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Summary Chemistry OCR A A level chapter 7-8
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Summary of chapter 7-8 of Module 3: Periodic Table and Energy of Chemistry OCR A A level. Great for Revision
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Module 3: Periodic table and energyChapter 7: Periodicity
7.1 The periodic table
The periodic table - then
Over 60 elements arranged by Mendeleev in order of atomic mass
Lined up elements in groups with similar properties
If properties did not fit, Mendeleev swapped elements and left gap
o Assuming atomic mass measurements were incorrect / elements yet to
discovered
Even predicted properties of the missing elements from group trends
The periodic table - now
As of 2014 PT has 114 elements arranged in 7 horizontal periods and 8 vertical groups
Most important organisation tool
First point of reference for chemists
Helpful to remember where common elements are positioned, their atomic number and
relative atomic masses
Arranging the elements
Position of elements linked to physical and chemical properties
Atomic number
o From left to right, elements are arranged in order of increasing atomic number
Groups
o Elements are arranged in vertical columns called groups
o Each element in a group has atoms with same no. of outer shell electrons and
similar properties
Periods and periodicity
o Elements are arranged in horizontal rows called periods
o No. of period gives the no. of the highest energy electron shell in an element's
atoms.
o Across periods, there is a repeating trend in properties - periodicity
o Most obvious periodicity in properties - trend from metals to nonmetals
o Periodicity of several properties:
Electron configuration
Ionisation energy
Structure
Melting point
Period trend in electron configurations
Chemistry of each element is determined by its electron configuration, particularly outer,
highest energy electron shell
Trend across a period
Each period starts with an electron in a new highest energy shell
o Across period 2, 2s subshell fills with 2 electrons, followed by the 2p subshell
with 6 electrons
o Across period 3, the same pattern of filling is repeated for the 3s and 3p subshell
o Across period 4, although the 3d subshell is involved, the highest shell number is
n = 4. From the n = 4, only the 4s and 4p subshell are occupied
o For each period, s- and p-subshells are filled in the same way
Trend down a group
Elements in each group have atoms with the same no. of electrons in their outer shell
, Elements in each group also have atoms with the same no. of electron in each subshells
o Element in the same group will have similar chemistry
Blocks
Elements can be divided into blockers corresponding to their hughes energy subshell
4 distinct blocks, s, p, d and f
Old numbers are the number used in IGCSE, group 1-7 and then 0
o This system is based on s- and p-blocks
, o Advantage of old numbering is that the group number matches the no of
electrons in the highest energy electron shell
New system run from 1-18, numbering each column in the s-, d-, and p-blocks
o New numbers were approved for use by IUPAC in 1988
o Periodic table use both number, with old numbers in bracket
7.2 Ionisation energy
Ionisation energy - measures how easily an atom loses electrons to form positive ions
o Attraction between nucleus and outer electrons
First ionisation energy - energy required to remove one electron from each atom in one
mole of gaseous atom of an element to form one mole of gaseous 1+ ions
o Example: Na(g) → Na+(g) + e- first ionisation energy = +496 kJmol-1
Factors affecting ionisation energy
Electrons are held in their shells from attraction of the nucleus
First electron lost will be in the highest energy level - least attraction
3 factors affect ionisation energy
o Atomic radius
Greater the distance the less nuclear attraction
Force of attraction falls off with inc distance, atomic radius - huge effect
o Nuclear charge
The more proton in the nucleus, the greater the attraction between the
nucleus and outer electrons
o Electron shielding
Electron are negatively charged - inner shell electron repel outer shell
This repulsion - shielding effect - reduced attraction
Successive ionisation energies
Elements have as many ionisation energies as there are electrons
Example: Helium has 2 electrons - 2 ionisation energies
o He(g) → He+(g) + e- first ionisation energy
o He+(g) → He2+(g) + e- second ionisation energy
Second ionisation energy of helium is larger than first
o After first is lost, single electron is pulled closer to nucleus
o Nuclear attraction on the remaining electron increases
o More ionisation energy is needed to remove the second
Successive ionisation energies = first ionisation energy
Second ionisation energy - energy required to remove one electron from each ion in one
mole of gaseous 1+ ions of an element to from one mole of gaseous 2+ ions
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