Week 2, Lesson 1
Chapter 3: The Periodic Table
Week 2
• Mendeleev developed the periodic table
based on discoveries by Boyle, Lavoisier,
Dalton and others.
• It organised gathered information based on
chemical properties and behaviour.
• It also became a helpful tool for making
predictions and guiding future developments.
• But…it did not adequately explain why
substances had certain properties and
behaved as they did.
Continued…
• The ideas from Chapter 1 and the
understanding discussed in Chapter 2 has
made the periodic table the most useful
reference available.
• It means that chemists do not have to
memorise facts about individual elements,
rather they can just identify patterns.
• It also provides a framework in which to
organise knowledge.
• Mendeleev discovered sets of elements with
similar chemical properties.
• He discovered that as he looked at the
elements in increasing atomic mass they
changed from metals, to non metals and back
to metals again.
• Similar patterns existed for other properties.
Mendeleev’s Periodic Law:
“The properties of elements vary periodically
with their atomic weights.”
• The number of protons in an atom is now
known to be the fundamental difference
between atoms of different elements, not
atomic weight.
• Because of this we know see the periodic
table organised based on ascending atomic
number.
Group 1 Elements
• Period 1 Elements are known as alkali metals.
• These elements all have similar properties and
are relatively soft metals that react highly with
water and oxygen.
• Have a look at their electron configurations.
–
–
–
–
–
Li = 1s22s1
Na = 1s22s22p63s1
K = 1s22s22p63s23p64s1
Rb = 1s22s22p63s23p63d104s24p65s1
Cs = 1s22s22p63s23p63d104s24p64d105s25p66s
• They all have similar outer-shell configurations
with their valence electrons being in the s
subshell.
Group VII or 17 Elements
• Are known as halogens.
• Again these elements have similar properties in
that they are all coloured and are also highly
reactive.
• They have the following electron configurations:
–
–
–
–
F = 1s22s22p5
Cl = 1s22s22p63s23p5
Br = 1s22s22p63s23p63d104s24p5
I = 1s22s22p63s23p63d104s24p64d105s25p5
• These elements all have the same valence
electron configuration, being s2p5.
Periodicity and Element Properties
• The arrangement of electrons in atoms is
responsible for the periodicity of element
properties.
• Properties such as melting temperature, electrical
conductivity, formulas of compounds formed
when two elements react, and many others
depend on the way the electrons are arranged,
particularly the outer shell configuration as these
are the electrons that are involved in bonding
Mendeleev’s Law Restated
“Variations of the chemical properties of
elements across a period and similarities
down a group are all associated with
electronic configurations of their atoms.”
Patterns in Electronic Structure
• Vertical columns (groups): contain elements
with similar outer-shell electron
configurations. They are numbered I-VIII or 118.
• Horizontal rows (periods): each period
contains elements with electrons in the same
outer shell. The number of the period is the
same as the number of the outer shell.
Blocks of Elements
• There are four main blocks of elements, with the
elements in each block filling the same type of subshell.
• s-block: contains elements in groups 1 and 2.
• p-block: contains elements in groups 13-18
• Although helium has an electron configuration of 1s2, it
is usually located in group 18 because its atoms are
unreactive.
• d-block: contains the transition metals.
• f-block: contains the lanthanides and actinides.
Lanthanides are a set of 14 elements with the atomic
numbers 58-71 and actinides have atomic numbers 90103.
Week 3, Lesson 1
Trends in Properties
• The periodic variation on the properties of
elements reflects the periodic variation in
their electron configurations.
• This is more clearly seen if you look at how
properties change from left to right across a
period.
• There are similarities between elements in a
group but there are also significant
differences.
Trends in Properties cont…
• Although elements in the same group have
the same number of electrons in their outer
shells, the atoms become larger as you move
down and this also effects the properties of
the elements.
• We can look at the properties of elements in
two ways:
– The properties that relate to individual atoms of
the element, ie electron configuration.
– The properties that reflect the way atoms or
groups of atoms interact with each other.
Atomic Properties
• Radius, ionisation energy and
electronegativity all depend on the strength of
the attraction between the valence electron
and the nucleus.
• This attraction will generally depend on:
– The positive charge that attracts the valence
electrons, or,
– The distance of these electrons from the nucleus.
Electrons and the Periodic Table
• The electrons of atoms of elements in the same period are
located in the same outer shell.
• As you move across a period the number of protons increases,
meaning the numbers of electrons also increases.
• The outer shell electron of lithium is attracted towards the
three protons. However it does not feel the full charge as the
two inner electrons act as a shield. The outer electron
experiences a core charge of +1.
• The outer shell electrons in fluorine experience a core charge
of +7 as the two inner electrons shield the outer electrons
from the 9 protons.
Electrons and the Periodic Table
cont…
• The outer shell of a sodium atom has a core charge of +1
because the 1 outer electron is attracted to the 11 protons
however the first two complete shells are shielding the
attraction.
• The core charge is the same moving down a group but the
number of electron shells increases. As such the valence
electrons are held less strongly.
• The electrons of elements in the same period are in the same
outer shell. As you move across the period the core charge
increases, therefore increasing the attraction towards the
nucleus.
Trends in the Periodic Table
Trend
Explanation
Atomic Radius increases down a group
Electrons occupy most of the volume of an
atom.
Atomic Radius decreases across a period
The size of atom decreases due to the
increasing positive charge of the nucleus
attracting the electrons.
First Ionisation Energy decreases down a group
The atoms get larger and as such valence
electrons are further away from the nucleus.
The energy needed to remove the outermost
electron decreases.
First Ionisation Energy increases across a group
The attraction to the nucleus is higher and as
such more energy is required to remove the
outer electrons.
Electronegativity decreases down a group
As the outer electrons become further away
from the nucleus, electrons are more weakly
attracted to an atom.
Electronegativity increases across a period
The electron-attracting ability of atoms
increases as the pull on the outer electrons
increases.
Defintions
• Atomic Radius: The size of the atom.
• First Ionisation Energy: The amount of energy
required to removed the first electron.
• Electronegativity: The ability of an atom to
attract an electron towards itself.
Metallic and Non-Metallic
Character
• Elements on the right hand side of the table
are non-metals, while the other elements are
metals.
• Moving from left to right across a period,
elements become less metallic and exhibit
more of the non-metal properties.
• There is also a variation in metallic character
within groups. Looking at group 14, Carbon is
a non-metal, where tin and lead are both
metals.
Metalloids
• Some metals like germanium, silicon, arsenic
and tellurium which are located across periods
display both metallic and non-metallic
properties.
• These are called metalloids.
Chemical Reactivity of Elements
Reactivity of Metals:
- The way in which metals react with water can
give us an indication of their relative reactivity.
- Those elements in group one are more
reactive than those in group 2.
- As you move down a group the reactivity of
the metal increases.
- Generally: Reactivity increases down and
group and decreases across a period.
Chemical Reactivity cont…
Reactivity of Non-Metals
- We can get an indication of the reactivity of
non-metals through reacting them with
aqueous solutions, for example potassium
iodide.
- This shows that the reactivity decreases down
a group. This is because the electronegativity,
or the ability to attract electrons is higher at
the top of a group.
Noble Gases
• Noble gases are all unreactive .
• These elements have very low melting and
boiling temperatures and all are gases at room
temperature.
• Their lack of reactivity is due to the electron
configuration. Each noble gas has an outer
shell that is considered ‘full’ or stable.
• As such they do not want to react with other
elements and unbalance their stability.
Week 3, Lesson 2
Compounds
• Both compounds and elements are considered
to be pure substances.
• Pure Substance: matter that always has the
exact same composition.
• Compounds are formed when atoms of two or
more elements chemically combine in fixed
proportions.
• Each compound has its own set of properties
and these can be quite different from the
elements that it is made up of.
Compounds cont…
• If we look at salt as an example…
– Salt, sodium chloride is made up of sodium and
chlorine.
– On the next slide is a table highlighting the
significant different characteristics of sodium,
chlorine and sodium chloride.
Sodium (Metal)
Chlorine (Non Metal)
Sodium Chloride
Melts at 98 degrees
Melts at -101 degrees
Melts at 801 degrees
Conducts electricity when
solid
Does not conduct
electricity when solid
Does not conduct
electricity when solid
Conducts electricity when
molten
Does not conduct
electricity when molten
Conducts electricity when
molten
Molecule
• A molecule is two or more non metal atoms
chemically combined, for example water.
• Both hydrogen and oxygen are both gases
under normal conditions, yet water is a liquid.
• Most compounds that are produced from two
chemically combined non-metals generally
have similar chemical properties to water.