Download Chapter 5 Notes

Chapter 5 Notes
The periodic table has not always looked as it does today. The first periodic table was
arranged by Dmitri Mendelev according to atomic mass. He noticed that physical and
chemical appeared to be in a repeating or PERIODIC pattern. He continued to arrange
his table in groups according to similar properties, but found that he had several empty
spaces between the elements. He proposed that undiscovered elements would fill the
empty spaces.
Working with Ernest Rutherford, Henry Mosley rearranged Mendelev’s periodic table
according to the number of protons (atomic number). This arrangement was consistent
with Mendelev’s groupings and lead to the modern PERIODIC LAW, the physical and
chemical properties of elements are a function of the atomic number.
The PERIODIC TABLE can be defined as the arrangement of elements in order of
atomic number so that elements with similar properties are grouped in columns.
The periodicity of the periodic table is as follows.
Group 1 elements
Li
Na
K
Rb
Cs
Fr
atomic number
3
11
19
37
55
87
difference between periods
11-3=8
19-11=8
37-19=18
55-37=18
87-55=32
Group 18 elements
He
Ne
Ar
Kr
Xe
Rn
atomic number
2
10
18
36
54
86
difference between periods
10-2=8
18-10=8
36-18=18
54-36=18
86-54=32
Notice how the difference in atomic numbers following the same repeating pattern
between groups 1 and 18.
The periodicity on the periodic table also explains how electrons fill orbitals. The
orbitals are filled in the following manner.
Period
1
2
3
4
5
6
7
# of elements
2
8
8
18
18
32
32
order of fill
1s
2s, 2p
3s, 3p
4s, 3d, 4p
5s, 4d, 5p
6s, 4f, 5d, 6p
7s, 5f, 6d, 7p
The groups on the periodic table can be broken down into blocks to help with filling
electron orbitals.
S- Block
The s-block elements consist of the elements in groups 1 and 2 of the periodic table.
Electrons from these elements fill the s orbital of each period.
Group 1 (ALKALI METALS) fill the s orbital with 1 electron. These elements are
considered to be reactive metals because they are not readily found in pure form in nature
and react violently with water. They have a silvery appearance and are soft.
Group 2 (ALKLINE EARTH METALS) fill the s orbital with the second electron. The
s-orbital is filled with group 2 elements. These elements are harder and more dense than
the alkali metals. They are also less reactive, but their reactivity is enough that they are
not readily found in nature in pure form.
HYDROGEN and HELIUM are part of the s-block elements, but their properties differ
from other s-block elements.
P-block
The p-block elements consist of the elements in groups 13-18. Electrons from these
elements fill the 3 orbitals of the p shell. There are 6 elements providing the 6 electrons
needed to fill the orbitals. Metals, metalloids, and non-metals make up the p-block
elements. Group 17 contain the HALOGENS, the most reactive non-metals. Halogens
will readily react with metals to form salts. Group 18 contain the NOBLE GASES.
Noble gases are unreactive because electrons from the gases fill the p orbitals.
The s-block and p-block elements make up the MAIN GROUP elements.
D-block
The metals in groups 3-12 make up the d-block. These metals are called TRANSITION
METALS. These metals have the properties associated with typical metals. They are not
especially reactive so they can be found in nature in pure form. They are good
conductors of energy. Many of these metals do not follow the general rules for electron
configuration.
F-block
The elements of the f-block can be found between the groups 3 and 4 in periods 6 and 7.
These elements are found in the insert at the bottom of the periodic table. The elements
in period 6 are called the LANTHANIDES. The lanthanides are similar to the alkaline
earth metals. The elements in period 7 are called the ACTINIDES. The actinides with
the exception of the first 4 elements are synthetic. All actinides are radioactive.
Many chemical properties will trend on the periodic table. These properties are atomic
radii, ionization energies, and electronegativity.
ATOMIC RADII is the distance
between the nuclei of identical
atoms that are bonded together.
Across a period the atomic radii of
elements tend to decrease from
group 1 to group 18. This decrease
in atomic radii is due to the
increasing positive charge created
by the addition of protons as you
move across the period. The more
protons the greater the positive
charge pulling the negatively
charged electrons closer to the
nucleus. The atomic radii do not
follow the same trend down a
group. The atomic radii of elements moving down a group will increase in size. The
increase in size is due to the addition of electron orbitals as the energy levels are added.
Electrons can be added or removed to atoms creating charged atoms called IONS. Ions
can carry either a positive or negative charge. The charge is created by the balance
between protons and electrons. An atom with more protons than electrons will carry a
positive charge. An atom with more electrons than protons will carry a negative charge.
IONIZATION is the process in
which electrons are removed
from an atom. The energy
required to remove the electron is
called IONIZATION ENERGY.
The greater the ionization energy
the more difficult it is to remove
an electron from the atom.
Ionization energy increases
across a period making it more
difficult to remove electrons from
non-metals. Electrons are more
easily removed from metals in groups 1 and 2. The ionization energy will decrease down
a group of main group elements due the electrons being located a further distance from
the nucleus.
The removal of an electron creates the situation where the protons exist in greater
numbers than electrons. The charge will be positive. Ions with a positive charge are
called CATIONS.
ELECTRONEGATIVITY is a measure of the ability of an atom to attract electrons. The
energy associated with gaining electrons is called ELECTRON AFFINITY. The greater
the electronegativity the more likely an atom is to accept an electron. Electronegativity
increases across a
period due to the filling
of the electron orbitals.
Elements in group 17,
halogens, are consider
to be the most reactive
of all non-metals
because they have 7 of
the 8 electrons needed
to fill the s and p
orbitals associated with
the period.
The acceptance of electrons by an atom creates a situation where the number of electrons
is greater than the number of protons. These atoms will carry a negative charge and are
called ANIONS.
The electrons that are lost, gained, or shared are called VALENCE ELECTRONS. These
electrons are involved in the formation of ions. The maximum number of valence
electrons is 8 regardless to where the element is on the periodic table. You can find the
number of valence electrons by counting the number of electrons in the s and p orbitals of
the highest energy level.