Inorganic Chemistry 2 - Atomic Orbitals, Electronic Configurations and the Periodic Table

Question 1

Electrons (like photons) display the properties of both ___ and ___.

Answer 1

particles and waves

Question 2

What is an atomic orbital?

Answer 2

An area within an atom with a high chance (>90%) of containing an electron/electrons

Question 3

Orbitals can hold a maximum of ___ electrons.

Answer 3

two

Question 4

What are the four types of orbitals?

Answer 4

s, p, d and f

Question 5

Electrons within atoms have fixed amounts of energy called ___.

Answer 5

quanta

Question 6

Each electron has ___ numbers that describe it.
What are they?

Answer 6

4 (n, l, ml, ms)

Question 7

No two electrons in one atom can have the ___ four numbers.

Answer 7

same

Question 8

n is called the…
What does n tell us?
n is somewhat related to the type of ___ containing the electron.

Answer 8

principle quantum number.
the energy level that the electron is on.
orbital

Question 9

What is the lowest value of n which can have a d orbital?

Answer 9

3

Question 10

L is called the…
L is always a value between _ and _.
What does L tell us?

Answer 10

angular momentum quantum number.
0 and n-1.
the type of orbital the electron is in

Question 11

Name the orbital represented by the following values of L.
3
1
2
0

Answer 11

f,
p,
d,
s

Question 12

ml is called the…
What does ml tell us?
ml is always a value between…

Answer 12

magnetic quantum number.
the orientation of the orbital.
-l and +l

Question 13

If l=0, how many orbitals and which type is it?
If l=1, how many orbitals and which type is it?
If l=2, how many orbitals and which type is it?
If l=3, how many orbitals and which type is it?

Answer 13

one, s.
three, p.
five, d.
7, f

Question 14

ms is called the…
What does ms tell us?
Name the possible values for ms.
How are these values represented?

Answer 14

spin magnetic quantum number.
the direction of spin of the electron.
+1/2, -1/2.
half arrows, pointing up or down

Question 15

What is a subshell?
What is an electron shell?

Answer 15

a bunch of degenerate orbitals of one type. (eg the 2p subshell is three p orbitals, the 3d subshell is five degenerate d orbitals, each with a different orientation, given by ml)

an electron shell is all of the subshells on an energy level (eg the 3rd level, n=3, has a 3s subshell, a 3p subshell, and a 3d subshell).

Question 16

State the Aufbau Principle.

Answer 16

electrons fill orbitals in order of increasing energy

Question 17

What is Hund’s Rule?

Answer 17

when degenerate orbitals are available, electrons fill singly, keeping their spins parallel until the subshell is half full, when pairing starts

Question 18

State the Pauli Exclusion Principle.

Answer 18

No two electrons in the same atom can have the same 4 quantum numbers (therefore orbitals can only hold 2 and they must have opposite spins)

Question 19

Name the following methods of representing electrons within atoms.
1s2, 2s2, 2p6, 3s2, 3p6, 3d7, 4s2.
The one with arrows in boxes.
[Ar] 3d2, 4s2

Answer 19

spectroscopic notation,
orbital box notation,
shortened spectroscopic notation (rounded down to the nearest noble gas + the remaining electrons)

Question 20

The periodic table is divided into _ blocks, corresponding to the ___ ___ ___ of the elements within these blocks.

Answer 20

4, outer electron configuration

Question 21

The anomalous elements on graphs of first (and subsequent) ionisation energies can be explained by the relative ___ of different ___ configurations.

Answer 21

stabilities, electron

Question 22

When are subshells most stable?

Answer 22

when they are half full or completely full

Question 23

When explaining that half full/completely full subshells are more stable, state the specific ___ when possible.

Answer 23

subshell (eg in Nitrogen it’s 2p)

Question 24

When an atom has a stable electron arrangement, its ionisation energy is ___.

Answer 24

higher

Question 25

An example of an element with an ‘anomalous’ first ionisation energy is nitrogen, as it has a higher ionisation energy than oxygen, despite…
Nitrogen’s electron configuration is 2, 5.
This means it has a ___ full p subshell, which is ___ stable than oxygen’s p subshell.

Answer 25

having an extra electron and proton.
half, more

Question 26

VSEPR theory can be used to predict what?
What does VSEPR stand for?

Answer 26

the shape of molecules and polyatomic ions.
Valence Shell Electron Pair Repulsion

Question 27

Arrange the following electron pairs in order of increasing strength.
Bonding : Bonding
Non-Bonding : Non-Bonding
Bonding : Non-Bonding

Answer 27

Weakest
Bonding : Bonding
Non-Bonding : Bonding
Non-Bonding : Non-Bonding
Strongest

Question 28

How do you find the number of electron pairs surrounding a central atom?
*There are four steps :)

Answer 28

1. Take the atom’s valence number/group number/number of outer electrons
2. Add one for every -ve charge and subtract one for every +ve charge
3. Add one for every bond (or bonding atom assuming they’re single bonds)
4. Divide by 2!

Question 29

Electron pairs are _vely charged, and ___ each other.

Answer 29

-vely, repel

Question 30

Electron pairs are arranged to maximise ___ and minimise ___.

Answer 30

separation, repulsion

Question 31

Name the shape that the following number of electron pairs will take around a central atom, and state the angle(s) between them.
Two
Three
Four
Five
Six

Answer 31

Two: linear (180 degrees)
Three: Trigonal Planar (120 degrees)
Four: Tetrahedral (109.5 degrees)
Five: Trigonal Bipyramidal (90 and 120 degrees)
Six: Octahedral (90 degrees)

Question 32

If all electron pairs are bonding, the bond angle will be (the same/different), compared to if all electron pairs are non-bonding.

Answer 32

the same.

Question 33

If some electron pairs in an atom are bonding, and others are non-bonding, the bond angles may ___ slightly.

Answer 33

change (eg a tetrahedral electron pair arrangement with one non bonding pair will have a 107 degree bond angle, instead of its typical 109.5)

Question 34

Non bonding electron pairs have a ___ repulsion force than bonding pairs.

Answer 34

stronger

Question 35

An increase in the number of (bonding/non-bonding) electron pairs causes the bond angle to decrease.

Answer 35

Non-bonding (stronger repulsion force)

Question 36

By how many degrees do bond angles decrease with each extra pair of non-bonding electrons?

Answer 36

2.5