A
Patterns Found in the Elements
Elements Known by the Ancients
The ancient civilizations knew of the following elements:
Au, Ag, Cu, Fe, Pb, Sn, Hg, S and C.
(gold, silver, copper, iron, lead, tin, mercury, sulfur and
carbon)
Elements Discovered by Alchemists in the Middle Ages
Alchemists experimented to turn base materials into precious
materials like gold or a potion for eternal life. As they did this
they discovered these new elements: As, Sb, Bi, P, Zn
(arsenic, antimony, bismuth, phosphorus, zinc).
Discovery of Electric Current
Alessandro Volta in 1800 discovered that electric current
could be made by using two different metals in a salt, acid or
base solution. He also discovered that the current could be
increased by hooking up several cells in a row (in series).
The combination of many cells in series was called a battery.
Batteries Used to Decompose Compounds into Elements
It was found that many pure substances (like water) could be
broken apart into its elements by using the electric current
that a battery provided. William Nicholson and Anthony
Carlisle in 1800 decomposed water into oxygen and hydrogen
gas.
Electric Current Used to Find New Elements
Humphry Davy and other experimenters (1807) used
electricity to break apart many substances into elements that
mankind had not seen until then (K, Na, Ca, Mg, B, Ba) .
Molten table salt (NaCl at 801o C)) was split apart into a
silvery, explosive metal (sodium) and an irritating yellowish
gas (chlorine).
Elements Discovered From 1765 to 1845
Electrical current enabled scientists to discover 42 new
elements from 1765 to 1845.
How to Classify the Elements?
Scientists began looking for patterns among the elements.
Since the atomic weights of these elements was known by
1845, chemists began to list the elements in order by weight
and then look for similarities or repetitions of characteristics.
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
The Octave of Elements
H
He
Li
Be
B
C
N
O
F
Ne
Na
Mg
In 1865, John Newlands published his “Law of Octaves”
which stated that every 8th element repeats the properties of
the element that was 8 before it, like a piano repeats a similar
tone every 8th key (an octave). It soon was found out that
this pattern just works for elements 1-20 but not beyond this.
Dmitri Mendeleev
In 1869, Dmitri Mendeleev published a chart of the elements
in which he noted recurring properties among the elements.
He formulated The Periodic Law which stated that if elements
are arranged in order by their weights (today by at. number),
then similar properties will recur periodically.
The
Periodic
Law:
Repeating
Properties
Among the
Elements
Mendeleev’s Chart Showing Patterns
As Mendeleev listed his
elements in order by
weight, he placed them
alongside other elements
with similar chemical and
physical properties to
show that they belonged
together in some way. He
referred to elements that
had similar properties as a
“family” of elements.
Gaps in Mendeleev’s Chart
In some places in his chart, Mendeleev found that the next
heaviest element had properties that fit with a group of elements
beyond the next family. When this happened, Mendeleev left a
space, placing this element with the family it belonged with.
Mendeleev predicted that an element would be discovered that
would have a weight and properties that would place it in the
space he had left.
Mendeleev’s Method of Predicting New Elements
Mendeleev predicted that elements to be discovered would have
chemical properties like the family they would be in. He predicted
the atomic masses, densities and other physical properties of the
elements to be discovered by averaging these from the elements
on either side of the blank space and from the elements above and
below the blank space. For example, Mendeleev had a blank space
after Ga because As had properties like P. To predict the atomic
weight of the element ?, Mendeleev averaged the weight top and
bottom of ? with the weights to the left and right of ? to get the
weight of 72.86 which was very close to 72.59, the weight of ? or
Ge.
Ga
Si
P
28.0855
30.973
?
As
69.72
74.92
Sn
118.71
The Success of Mendeleev’s Predictions
Mendeleev’s predictions of several elements to yet be discovered
were quickly confirmed by other chemists and his predictions
about their properties proved very accurate. This made chemists
realize that the idea of chemical families and Mendeleev’s periodic
chart demonstrated some underlying structure among atoms still
to be discovered.
The Reason for Chemical Families Revealed
In 1913, Niels Bohr presented his model of the atom which
helped explain why elements repeat their properties as they
are ordered from lightest to heaviest. Bohr’s model showed
that elements in a chemical family all have the same number
of outermost electrons.
The Quantum Atomic Model Also Explains Periodicity
Elements in the same chemical family have the same electron
configuration, just at increasing energy levels.
The Periodic Chart Reflects Quantum Atom Structure
The ordering of the elements follows the Aufbau Principle of
the Quantum Model of the Atom.
The Modern Periodic Chart
The modern Periodic Chart of the Elements arranges
elements into vertical columns and horizontal rows. Vertical
columns are called chemical families and elements in vertical
columns have similar chemical and physical properties.
The Alkali Metal Family
The alkali metal elements are silvery, have a low density (float
on water), react very strongly with water to form hydrogen
gas and very strong bases (alkalis).
The Alkaline Earth Family
The alkaline earth metal family is less reactive than the alkali
metals but alkaline earth metals burn brightly in air, some
very colourful (strontium is used in red fireworks). Alkaline
earth elements also react with water and are very common
elements in earth rocks (calcium and magnesium especially).
The Halogen Family
The halogens are very reactive and corrosive, especially with
metals (Na + Cl  NaCl {table salt}). At room temperature,
fluorine and chlorine are gases while bromine is a liquid and
iodine is a solid. Chlorine, bromine and iodine are very good
anti-bacterial agents.
The Noble Gas Family
The noble gases are non-reactive and stable by themselves.
They are all colourless and odourless gases. They glow with
distinctive colours (Ne is reddish) when electricity is passed
through them. Their ion charges of zero indicate that they do
not form charged ions.
A Major Pattern Within The Elements
The elements can be classified as either metals or nonmetals. Metals
are shiny, malleable, ductile, good conductors of heat/electricity, and
form + ions. Nonmetals are dull, brittle/crumbly, poor conductors, and
form – ions. Metalloids or semi-metals have characteristics inbetween
metals and nonmetals.
Some Periodic Properties of Atoms
Atomic and Ionic size, Ionization energies and Electron
Affinities are properties which can be predicted using the
Quantum Model of the atom.
Atomic and Ionic Size
An important atomic property is its size relative to other
atoms. An atom’s size changes when it becomes an ion.
Ionization Energy
Ionization energy is the amount of energy needed to remove
the outermost electron from an atom of a gaseous element.
Electron Affinity
Electron Affinity is the amount of energy released when an
electron is added to a neutral atom in the gaseous state.
Factors Used to Predict Atomic Properties
1. The pel (principal energy level) of the outermost electrons.
This factor relates to the distance that the outermost
electrons are from the nucleus. It is the primary factor for
determining many atomic properties. Check the highest pel
first!
Factors Used to Predict Atomic Properties
2. The size of the nuclear charge is the second most
important factor in determining atomic properties. A greater
nuclear charge (+) pulls electrons (-) closer because of the
attraction of opposite charges. Electrons are held more
tightly (attracted more strongly) by the nucleus when they are
closer to the nucleus.
Factors Used to Predict Atomic Properties
2. The size of the nuclear charge is the second most
important factor in determining atomic properties. A greater
nuclear charge (+) pulls electrons (-) closer because of the
attraction of opposite charges. Electrons are held more
tightly (attracted more strongly) by the nucleus when they are
closer to the nucleus.
Factors Used to Predict Atomic Properties
3. The screening effect of inner electrons between outer
electrons and the nucleus is the third most important factor
(after nuclear charge and pel) in determining atomic
properties. The layers of inner electrons affect the attraction
force that the nucleus has for its outer electrons by
weakening or diluting the effective nuclear attraction as these
layers increase.
Factors Used to Predict Atomic Properties
4. The repulsive effect of electrons within the same orbital is
the fourth factor affecting atomic properties. The repulsive
effect only becomes significant in the rare occasions when,
for 2 situations, the effects of pel, nuclear charge and
screening are generally equivalent. Repulsive effects tend to
move electrons slightly apart from each other due to their
identical charge.
Using the 4 factors to Predict Size
Within the alkali metal family, the size of atoms increases
going from Li to Na to K because in this progression, a new
pel is being added with each successive element. A new pel
being filled means successive outer electrons are farther
from the nucleus.
Using the 4 Factors to Predict Sizes
For the atoms, Mg and Cl, their outer electrons are on the
same pel but the nuclear charge of Cl is greater. Thus an
atom of Cl will be smaller than an atom of Mg due to its
stronger attraction for electrons on the same pel.
Atomic Size Trends in the Periodic Chart
Atoms increase in size from top to bottom due to the addition
of new pels. Atoms decrease in size from left to right due to
increasing nuclear charge, increasing effective nuclear
charge on outer electrons.
Predicting Ion Sizes Relative to Atom Sizes
How would the sizes of the following species compare : 9F,
+
2+ ?
9F , 10Ne, 11Na, 11Na , 12Mg, 12Mg
Write the electron configuration for each species:
: 1s22s22p5
9F
: 1s22s22p6
9F
2
2
6
10Ne : 1s 2s 2p
2
2
6
1
11Na : 1s 2s 2p 3s
+
2
2
6
11Na : 1s 2s 2p
11Na
12Mg
2
2
6
2
12Mg : 1s 2s 2p 3s
2+
2
2
6
12Mg : 1s 2s 2p
Based on the above, neutral sodium and magnesium atoms
would be the largest since they have a 3rd pel (1st factor). A
magnesium atom would be smaller than a sodium atom
because it has more nuclear charge (2nd factor).
Predicting Ion Sizes Relative to Atom Sizes (Cont.)
The remaining species all have the same outer pel. Thus the
second factor (nuclear charge) is checked to determine size.
22s22p5
F
:
1s
9
: 1s22s22p6
9F
2
2
6
10Ne : 1s 2s 2p
+
2
2
6
11Na : 1s 2s 2p
2+
2
2
6
12Mg : 1s 2s 2p
9F
9F
The smallest nuclear charge (the smallest or weakest
effective attraction) is found in 9F and 9F-, making these the
largest. Of these two, both with same pel and nuclear
charge, 9F- has one more electron which causes it to be the
larger of the two due to factor 4, the repulsion effect.
Predicting Ion Sizes Relative to Atom Sizes (Cont.)
In the remaining species which have the same outer pel, the
second factor, nuclear charge, is checked for differences.
22s22p6
Ne
:
1s
10
+
2
2
6
11Na : 1s 2s 2p
2+
2
2
6
12Mg : 1s 2s 2p
10Ne
+
11Na
2+
12Mg
The smallest to largest nuclear charge, the weakest to
strongest effective electron attraction, hence the largest to
smallest sized, is 10Ne, 11Na+, 12Mg2+ .
Relative Sizes Compared
22s22p5
F
:
1s
9
: 1s22s22p6
9F
2
2
6
10Ne : 1s 2s 2p
22s22p63s1
Na
:
1s
11
+
2
2
6
11Na : 1s 2s 2p
2
2
6
2
12 Mg : 1s 2s 2p 3s
2+ : 1s22s22p6 A
Mg
12
11Na
12 Mg
9F
9F
10Ne
+
11Na
2+
12Mg
Periodic Ionization Energy Patterns
Ionization energy (The energy needed to remove the
outermost electron) increases across a period due to
increasing nuclear charge which attracts electrons more
strongly, making it more difficult to remove them. Having
fewer pels means outer electrons are closer to the nucleus,
making it more difficult to remove them.
Explanation for Ionization Energy Patterns
Within each period, the noble gases have the largest nuclear
charge so they attract their outer electrons most strongly so
that their ionization energies are at the top in the graph
above. In contrast to this, in each period, the alkali metals
have the lowest nuclear charge which means they have the
least attraction for their outer electrons and are at the bottom
of the graph
Ionization Energy Patterns Reflect Quantum Atom
There is an ionization energy drop from Be to B despite the
fact that B has a greater nuclear charge and same pel. This is
explained by noting that in the quantum atom, B’s outermost
electron is in the second minor energy level (p) which makes
it slightly farther from the nucleus and thus easier to remove.
Ionization Energy Patterns Reflect Quantum Atom
Oxygen has a lower ionization energy than nitrogen despite
the fact that oxygen has greater nuclear charge and the same
pel. This is explained by considering that oxygen’s
outermost electron is the first paired p electron. Being in the
same orbital as another electron, this outermost electron
experiences repulsion force and thus is easier to remove.
Electron Affinity Trends
Electron Affinity is the energy an atom gives off when an
electron is added to it. A large electron affinity (attraction for)
means that an atom has a strong “desire” to get an electron.
The nonmetals tend to have higher electron affinities because
they have higher nuclear charge on a pel and getting electrons
gives them electron configurations like the noble gases.
Summary of Atomic Property Trends
End of Presentation