Transcript File - Science at St. Dominics

Electronic Structure of Atoms
atomic radii,
ionisation energies
and trends within groups
The size of an atom
• It is difficult to measure the size of an atom directly.
• The exact position of the furthest electron from the
nucleus cannot be found!
However.. Scientists can
measure the length of a
covalent bond between two
atoms using various methods
The distance is .074 nm
Atomic radii
The atomic radius is half the distance between two atoms of
the same element joined together by a single covalent bond.
It only works if it is the same
atoms.. And a single covalent bond
The atomic radii of the elements
A pattern in atomic radii
Descrease going across a period
Ìncrease going down a group
3
11
19
Potassium
As you go down a
group…
1. the atoms have extra electrons which
have to go into a new shell further away
from the nucleus.
2. the atoms have a bigger nuclear charge.
You might expect that this would pull the
electrons closer to the nucleus but the
electrons in the extra shells are being
shielded from the extra nuclear charge
by the inner shells of electrons so this
doesn’t happen!
“The shielding effect”
The atomic radii increases..
3
Lithium
4
Beryllium
5
6
Aluminium
Carbon
As you go across a period..
1. the nuclear charge in the atom increases. The increasing
nuclear charge pulls the electrons closer to the nucleus
2. the atoms have extra electrons but they do not go into new
shells! There is therefore no extra shielding effect as you go
across a period!
The atomic radii decreases..
Higher level only
Ionisation energy
Ionisation energy is the minimum energy needed to
remove the most loosely bound electron from a neutral
gaseous atom.
Check the ionisation energy
• What is the pattern
• 1) as you go down a group
• 2) as you go across a period
Higher level only
Ionisation energy – the pattern
Decreases as you go across any period
Increases as
you go down
any group
Higher level only
Ionisation energy – explaining the trend
Increases as
you go down
any group
3
11
19
As you go down any
group…
1. The atomic radii increase so the
furthermost electron gets further away
from the nucleus
2. The atoms have a bigger nuclear
charge. But the furthermost electron is
shielded from the extra nuclear charge by
the inner shells of electrons
“The shielding effect”
The ionisation energy decreases.. It
gets easier to pull the electron away
from the nucleus
Ionisation energy – explaining the trend
Higher level only
Decreases as you go across any period
3
Lithium
4
Beryllium
5
6
Aluminium
Carbon
As you go across a period..
1. the nuclear charge in the atom increases. The increasing
nuclear charge pulls the furthermost electron tighter to the
nucleus
2. The atomic radii decrease so the furthermost electron is closer
to the nucleus
The ionisation energy increases.. It gets harder to pull the
electron away from the nucleus
Plot of atomic number against
ionisation energy
The second period
Higher level only
Higher level only
Ionisation energies across the second period
Generally they go up - but there are exceptions…
Beryllium and Nitrogen have higher ionisation energies
than you would expect.
Higher level only
The electronic configuration of:
Lithium
It has three electrons.
The electronic configuration
is 1s2 2s1
Higher level only
The electronic configuration of:
Beryllium
It has 4 electrons.
The electronic configuration:
1s2 2s2
It has a full 2s sublevel
and this gives it extra
stability!
It would be extra hard to
pull the electron away so
the ionisation energy is
especially high.
Higher level only
The electronic configuration of:
Boron
It has 5 electrons.
The electronic configuration:
1s2 2s2 2p1
Higher level only
The electronic configuration of:
Carbon
It has 6 electrons.
The electronic configuration:
1s2 2s2 2p2
Higher level only
The electronic configuration of:
Nitrogen
It has 7 electrons.
The electronic configuration:
1s2 2s2 2p3
It has a half full 2p
sublevel and this gives it
extra stability!
It would be extra hard to
pull the electron away so
the ionisation energy is
especially high.
Higher level only
The electronic configuration of:
Oxygen
It has 8 electrons.
The electronic configuration:
1s2 2s2 2p4
Plot of atomic number against
ionisation energy
The third period
Higher level only
Higher level only
Ionisation energies across the third period
Generally they go up - but there are exceptions…
Magnesium and Potassium have higher ionisation
energies than you would expect.
Higher level only
The electronic configuration of:
Magnesium
It has 12 electrons.
The electronic configuration:
1s2 2s2 2p6 3s2
It has a full 3s sublevel
and this gives it extra
stability!
It would be extra hard to
pull the electron away so
the ionisation energy is
especially high.
Higher level only
The electronic configuration of:
Potassium
It has 15 electrons.
The electronic configuration:
1s2 2s2 2p6 3s2 3p3
It has a half full 3p
sublevel and this gives it
extra stability!
It would be extra hard to
pull the electron away so
the ionisation energy is
especially high.
Check your learning..
•
•
•
•
•
What is an energy level?
What is an orbital?
What is an energy sublevel?
Define atomic radius
Define ionisation energy
Today’s objectives
• Second and successive ionisation
energies and evidence for energy levels
• The chemical properties of the elements
based on atomic radii and ionisation
energies
Higher level only
Second ionisation energies
• Def the energy needed to remove the second electron from
a singly charged positive ion.
First ionisation energy
needed for:
Second ionisation energy
needed for:
Be
Be+
Be+ + e
Be2+ +e
Higher level only
Beryllium: Successive ionisation energies
+
4
Beryllium
(neutral atom)
4
Beryllium+
(positive ion)
The second ionisation energy is always higher than the first because in
an ion the atomic radius decreases and the electrons are closer to
the nucleus – harder to take away!
Higher level only
Beryllium: Successive ionisation energies
Beryllium
There is a dramatic increase in ionisation energy
needed to remove the third electron!
Higher level only
Beryllium: Successive ionisation energies
There is a large
ionisation energy needed
because the electron is
removed from a new
energy level!
The electronic configuration:
1s2 2s2
Higher level only
Ionisation values provide evidence for
the existence of energy levels:
• The ionisation energy required to remove an electron from
a new shell jumps dramatically! This is because it is much
harder to remove an electron from an inner shell as it is
closer to the nucleus and has less shielding from the
nuclear charge than before.
Chemical Properties of the elements
The chemical properties of each element depends on its
electronic structure.
Elements in the same group in the Periodic Table have
similar electronic structures – they all have the same number
of electrons in their outermost shell - and so have similar
chemical properties.
Periodic table
Group 1- The Alkali metals
Why do they always react to
lose an electron?
Can you explain why
reactivity increases going
down the group?
They all have one electron in their outermost
shell, which they have a tendency to lose.
Potassium
3
Group 1 – the Alkali metals
• Going down the group the atomic
radius increases
11
• Going down the group the
nuclear charge also increases,
but screening also increases so
this effect is cancelled out.
19
Potassium
• The result is that as you go
down the group the outermost
electron becomes easier to
remove, and the element is
therefore more reactive!
Alkali metals reacting with water
Sodium
Lithium
Potassium
Periodic table
Group 17- The Halogens
9
17
24
Why do these elements always gain an electron?
Can you explain why reactivity decreases as you go down
this group?
They all have 7 electrons in their outermost shell and
have a tendency to gain one electron in chemical
reactions
9
Group 17- The Halogens
• A high nuclear charge and small atomic
radius attracts electrons to the halogens
17
• Going down the group the atomic radius
increases
• Going down the group the nuclear
charge also increases, but screening
also increases so this effect is cancelled
out.
24
• The result is that as you go down the
group the ability to gain an electron
decreases… and the element is
therefore less reactive!