Topic 4: Periodic trends and molecular orbital theory Flashcards
What is nuclear charge and effective nuclear charge?
- Nuclear charge: number of proteons in nucleus
- Effective nuclear charge: partial nuclear charge felt by electrons in multi-electron atoms
Compare the effective nuclear charge on electrons that are closer and further from nucleus
Electrons closer to the nucleus are more strongly attracted to the nucleus, thus higher Zeff
Electrons further away from the nucleus or in outer orbitals are shielded by those closer –> weaker attraction and lower Zeff
Identify and explain the trend of effective nucler charge and atomic radii in the periodic table
- Across the period:
+ same quantum number but the number of electrons in the outer orbital increases
+ orbitals of the same quantum number don’t shield effectively
–> more affected by nuclear charge –> attracted more strongly to nucleus
–> decreasing radii and increasing Zeff - Down a group:
+ increasing quantum number with same number of electrons in the outer orbital
+ orbitals of smaller quantum number shield those with higher quantum number effectively
–> less affected by nuclear charge –> attracted less strongly to nucleus
–> increasing radii and decreasing Zeff
Compare anionic and cationic radii with neutral atoms
- Anionic radii:
+ anions are usually formed with electrons added to the same orbital
+ same Z and same Zeff
–> more electrons means more electron repulsion
–> larger radii (-2 also > -1) - Cationic radii:
+ electrons removed from outer orbital
–> valence electrons closer to nucleus –> stronger Zeff
–> smaller radii than neutral atoms
–> cations usually are smaller than anions in ionic compounds
Explain ionisation energy (Ei) and compare Ei of atoms with larger or smaller Zeff
Ionisation energy is energy required to remove an electron completely from an atom.
–> energy input –> always positive
Larger Zeff –> strongly attracted to the nucleus
–> more energy required to remove electron
–> larger Ei
Identify and explain the trend of ionisation energy in the periodic table
- Across a period: Zeff increases
–> electrons tightly bound to nucleus –> more energy input to remove –> increasing Ei - Down a group: Zeff decreases
–> electrons less tightly bound to nucleus –> less energy input to remove –> decreasing Ei
What is first ionisation energy and how is it compared to second ionisation energy?
Energy required to remove the first electron from a neutral atom
When 1 electron is removed, if there are still electrons in that orbital, then the electron repulsion is still less –> a bit more affected by Z –> higher 2nd Ei.
If after the 1st electron being removed, the outer orbital is empty, then outer electrons are closer to nucleus –> higher 2nd Ei
Explain electron affinity (Eea)
Energy change when an electron is added to an atom in the gas phase
What does it mean when energy is required or released when electron is added to neutral atoms?
- Require energy input means the added electron needs to occupy higher energy level or it disrupts half-filled orbital
–> atoms not have an affinity for electrons - Release energy output means the added electron occupy the same orbital with the valance electrons or fulfill the valence shell
–> more stable
–> the atom has affinity for electrons and has a negative Eea
Compare the ionisation energy and electron affinity between metals and non-metals
- Metals have a very low ionisation energy as they want to lose rather than gain electrons
- Non-metals have an affinity for electrons as they readily gain electrons to fulfill valence shell.
Explain how electron sharing leads to lowering of energy in terms of
i) de Brogile wavelength
ii) attraction and energy level
i) When atoms come together, electrons are shared in between the nuclei -> increased wavelength -> decreased energy
ii) When 2 atoms come together, they share electrons in the area in between the two nuclei. This is a covalent bond so the attraction between nucleus and the new electron is stronger than the repulsive force. As electrons are now affected by more electrostatic attarction than usual, their energy lower down.
Explain how the distance in between 2 atoms affect their interaction, overall energy and stability
Explain bond length and bond energy
- Atoms of different elements will have different distances at which they display no interaction or start attracting.
- When 2 atoms are too far apart, they exert no attraction towards each other –> total enegry is the sum of 2 individual atoms
- Bond energy:
+ the energy difference between individual atoms and the molcule; or
+ the energy required to break the bond of the molecule - As the distance gets smaller, the 2 atoms start to attact, energy level decreases and the atoms become more stable
- Bond length: the distance at which
+ the molecule has the most stable balance between attractive and repulsive forces; if lower, repulsive forces overwhelm
+ the molecule has the most energetic advantage over individual atoms
+ the new molecule is at the lowest energy position
What is the unit of distance when discussing atomic distances?
What is the equilibrium bond length?
What is Bohr radius?
- Unit: Ångström (Å)
- Equilibrium bond length: ~1Å
- Bohr radius: the most like distance between a proton and an electrons in a H atom - 0.53Å
How to determine the number of molecule orbitals in general?
Total number of atomic orbitals in component atoms
= number of molecular orbitals
Explain how constructive and destructive interaction between orbitals result in molecule with higher and lower energy level compared to individual atoms
- Same phase –> constructive interaction
–> larger amplitude –> higher electron density in between the nuclei –> minimize internuclear repulsion
–> Bonding molecular orbital - Different phase –> destructuve interaction
–> smaller amplitude or cancel out
–> lower or NO electron density in between the nuclei
–> nodes appear and internuclear repulsion increases
–> Anti-bonding molecular orbital with nodes
