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Chemistry - Module 3 - Periodic table - Part 1 - Coggle Diagram
Chemistry - Module 3 - Periodic table - Part 1
Structure, bonding and properties
giant covalent lattices and carbon allotropes
general: giant networks of strong covalent bonds require large energy to break
diamond
very high melting point, extremely hard, good thermal conductor, electrical insulator, insoluble
silicon forms a similar crystal lattice
each carbon bonded to four carbons in tetrahedral lattice
graphite
layers held by weak forces that allow sliding, making it slippery and useful as lubricant and in pencils
high melting point, insoluble, lower density than diamond, used for lightweight strong materials
each carbon bonded to three others, delocalised electron allows electrical conductivity
graphene
delocalised electrons give excellent electrical conductivity and strengthen bonds
transparent, extremely strong, light, flexible, useful for electronics and touchscreens
single-atom-thick sheet of hexagonally bonded carbon
metallic bonding and properties of metals
metals form lattices of positive ions in a sea of delocalised electrons
electrostatic attraction between cations and electrons is metallic bonding
properties explained by metallic bonding
malleability and ductility from ability of ions to slide past each other
good thermal and electrical conductivity via delocalised electrons
melting point depends on number of delocalised electrons per atom, ion size, and lattice structure
generally insoluble except in liquid metals
simple molecular structures and london forces
small molecules with strong covalent intramolecular bonds but weak intermolecular london forces
low melting and boiling points compared to giant structures
larger molecules have stronger london forces and higher melting/boiling points
noble gases are monatomic with very weak london forces and lowest melting/boiling points in periods
bond-type changes across a period and related property trends
across a period, bonding type changes from metallic to giant covalent to simple molecular and to noble gases
metallic region: melting/boiling points increase as metallic bonds strengthen with decreasing ionic radius and more delocalised electrons
giant covalent region: very high melting/boiling points due to network covalent bonds
simple molecular region: low melting/boiling points due to weak intermolecular forces
noble gases: lowest melting/boiling points due to monatomic nature
Ionisation energies
definition and measurement notes
first ionisation energy is energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms
process is endothermic and measured for gaseous atoms
equations remove only one electron per equation and refer to one mole of atoms
factors affecting ionisation energy
nuclear charge
more protons increase attraction for electrons
atomic radius
greater distance reduces nuclear attraction sharply
shielding
inner electrons reduce attraction felt by outer electrons
trends down groups and across periods
down a group, ionisation energies generally decrease
extra electron shells increase atomic radius and shielding, making electrons easier to remove
across a period, ionisation energies generally increase
increasing proton number pulls electrons closer, reducing atomic radius
extra protons are added to same outer level so no extra shielding
exceptions to period trend
drop between groups 2 and 3
outer electron in group 3 removed from a p orbital, which is slightly higher energy and more shielded by s electrons
drop between groups 5 and 6
group 5 removal from singly occupied orbital; group 6 removal from doubly occupied orbital where electron-electron repulsion makes removal easier
successive ionisation energies and shell evidence
multiple ionisations produce successive ionisation energies that increase within a shell
increase due to removal from increasingly positive ion and reduced electron repulsion
large jumps in successive ionisation graph indicate breaking into inner shell
number of electrons removed before big jump indicates group number
spacing of points can predict electronic structure by counting electrons per shell
The Periodic table
early classification and modern arrangement
two early 1800s classification methods: by physical and chemical properties, and by relative atomic mass
at that time, protons and electrons were unknown
modern table arranged by increasing atomic number (proton number)
elements placed into periods (rows) and groups (columns)
periodicity concept
elements in same period share the same number of electron shells
elements in same group share the same number of outer-shell electrons and similar chemical properties
Periodic table structure
periods and groups
periods indicate number of electron shells
repeating trends in properties across a period is called periodicity
groups indicate number of electrons in outer shell
similar chemical behaviour within a group
blocks of the table
s-block: groups 1 and 2, outer electron in s subshell
d-block: transition elements, outer electron in d subshell
p-block: groups 3 to 0, outer electron in p subshell