Physical Properties
Syllabus statements
• 3.2.1 Define the terms first ionization energy, and
electronegativity.
• 3.2.2 Describe and explain the trends in atomic radii, ionic
radii, first ionization energies, and electronegativities and
melting points for the alkali metals, and the halogens.
• 3.2.3 Describe and explain the trends in atomic radii, ionic
radii, first ionization energies, and electronegativities for
the elements across period 3.
•
• 3.2.4 Compare the relative electronegativity values of two
or more elements based on their positions on the
periodic table.
• The first ionisation energy of an atom tells us how
difficult it is to remove an electron from that
atom.
• We only talk about first ionisation energy when
we make a positive ion – this will be important to
remember when we study energetics!
• A definition (learn it!)
• The first ionisation energy of an element is the
energy required to remove one mole of electrons
from one mole of atoms in the gaseous state.
• Every part of the definition is important!
• First ionization energies are measured in
kJ mol-1
• You don’t need to remember the numbers –
they are given in the data booklet.
• You DO need to remember the trends.
Trends in Properties
• We will look at the chemical properties later;
first we need to think about some of the
physical properties:
Atomic radius
• The atomic radius is the distance from the
nucleus of an atom to its outermost electron.
• BUT
• Electrons don’t stand still (and that’s without
even mentioning the Heisenberg Uncertainty
Principle!)
• So atomic radius is sometimes defined as half
the distance between neighbouring nuclei
For group 1 elements:
Element
Lithium Li
Sodium Na
Potassium K
Rubidium Rb
Cesium Cs
Electron
arrangement
2,1
2,8,1
2,8,8,1
2,8,8, . . . ,1
2,8,8 . . ., . . ., 1
Atomic radius
(10-12 m)
152
186
231
244
262
• The numbers are in the data booklet – you
don’t need to learn them.
• The atomic radius increases down a group as
the number of occupied electron shells
increases.
• The data booklet doesn’t give atomic radii for
the noble gases.
• Why?
• All the group 1 elements lose a single electron
when they are ionized.
• As we would expect, positive ions are smaller
than the atoms from which they were formed.
• Positive ions have fewer occupied electron
shells, and greater electrostatic attraction
than the atoms from which they were formed.
• There is a similar trend in atomic radii going
down group VII.
Element
F
Cl
Br
I
Atomic Radius
10-12m
64
99
114
133
• Fluorine has more electrons than Li.
• So why does it have a smaller atomic radius?
• We will answer that question in a minute!
• All the group VII elements gain one electron
when they form ions.
• The ions have more electrons than the atoms
from which they were formed.
• Hence negative ions are larger than the atoms
from which they were formed.
• Cations contain _______ electrons than
protons, so they are ______ than their parent
atoms;
• Anions contain _______ electrons than
protons, so they are ______ than their parent
atoms.
• Cations contain _fewer_ electrons than
protons, so they are smaller than their parent
atoms;
• Anions contain __more_ electrons than
protons, so they are bigger than their parent
atoms.
• Now consider period 3
• This is the row which contains:
• Na Mg Al Si P S Cl Ar
• How does atomic radius change between Na
and Cl?
• It gets smaller, even though Cl has more
electrons.
• Why?
• All the elements in period 3 contain electrons in
the 3rd shell.
• I.e. they all have the same outer shell.
• But, as we go across period 3 each element has
one more proton than the previous element.
• This extra positive charge pulls the outer shell of
electrons slightly.
• As we go across a period, atomic radius decreases
because of increased electrostatic attraction
between the nucleus and the outermost
electrons.
• The situation is a little harder for ionic radii.
• We need to consider positive ions (cations)
and negative ions (anions) separately.
• All the ions across a period contain the same
number of electrons
• They are “isoelectronic” with a noble gas
• For cations:
• As we go across a period, we add protons.
• The increased nuclear charge pulls the outer
electrons in closer.
• Ionic size decreases across a period.
• When we start forming negative ions (anions)
• There is suddenly a large increase in size.
• This is because electrons have been added to
the outer shell, resulting in increased
repulsion with the nucleus.
• BUT as we continue across the group we add
further protons
• (ie the number of protons gets closer to the
number of electrons)
• Hence the anions get smaller as we add more
protons.