The Periodic Table Flashcards
What is the first ionisation energy?
The energy required to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous (1+) ions
What factors affect ionisation energy?
~Atomic radius- the greater the radius, the smaller the attraction of the nucleus on the outer electrons
~Nuclear charge- the greater the charge, the greater the attraction on electrons
~Electron shielding (on screening)- inner shells of electrons repel outer electrons. The more inner shells the greater the shielding effect and the smaller the attraction on outer electrons
What are the trends of ionisation energy down a group:
~the nuclear charge increases as there are more protons
~the number of shells increases as there are more electrons
~the atomic radii increases as there are more protons
~shielding increases as there are more shells
~the nuclear attraction on electrons decreases as there are more shells
~ionisation energy decreases because of all these factors
What trends are there with ionisation energy as you go across a period:
~the shells stay the same
~the shielding stays the same as there are the same amount of shells
~the nuclear charge increases as there are more protons
~nuclear attraction on electrons increases as there are more protons and still the same shells
~atomic radii decreases as there is a greater attraction between protons and electrons so the nucleus is tighter decreasing in size
~ionisation energy increases because of these factors
What is successive ionisation energy?
Energy require to remove each electron in turn from an ion in 1 mole of gaseous ions to form 1 mole of gaseous ions
Why does the second ionisation energy require more energy than the first?
Because when it loses electrons it becomes an ion so it requires more energy to get rid of the electron
Electron shells:
~A shell contains sub-shells
~Sub-shells contain orbitals
~Electrons occupy orbitals
~Each orbital can hold 2 electrons
~Electrons in orbitals are represented by boxes
~Electrons have a ‘spin’ which is represented by an arrow
~Electrons in orbitals have opposite spin ‘up and down’
~Electrons don’t like pairing in orbitals so they avoid it by filling up the lowest available energy levels then the next available energy level is filled
~Electrons occupy sub-shells in order of increasing energy sales
~The 4s orbital fills before the 3D orbital
What are the different types of orbital?
~s
~p
~d
~f
Information about the sub-shells:
~Sub-shell-S
number of orbitals in sub shell:1
max number of electrons in sub shell:1x2=2
~Sub-shell-P
number of orbitals in sub shell:3
maximum number of electrons in sub shell:3x2=6
~Sub-shell-D
number of orbitals in sub shell:5
maximum number of electrons in sub shell:5x2=10
How can you simplify notation of electron configuration?
~Use the noble gas configuration ~e.g. 1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 4s^1 Simplify using Ar= 1s^2, 2s^2, 2p^6, 3s^2, 3p^6 To [Ar]4s^1 ~F[He]2s^2, 2p^5 ~Ca[Ar]4s^2 ~As[Ar]4s^2, 3d^10, 4p^3
What is the order of filling orbitals?
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
Why is each successive ionisation energy higher than the one before?
~As each electron is removed there is less repulsion between the remaining electrons and each shell will be drawn in slightly close to the nucleus
~The positive nuclear charge will outweigh the negative charge every time an electron is removed
~As the distance of each electron from the nucleus decreases slightly the nuclear attraction increases. More energy is needed to remove each successive electron.
Why does the ionisation energy increase across a period?
~Number of protons in the nucleus increases, so there is higher attraction on the electrons
~Electrons are added to the same shell so the outer shell is drawn inwards slightly
~There is the same number of inner shells, so electron shielding will hardly change
Why is there a decreases in atomic radium across a period?
The increased nuclear charges pulls the electrons in towards it
What is bond enthalpy?
The amount of energy required to break a bond
Why is the ionisation energy high for noble gases?
The atoms have a full outer shell of electrons and a high positive attraction from the nucleus so the ionisation energy values are large
What is the trend of ionisation values across a period?
There is a general increase across each period because:
•the number of protons in the nucleus increases, so there is a higher nuclear charge therefore a higher attraction on the electrons
•There is the same amount of shells so the atomic radii is smaller as there is a higher nuclear attraction on the electrons as the charge increase meaning the outer shell is drawn inwards slightly
•Same amount of inner shells so shielding stays the same
These factors mean more energy is needed to remove an electron so the first ionisation energy increases
Why is there a decrease in ionisation energy between group 2 and 13 elements?
As group 13 elements have the outermost electron in a p-orbital whereas group 2 elements have theirs in an s-orbit also they are marginally further away from the nucleus so the electrons in these orbitals are slightly easier to remove so the elements have lower ionisation energies
Why is there a decrease in ionisation energy from group 15 and 16?
As you move from group 13 towards group 18 outer electrons are found in p-orbitals. In groups 13,14 and 15 each of the p-orbitals contains only a single electron. In group 16 however the outermost electron is now spin paired in the p-orbitals. Electrons that are spin paired experience some repulsion making the first outer electron slightly easier to remove so a slightly lower first ionisation energy is observed
