Topic 1: Atomic Structure + Periodic Table Flashcards
What is meant by ‘Atomic number’
Tells us the number of protons in an atom
What is meant by ‘mass number’
Tells us the number of protons and neutrons in the nucleus
What is an isotope?
Elements with the same number of protons and electrons, but different numbers of neutrons.
What is an ion?
An atom which has either gained or lost electrons, to become charged
Define ‘Relative atomic mass’
The weighted mean mass of an element, relative to 1/12 the mass of a carbon 12 atom
Define ‘Relative isotopic mass’
The mass of an atom of an isotope compared to 1/12 the mass of a carbon 12 atom
Define ‘Relative molecular mass’
The mean mass of a molecules, compared to 1/12 the mass of a carbon 12 atom
How to calculate Relative atomic mass
(Abundance of A x m/z) x
(Abundance of B x m/z)
————————————-
Total Abundance
Define ‘First ionization energy’
The energy required to remove one electron from each atom in one mole of gaseous atoms, forming one mole of gaseous 1+ ions
Explain the trend in ionization energy along a period
Ionization energy increases along a period.
Shielding is similar
This is because the proton number increases.
So nuclear charge increases
Pulling outer electron in closer to the nucleus
So more energy is required to remove it
Explain why there is a difference in ionization energy between groups 2 - 3
Ionization energy is less for G3. This is because the outermost electron is in a higher energy subshell - further from the nucleus.
Therefore more shielding, less attraction to nucleus, and lower ionization energy
Explain why there is a difference in ionization energy between groups 5 - 6
Ionization energy is less for G6. This is because the outer electrons in the P orbital (of group 6) have a pair. The pair experience opposite spin, and repulsion, meaning less energy is required to remove an electron
Explain why ionization energy decreases down a group
The atomic radius increases as we go down the group
–> So outer electrons are further from the nucleus
Shielding increases as we go down the group
–> So more shells between the nucleus and outer shell
The attractive force between the nucleus and outer shell is weaker, so less energy is required to remove the outer electron
Successive ionization energy
The removal of more than 1 electron from the same atom
Explain the trend in Atomic radius along a period
Decreases along the period
Electrons go into the same shell - so similar shielding effect
Proton number increases by 1, nuclear charge increases
-> Pulls outer shell towards nucleus
Explain the trend in Atomic radius down a group
Increases down the group
Due to additional electron shells added.
Trend in melting points: Group 1 - 3
Melting point increases from Groups 1- 3. Metal ions have an increasing positive charge.
Increasing number of delocalised electrons.
Smaller ionic radius
Therefore a stronger metallic bond + higher melting point to overcome this
Trend in melting points: Group 4
Has the highest melting point - due to the giant covalent structure.
Many strong covalent bonds hold the atoms together.
Therefore a large amount of energy is needed to overcome this.
Trend in melting points: Group 4 - 5 - 6
Large decrease from groups 4 - 5. Group 5 elements have a simple covalent molecular structure. With weak London forces which require little energy to overcome.
Slight increase from groups 5 - 6. Group 6 elements still have a simple molecular structure, yet have more electrons, so have larger London forces.
Trend in melting points: Group 6 - 7 - 8
Decrease from group 6 - 7. Due to the smaller simple covalent molecular structure and therefore smaller London forces.
Group 8 has the lowest melting point, as it exists as a single atom, so has very small London forces.
Define an orbital
a region within an atom that can hold up to two electrons with opposite spins
