When comparing the properties of the chemical elements, recurring ('periodic') trends are apparent. This led to the creation of the periodic table as a useful way to display the elements and rationalize their behavior. When laid out in tabular form, many trends in properties can be observed to increase or decrease as one progresses along a row or column.
These periodic trends are distributed among 3 different properties namely, physical properties, chemical properties and on the basis of chemical reactivity. In chemical properties, it is classified on the basis of two i.e.periodicity of valence or oxidation states, anomalous properties of second period elements. Here, we are going to discuss about the periodic trends with respect to their physical properties.
- electronic configuration
- atomic radius
- ionization energy
- electronegativity
- electron affinity
- metallic character
1. ELECTRONIC
CONFIGURATION
·
The electronic configuration of
an atom is the representation of the arrangement of electron distributed among the shell and
subshells
·
Electron fill orbital in a way
to minimize the energy of atom
·
The electron in an atom fill
the energy level in order increasing energy
·
In order such as 1s, 2s, 3s,
3p, 4s, 3d, 4p, 5s,4d, 5p, 6s, 4f, 5d,6p,7s,5f, 6d, 7p
RULES ASSIGN ELECTRON ORBITAL
PAULI EXCLUSION PRINCIPLE
·
No two electron in an atom can have
the same set of four quantum number n, l, ml are
fixed
·
To assign different ms values
to electron
·
No two electrons in the same atom can
have exactly the same energy
|
HUND’S
RULE
·
every orbital in a subshell is singly
occupied with one electron before any one orbital is doubly occupied, and all
electrons in singly occupied orbitals have the same spin.
|
For
example ;
WRITING
ELECTRON CONFIGURATION
ORBITAL
DIAGRAM
·
A visual way to
reconstruct the electron configuration by showing each of the separate orbital
and the spins on the electron by determining the shell , s , p ,d , f
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ELECTRON ARRANGE IN ORBITAL DIAGRAM
SPDF NOTATION
·
The total number of
electron in each energy level is described by the superscript follow relating
energy level.
For
example :
2. ATOMIC RADIUS
DOWN THE GROUP
·
atomic radius increase
·
no electron increase
·
valence electron far from the nucleus
·
the attraction between the valence
electron and nucleus become weaker
·
atomic radii increase
|
ACROSS PERIOD
·
atomic radius decrease
·
effective nuclear charge ,
Zeff
increase
·
valence electron close to the nucleus
·
the attraction between nucleus and valence
electron stronger
·
atomic radii decrease
|
Energy
level from the nucleus
3. ionization energy
First ionization energy
The
minimum energy required to remove one mole of electron from one mole of gaseous
atom in its ground state
Second ionization
energy
The
minimum energy required to remove one mole of electron from one unpositive ion
in the gaseous state
size
|
stability
|
Size increase, IE decrease
|
Half-filled
Fully-filled
|
First ionization
energy across the period
*
First
I.E increases
WHY? - As you move
across a period, the atomic radius decreases, that is, the atom is
smaller. The outer electrons are closer to the nucleus and more strongly
attracted to the center. Therefore, it becomes more difficult to remove
the outermost electron
1 1. Mg > Al
Look at their electronic configurations:
Magnesium: 1s2 2s2 2p6 3s2 ... and ... aluminium: 1s2 2s2 2p6 3s2 3p1
The outer electron in aluminium is in a p
sub-level. This is higher in energy than the outer electron in magnesium, which
is in an s sub-level, so less energy is needed to remove it.
2. P > S
Look at their electronic configurations:
Phosphorus: 1s2 2s2 2p6 3s2 3p3 ... and ... sulphur: 1s2 2s2 2p6 3s2 3p4
The 3p electrons in phosphorus
are all unpaired. In
sulphur, two of the 3p electrons are paired.
There is some repulsion between paired electrons in the same sub-level. This
reduces the force of their attraction to the nucleus, so less energy is needed
to remove one of these paired electrons than is needed to remove an unpaired
electron from phosphorus.
First ionization
energy down the group.
*
The I.E
will decrease when down the group
§ Down the group, the shelding effect increase,
the atomic radii increase
§ So, valance electron have low attraction toward
the nucleus
§ Thus, lower energy need to remove the outermost
electron
4. electronegativity
*
What? :
Ability of an atom to attract electron to itself
*
Trend in
periodic table
From left to right across a period of
elements, electronegativity increases.
If the valence shell of an atom
is less than half full, it requires less energy to lose an electron than to
gain one. Conversely, if the valence shell is more than half full, it is easier
to pull an electron into the valence shell than to donate one.
From top
to bottom down a group, electronegativity decreases.
This is because atomic number
increases down a group, and thus there is an increased distance between the
valence electrons and nucleus, or a greater atomic radius.
5.Electron Affinity
Electron
Affinity is the energy associated with the addition of an electron to a gaseous
atom.
Example:
Cl(g) + e- → Cl-(g)
|
E.A. = -349 kJ/mole
|
Notice
the sign on the energy is negative. This is because energy is usually released in
this process, as opposed to ionization energy, which requires energy.
A more negative electron affinity corresponds to a greater attraction
for an electron. (An unbound electron has an energy of zero.)
As with ionization energy, there
are two rules that govern the periodic trends of electron affinities:
Electron affinity becomes less
negative down a group.
As the principal quantum number increases, the size of the orbital
increases and the affinity for the electron is less. The change is small and
there are many exceptions.
Electron affinity decreases or
increases across a period depending on electronic configuration.
This occurs because of the same
sub shell rule that governs ionization energies.
Down the group
|
Across
the period
|
|
|
6.metallic character
The metallic character of an element can be defined as how
readily an atom can lose an electron. Metallic character relates to the ability
to lose electrons, and nonmetallic character relates to the ability to gain
electrons.
Down the group
|
Across the period
|
|
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Another easier way to remember the trend of metallic
character is:
- moving left and down toward the bottom-left corner of the periodic table, metallic character increases toward Groups 1 and 2, or the alkali and alkaline earth metal groups.
- Likewise, moving up and to the right to the upper-right corner of the periodic table, metallic character decreases because you are passing by to the right side of the staircase, which indicate the non-metals. These include the Group 8, the noble gases, and other common gases such as oxygen and nitrogen.
In other words:
- Move left across period and down the group: increase metallic character (heading towards alkali and alkaline metals)
- Move right across period and up the group: decrease metallic character (heading towards non -metals like noble gases)
summary of periodic table trends
Moving Left → Right (across the period)
Atomic Radius Decreases
Ionization Energy Increases
Electronegativity Increases
electron affinity
increases
metallic character
increases
Moving Top → Bottom
Atomic Radius Increases
Ionization Energy Decreases
Electronegativity Decreases
electron affinity decreases
metallic character decreases