FOUNDATIONS
IN CHEMISTRY
Atomic Structure
Atoms
Atoms a re made of protons, neutrons and electrons. These subatomic particles have different
properties.
·
Units
Li Subatomic c onventional S I useful.
particles a re so small that u n its aren't
very
↳ We define certain units for use with atoms.
↳ The atomic unit (amul is defined as twelfth (2) the of
single
mass one m a ss a
carbon-12 atom.
(i) The
elementary charge ( e)
unit is equal ttoh e charge
on an electron.
·
Protons
LD Has
charge of1.0073
a
amu.
Lix
charge of t h e
Have a
·
Neutrons
↳ Has a mass of1.0087 amu
↳N o electric
charge.
↳
Slightly heavier than a
proton, by the difference is so small we often t a ke the masses
to be the and to namu.
same
equal
·
Electrons
LD Has
charge of -he.
a
11Mass is
small it
usually approximated
so
is to ze ro
↳
Roughly 0.00055
amu.
Atoms a re made of a nucleus and shells of
e lectrons.
·
Nucleus
↳ Contains
protons and neutrons.
Li Most
ofthe mass of
a n atom is i n the nucleus.
Li
Positively charged.
·
Electron Shells
↳ Further
spliti n t o sub-shells.
↳Each sub-shell has levels.
slightly different energy
↳
Occupy most of the space in an atom.
,Isotopes
An atom can be identified byits m a ss and
proton number.
·
Mass Number
Lis The number of the number of neutrons.
protons plus
↳Often by
the symbol A.
given
·
Proton Number
4 The number of
protons in a nucleus.
↳ Often by
given the symbol.
Isotopes a re atoms with the same
proton number, but a different m a ss number. This means
that
they have a different
n umber of neutrons.
Relative Masses
Relative Atomic Mass Ar is the of element compared to twelfth the ofa
avarge m a ss an one mass
single carbon-12 atom.
↳I n name
other words, is one-twelfth of t h e m a ss of a
single carbon-12 a to m .
↳Ar is the of in
element units of
average
mass an amu.
Average Mass
·
↳Almost element has multiple stable
every isotopes meaning some will
weigh m o re than
others.
↳ The of
a n element's
isotopes weighted
totheir
abundance.
average according &
Relative Molecular Mass Mr i s the m a ss of molecule compared to one twelfth the mass of
average
a
single carbon-12 atom.
a
Li
Simply add Ar values for all the atoms i n a molecule to calculate the Mr.
Mass
Spectrometry
Once a
sample has
passed through a
spectrometer we c a n
analyse the data to
identify the
molecule.
·
Spectrum Produced
Lis The x- a x i s is
mass/charge ratio m/2.
↳The is the % abundance.
y-axis
·
Main Peak
↳ Also called the molecular ion
peak, the one with the
greatestm a ss (charge ratio.
·
Isotopes
↳ Smaller cluster around the main
peaks will peak.
↳ These from the w i td
molecules but hifferent
isotopes in them.
a re
same
Lis The molecules have different
masses and different
isotopic so
mass/charge ratio.
·
Fragmentation
↳ Smaller and
significantly lighter
peaks in the
spectrum a re because of
fragmentation.
↳ Molecules can
fragment in the spectometer.
, Once we have the m a ss
spectrum we can calculate the relative atomic mass.
[ Isotopepercentage IstoPass the
"Ar=
x
900
The abundance be quoted in
many different
can units.
Percentage
·
↳ If abundance expressed becomes
is
percentage, calculation
t he
slightly easier.
as a
Example -
Boron
·
Boron has 2
isotopes "OB and "B.
↳"B -20% and "B-80%
(20x10) 180xx)
8
x
+
=
100
IN CHEMISTRY
Atomic Structure
Atoms
Atoms a re made of protons, neutrons and electrons. These subatomic particles have different
properties.
·
Units
Li Subatomic c onventional S I useful.
particles a re so small that u n its aren't
very
↳ We define certain units for use with atoms.
↳ The atomic unit (amul is defined as twelfth (2) the of
single
mass one m a ss a
carbon-12 atom.
(i) The
elementary charge ( e)
unit is equal ttoh e charge
on an electron.
·
Protons
LD Has
charge of1.0073
a
amu.
Lix
charge of t h e
Have a
·
Neutrons
↳ Has a mass of1.0087 amu
↳N o electric
charge.
↳
Slightly heavier than a
proton, by the difference is so small we often t a ke the masses
to be the and to namu.
same
equal
·
Electrons
LD Has
charge of -he.
a
11Mass is
small it
usually approximated
so
is to ze ro
↳
Roughly 0.00055
amu.
Atoms a re made of a nucleus and shells of
e lectrons.
·
Nucleus
↳ Contains
protons and neutrons.
Li Most
ofthe mass of
a n atom is i n the nucleus.
Li
Positively charged.
·
Electron Shells
↳ Further
spliti n t o sub-shells.
↳Each sub-shell has levels.
slightly different energy
↳
Occupy most of the space in an atom.
,Isotopes
An atom can be identified byits m a ss and
proton number.
·
Mass Number
Lis The number of the number of neutrons.
protons plus
↳Often by
the symbol A.
given
·
Proton Number
4 The number of
protons in a nucleus.
↳ Often by
given the symbol.
Isotopes a re atoms with the same
proton number, but a different m a ss number. This means
that
they have a different
n umber of neutrons.
Relative Masses
Relative Atomic Mass Ar is the of element compared to twelfth the ofa
avarge m a ss an one mass
single carbon-12 atom.
↳I n name
other words, is one-twelfth of t h e m a ss of a
single carbon-12 a to m .
↳Ar is the of in
element units of
average
mass an amu.
Average Mass
·
↳Almost element has multiple stable
every isotopes meaning some will
weigh m o re than
others.
↳ The of
a n element's
isotopes weighted
totheir
abundance.
average according &
Relative Molecular Mass Mr i s the m a ss of molecule compared to one twelfth the mass of
average
a
single carbon-12 atom.
a
Li
Simply add Ar values for all the atoms i n a molecule to calculate the Mr.
Mass
Spectrometry
Once a
sample has
passed through a
spectrometer we c a n
analyse the data to
identify the
molecule.
·
Spectrum Produced
Lis The x- a x i s is
mass/charge ratio m/2.
↳The is the % abundance.
y-axis
·
Main Peak
↳ Also called the molecular ion
peak, the one with the
greatestm a ss (charge ratio.
·
Isotopes
↳ Smaller cluster around the main
peaks will peak.
↳ These from the w i td
molecules but hifferent
isotopes in them.
a re
same
Lis The molecules have different
masses and different
isotopic so
mass/charge ratio.
·
Fragmentation
↳ Smaller and
significantly lighter
peaks in the
spectrum a re because of
fragmentation.
↳ Molecules can
fragment in the spectometer.
, Once we have the m a ss
spectrum we can calculate the relative atomic mass.
[ Isotopepercentage IstoPass the
"Ar=
x
900
The abundance be quoted in
many different
can units.
Percentage
·
↳ If abundance expressed becomes
is
percentage, calculation
t he
slightly easier.
as a
Example -
Boron
·
Boron has 2
isotopes "OB and "B.
↳"B -20% and "B-80%
(20x10) 180xx)
8
x
+
=
100
