Module 3, Chapter 1: Periodicity

First Ionisation Energy
the energy needed to remove 1 mole of electrons from 1 more of gaseous atoms
in the 1800s, how were elements grouped?
by atomic mass
What were the 2 ways to categorise elements in the 1800s?
- by their physical and chemical properties

- by their relative atomic mass
How is the Modern Periodic Table is arranged
by proton number
What did John Dobereiner attempt to do?
group similar elements, these groups were called Dobereiner's Triads
Who made the first table of elements in 1963?
John Newlands
How did John Newlands arrange the elements?
by mass
Law of Octaves
the same properties appear every eighth element when the elements are listed in order of their atomic masses
How did Mendeleev organize his periodic table?
- he arranged the elements in his periodic table in order of increasing atomic mass

- left gaps

- similar properties in the same groups

- also predicted undiscovered elements that would go in the gaps
when electrons have been removed
What do you have to do to ionise an atom?
you have to put energy in
The Lower the Ionisation Energies.....
the easier it is to form an ion
Rules for Ionisation Energies
- you must use the gas state symbol (g), because ionization energies are measured for gaseous atoms

- always refer to 1 mole of atoms, rather than a single atom
Factors Affecting Ionisation Energy
- nucleus charge

- atomic radius

- shielding
How does Nuclear Charge Affect Ionisation Energy
- the more protons there are in the nucleus

- the more positively charged the nucleus is

- the stronger the attraction for the electrons
How does Atomic Radius Affect Ionisation Energy
- attraction falls of very rapidly with distance

- an electron close to the nucleus will be much more strongly attracted than further away
How does Shielding Affect Ionisation Energy
- as the number of electrons between the outer electrons and the nucleus the outer electrons and the nucleus increases, the outer electrons feel less attraction towards the nuclear charge

- this lessening of the pull of the nucleus by inner shells of electrons is called shielding
High Ionisation Energy
there's a strong electrostatic attraction between the electron and the nucleus, so more energy is needed to overcome the attraction and remove the electron
Pattern between Ionisation and Groups
- Ionisation Energy Decreases down a group
Why do Ionisation Energies Decreases down a group?
- elements further down a group have extra electron shells compared to the ones above

- extra shell = atomic radius is larger

- outer electrons further away from the nucleus, which greatly reduces their attraction to the nucleus

- extra inner shells shield the outer electrons from the attraction of the nucleus
What provides evidence that electron shells exist?
ionisation energies
Why does the fact that the positive charge of the nucleus does increase as you gown a group (due to extra protons) not affect the ionisation energy?
the effect is overridden by the effect of the extra shells
Relationship between Ionisation Energies and periods?
- Ionisation Energy Increases ACROSS period
Why do ionisation energies Increases across a period?
- across period = harder to remove outer electrons

- number of protons increasing as positive charge of nucleus increases = electrons are pulled closer to the nucleus = atomic radius smaller

- extra electrons that the elements gain across a period are added to outer energy level (don't provide extra shielding)
Exceptions to Ionisation energy
first ionisation energy decrease between groups 2 & 3 and between 5 & 6
Why is there a drop in ionisation energy between group 2 and 3?
- is due to sub - shell structure

- outer electron in group 3 is in a p orbital rather than an s

- a p orbital has a slightly higher energy than an s orbital in the same shell, average to be found further from the nucleus

- the p orbital also has additional shielding provided by the s electron
Why is there a drop in ionisation energy between group 5 and 6?
- in the group 5 elements, the electron is being removed from a singly - occupied orbital

- in the group 6 elements, the electron is being removed from an orbital containing 2 electrons

- the repulsion between 2 electrons in an orbital means that electrons

- the repulsion between 2 electrons in an orbital means that electrons are easier to remove from shared orbitals
Succesive Ionisation Energies
- involves removing additional electrons

- you can remove all the electrons from an atom, leaving only the nucleus

- each time you remove an electron, there's a successive ionisation energy
Successive Ionisation Energies within each Shell
Why do Successive Ionisation Energies within each Shell increase?
- this is because electrons are being removed from an increasingly positive ion

- less repulsion amongst the remaining electrons

- electrons are held more strongly by the nucleus
Function of Graphs of Successive Ionisation Energies
- provides evidence for the shell structure of atoms

- they tell you which group of the periodic table an element belongs to

- these graphs also predict the electronic structure of an element
How to Graphs of Successive Ionisation Energies tell you which group of the periodic table an element belongs to?
count how many electrons are removed before the first big jump to find the group number
How to Graphs of Successive Ionisation Energies predict the electronic structure of an element?
working from right to left, count number of protons there are before each big jump to find how many electrons are in each shell
Giant Covalent Lattice Structure
huge networks of covalently bonded atoms (macromolecular structures)
Macromolecular Structures
huge networks of covalently bonded atoms
Why can carbon atoms form macromolecular structures?
because they can each form four strong covalent bonds
different forms of the same element in the same state
How many carbon atoms are each carbon atom bonded to in Diamond?
How do Carbon atoms arrange themselves in Diamond?
in a tetrahedral shape - Crystal Lattice Structure
Melting Point of Diamond
- very high

- strong covalent bonds
Thermal Conductivity of Diamond
- vibrations travel easily through stiff lattice

- good thermal conductor
Why is diamond hard?
each carbon atom is bonded to four others via strong covalent bonds
Electrical Conductivity of Diamond
- can't conduct electricity

- all the outer electrons held in localised bonds

- no electrons to move and carry charge
Solubility in Solvent of Diamond
won't dissolve in any solvent
Silicon Bonding
- each silicon atom is able to form 4 strong, covalent bonds
Bonding in Graphite
- each carbon atom covalently bonded with 3 bonds each

- sheets of hexagons are bonded together by weak induced dipole dipole forces

- 4th outer electron of each carbon atom is delocalised between the sheets of hexagons
Why is graphite slippery?
- weak forces between layers in graphite are easily broken

- sheets slide over each other

- - feels slippery

- - used as dry lubricant and in pencils
Electrical Conductivity of Graphite
- the delocalised electrons in graphite aren't attached to any particular carbon atom

- free to move

- electric current can flow
Why is graphite less dense?
layers are far apart compared to length of covalent bonds
Why does graphite have a very high melting point?
because of strong covalent bonds in the hexagon sheets
Why is graphite insoluble in any solvent?
covalent bonds are too strong to break
How are the sheets of hexagons in graphite bonded together?
by weak induced dipole dipole forces
- one layer of graphite

- 2D

- one atom thick
Why is Graphene strong?
- delocalised electrons strengthen the covalent bonds between the carbon atoms
Can graphite conduct electricity?
- delocalised electrons in graphene free to move along the sheet

- no layers = move quickly above and below the sheet

- making graphene the best known electrical conductor
Why do metals conduct electricity?
The bonding in metals creates delocalised electrons (free to move and pass on electrical current)
Structure of Metals
- giant structure

- very strong forces of electrostatic attraction between positive metal ions and negative electrons
Why do metals have high melting points?
- more delocalised electrons

- stronger bonding

- higher the melting point
What affects the melting points of metals ?
- the number of delocalised electrons

- the size of the metal ion

- the lattice structure
How does the ionic radius affect the melting point?
smaller ionic radius will hold the delocalised electrons closer to the nuclei
Why are metals malleable and ductile?
no bonds holding specific ions together, the metal ions can slide past each other when the structure is pulled
Why are metals good thermal conductors?
the delocalised electrons can pass kinetic energy to each other
Why are metals insoluble?
because of the strength of the metallic bonds
Bonds between the atoms in Simple Molecular Structures
strong covalent bonds
What do the melting and boiling points in Simple Molecular Structures depend on?
induced dipole dipole forces between their molecules
Why do Simple Molecular Structures have low melting points?
- induced dipole dipole forces

- weak and easily overcome
Why do noble gases have very low melting and boiling points?
- exists as individual atoms

- very weak induced dipole dipole forces
Boiling Point Pattern of Metals across the Period
- boiling points increase across the period

- this is because the metallic bond gets stronger

- as the ionic radius decreases and the number of delocalised electrons increases
Bond Strength of Giant Covalent Lattice Structures
- strong covalent bonds which require a lot of energy to break
Bond Strength of
Reactivity Trend Down Group 2
- as you gown the group, ionisation energies decreases

- this is due to the increasing atomic radius and shielding effect
Lower First and second ionisation Energies mean?
that it's easier to lose electrons
Group 2 Metal + Metal --->
metal hydroxide + hydrogen
Group 2 Metal + Oxygen --->
Solid White Oxides
Group 2 Metal + Dilute Acid
Salt (METAL CHLORIDE) + Hydrogen
Examples of Bases
- group 2 oxides

- hydroxides are bases
Hydroxide trend down group
oxides form more strongly alkaline solutions as you go down the group
Alkaline Metals
Group 2 metals
What is Calcium Hydroxide used for?
used in agriculture to neutralise acidic soils
What is Magnesium hydroxide and Calcium Carbonate used for?
used in some indigestion tablets as antacids
Colour of Fluorine
Colour of Chlorine
Colour of Bromine
Colour of Iodine
Melting Point Trend in Group 1
- increase down the group

- this is due to the increasing strength of the London forces

- the size and relative mass of the atoms increases
Volatility Pattern down Group 1
- volatility decreases down the group
Reactivity Down Group of Halogens
- as you go down the group, the atomic radii increases so the outer electrons are further from the nucleus

- outer electrons are also shielded more from the attraction of the positive nucleus, because there are more inner electrons

- making it harder for larger atoms to attract the electron needed to form an ion

- despite increases charge of the nucleus

- so larger atoms are less reactive
Displacement Reactions
- halogens displace less reactive halide ions from solution
Colour of Chlorine Water
Colour of Bromine Water
Colour of Iodine Water
orangey - brown
Test for Halides
1. Add dilute nitric acid

(this removes ions that might interfere with the test)

2. Add silver nitrate solution

3. Precipitate formed

4. To be sure of results, add ammonia solution
each ion has a different solubility in ammonia

- the larger the ion, the more difficult it is to dissolve
Chloride Ion in Dilute Ammonia Solution
WHITE precipitate
Bromide Ion in Dilute Ammonia Solution
CREAM precipitate
Iodide Ion in Dilute Ammonia Solution
YELLOW precipitate
Disproportionation of Halogens
- halogens undergo disproportionation with alkalis

- halogen is simultaneously oxidised and reduced
chlorine and sodium hydroxide make bleach
Chlorine in beach
- oxidation number of Cl goes UP and DOWN = disproportionation
Uses of Bleach
- water treatment

- bleach paper and textiles

- cleaning
Chlorine + Water
- forms a mixture of HCl and Chloric (I) Acid (also called hypochlorous acid)

- chlorine undergoes disproportionation
Pros of Chlorine in Water
- kills diseases causing microorganisms

- some chlorine remains in the water and prevents reinfection further down the supply

- prevents growth of algae - eliminating bad tastes and smells

- removes discoloration caused by organic compounds
Cons of Chlorine in Water
- chlorine gas is very harmful if breathed it (irritates respiratory system)

- liquid chlorine on the skin or eyes causes severe chemical burns

- accidents involving chlorine could be really dangerous / fatal

- water contains a variety of organic compounds e.g. from the decomposition of plants
Ethical Choices of Chlorine in Water
- we don't get a choice about having our water chlorinated

= Forced 'Mass Medication'
Alternatives to Chlorine
- ozone

- UV Light
- strong oxidation agent, which makes it great at killing microorganisms

- expensive to produce

- its short half life in water means that treatment isn't permanent

- won't stop water from being contaminated further down the line
UV Light
- kills microorganisms by damaging their DNA

- inefffective in cloudy water

- won't stop water from being contaminated further down the line
Test for Carbonates
1. Add dilute acid

- carbon dioxide should be released
lime water goes cloudy if CO2 is present
Test for Sulfates
1. Add dilute HCl

2. Add barium chloride solution

- white precipitate forms
barium sulfate is insoluble
Silver Chloride colour
white precipitate
Silver Bromide colour
cream precipitate
Silver Iodide colour
yellow precipitate
Why doesn't silver fluoride give any precipitate?
it's soluble
Which order should you do tests?
1. Test for carbonates

(no CO2)

2. Test for sulfates

(no precipitate)

3. Test for Halides
UV Light
UV Light
- kills microorganisms by damaging their DNA

- inefffective in cloudy water

- won't stop water from being contaminated further down the line
Test for Carbonates
1. Add dilute acid

- carbon dioxide should be released
lime water goes cloudy if CO2 is present
UV Light