Group 2 Elements

• Elements in Group 1 and Group 2 are known as: s-block elements because their valence [bonding] electrons are in the s orbitals.

M  → M2+

+

2e-

• These elements give away 2 electrons when they react.
• As the Group 2 elements cause the reduction of other compounds or elements (as it gives two electrons to another compound), we say it is a good Reducing agent.
• Reactivity increases as you go down the Group.  This means they lose their electrons more readily.
• This means as you go down Group 2, they become better Reducing agents.

Physical Properties

• High melting and boiling points
• Low density metals
• Form colourless compounds/white when solid

Atomic Radius

The atomic radius increases as you go down the group due to:

• The increased number of electron shells
• The less effective nuclear attraction

1st Ionisation energy

The first ionisation energy decreases as you go down the group due to:

• despite the increased nuclear charge
• there is increased electron shielding
• the atomic radii increases
• Overall the effective nuclear attraction decreases.
• Become more reactive as you go down group 2.

Group 2 elements reacted with oxygen

• Group 2 metals react vigorously with oxygen to form simple ionic oxide

e.g. 2Mg (s) + O₂ (g) → MgO (s)

• It reacts with increasing vigour as you go down the group

Group 2 elements reacted with water

• Beryllium does not react with water
• The rest of the group 2 metals react with increasing vigour as you go down the group to form a metal hydroxide, M(OH)₂ and hydrogen gas

e.g. M (s) + 2H₂O (l) → M(OH)₂ (aq) + H₂ (g)

• These hydroxides have increasing solubility in water to form alkaline solutions.

Group 2 oxides and hydroxides
 

• Group 2 oxides and hydroxides are bases
• They are neutralised by acids to from a salt and water

e.g. MgO (s) + 2HCl (aq) → MgCl₂ (aq) + H₂O (l)
       Ca(OH)₂ (s) + 2HCl → CaCl₂ (aq) + H₂O (l)
 
Oxides

• Group 2 oxides react with water to form a solution of the metal hydroxide.
• These solutions usually have a pH of 10-12

e.g. MgO (s) + H₂O (l) → Mg(OH)₂ (aq)
 
Hydroxides

• Group 2 hydroxides dissolve in water to form alkaline solutions.

e.g. Ca(OH)₂ (s) + aq → Ca²⁺ (aq) + 2OH^- (aq)

• The solubility of the hydroxides in water increases as you go down the group.

Uses of group 2 hydroxides?
 
- Calcium hydroxide, Ca(OH)₂ is used by farmers and gardeners to neutralise acidic soils
- Magnesium hydroxide, Mg(OH)₂ is used in 'Milk of magnesia' to relieve indigestion. It works by neutralising any excess acid in the stomach.
 
Group 2 Metal Carbonates
 
THERMAL DECOMPOSITION is the breaking of a chemical substance using heat into at least 2 smaller substances.

• The group 2 carbonates undergo thermal decomposition to form the metal oxide and carbon dioxide gas.

 e.g. MgCO₃ (s) → MgO (s) + CO₂ (g)

• The group 2 carbonates decompose at a higher temperatures as you go down the group.

• e.g. BaCO₃ needs the most energy to break the bonds in the compound.

 

 

Calcium Compounds

 

 



 
 
- Most calcium is found as calcium carbonate in limestone
 
Uses of calcium compounds

• Limestone - Calcium carbonate CaCO₃ (s) - making cement
• Quicklime Calcium Oxide CaO (s) - iron purification
• Slaked Lime - Solid calcium hydroxide Ca(OH)₂ (s) - Soil Treatment
• Lime water - Aqueous calcium hydroxide Ca(OH)₂ (aq) - testing for CO₂

 

Intermolecular Forces

What is an intermolecular force?

An intermolecular force is an attractive force between neighbouring molecules.

Bond Type [Intermolecular forces]                     Relative Strength
Ionic and covalent bond                                    1000
Hydrogen bonds                                                 50
Dipole-dipole forces                                           10
Van-der-Waals forces                                         1

• Van-der-Waals forces exist between all covalent molecules, whether polar or non-polar.

What causes Van-der-Waals forces?

• The uneven distribution of electrons creates an instantaneous dipole on one atom.
• This induces dipoles on nearby molecules.
• Molecules are now attracted to each other by weak forces
• The greater the number of electrons, the stronger the Van-der-Waals force.
• The more Van-der-Waals forces, the higher the melting and boiling point.

 

 

Permanent Dipole

 

Permanent dipole-dipole forces: is a weak attractive force between permanent dipoles in neighbouring POLAR molecules.

 

 

 

• There is a dipole-dipole interaction between a partially positive atom and a partially negative atom of another molecule.
• This example shows the permanent dipole between two hydrochloric acid molecule



Hydrogen Bonds

A hydrogen bond is a strong dipole-dipole attraction between:

• An electron deficient hydrogen atom
• And a lone pair of electrons on a highly electronegative atom on a different molecule          (e.g. N, O, F)

It is the strongest intermolecular force.

 

 

 

 

 

 

Ice Lattice (open network of water molecules)

• Hydrogen bonds hold the H2O molecules apart
• Each oxygen atom has 4 bonds (2 covalent, 2 hydrogen bonds)
• Hydrogen bonds are slightly longer
• Open structure made up of rings of 6 oxygen atoms
• Covalent bonds are stronger than the hydrogen bonds.

Q: Suggest why ice has a higher melting point than solid ammonia. (2)

A:

• Ice has stronger hydrogen bonds than ammonia
• O has two lone pairs/N has one OR YOU COULD SAY - O is more electronegative than N

Here is the electronegativity's of the elements in the periodic table:

 

 

Q: Nitrogen can form a fluoride, NF₃ which has a permanent dipole. Explain why NF₃ has a permanent dipole. (2)
A:

• It has a permanent dipole because F is more electronegative than N.
• Also the shape of the molecule is non-symmetrical as NF₃ is non-symmetrical as it is trigonal pyramidal.
• Therefore the dipoles do not council each other out so it has a permanent dipole.



Crude Oil and Hydrocarbons

Definitions:

Hydrocarbon: Is a compound or a molecule made up of carbon and hydrogen ONLY

Saturated: Contains only C-C bonds

Unsaturated: Contains a C=C bond

Crude Oil: A mixture of hydrocarbons

1. Crude oil can be separated into fractions because the fractions have different boiling points.
2. Small size hydrocarbon molecules have less surface contact and so less van-der-Waals forces between the molecules, so less energy is needed to break them.
3. Branched chain hydrocarbons have lower boiling points than straight chain hydrocarbons.
4. Branched chains have less surface contact, less Van-der-Waals forces between molecules, less energy needed to overcome them.

Fractional Distillation Process

• Fractions are separated because of their differences in boiling point.
• Crude oil vapour is fed into the fractionating tower.
• The tower has a temperature gradient.
• The upper part has lower temperatures and the temperature in the tower increases as you go down.
• Small size hydrocarbon fractions condense from the top.
• Larger sized hydrocarbon fractions condense from the bottom.

 

• Petrol is of highest demand.
• The amount of petrol produced from fractional distillation is not enough to meet the demand..
• Petrol has to be refined to increase the quality/octane number.

Here are the three ways to increase the octane number:

Alkanes

 

Alkanes General Formula:  C_nH_2n+2

 

Reactions of Alkanes

 

Alkanes are extremely unreactive. Here are the 3 reasons why:

 

1. C-C and C-H bonds are non-polar and cannot be attracted to nucleophiles.

 

2. C-C has a low electron density (does not have an electron rich centre) so it is NOT attracted to electrophiles.

 

3. The bond enthalpy of C-C and the C-H bonds are TOO HIGH.

 

Therefore... Alkanes undergo: Free Radical Substitution.

 

Definition of a free radical: A species with an unpaired electron.

 

The mechanism for the reaction of alkanes with a free radical is shown below:

 



 

 

 

Why can a mixture of products be formed?

 

1. Because different products can be formed in the termination step.

 

2. Several isomers of a product can be made (e.g. the halogenoalkane)

 

3. Multi-substitution of the hydrogen, e.g. in CH₄.

 

 

 

Alkanes can also react with halogens to form halogenoalkanes.

 

E.g.

C₄H₁₀ + I₂ → C₄H₉I + HI

C₆H₁₄ + Br₂ → C₆H₁₃Cl + HCl

 

 

Quick Quiz:

 

1. What is a nucleophile? (1)

2. Why does the boiling point increase as the Mr increases in the straight chain alkanes? (1)

3. What type of fission takes place when the bond in Cl₂ breaks? (1)

4. State the condition necessary for the fission of the bond in Cl₂? (1)

5. Write a balanced equation for the complete combustion of propane? (1)

 

Answers:

 

1. An electron pair donor.

2. There are more Van-der-Waals forces because there is more surface contact.

3. Homolytic fission

4. UV light

5. C₃H₈ + 5O₂ → 3CO₂ + 4H₂O

 

 

 

 

 

F322 Mechanisms

 

The F322 exam has quite a lot of mechanisms and equations. Therefore I decided to make a revision poster containing all of them! I then stuck it to the back of my door. Today I want to share my poster with you, and I advise any one doing this exam to make a similar poster... and don't forget to use colours!

 

 

 

 

Periodicity: Ionisation energies and Atomic Radii

Key Words:

o   Period

Is a horizontal row of elements in the Periodic Table. Elements show trends in properties across a period.

o   Group

Is a vertical column in the periodic table. Elements in a group have similar chemical properties and their atoms have the same number of outer shell electrons.

o   Periodicity

 

Is a regular periodic variation of properties of elements with atomic number and position in the Periodic table.

o   First Ionisation energy

 

   The first ionisation energy of an element is the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.

 

o   Electron shielding

 

    Is the repulsion between electrons in different inner shells. Shielding reduces the net attractive force from the positive nucleus on the outer-shell electrons.

Periodicity: Ionisation energies and atomic radii

Recap:

• In a vertical group they have similar electron configurations.  They have 1) The same Nº of electrons in the outer shell and 2) the same type of orbitals
•  á Nuclear charge, greater the attractive force on the outer electrons, á I.E
• Greater the atomic radius, Smaller the nuclear attraction experienced by the outer shells, â I.E
•  Electron Shielding= á inner shells, larger shielding effect, smaller nuclear attraction experienced by outer electrons, â I.E

Across a period

§  Ionisation energy increases

§  Most important factor= Increased nuclear charge

§  Attraction between the nucleus and outer electrons increases → More energy is needed to remove an electron

§  á Number of protons → á attraction acting on electrons

§  Electrons added to the same shell → Outer shell is drawn inwards slightly

§  Same Nº of inner shells → electron shielding hardly changes

§  â Atomic radius as increased nuclear charge pulls the electrons in towards it.

 

There is a sharp decrease in ionisation energy between one end of a period to the start of the next:

« Addition of new shell→ further from the nucleus

« Increased distance of outermost shell from the nucleus

« Increased electron shielding of the outermost shell by inner shells.

Down a group

µ First Ionisation energy decreases

µ Most important factor= Increased distance and shielding

µ Nº of shells á, distance of the outer electrons from the nucleus increases, weaker force of attraction on the outer electrons

µ More inner shells, shielding effect on the outer electrons from the nuclear charge increases, less attraction

µ Nº of protons in the nucleus á, but outweighed by the increase in distance and shielding

µ Atomic radius increases down the group, less attraction so electrons are not pulled as close to the nucleus.

Conclusion:

Ionisation energy shows a general increase across each period.                                                                       

Across each period, the number of protons increases, so there is more nuclear attraction acting on the electrons.

Electrons are added to the same shell, so the nuclear attraction draws the outer shells inwards slightly.

There is the same number of inner shells and so electron shielding will hardly change.

The increased nuclear charge is the significant factor.

______________

Ionisation energy shows a general decrease down each group.
The number of shells increases.

The distance of the outer electrons from the nucleus increases, increasing the atomic radius. There is a weaker force of attraction on the outer electrons.     

 

There are more inner shells, so the electrons are more effectively shielded from the nuclear charge. Again, there is less attraction.                                                                                             

The number of protons in the nucleus also increases, but the resulting increased attraction is far outweighed by the increase in distance and shielding.

This is a video that really helped me to grasp this topic: