ORBITAL - can hold up to 2 electrons,
each sub-shell has different orbitals
s,p,d,f
s orbital has a spherical shape
P orbital has a 3D
dumbbell shape at
right angles to one
another
orbits are regions
around the nucleus that
have electron density
electron
repel each
other and
have a
property
called a
spin
in orbitals they
have opposite
spins to
counteract the
repulsion
carbon- 12= international standard
the mass of carbon-12 is exactly 12
the mass of 1/12 of an atom of carbon 12 is exactly 1
RELATIVE ISOTOPIC MASS-
the mass of an atom of an
isotope compared to 1/12 of the
mass of an atom of carbon-12
RELATIVE ATOMIC MASS-
is the weighted average
mass of an atom of an
element compared with 1/12
of the mass of am atom of
carbon-12
RELATIVE MOLECULAR
MASS- the weighted
average mass of an
molecule of a compound
compared with 1/12 of the
mass of an atom of
carbon-12
ISOTOPES-atoms of the
same element with
different masses with
the same number of
protons and neutrons
RELATIVE FORMULA
MASS- the average mass of
the formula unit of a
compound compared with
1/12 the mass of an atom of
carbon-12
IONISATION ENERGY
the first ionisation energy is
the energy required to
remove 1 electron from each
atom in 1 mole of gaseous
atoms to form 1 mole of a
gaseous 1+ ion
1st ionisation energy == Na(g)-->Na+(g) + e-
2nd i.e==Na+(g)-->Na2+(g) + e-
factors affecting ionisation energy
Atomic radius-the greater the distance
between the nucleus and the outer
electron, the less the attractive force
Nuclear charge - the greater the
number of protons in the nucleus the
greater the attractive force
Electron shielding- the outer electron shells
are repelled by any inner shells between
electrons and the nucleus
the second ionisation energy
is the energy required to
remove 1 electron fro each
atom in 1 mole of a gaseous
1+ ion to form a mole of
gaseous 2+ ions
trends in 1st I.E-- There is a
increase in 1st I.E across a period
because the there is a increase in
nuclear charge as electrons are
added ti the same shell across a
period.
there is a sharp decrease in 1st I.E
between the end of one period and
the start of the next period, this is
because there is a new outer shell
which is due to distance and shielding
A decrease in 1st I.E down a
group due to the presence of
extra shells down the group
redox
oxidation- the loss of electrons
reduction is the gain of electrons
half equations-
Mg + Cl2 -->MgCl2
electron transfer
Mg-->Mg2+ + e-
Cl2--> 2e- + 2Cl-
Mg is the reducing
agent, it has reduced
Cl2
Cl2 is the oxidising
agent , it has oxidised
Mg to Mg 2+
oxidation agent accepts electrons from other reactants
non-metals are oxidising agents==
F2,Cl2,O2
A reducing agent donates
electrons to another reagent
metals are reducing
agents==Na, Fe, Zn
Acids bases and Salts
All acids release H+ ions (aq)
All bases accept H+ ions (aq)
An alkali is a soluble
base that dissolves in
water, releasing OH-
ions (aq)
SALT- a compound formed from an acid,
when H+ ion from the acid has been
replaced by a metal ion or a another
positive ion, such as the ammonium ion
(NH3+)
water of crystallisation
hydrated- containing water molecules
anhydrous- without water
the molar proportion of water in hydrated crystals
Bonding
structure and
the periodic
table
ionic
electrostatic attraction
between oppositely
charged ions
ionic lattices
structure compromising of hundreds of
thousands of ions
an ion is surrounded by oppositely
charged ions, forming a giant ionic
lattice
properties of ionic compounds
high melting and boiling point due to lot of energy needed to
break the strong electrostatic forces holding the ions rigidly in the
solid lattice
electrical conductivity- (solid lattice)- ions are in a fixed position and there are no MOBILE
charge carriers so ionic compounds are a non conductor of electricity in the solid state.
When melted or dissolved in water the solid lattice breaks down and the ions are free to
move as mobile charge carriers therefore ionic compounds are electricity conductors in a
liquid or aqueous state
An ionic lattice often dissolves in polar solvents (e.g) water
covalent
a shared pair of electrons
between atoms of non
metals
a molecule is the smallest
part of a covalent compound
that can take part in a
chemical equation
DATIVE COVALENT
BOND FORMS WHEN
THE SHARED PAIR OF
ELECTRONS COMES
FROM ONE BONDED
ATOM
e.g- and ammonium ion
NH4+ (containing3
covalent bonds and one
dative bond
an electron that takes part
in forming a chemical bonds
is called a valence electron
simple covalent
small molecules with weak intermolecular forces
low melting and boiling
points due to less energy
required to break the
weak intermolecular
forces
no mobile charged
particles so the
structures are
non-conductors of
electricity
van der waals
forces form between
a simple molecular
structure and a
non-polar solvent
which weakens the
structure simple
molecular structures
are often soluble in
non polar solvents
giant covalent
e.g- diamond, graphite and SiO2
high melting and boiling point due to
lots of energy required to break the
strong covalent bonds in the lattice
except fro graphite there are no mobile charged
particles so they are non conductors
the strong covalent bonds in a lattice are
too strong to be broken by polar or
non-polar solvents
metallic
the electrostatic attraction
between the positive
metal ions and
delocalised electrons
the delocalised
electrons in metals are
able to move
throughout the
structure can cannot
assign to s positive
metal ion
shapes of molecules
the shape of the molecule
depends on the number of
electron pairs surrounding the
central atom
electron pairs repel
each other and move
as far apart as
possible
molecules with bonded pairs
2- linear
3- trigonal planar
4- tetrahedral
6- ocrahedral
SF6
90
CH4
109.5
BF3
120
BeCl2
180
molecules with lone pairs
0- tetrahedral
1- pyramidal
2- non-linear
H20
104.5
NH3
107
CH4
109.5
electronegativity, polsrity and polarisation
ELECTRONEGATIVITY- the
measure of the attraction of an
atom in a molecule for the pair of
electrons in a covalent bond
generally the smallest atoms are
the most electronegative and the
most electronegative atoms are
those of highly reactive non-metallic
elements- e.g.. O, Fe and Cl
Polar and non-polar molecules
polar bonds
the bonded electrons
are shared equally
between both atoms,
and the bonded atoms
have similar
electonegativities
non-polar bonds
the bonded electrons are
shared equally between both atoms and the
bonded atoms have similar
electronegativities
symmetrical and
unsymmetrical molecules
each bond is polar but
the dipoles act in
different directions, the
overall effect of the
dipoles to cancel out
Intermolecular forces
Van der waals forces
weak intermolecular forces that act between molecules
caused by the uneven distribution of electron in
molecules which inures a dipole in an neighbouring
molecule which is induces more dipoles
the greater the number of electrons in
each molecule the larger the oscillating
and induced dipoles. the greater the
attractive forces between molecules
and the greater the van der waals
forces
van der waals
forces increase
in strength with
increasing
number of
electrons
Permanent dipole-dipole
the delta positive/negative
charges on a polar molecule
attract oppositely charged dipoles
on another polar molecule
this gives a weak intermolecular
force called a permanent
dipole-dipole interaction- eg- HCL
stronger than van der waals
forces although they are both
weak
Hydrogen bonds
strong intermolecular bond between polar molecules
special properties of water due to hydrogen bonds are -that ice is
less dense than liquid water this is because particles in solids are
usually packed closely together than in liquids however hydrogen
bonds hold water molecules apart in an open lattice structure -ice
has a high melting point and water has a high billing point
because of the strong hydrogen bonds between water molecules