidation state of elements in covalent compounds.

Hydrogen Chloride - HCl

oxidation state HCl

Cl is more electronegative than H and will pull the electron pair closer to itself in the covalent bond.

When the H-Cl bond is broken both electrons will go to Cl, hence Cl will gain a -1 charge as an ion.

Therefore oxidation state of Cl in HCl is -1.

H loses its electron hence will gain a +1 charge and oxidation state.

Methane - CH4

oxidation state CH4

C is more electronegative than H hence when C-H bond is broken, C will gain -1 charge and H gains +1 charge.

When all the four C-H bonds are broken, C will gain a total charge and oxidation state of -4; H will have charge and oxidation state of +1.

Chlorine element - Cl2

oxidation state Cl2

There is no difference in electronegativity in Cl-Cl bond, hence it will break equally to give Cl atoms with zero charge.

Therefore the oxidation state of Cl in Cl2 element is zero.

This is consistent with what we memorised at secondary level where elements have zero oxidation states.

Topic: Redox Reactions, Physical Chemistry, A Level Chemistry, Singapore

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Actual Oxidation State versus Average Oxidation State

In this JC1 webinar we want to compare Actual Oxidation State versus Average Oxidation State.

Let's use S4O62- as an example and calculate the oxidation state of sulfur.

actual oxidation state calculate OS of sulfur

Using oxidation state of -2 for oxygen as a reference, we can determine sulfur to have oxidation state of +2.5.

It might not seem unusual but once we consider the definition of oxidation state, we will realise that it is not possible.

Oxidation state of an atom is the charge that it carries as a monatomic ion for ionic compounds, or the hypothetical charge it would have if it were an ion for covalent molecules.

Check out this video for a more detailed discussion of the meaning of oxidation state.

So for sulfur to have an oxidation state of +2.5, it must lose 2.5 electrons to form an ion of +2.5 charge.

This is not possible as we cannot lose half an electron.

Hence what we are calculating is the average oxidation state of sulfur, which is good enough for us to handle most redox reactions and questions.

So how do we determine the actual oxidation state of sulfur then?

We would need to consider the lewis structure of S4O62- and break all the covalent bonds to determine the hypothetical charge of sulfur if it were an ion.

actual oxidation state lewis structure of S4O62minus

Oxygen is more electronegative than sulfur hence:
- when O-S single bond is broken, O will gain -1 charge and S will gain +1 charge;
- when O=S double bond is broken, O will gain -2 charge and S will gain +2 charge

There is no difference in electronegativity in S-S bond hence no charge is acquired when S-S bond is broken.

Therefore we can deduce that:
- the yellow sulfur that is directly bonded to 3 oxygen atoms will acquire a charge of +5 when it forms an ion. Hence oxidation state of that yellow sulfur is +5.
- the blue sulfur will not acquire any charge when all bonds around it are broken. Hence oxidation state of that blue sulfur is zero.

Notice if we sum up all the actual oxidation states of sulfur (5 + 0 + 0 + 5 = 10) and divide by the number of sulfur atoms (4), we have the average oxidation state of +2.5.

We need to know this technique of determining actual oxidation state of a specific atom in a compound as it will be useful in Organic Chemistry.

Topic: Redox Reactions, Physical Chemistry, A Level Chemistry, Singapore

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Three Definitions of Acids and Bases

In this JC1 webinar we want to discuss the three definitions of acids and bases in A Level Chemistry.

The balanced equation for reaction between sulfuric acid and ammonia is as follows:

H2SO4 + 2NH3 → (NH4)2SO4

This is clearly an acid base reaction, but where is the water of neutralisation?

In order to understand this, we need to go through the different definitions of acids and bases.

1. Arrhenius Theory

This is the definition used in secondary schools and O Levels where:

Arrhenius acid - dissolves in water to give H+(aq)
Arrhenius base - dissolves in water to give OH-(aq)

Therefore any acid base reaction will simply be the reaction between H+ and OH- to give water.

3 definition acid base arrhenius acid base theory

This definition is quite restrictive as the acid and base must be soluble in water.

Hence this will not apply to acid base reactions in gaseous systems or organic compounds which are not soluble in water.

2. Bronsted-Lowry Theory

This is the default definition used in A Levels where:

Bronsted acid - H+ or proton donor
Bronsted base - H+ or proton acceptor

Bronsted acid and base need not be soluble in water, hence will be more general than Arrhenius theory.

The major difference is Bronsted base just need to be a H+ acceptor and forming OH- is not required.

3 definition acid base bronsted lowry acid base theory

For the example given earlier, we can see that H2SO4 is the proton donor hence Bronsted acid, while NH3 is the proton acceptor hence Bronsted base.

Bronsted acid base reaction is just a transfer of H+ from acid to base so water need not be formed.

3. Lewis Theory

We use this definition sparingly in A Levels as it is the most abstract.

Lewis acid - electron pair acceptor
Lewis base - electron pair donor

Notice there is no need for the Lewis acid to form H+, hence it is harder to link this to Arrhenius theory.

3 definition acid base lewis acid base theory

One way to remember is to use NH3 as an example of Lewis base.

Since nitrogen in NH3 has a lone pair of electrons to donate, a Lewis base is an electron pair donor.

Hence Lewis acid will be an electron pair acceptor.

Topic: Ionic Equilibria, Physical Chemistry, A Level Chemistry, Singapore

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