Atomic number: The atomic number is the total number of protons in the nucleus of an atom. Mass/Nucleon number: Mass/Nucleon number refers to the total number of protons and neutrons in the nucleus of an atom. Subatomic particles: Electrons, Protons and neutrons are called subatomic particles.
Ions
Ions have an unequal number of protons and electrons.
Ions are formed due to the loss and gain of electrons.
Ions can be identified in two respects:
Cations:
Positively charged ions are called Cations
Cations are formed due to the loss of electrons.
Cations have more protons than electrons.
The number of charge on the cations indicates the number of electrons donated.
For example:
Anions:
Negatively charged ions are called Anions.
Anions are formed due to the gain of electrons.
Anions have more electrons than protons.
The number of negative charges on the Anions indicates the number of electrons taken.
For example:
Isotopes
Isotopes are the atoms of the same element that have the same number of protons but a different number of neutrons.
Isotopes have the same:
number of protons
atomic radius
electronic configuration
Similar chemical properties
On the other hand
Isotopes have different:
number of neutrons
nucleon number
physical properties
relative isotopic mass
mass number
Isoelectronic Ions
Ions that have the same number of electrons.
Example - Na+ , Mg2+ , O2- , F-
All of the ions above have the same number of electrons (10 electrons)
Sodium ion and magnesium ions are isoelectronic cation(as they are positively charged).
Oxide ion and fluoride ion are isoelectronic anions(as they are negatively charged.)
Ionic radius It is the distance between the nucleus and the outermost shell of an ion.
For isoelectronic cations, the ionic radius of the ion decreases with the increase of charge on the cation.
For isoelectronic anions, the ionic radius of the ion increases with the increase of charge on the cation.
Subatomic particles in an electric field
Protons are positively charged and will be deflected away from the positive plate and towards the negative plate.
Neutrons have no charge and will continue moving in a straight line.
Electrons are negatively charged and will be deflected away from the negative plate and towards the positive plate.
Electronic configuration Principal quantum shell = number of energy shells. n = number of energy shells. Maximum number of electrons in an energy shell = 2n²
The number of electrons present in the outermost shell of an atom represents the group number.
The number of energy shells of an atom represents the period number.
Metallic elements (group 1,2,3) donate electrons.
Non-metallic elements (4,5,6,7) gain electrons.
A principal quantum shell has sub-shells.
Sub-shells have orbitals.
Electrons are present as "clouds" in these orbitals.
Writing electronic configuration.
Examples
Electronic configuration box diagram
Examples:
Na -
P-
**The shielding effect can be defined as a reduction in the effective nuclear charge on the electron.
Exception
Half filled d-subshell or fully filled d- subshell is more stable than a partially filled d-subshell.
As a result, electronic configurations of Copper (Cu) and Chromium(Cr) are different.
During the formation of a cation, 3d-block elements donate electrons from the 4s subshell first.
Shapes of orbitals
The shape of a 1s orbital
The shape of a 2s orbital
The shape of a 2s orbital is larger than 1s .
The shape of a p orbital
Ionisation energy
It is the amount of energy required when one electron is removed from each atom in one mole of gaseous atoms to form one mole of gaseous ion.
Second ionisation energy is the amount of energy required when one electron is removed from each ion of +1 gaseous ion to form one mole of gaseous +2 ion.
Variation of ionisation energy
Down the group ionisation energy decreases.
Atomic radius increases.
The distance between the nucleus and outer electron increases.
The outer electron moves further away.
The shielding effect increases.
The attraction between the nucleus and the outer electron increases.
Although the nuclear charge increases, less energy is needed to remove the electron.
Across the period the ionisation energy increases.
The number of protons in the nucleus increases.
The shielding effect remains almost constant.
The distance between the outer electrons and nucleus decreases.
The nuclear attraction increases for the outer electron.
The atoms gradually get smaller.
As a result more energy is needed to remove the outer electron.
The first ionisation energy of Aluminium is less than Magnesium
The outer electron of Magnesium is present in the 3s orbital.
The outer electron of Aluminium is present in the 3p orbital.
3p orbital is much further away from the nucleus compared to 3s.
The outer electron is more shielded.
The electron is easier to remove so less energy is required.
The first ionisation energy of Phosphorus is more than Sulfur.
The 3p orbitals of phosphorus are singly filled.
One of the 3p orbital of sulfur has paired electrons.
The paired electrons repel.
Electron is easier to remove.
Successive ionisation energy data can be used to determine the group number of an unknown element
The number of electrons removed from the atom before the first large energy gap indicates the group number.
In the example above the group number of this atom is 2 as two electrons were removed before the large energy gap.
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