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The structure of the periodic table
Periodicity is the repeating pattern in chemical and physical properties
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The periods show repeating trends in chemical and physical properties
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g
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Groups have similar chemical properties
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g
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Periodic trend in electronic configuration and ionisation energy
Similarities in electron configuration reflect similarities of chemical reactions
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Be: Be → Be²⁺ + e¯
[He] 2s² → [Ne]
Mg: Mg → Mg²⁺ + e¯
[Ne] 3s² → [Ar]
Ca: Ca → Ca²⁺ + e¯
[Ar] 4s² → [Kr]
Sr: Sr → Sr²⁺ + e¯
[Kr] 5s² → [Xe]
Electron configuration can be based on the previous noble gas, so only outer-shell electrons, responsible for reactions, can
be shown
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The first ionisation energy increases overall across period 2 and 3 because
the number of protons increase so there is greater nuclear charge
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They have the same number of inner
shells so the outermost electrons experience the same shielding
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The first ionisation energy decreases as you move from period 2 to 3
because the atomic radius increases as there is one more inner shell and
the electron shielding increases as the number of outer electrons increase
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There is a small decrease in first ionisation energy between Be and B because the outer electrons move from an s-orbital
to a p-orbital
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There is another small decrease in first ionisation energy between N and O due to p-orbital electron repulsion
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The first ionisation energy decreases down a group because
the number of inner shells increases, so the atomic radius
increases and the electron shielding increases
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This means there is less nuclear attraction and it is easier to
remove an electron from the nucleus
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Successive ionisation energy allows us to
predict the number of electrons in each
shell of an atom and the group of an
element as big jumps in ionisation energy
happen when an electron is being removed
from a shell that is closer to the nucleus
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The first
electrons are removed from the outer shell,
working towards the inner shells
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The number of electrons counted
equals the group
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Metals have a giant metallic lattice structure held together by metallic bonds or forces
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They have high melting and boiling points because they have delocalised electrons which can move through the structure
and positive ions (cations) which remain where they are
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They are insoluble in water, except liquid metals, because of the strong electrostatic attraction between the cations and
delocalised electrons which can’t be interfered by the hydrogen bonds
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They are good thermal conductors because the delocalised electrons can move freely through the metal lattice and carry
kinetic energy throughout the metal lattice
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Diamond and graphite form gaseous atoms of carbon when they are ionised
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It’s a poor electrical
conductor because there are no delocalised electrons as all the outer-shell electrons are used for covalent bonds
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They are insoluble in water
because of the strong covalent bonds
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They are good electrical conductors because it has delocalised
electrons between the layers which are mobile
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It’s soft and slippery because there are weak van der Waals’ forces between the
layers allowing it to slide
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It’s very thin, so it is transparent
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It has a low density but is very strong, making it useful in high-speed
electronic and aircraft technology
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It doesn’t
conduct electricity because it doesn’t have any free electrons and it is hard due to the strong covalent bonds
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The high melting and boiling
points are due to the strong metallic bonds between the metals which get stronger because the metal ions have a greater
charge, an increasing number of delocalised electrons and a decreasing ionic radius
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The fourth element has a giant covalent structure which is held together by
covalent bonds
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Element five onwards are simple molecules
which have strong covalent bonds but have weak induced dipole-dipole forces which are easily overcome
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