Chapter 7 - Periodicity Flashcards
Define periodicity
• Across periods, repeating trend in properties of elements
How did Mendeleev order the periodic table
arranged them in order of atomic mass
• Lined up elements in groups of similar properties
• He swapped some around or left gaps – assuming some atomic masses were wrong and some undiscovered
• Predicted properties of missing elements from group trends
How are elements arranged now
In increasing atomic number
If an element is in the P block what does that mean
The elements highest energy subshell is a P shell
What is the trend across periods
o Across period 2, the 2s sub-shell fills with two electrons, followed by the 2p sub-shell with six electrons.
o Across period 3, the same pattern of filling is repeated for the 3s and 3p sub-shells,
Define first ionisation energy
energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.
What is the first ionisation energy of sodium
Na (g) → Na+ (g) + e-
What are the factors effecting ionisation energy
Atomic radius
Nuclear charge
Electron shielding
How does atomic radius effect ionisation energy
o electrons in shells that are further away from the nucleus are less attracted to the nucleus
- so the further the outer electron shell is from the nucleus, the lower the ionisation energy
How does nuclear charge effect ionisation energy
nuclear charge increases with increasing atomic number,
which means that there are greater attractive forces between the nucleus and outer electrons,
more energy is required to overcome these attractive forces when removing an electron
How does electron shielding effect ionisation energy
electrons in full inner shells repel electrons in outer shells preventing them to feel the full nuclear charge
so the greater the shielding of outer electrons by inner electron shells, the lower the ionisation energy
Trend in ionisation energy down a group
Decrease
Why does ionisation energy decrease down a group
o The atomic radius increases
o The shielding (by inner shell electrons) increases
o Therefore, the attraction between the nucleus and the outer electrons decreases
But nuclear charge does increase down a group?
Effect is outweighed by increase radius
How does ionisation energy change across a period
Increase
Why does ionisation energy increase across a period
o Across a period, the nuclear charge increases
o The distance between the nucleus and outer electron remains reasonably constant
o The shielding by inner shell electrons remains the same
Graph of first ionisation energy against atomic number
Why does boron have a lower ionisation energy than beryllium
the fifth electron in boron is in the 2p subshell, which is further away from the nucleus than the 2s subshell of beryllium
Why does oxygen have a lower ionisation energy than nitrogen
the paired electrons in the 2p subshell of oxygen repel each other, making it easier to remove an electron in oxygen than nitrogen.
Define second ionisation energy
energy required to remove one electron from each ion in on mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions.
Do successive ionisation energies increase or decrease
Increase
Why do successive ionisation energies increase
as removing an electron from a positive ion is more difficult than from a neutral atom
• As more electrons are removed the attractive forces increase due to decreasing shielding and an increase in the proton to electron ratio
What do jumps in successive ionisation energy suggest
A change in shell
How are metals arranged
tightly packed together in lattice structures
Bonding in metals
strong electrostatic forces of attraction between the positive metal centres and the ‘sea’ of delocalised electrons
Define covalent bonds
bonds between nonmetals where there is a shared pair of electrons between the atoms
Structure of all metals
Giant metallic lattice
Three forms of carbon
Diamond
Graphite
Graphene
How many carbons are bonded in diamond
4
Arrangement of carbon atoms + bond angle in diamond
Tetrahedral + 109.5
Structure of diamond
Giant covalent lattice
How many carbons bonded in graphite
3
Structure of graphite
Giant covalent lattice
Bond angle in graphite
120
Bonding in graphite
• All atoms in same layer held together by strong covalent bonds
• Layers held together by weak intermolecular forces = allows to slide
Conductivity of graphite
Yes = delocalised electrons
Structure of graphene
Giant covalent lattice
How many carbon atoms are bonded in Graphene
3
Graphene = layered?
No one layer of carbon atoms
Structure of silicon
networks of atoms bonded by strong covalent bonds
Same structure as diamonds
Properties of metallic substances
• Due to delocalised sea of electrons…
• They have high M.P/B.P = lot of energy required to overcome strong electrostatic forces of attractions between positive ions and delocalised electrons
• Solubility = metals don’t dissolve
• Electrical conductivity = conduct electricity in solid + liquid state
Do metals dissolve
No
Properties of giant covalent substances
• High M.P / B.P = large number of covalent bonds linking the whole structure = lot of energy required to overcome
• Can be hard or soft
o Graphite = soft = intermolecular forces between layers are weak
o Diamon + SiO = hard = difficult to break 3D network of covalent bonds
o Graphene = strong + flexible + transparent = useful
• Most compounds insoluble with water
• Diamond and silicon = no conducting electricity = 4 carbons
• Graphene + graphite = yes conducting = 3 carbons
Does melting point increase or decrease across period 2 /3
Both = increases till group 4 then decreases
Why does melting point increase across period 2/ 3
o Groups 1 to 3 (13) have metallic bonding which increases in strength due to increased forces of attraction between more electrons in the outer shell that are released to the sea of electrons and a smaller positive ion
o Group 4 (14) has a giant covalent structure with many strong covalent bonds requiring a lot of energy to overcome
Why does the melting point decrease from group 4 to 0
o Groups 5 (15) to 0 (18) have simple molecular structures with weak London forces between molecules requiring little energy to overcome
