Inorganic Chemistry - Module 2 Flashcards
Define ionisation energy.
The energy needed to form positive ions.
What is the first ionisation energy of an element?
The energy needed to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.
Do outer-shell electrons require the most or the least ionisation energy? Explain your answer.
They require the least ionisation energy. This is because they are further away from the nucleus and experience less nuclear attraction.
Which three factors affect the nuclear attraction experienced by an electron?
Atomic radius, nuclear charge, and electron shielding/screening (inner shells of electrons repel outer-shell electrons).
What is the second ionisation energy of an element?
The energy needed to remove an electron from a 1+ ion to form a 2+ ion.
Define principle quantum number.
A number representing the relative overall energy of each orbital, which increases with distance from the nucleus.
Define atomic orbital.
v
What is the formula to work out how many electrons shells hold?
2n2
Within a shell, state the number of s-, p-, d-, and f-orbitals.
- s = 1
- p = 3
- d = 5
- f = 7
How many electrons can an orbital hold?
2
Define sub-shell.
Give the electron configuration in terms of sub-shells for Si (atomic number of Si = 14).
1s2, 2s2, 2p6, 3s2, 3p2
Give the electron configuration in terms of sub-shells for N3- (atomic number of N = 7).
1s2, 2s2, 2p6
Why are noble gases so unreactive?
They have full outer shells.
Which type of compound does ionic bonding take place in?
Metal-Non Metal.
Which type of compound does covalent bonding take place in?
Non metal-Non metal.
Define ionic bond.
The electrostatic attraction between oppositely charged ions.
Write the electron configuration for K and K+ (atomic number of K = 19)
K: 1s2, 2s2, 2p6, 3s2, 3p6, 4s1
K+: 1s2, 2s2, 2p6, 3s2, 3p6
Define covalent bond.
A bond formed by a shared pair of electrons.
Define a dative covalent/co-ordinate bond.
A shared pair of electrons which has been provided by one of the bonding atoms only.
What is the bond angle in a water molecule?
104.5
What is the bond angle in an ammonia molecule?
107
Define permanent dipole.
A small charge difference across a bond due to the difference in electronegativities of the bonded atoms.
Across the period table, does electronegativity increase or decrease?
Increases.
What are the strongest bond types?
Ionic and covalent bonds.
What is the weakest bond type?
Van der Waals’ forces.
Define Van der Waals forces.
Attraction between induced dipoles of neighbouring molecules.
Where do hydrogen bonds form?
Between an electron-deficient H atom on one molecule and the lone pair of electrons on an O or N atom of another molecule.
Give a diagram for hydrogen bonding.
Define metallic bonding.
The electrostatic attraction between positive metal ions and delocalised electrons.
Why do ionic compounds have high melting and boiling points?
Because of the strong electrostatic forces holding the oppositely charged ions together in a lattice structure.
Why can’t ionic compounds conduct electricity when solid, but can when molten or dissolved?
When solid, the ions are in a fixed position and can’t move. However, when molten or dissolved, the ions can move, allowing them to conduct electricity.
