The Periodic table trends Flashcards
What do elements in a period show?
A trend in physical and chemical properties
Why do elements in a period show a trend in physical and chemical properties?
They have different electron configurations
What is a group?
A vertical column in the periodic table
What is a period?
A horizontal row in the periodic table
Do elements in the same group have the same or different number of electrons in their outer shell?
Same
Do elements in the same group have the same or different types of orbital in the outer shell
Same
Do elements in the same group have outer electrons in the same or different quantum shell?
Different
How is the periodic table arranged?
In order of increasing atomic number
Do elements in the same period have the same or different numbers of electrons in their outershells?
Different
Do elements in the same period have the same or different types of orbitals in their outer shells?
Different
Define periodicity
A repeating trend in properties of the elements across each period of the periodic table
What do elements in the same group show?
Similar physical and chemical properties
Why do elements in the same group have similar chemical and physical properties?
They have a similar electronic configuration
Define 1st ionisation energy
The amount of energy required to remove one electron from EACH ATOM in one mole of gaseous atoms to form one mole of gaseous 1+ ions
What is the general trend in 1st ionisation energy across a period?
It increases
What is the general trend in 1st ionisation energy down a group?
It decreases
What is the trend observed between group 0 1st of one period and the group 1 element of the next period?
There is a sharp drop
Why is there a sharp drop in 1st IE observed between group 0 of one period and group 1 element of the next?
- electrons are being removed from a new electron shell
- this shell is further from the nucleus and experiences more shielding despite the increase in nuclear charge
- the overall nuclear attraction is weaker
- therefore less energy is required to remove the outer electron
What 4 factors affect the value of ionisation energy?
- number of electron shells
- atomic radius
- number of protons
- nuclear attraction
Explain the general trend in 1st ionisation energy across a period
- 1st ionisation energy increases
- nuclear charge increases across a period
- shielding does not change across a period
- atomic radius gets smaller across a period
- therefore attraction between the nucleus and outer electron becomes greater
Explain the general trend in 1st ionisation energy down a group
- 1st IE decreases down a group
- shielding increases down a group
- atomic radius increases down the group
- although nuclear charge increases, the other factors outweigh this
- there is a decrease in nuclear attraction
Some points do not obey the general trend in 1st ionisation energy- where are these?
- between groups 2 and 3
- between groups 5 and 6
Explain the difference in 1st ionisation energy between Be and B
- the fall in IE marks the start of filling the p subshell
- in Be the outer electron is in a 2s orbital
- in B the outer electrons is in a 2p orbital
- the 2p orbital in B is of higher energy and experiences more shielding from inner electrons as it is also shielded by the electrons in the 2s subshell
- despite the fact the nuclear charge increasing by +1 from Be to B the difference in shielding and energy means the nuclear attraction is weaker in B
- less energy required to remove the electron
Explain difference in 1st ionisation energy between N and O
- the fall in IE between groups 5 and 6 marks the start of electron pairing in the p-orbitals in p-subshell
- In N the outer electron is unpaired in a 2p orbital
- In O the outer electron is paired in a 2p orbital
- The paired electrons in O repel slightly
- This makes it easier to remove an electron from O than N
What is the trend in melting points across periods 2 and 3?
- Groups 5-8 have much lower melting points
- Group 4 elements occupy the peaks of melting points
- melting points increase from group 1 to group 2 to group 3
Why do group 4 elements occupy the peaks of melting points?
- Carbon and Silicon have giant covalent structures
- each atom makes 4 covalent bonds
- all of these bonds must be broken which requires a lot of energy
- whereas Boron only makes 3 covalent bonds
Why do elements with diatomic molecules or single atoms have the lowest melting points?
- simple covalent structure
- only intermolecular forces must be overcome, no bonds
- the strength of London forces is related to the no. electrons that can take part in instantaneous induced dipole-induced dipole interactions
- smaller the molecule the smaller the London forces the lower melting point
Why do melting points increase from group 1 to group 2 to group 3?
- metallic bonding present
- metallic bonding increases as the number of electrons present that can be delocalised increases and as nuclear charge increases
- The charge on the metal ion increases from group 1 (1+) to group 2 (2+) to group 3 (3+)
- therefore there are more delocalised electrons in the structure
- therefore the strength of electrostatic attraction between positive ions and delocalised electrons increases along the group
- More energy is needed to overcome these attractions
What is the trend in electrical conductivity across period 3?
Electrical conductivity increases from sodium to magnesium to aluminium
Why does electrical conductivity increase from sodium to magnesium to aluminium?
- the charge on ions increases
- so the number of delocalised electrons within the metallic bond increases
- delocalised electrons act as mobile charge carriers so conduct electricty
- more delocalised electrons means greater electrical conductivity
Why do elements with molecular structures have poor electrical conductivity?
-because they do not contain any mobile charge carriers
What is the trend in atomic radius across periods?
-atomic radius decreases
Why does atomic radius decrease across periods?
- across a period there is little change in shielding
- there is a regular increase in nuclear charge
- the nuclear attraction increases
- electrons are pulled more strongly towards the nucleus
What the trend in atomic radius down the group?
It increases
Why does atomic radius increase down the group?
- there are increasing numbers of electron shells
- therefore shielding increases
- the attraction of electrons by the nucleus is less
