Transcript 1-4 ppt - BSCS

1.2.1
Evidence for shells
12-Apr-16
Lesson objectives
• Define first ionisation energy and successive
ionisation energy.
• Explain the factors that influence ionisation energies.
• Predict the number of electrons in each shell as well
as the element’s group, using successive ionisation
energies.
Evidence for shells
Using your knowledge from GCSE draw the electronic
structures for the following atoms:
Calcium
Chlorine
Aluminium
Sodium
Ions
• Draw electronic configurations for the
ions
Ca2+
Al3+
Cl-
O2-
Forming Ions
The first ionisation energy (1st I.E.) of an element is
the amount of energy required to remove one electron
from each atom in a mole of gaseous atoms to form one
mole of gaseous 1+ ions.
X (g)
Example
X+(g)
Na (g)
+ eNa+(g)
+ e1st I.E. = + 496 kJ mol-1
Don't forget state symbols
This is what happens inside a
plasma TV screen
• The screen consists of 100’s of 1000’s of tiny
cells which each contain a mixture or neon and
xenon between two plates of glass.
• The gas is electrically ionised into a mixture
of +ve ions and –ve electrons.
Xe (g)
Xe+(g)
+ e-
• The formation of these ions required energy .
The mixture is called a plasma which causes
the screen to emit light
Forming Ions
Consider the atomic
structure and discuss in
pairs what factors you
think will affect ionisation
energies.
Ionisation energy is affected by:
1 - Atomic Radius
Attraction falls off very
rapidly with distance. An
electron close to the nucleus
will be much more strongly
attracted than one further
away.
2 - Nuclear Charge
Hydrogen
1st I.E = 1310 KJ mol-1
Helium
1st I.E = 2370 KJ mol-1
Why is the 1st IE greater for helium than for hydrogen?
The greater the nuclear charge, the greater the attraction of the
nucleus for the outer shell electron, therefore it requires more
energy to remove it.
3 - Electron Shielding
Sodium
1st I.E = 494 KJ mol-1
Potassium
1st I.E = 418 KJ mol-1
There is an increase in nuclear charge from Na to K. However, for
potassium there is an additional shell of electrons which shield the
outer electron from the attraction of the nucleus. So the attraction of
the nucleus for the outer shell electron is less, so is held less strongly.
This is called electron shielding.
Key definition
• Electron shielding is the repulsion
between electrons in different inner
shells. Shielding reduces the net
attractive force from the positive
nucleus on the outer-shell electrons.
True or false?
• I played rugby at University and in my
final year was awarded half colours
(Palatinate – Durham)
• One of my hobbies is breeding pythons.
Currently I have 4 rock pythons, 10 eggs
incubating which are due to hatch at the
end of September
Successive Ionisation Energies
This is a measure of the energy required to remove
each electron in turn.
Example,
Li (g)
Li (g)
+
e-
1st I.E. = + 520 kJ mol-1
Li + (g)
Li
(g)
+
e-
2nd I.E = +7298 kJ mol-1
Li 2+ (g)
Li 3+ (g)
+
e-
3rd I.E. = +11815 kJ mol-1
+
2+
Why does it take more each energy for each successive
ionisation?
Three successive ionisation energies of
lithium
Successive ionisation energies of nitrogen
1. Write an equation to represent the 4th ionisation energy of
chlorine
2. The successive ionisation energies for carbon are shown below.
Draw its electronic configuration and use this information to
explain any trends shown.
2nd = 235
6th = 4730
3. The graph below shows the
successive ionisation energies
of sodium, explain the trends
shown.
4. How does the graph
confirms the suggested simple
electronic configuration for
sodium of (2,8,1)
3rd = 462
4th = 622
(to the nearest 10 kJ mol-1)
5.5
LOG ionisation energy (kJ/mol)
1st = 109
5th = 3780
5
4.5
4
3.5
3
2.5
2
0
2
4
6
number of electrons removed
8
10
Lesson Objectives
• State the number of electrons that can fill the first
four shells of an atom.
• Define an orbital.
• Describe the shapes of s- and p-orbitals.
At GCSE : electrons in shells
Lowest shell holds ____ electrons
Higher levels hold ____ electrons
Fill from the centre outwards
At A level . . .
• Electrons travel far from the nucleus, but there are
main areas where they are commonly found. These are
the principal shells and have distinct energy levels
(represented by rings at GCSE). The innermost ring is
level 1 and contains 2 electrons. The next level
contains 8 electrons and so on. The levels are given
numbers 1, 2, 3 etc which are also known as quantum
numbers.
Energy levels or shells
• Electrons are constantly moving, and it is
impossible to know the exact position of an
electron at any given time. However,
measurements of the density of electrons as
they move around the nucleus show us there
are areas where it is highly probable to find
an electron. These regions of high probability
are called ORBITALS. Each orbital can hold 2
electrons.
• There are 4 different types of orbitals -s, p,
d and f and they all have a different shape!!
Key definition
• ORBITAL – an atomic orbital is a
region within an atom that can hold 2
electrons, with opposite spins
s-orbitals
• An s-orbital is spherical in shape
• Each shell has an s orbital
• This gives a total of 1 x 2 = 2s electrons in each shell
p-orbitals
• From n=2 upwards (ie.
the second shell), each
shell contains 3 porbitals, px, py, and pz
• This gives a total of 3
x 2 = 6p electrons
d-orbitals and f-orbitals
• These are more complex . . .
• From n=3 upwards each shell has five dorbitals. This gives 5 x 2 = 10 d
electrons
• From n=4 upwards each shell has seven
f-orbitals. This gives 7 x 2 = 14 f
electrons
Let’s revisit our lesson objectives
• State the number of electrons that can fill the first
four shells of an atom.
• Define an orbital.
• Describe the shapes of s- and p-orbitals.
With orbitals having different shapes, chemists
often represent an orbital as a box. We call
this - Electrons in boxes
Electrons within an orbital will repel each other. An
electron has a property called spin and each electron in
an orbital will always have opposite spins.
Lesson objectives
•
State the number of orbitals making up s, p and d sub-shells.
•
State the number of electrons that occupy s, p and d sub-shells.
•
Describe the relative energies of s-, p- and d-orbitals for shells 1, 2
and 3.
•
Deduce the electron configurations of atoms for the first two periods.
Shells, sub-shells and energy levels
Filling the 2p-orbital
• Describe the relative energies of s-, p- and d-orbitals
for the shells 1, 2, 3 and of the 4s- and 4p-orbitals.
• Deduce the electron configuration of atoms and ions
up to Z = 36.
• Classify the elements into s-, p- and d-blocks.
Overlap of 4s- and 3d sub-shells
Filling the orbitals in a potassium atom
The Periodic Table, sub-shells and blocks