Module 7.1 - Periodicity (3.1.1 - Spec reference) Flashcards
Who crated one of the first periodic tables with gaps in it?
Dmitri Mendeleev
What did the gaps in the table represent from Mendeleev?
They represented elements that he predicted would have specific properties
These predictions were based on other elements in the same group
How many elements were in Mendeleev’s table?
63
What did Mendeleev call Germanium before it was discovered?
Eka - silicon
It was next to silicon in the table
How did Mendeleev arrange his table?
By Relative atomic mass(Mr)
How are elements arranged in the table today?
By increasing atomic number
What is a group?
What does it show?
Vertical columns of elements
Elements in groups have the same number of outer shell electrons
Group number gives the number of outer shell electrons
What is a period?
What does it show?
Horizontal rows in the periodic table
Number of the period gives the number of the highest energy electron shell of the element’s atom
Periodicity def
A repeating trend in properties of elements
Examples of periodicity in elements
Electron Configuration
Ionisation energy
Structure
Melting points
What is the trend across a period?
For each period, the s - and p - sub shells are filled in the same way
A periodic pattern
Trend down a group
All elements in that group have the same number of outer shell electrons
What are the old and new group numbers now for the table?
Which groups do they represent?
(Old)1 - 1 - alkali metals (Old)2 - 2 - alkaline earth metals Groups 3-12 are transition metals (Old)7 - 17 - halogens (Old)8- 18 - Noble gases
Definition of first ionisation energy
The energy required to remove one electron from each atom in one mole of gaseous atoms of an element to from one mole of gaseous 1+ ions
Factors affecting ionisation energy
Atomic radius
Nuclear Charge
Electron Shielding
How atomic radius affects ionisation energy
Greater the distance between nucleus and outer shell electrons the less the nuclear attraction
(Force of attraction falls off sharply with increasing distance, so atomic radius has a large effect)
How nuclear charge affects ionisation energy
More protons in nucleus gives greater attraction between nucleus and outer shell electrons
How electron shielding affects ionisation energy
Electrons are negatively charged, so inner shell electrons repel outer shell electrons
This reduces attraction between nucleus and outer shell electrons
Difference in first and second ionisation energies
Second ionisation energy is larger than first
When one electron is lost, other electrons have a stronger nuclear attraction to unchanged number of protons
Therefore electrons pulled closer to nucleus, so atomic radius decreases
So more attraction to nucleus so more energy needed to remove the remaining electrons
Def of second ionisation energy
Amount of energy required to remove on electron from each atom of one mole of gaseous 1+ ions to form one mole of gases out 2+ ions
What do large spikes in ionisation energies in successive ionisations show?
Suggest that electron must have been removed from a new electron shell
Due to atomic radius being smaller, so nuclear attraction increases
What can successive ionisations show?
Number of electrons in the outer shell
Group of the element in the periodic table
The identity of an element
Trend in first ionisation energy down a group
Atomic radius increases
More inner shells so shielding increases
Nuclear attraction on outer shells decreases
Therefore first ionisation energy decreases
Trend in first ionisation energy across a period
Nuclear charge increases(same no. of protons whilst electrons increases) Same shell:similar shielding Nuclear attraction increases Atomic radius decreases So first ionisation energy increases
Sub shell trends in first ionisation energy
First ionisation energy decreases when a new sub shell is filled
E.g. Be and B marks 2p sub shell being filled
2p sub shell has a higher energy than 2s sub shell
So 2p electron is easier to remove than one of 2s electrons in Be
So first I.E in B is less than Be
Nitrogen and oxygen sub shell trends
H
Structure of metallic bonding
Each atom donates its negative outer-shell electrons to a shared pool of electrons, which are delocalised(spread out) throughout the whole structure
The cations left behind consist of the nucleus and the inner electron shells of the metal atom
Def of metallic bond?
The strong electrostatic attraction between cations and delocalised electrons
Structure of metals
Cations are fixed in position, maintaining the structure and shape of the metal
Delocalised electrons are mobile and able to move throughout the structure - only the electrons move
Properties of metals
Strobe metallic bonds(attraction between positive ions and delocalised electrons
High electrical conductivity
High MP/BP
Why metals have electrical conductivity
They conduct electricity in solid and liquid states
When voltage applied across a metal, delocalised electrons can move through structure carrying a charge
(In comparison, ionic compounds have no mobile charge carriers)
Why metals have high MP/BP
Strong electrostatic attraction between the cations and sea of delocalised electrons
So more energy needed to break these forces
Solubility in metals
Metals do not dissolve
Any interactions lead to a reaction, rather than dissolving
E.g. sodium and water
Properties of giant covalent structures
High MP/BP
Insoluble in most solvents
Don’t conduct electricity
High MP of giant covalent
Contain strong covalent bonds
So more energy needed to break these strong covalent bonds
Solubility of giant covalent structures
Insoluble in most solvents
Covalent bonds holding atoms together in lattice too strong to be broken by interactions with solvents
Electrical conductivity in giant covalent structures
Don’t conduct electricity
E.g. in diamonds, all four outer shell electrons involved in covalent bonding, so none are available for conducting electricity
Exceptions to conductivity in giant covalent structures
Graphene
Graphite
Both made of carbon
Features of graphene
A single layer of graphite
Shape - hexagonal planar layer
Bond angles - 120 - by electron-pair repulsion
Composed of hexagonally arranged carbon atoms linked by strong covalent bonds
Very strong material
Same electrical conductivity as copper
Structure of graphite
Bond angles - 120
Compose of layers of hexagonally arranged carbon atoms, like a stack of graphene layers
Layers bonded by weak London Forces
Bonding in the hexagonal layer only uses three of carbon’s four outer shell electrons.
Spare electron is delocalised between layers
This means electricity can be conducted as in metals as has a mobile charge carrier
Periodic trend in melting points across periods 2 and 3
Melting point increases from group 1 to group 4
Sharp decreases in melting point between group 4 and group
Melting pints are comparatively low from group 5 to group 0
What does the sharp decrease in melting points show?
Marks change from giant to simple molecular structures. Shows start of divide between metals and non-metals
Difference between giant and simple molecular structures
Giant covalent structures have strong forces to overcome so have high melting points.
Simple molecular structures have weak forces to overcome, so have much lower melting points
Where does the trend in melting points across a period continue
Across period 2 it is repeated across period 3, and continues across the s- and p- blocks from period 4 downwards
