Topic 1: Atomic Structure And The Periodic Table Flashcards
Define isotope
Elements with the same number of protons but different number of neutrons
Define relative atomic mass
The weighted mean mass of an atom of an element, compared to 1/12th mass of an atom of carbon 12
Define relative isotopic mass
The mass of an atom of an isotope, compared to 1/12th of the mass of an atom of carbon 12
What must your abundance add up to?
100%
Formula for relative atomic mass
(Abundance A x m/z of A) + (Abundance B x m/z of B) all over total abundance
How do you calculate isotopic mass?
Do relative atomic mass and then solve for x
Why are arrows in electronic figuration in different directions?
Spin pairing. When 2 electrons occupy one orbital, they ‘spin’ in opposite directions
Why do you fill orbitals singly first then pair them up?
Electron repulsion
Why would a transition metal with an electronic configuration of 3d5 4s1 form a unipositive cation with an electronic configuration of 3d5 4s1 and not 3d4 and 4s2?
Lose from the 4s orbital first then from 3d as 4s has higher energy when filled but lower energy when empty.
What is n=1 in the atomic emission spectra?
Ground state
Why are line spectra used?
To identify elements and it is evidence for quantum shells
What is a series in atomic emission spectra?
A group of lines
Why do lines in atomic emission spectra get closer together?
The energy and frequency increases
What does the line spectrum show?
The frequency of light in coloured bands
Where must an electron fall to for a line to be produced at uv?
Ground state, n=1
Where must an electron fall for a line to appear at visible light in atomic emission spectra
N=2
Where will the line in atomic emission spectra appear if n=3?
Infrared
Define first ionisation energy
The minimum amount of energy required to remove 1mol of electrons from 1mol of atoms in the gaseous state to form one mole of unipositive gaseous cations
Ionisation energy trends down a group
Decreases down groups. Atomic radius increases. Outer electrons are further from nucleus, so there is a weaker attractive force, therefore less energy is required. Shielding increases so less energy is required to remove electrons.
What is successive ionisation
The removal of more than 1 electron from the same atom
What is the general trend of successive ionisation
There is a general increases in energy as removing electrons causes the atom to become increasingly more positive
What happens to the atomic radius across period 3?
It decreases. There is an increased nuclear charge because there is an increase in protons
What happens to the atomic radius down a group?
Increases due to extra electron shells being added
What is the trend of ionisation energies across periods
Increases due to increasing number of protons and nuclear attraction increases. Shielding is similar and distance marginally decreases. More energy is required to remove outer electrons.
Why is there a decrease at the 3rd element in the ionisation energies across periods
The element sits in a higher energy subshell (p instead of s) so it is slightly further from the nucleus and there is some shielding from s, the p orbitals are higher in energy so less enegry needed to remove an electron
Why is there a decrease at the 6th element in ionisation energies across periods
Electronic configuration p3 to p4 which means the outer electron is paired in its orbital. This means there is more repulsion between the electrons and so they are higher in energy and therefore require less energy to remove.
What is the general trend of melting points across period 3 metals
Increases as metal ions have a positive charge, increasing the number of delocalised electrons and a smaller ionic radius
Why does silicon have the highest melting point in period 3
Due to its giant covalent structure
Why is argon the lowest melting point in period 3
Argon only exists as atoms
