Module 3: Chapter 7 (Periodicity) Flashcards
How is the periodic table arranged?
- by increasing atomic (proton) number
- in periods showing repeating trends in physical and chemical properties (periodicity) -> number of period gives the highest energy electron shell
- in groups having similar chemical properties -> same number of electrons in the outer shell
Describe the development of the periodic table?
- Dimitri Mendeleev published his periodic table in 1865 with 60 elements
- he arranged them in order of atomic mass and lined up elements in groups with similar properties
- swapped elements around if they didn’t fit properties and left gaps for elements that weren’t discovered yet
- today the period table has 118 elements arranged by increasing atomic number
What is periodicity?
A repeating pattern across periods is called periodicity
Define first ionisation energy?
The energy to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions
How does atomic radius affect ionisation enthalpy?
The greater the distance between the nucleus and the outer electron the weaker the attraction
How does nuclear charge affect ionisation enthalpy?
the more protons in the nucleus the stronger the attraction
What is electron shielding?
the shielding effect is the repulsion between the inner and outer shell electrons as electrons are negatively charged
How does electron shielding affect ionisation enthalpy?
shielding increases when the number of inner shells increases, reduces the attraction between the nucleus and outer electrons
Describe the trend in successive ionisation energies
As each electron is removed the remaining electrons
are pulled closer to the nucleus - nuclear attraction
increases
Describe the periodicity in first ionisation energy
increases as you move across the group
Why does the first ionisation energy decrease as you go down the group?
This is because the atomic radius increases and shielding increases, therefore weaker attraction between outer electron and nucleus.
Why is there a general increase in first ionisation energy across a period?
- Electrons are in the same shell so shielding is similar.
- Nuclear charge increases (more protons in nucleus)
- Nuclear attraction decreases (between nucleus and outer electrons) so atomic radius decreases
- First ionisation energy increases as attraction between outer electrons and nucleus increases
Explain why the first ionisation energy of boron is less than of beryllium?
In B, electron is removed from a 2p orbital rather than
2s orbital in Be. The 2p sub-shell is at higher energy
and its electron is easier to remove
Why is there a drop in ionisation energy between nitrogen and oxygen?
In O, one of the 2p orbitals contains paired electrons
whereas in N, all three orbitals are singly occupied
The paired electrons in O repel and electron is easier to
remove
Describe the structure for metallic bonding
In a SOLID…
• a regular repeating pattern of fixed cations (maintain shape)
• delocalised electrons spread out in the whole structure (each atom donated one outer shell electron), act as mobile charge carriers
Describe the bonding of metallic bonds
Electrostatic attraction between cations and delocalised electrons
What are the properties of metals?
- conduct electricity in a solid or liquid state as there as mobile charge carriers
- high melting and boiling points because of the strong electrostatic attraction
- insoluble in all solvents
Describe the properties of giant covalent structures
- high melting and boiling points as there are very strong covalent bonds
- insoluble in almost all solvents, as they’re held in a lattice and bonds are too strong to be broken down by interactions with solvents
- non-conductors of electricity except for graphene and graphite as there are delocalised electrons between layers as carbon only bonds 3 times
What is the structure of giant covalent structures?
- billions of atoms held together by strong covalent bonds, forms a giant covalent lattice
What is structure and bonding of simple molecular structures?
- strong covalent bonds between atoms within the molecules with only weak attractive forces between molecules
- for elements the molecules will always be non-polar so the intermolecular forces are London Forces
What are the properties of simple molecular structures?
- low melting and boiling points (not as much energy is required to overcome the weak intermolecular forces)
- the more electrons in the molecule the stronger the London forces, and the higher the melting and boiling points
Describe the giant covalent structure diamond
- Each carbon forms 4 single covalent bonds to another C atom (tetrahedral shape)
- a non-conductor of electricity
Describe the giant covalent structure graphene
- one layer of graphite
- each carbon only uses 3 of its outer shell electrons to form a planar hexagonal layer
- remaining electrons are delocalised in the layer, making them excellent conductor
Describe the giant covalent structure graphite
layers of graphite stacked together, attracted by weak forces
State and explain the differences of melting points between solid iodine and graphite? (2 marks)
- Graphite has a giant covalent lattice and iodine has a
simple molecular lattice - On melting, covalent bonds are broken in graphite
- London forces are broken between iodine molecules
- Covalent bonds are much stronger
- Covalent bonds require more energy input to break and
graphite has a much higher melting point
State and explain the difference in electrical conductivity between solid iodine and solid graphite? (2 marks)
- Graphite contains delocalised electrons between its
layers - The delocalised electrons can move allowing graphite to conduct
- Iodine has no mobile charge carriers and cannot
conduct
Explain why there is a drop in first ionisation energies in group 0?
Down a group, electrons are added to a new shell,
further from the nucleus - atomic radius increases
There are more inner shells between the outer electrons
and the nucleus, increasing the shielding
Attraction between nucleus and outer electrons
decreases
