Atomic Structure Flashcards
Characteristics of S orbitals
Spherical in shape, non directional , size of s orbital increases with principal quantum number
Characteristics of P orbital
• Dumb-bell shaped
• Directional
• Three p orbitals have different orientation in space
Size of the p orbitals increases with principal quantum number
Characteristics of D orbitals
• Butterfly shapes for dxy, dyz, dxz, dx^2−y^2
• Doughnut shape for dz^2
• Directional
• Different orientation in space
Size of d orbitals increases with principal quantum number
Why are 4s orbitals filled before 3d orbitals
4s has a lower energy level than 3d subshell when it is not occupied by electrons. Once the 4s subshell is occupied by electrons, 4s subshell will have a higher energy level than 3d subshell due to repulsion and close energy gap between 4s and 3d subshell
What are the exceptions of the 4s orbital before 3d orbital rule
Copper and chromium
Cr atom gains extra stability with half filled 3d subshell
Cu atom gains extra stability with a fully filled 3d subshell
What is isoelectronic, isotopic, isotonic?
Atoms with the same number of electrons are isoelectronic
Atoms with the same number of protons are isotopic
Atoms with the same number of neutrons are isotonic
Why higher IE, more difficult to remove electron?
The higher the ionisation energy of an element, the more difficult it is to remove an electron.
2nd I.E. is more than 1st I.E. because more energy is required to remove an electron from a positively charged ion compared to a neutral atom due to stronger net electrostatic attraction
Why does ionisation energy increases across period
Across a period, the number of proton increases , nuclear charge increases. The number of electron increases but they are added to the same outermost shell, hence shielding effect remains relatively constant. Effective nuclear charge increases, resulting in stronger electrostatic attraction between the nucleus and the valence electrons. Hence more energy is required to remove the valence electron and first ionisation energy generally increases across the period.
Deviations from general trend
For period 3 elements
Why 1st I.E. of group 13 element < 1st I.E. of Group 2 element
E.g. 1st I.E. of Al < 1st I.E. of Mg
The 3p valence electron to be removed from Al has higher energy than the 3s valence electrons in Mg
Hence less energy is required to remove the 3p electron in Al than the 3s electron in Mg, and first ionisation energy of Al is lower than that of Mg.
Deviation from general trend
Why 1st I.E. of Group 16 elements < 1st I.E. of Group 15 elements?
E.g. 1st I.E. of S < 1st I.E. of P
Inter-electron repulsion is present between the paired electron in 3p orbital of S atom. Hence less energy is required to remove the valence electron from S, and first ionisation energy of S is lower than that of P.
Why Ionisation energy decreases down the group?
Down the group, the number of proton increases and hence the nuclear charge increases. However the number of electron shell and shielding effect also increases. Hence the valence electron experience weaker electrostatic attraction to the nucleus. Less energy is required to remove the valence electron, and first ionisation energy decreases down the group.
*increase in shielding effect outweighs increase in nuclear charge
