Chemistry UNIT 2- Electron configuration and bonding Flashcards
What is the Heisenberg Principle?
You can’t know with certainty both where an electron is and where an electron will travel to next.
This makes it impossible to plot the orbit and location of an electron around a nucleus.
What is the arrangement of electrons in their subshells up to 4D^10
1S^2 2S^2 2P^6 3S^2 3P^6 3D^10 4S^2 4P^6 4D^10
What is the aufbau principle?
Electrons fill the lowest energy level subshell first before moving to the next level.
What is Hund’s rule
The most stable arrangement of electrons is that with the maximum number of unpaired electrons all with the same spin direction. (one electron goes into a orbital at a time before doubling up)
What is the Pauli Exclusion principle?
No two electrons in a atom can have the same set of four quantum numbers in an orbital and they must have opposite spins.
What is ionisation energy?
The energy required to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mol of gaseous 1+ ions.
What are the main key points for electron configuration?
- Electrons arranged in energy levels (shells) which also themselves have subshells.
- Each subshell consists of 1+ electron orbital
- Each orbital can hold 2 electrons with opposite spins
- Electrons fill lowest energy level orbitals first.
- Exceptions are Cr & Cu have 4S electrons promoted to 3D
- Electrons occupy orbitals singularly and only pair up when no empty orbitals of the same energy level are available.
What is successive ionisation energy?
To form a positive ion, we need energy to overcome the electrostatic attraction from the nucleus.
A each electron in removed from an atom the remaining ions gain a more positively charged structure so more energy is required.
Complete the first and second ionisation equation for potassium.
1st IE: K(g) —-> K+ + e-
2nd IE: K+(g) —-> K2+ + e-
Each time you do a successive ionisation equation, start from the previous IE
What are the main factors affecting Ionisation energy?
- Atomic radius- Greater the atomic radium, smaller nuclear attraction experienced by the outer electrons
- Nuclear charge- Greater nuclear charge, greater the attraction force on the outer electrons.
- Electron shielding- Inner shells of electrons repel the outer shell electrons.
What is the general trend in ionisation energy down a group?
- Negative trend due to a greater number of orbitals creating a larger shielding effect which means that IE energy is required.
- Larger atomic radius so there is less attraction between the positive nucleus and negative electrons.
- Nuclear charge is greater due to an increase in proton number however more orbitals means more shielding therefore is less effective.
What is the general trend in ionisation energy across a period?
The general trend increases the first IE as you go across the period as the is a higher nuclear charge and no extra orbitals causing electron shielding.
What is the explanation across a period for the difference in Group 3 elements?
There is a drop in the value of the element. This is because the extra electron has gone into one of the 2P orbitals. This increases the shielding making the electron easier to remove.
What is the explanation across a period for the difference in Group 6 elements?
Drop in the value of the element due to the extra electron being paired up with one of the electrons already in one of the 2P orbitals. The repulsive force between the two paired up electrons means that less energy is required to remove one of them.
Why do transition metals have similar ionisation energies bar zinc?
All of the elements have an electron structure of [Ar] 3D^n 4S^2. The electron being lost always comes from the 4S orbital.
Which would have the higher value, the 1st IE of sodium or the second IE of magnesium?
Magnesium would have the higher IE even though they have the same electronic structure. Magnesium has 12 protons and sodium has 11 therefore there is a greater nuclear charge attracting the electron to be removed in magnesium so a higher ionisation energy is required.
What is the trend for the size of the atomic radius across a period?
Number of protons increases which increases the “pull” on each electron, therefore pulling them in closer and reducing the atomic radius.
There is no increase in the number of shells/ energy levels to cause any effective shielding.
What is the trend for the size of the atomic radius down a group?
Increase in the number of energy levels/shells which causes a significant shielding effect.
There is a greater number of protons as you go down the group however these protons do not have a significant enough effect to overcome shielding.
What are the factors effecting effective nuclear charge?
Number of protons.
Shielding from inner electrons/shells
What is the ionic radius trend in metals and non-metals?
Metals- lose electrons when gaining an ion so the electrons left in the atom will feel a greater effective nuclear charge and so the ionic radius is always smaller than the atomic radius.
Non-metals- Gain electrons so the ionic radius becomes larger than the atomic radius.
