chemistry is easy!: periodic table trends

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

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
Smaller size, zeff increase More strong te electron valence held by nucleus More energy required to remove the electron Size increase, IE decrease	More energy required to remove an electron from higher stability 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

Electron affinity decrease since the additional electron is entering an orbital farther away from the nucleus.
Since this electron is farther away, it should be less attracted to the nucleus and release less energy when added.
However, this trend applies only to Group-1 atoms.
Electron affinity follows the trend of electronegativity: fluorine (F) has a higher electron affinity than oxygen (O), and so on.

Electron affinity increases across a period (row) in the periodic table, due to the filling of the valence shell of the atom.

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
Metallic character increases down a group because the atomic size is increasing. When the atomic size increases, the outer shells are farther away. The principle quantum number increases and average electron density moves farther from nucleus. The electrons of the valence shell have less attraction to the nucleus and, as a result, can lose electrons more readily.	metallic character increases because the attraction between valence electron and the nucleus is weaker. enabling an easier loss of electrons.

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)