Download the modern periodic law

HISTORY OF THE PERIODIC TABLE
Early in the 19th century, scientists began to seek ways to classify the elements.
1) Johann Döbereiner (German chemist):
- in 1817, he noticed similarities among certain elements
- similar elements seemed to appear in groups of three: [S, Se, Te] [Cl, Br, I] [Li, Na, K]
- he called these groups triads
2) John A. R. Newlands (English chemist):
- in 1863, he arranged the known elements in order of increasing atomic mass
- he noticed that similar chemical and physical properties occurred every eight element: ex: atoms 2, 9 and 16
(Li, Na and K) resembled each other chemically
- he classified the 49 known elements in seven groups of seven elements
- his classification was known as law of octaves
- one flaw in his scheme: his proposed table had no gaps and therefore left no room for the discovery of new
elements
3) Dmitri Mendeleev (Russian chemist) and Lothar Meyer (German chemist):
- they both discovered the periodic law in 1869, although Mendeleev is given priority since he published his ideas
first
- periodic law: when elements are arranged in order of increasing atomic mass, elements with similar properties
recur at regular intervals
- the periodic law clearly proved that the classification of elements can be done in a logical and natural way
- it established a relation between the periodicity of elements and the regular increase of their atomic masses and
numbers
- Mendeleev's table allowed gaps for the discovery of new elements, and their properties
- approximately 65 elements were known at the time, and Mendeleev had left empty spaces to respect similarities
within the columns
- Mendeleev predicted with remarkable accuracy the properties and atomic masses of elements that were not
discovered at the time
- he called an undiscovered element ekasilicium, to be placed below silicium and above tin; Mendeleev was able
to predict its properties with accuracy; this element was discovered by Clemens Alexander Winkler in 1886 and
was named germanium (Ge)
- Mendeleev's table has stood the test of time with few alterations brought by the discovery of the first noble gas,
argon, by Lord Rayleigh and Sir William Ramsay, in 1894
- see Mendeleev's periodic table on the transparency
THE MODERN PERIODIC LAW
1- The development of the modern periodic table:
1) A. van den Broek (Dutch physicist):
- with the discovery of the noble gases, a new group had to be added to Mendeleev’s table
- in 1911, after the discovery of the electron and proton, and the development of the nuclear atom, A. van den
Broek suggested that several minor inconsistencies would be removed if the table were arranged according to
atomic number instead of atomic mass
2) Henry Moseley (English physicist):
- in 1913, while studying X-rays produced by elements, he showed that the nucleus of each element had a positive
charge equal to a whole number, which happened to be its atomic number
- Moseley’s experiment confirmed van den Broek’s suggested classification: Mendeleev’s exceptions disappeared
when elements were placed in order of increasing atomic numbers
- his work confirmed conclusively that the properties of elements depended on their atomic number rather than
their atomic mass
Statement of the modern periodic law:
When the elements are arranged in order of increasing atomic number, elements with similar properties occur at
regular intervals.
The physical and chemical properties of the elements are a periodic function of the atomic numbers.
2- Features of the modern periodic table:
1) Groups:
- 18 vertical columns containing elements with similar chemical and physical properties, also called families
- groups 1 and 2, on the left-hand side of the periodic table, and Groups 13 to 17, on the right-hand side, together
constitute the main-group elements
- groups 3 to 12 are called the transition metals (their similarities are found within their horizontal periods rather
than in their groups)
- group 1: alkali metals
- group 2: alkaline earth metals
- group 13: earth metals
- group 17: halogens
- group 18: rare gases / noble gases / inert gases
2) Periods:
- 7 horizontal rows
- 1st period contains 2 elements: hydrogen and helium
- in 1940, the 7th period had 6 elements
- the two bottom rows represent two series and are called the lanthanide series and the actinide series
3) Metals:
- more than ¾ of the elements known today are metals
- on the left of the zigzag line on the table
- good conductors of heat and electricity
- characteristic lustre
- malleable (can be hammered into sheets)
- ductile (can be drawn into wires)
- high densities
- high melting and boiling points
- mercury is the only liquid metal at room temperature
- gold, silver and copper are called coinage metals because they are used to make coins
4) Non-metals:
- on the right of the zigzag line
- do not have lustre
- brittle
- non-ductile, non-malleable
- poor conductors of heat and electricity
- iodine, phosphorus and carbon are solid non-metals
- bromine is the only non-metal that is a liquid at room temperature
- gases such as chlorine, oxygen and helium are non-metals
5) Metalloids:
- on each side of the zigzag line on the table
- no sharp dividing line between metals and non-metals
- a metalloid is an element that exhibits properties of both metals and non-metals
- boron, silicon, germanium, arsenic, antimony, tellurium, polonium, and astatine fall in this category
6) Properties of noble bases, alkali metals, halogens and hydrogen:
The first rare gas was discovered by Ramsay around 1892, while he was studying atmospheric nitrogen. He
successfully extracted argon and demonstrated that it was inert. He also discovered helium, neon, krypton, and
radon. All these gases are colourless, odourless and chemically inert. For over 50 years, scientists believed that
rare gases could not combine with any other element, but now we know that xenon can combine with fluorine to
form many compounds. The same goes for krypton and radon.
Helium is approximately 7 times lighter than air; it has a considerable lifting power. Because it does not react
chemically, it is the gas used in trial balloons (for weather and environmental tests), and toy balloons. Helium is
often used for research on low temperatures because its boiling and freezing points are lower than any other
substance.
Argon is an excellent heat conductor. It is used to fill electric and electronic light bulbs because it disperses the heat
given off by the filament and prevents it from burning. The rare gases are also used in neon signs.
The alkali metals are those of group 1. They all have metallic properties and are chemically active. They react
vigorously with water to release hydrogen from water. Sodium is so active that it must be kept under an oil layer;
this prevents it from any contact with the humidity and the oxygen contained in the air. All alkali metals react with
chlorine to form colourless compounds which crystallize in cubic shapes and have similar formulas: LiCl, NaCl,
KCl, RbCl, CsCl, and FrCl.
Halogens are elements of group 17. The word halogen means salt-former. All these elements have non-metallic
properties. They react with hydrogen to form compounds that dissolve in water to make an acidic solution. They
also react with metals to form salts. In their natural state, these elements are diatomic: F 2, Cl2, Br2, I2.
Hydrogen is unique with its atomic number 1. Sometimes, it is placed in group 1 of the periodic table, although it
is not a metal. However, it can form compounds with halogens, just as the metals of group 1. It could easily be
placed in group 17 with the halogens because it can also react with metals, such as sodium, and non-metals, such as
nitrogen, in the same way as elements of group 17. Because of its unique characteristics, hydrogen does not have a
definite position in the periodic table - although it is found at the top of group 1 in most periodic tables.
3- Metallic trends in the periodic table:
The most metallic elements are located in group 1.
The metallic character of the elements decreases from left to right in a row and increases from top to bottom in a
group. Francium is the most metallic element in the periodic table.
The most non-metallic elements are located in group 17.
There is a gradual transition from metallic to non-metallic properties as one goes from left to right within a period.
Fluorine is the least metallic element (or the most non-metallic element) in the periodic table.
THE PERIODIC LAW AND ENERGY LEVELS
1- Valence electrons and the periodic table:
Valence electrons are the electrons located in the outermost shell. These are the ones that are lost, gained or shared
during the formation of a chemical bond.
According to their placement on the periodic table, the following two trends apply to the elements in groups 1, 2, 13 to
18:
1) The group number indicates the number of valence electrons.
ex:
group 1  1e- on the outermost level
group 13  3e- on the outermost level
2) The period number indicates the outermost level where the valence electrons can be found.
ex:
period 1  valence electron(s) are on level 1, or 1st electron shell
period 7  valence electron(s) are on level 7, or 7 th electron shell
2- Chemical properties and outer energy level electrons:
Let us look at similarities between elements belonging to the same group.
a) group 1: Li, Na, K
- alkali metals
- each one has 1 electron in its outer energy level
- all three have similar chemical properties: they have a tendency to lose that electron when they combine
with another element to form a compound
b) group 2: Be, Mg, Ca
- alkaline earth metals
- each one has 2 electrons in its outer energy level
- all three have similar chemical properties: they have a tendency to lose these two electrons when they
combine with another element to form a compound
c) group 17: F, Cl, Br
- halogens
- each one has 7 electrons in its outer energy level
- all three have similar chemical properties: they have a tendency to gain an electron when they combine with
another element to form a compound
d) group 18: He, Ne, Ar
- noble/inert/rare gases
- each one has 8 electrons in its outer energy level
- all three have similar chemical properties: they are chemically stable, i.e. they do not lose or gain electrons
to form compounds, therefore the name inert gases

The similarities between elements in a group are related to the number of electrons in the outer energy level.
The similarity of this listing in the periodic table is not accidental. The periodic table was first constructed to group
elements according to their chemical properties. However, it can also be constructed by grouping elements according to
their electron configurations.

Chemical properties are related to the electron configuration of an element.
Example:
Consider elements 9, 17 and 35:
The electron configurations of these elements all end in s2 p5. Therefore, these elements have similar properties and
belong to the same group in the periodic table.
PERIODIC PROPERTIES
1- Atomic radius:
 It is the distance between the nucleus of an atom and the atom's outer limits, usually expressed in picometres or
nanometres.
a) from left to right of the periodic table:
The radii of atoms decrease from left to right in the periodic table. As you move from the left to the right of a
period, the number of electrons in the atoms is increasing, and you might expect the radii of atoms to increase.
However, this does not happen.
Explanation:
As you go from the left of a period to the right, each additional electron is going into a sublevel of the same energy
level. Electrons having approximately the same energy would be expected to be approximately the same distance
from the nucleus. However, as you go from the left to the right side of a period, protons are also being added to the
nucleus, increasing the positive nuclear charge. As this happens, electrons are being pulled closer to the nucleus, and
the size of the atom decreases.
b) from top to bottom of the periodic table:
The atomic radii increase from the top of a group of the periodic table to the bottom of the same group.
Explanation:
As you move from the top of a group to the bottom, the extra electrons are going into higher energy levels that place
the electrons farther from the nucleus. Even though the nuclear charge increases from the top to the bottom of a
group, it is less effective at attracting electrons which are much farther away from the nucleus. Since the principal
quantum number n increases, the size of the orbitals increase as well as the size of the atom.
A positive ion (cation) is always smaller than the atom from which it is derived because of the extra nuclear charge
which pulls the electrons closer to the nucleus. Remember that a positive ion is an atom that has lost one or more
electrons, so it has one or more extra protons to attract the electrons that are left much more strongly.
radius (X+)  radius (X°)
A negative ion (anion) is always larger than the atom from which it is derived because of the extra negative charge.
Since electrons repel each other, they will increase the size of the atom.
radius (X-)  radius (X°)
2- Ionization energy:
 The energy required to remove an electron completely from an atom is called ionization energy or the first ionization
energy of the atom, usually expressed in kJ/mol. Here, we mean a valence electron in its ground state belonging to
an atom in its gaseous state.
X(g)
+
IE1

X+(g)
+
1e-
If the ionization energy is low, it is easy to remove an electron.
If the ionization energy is high, it is difficult to remove an electron.
The easiest electron to remove is the one occupying the highest energy level, i.e. the furthest from the nucleus.
For any given atom, it is theoretically possible to have as many ionization energies as there are electrons in that
atom. We refer to those as second, third, fourth, etc. ionization energy.
The second ionization energy is shown by this equation:
X+(g)
+
IE2

X2+(g)
+
1e-
Therefore, to remove 2 electrons from a neutral atom, we must supply IE = IE 1 + IE2
X(g)

X+(g)
+
1e-
X+(g)  X2+(g) + 1e________________________________
X(g)  X2+(g) + 2e-
IE1
IE2
IE1 + IE2
a) from left to right:
The ionization energy increases as one goes from the left side to the right side of a period.
Explanation:
From left to right in the periodic table, the nuclear charge increases while the energy level n remains the same. It is
more difficult to pull electrons away when they are attracted closer to the nucleus.
b) from top to bottom:
The ionization energy decreases as you go from the top of a group to the bottom.
Explanation:
From the top of a group to the bottom, the number of energy levels increases and the outermost electrons are farther
away from the nucleus, hence easier to be removed.
3- Electron affinity:
 It is the energy given off when an electron is added to an atom, usually expressed in kJ/mol.
X(g)
+
1e-

X-(g)
+
EA
IE = tendency for atoms to lose 1eEA = tendency for atoms to gain 1e-
Attention:
Some atoms have a tendency to pick up additional electrons. This process results in a lowering of energy of the
system; i.e. it ichieves a more stable state for the added electrons. Energy will be given off when an added electron
is moved into the positive field of the nucleus. This energy is the electron affinity.
a) from left to right:
Electron affinity increases from left to right within a period.
Explanation:
From left to right, the size of the atoms decreases and the electrons are pulled closer to the nucleus. More energy
(EA) is given off as a smaller atom accepts an additional electron to become a negative ion (anion).
b) from top to bottom:
Electron affinity decreases as you go from the top to the bottom within a group.
Explanation:
The elements at the top of a group generally are smaller and have a greater tendency to accept an electron to become
an anion. The smallest halogen, fluorine, is the most reactive element known.
Shielding effect:
Any given electron is subjected to attractive forces from the nucleus as well as repulsive forces from the surrounding
electrons; this effect is particularly noticeable when many electrons are located between the nucleus and the given
electron. In this case, the attractive force felt by this electron is greatly reduced.
Example:
13 Al
The electron on 3p orbital is attracted by the nucleus but also repelled by the other 12 electrons.
The internal electrons block the positive charges from the nucleus, hence forming a negative shield which diminishes the
attractive force felt by the external electron. This is called shielding effect and it increases as the number of electrons
located in the internal orbitals increase.
This is why the valence electron in the potassium atom is not as strongly attracted by its nucleus as the valence electron
in the lithium atom, although the potassium nucleus holds more positive charges.
4- Electronegativity:
 It is tendency for an atom to attract electrons to itself when it is chemically combined with another element.
Electronegativities have been calculated for the elements. They are expressed in arbitrary units on the Pauling
electronegativity scale. (see p. 285 of Addison-Wesley). This scale is based on a number of factors including the
electron affinity and ionization potential of the atoms.
Electronegativity follows the same periodic trend as ionization energies:
a) from left to right:
Electronegativity increases from left to right within a period.
b) from top to bottom:
Electron affinity decreases as you go from the top to the bottom within a group.
A few more things to note about electronegativity:
 The largest electronegativity values belong to the nonmetals, specifically the nonmetals of group 17.
Fluorine is the most electronegative element of the periodic table.
 The smallest electronegativity values belong to the most active metals of group 1. Francium and
cesium are the least electronegative elements of the periodic table.
 Noble gases do not have electronegativity values.