Summary Chemistry OCR A A level chapter 7-8

Module 3: Periodic table and energy
Chapter 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