Quantum Theory Flashcards Preview

Advanced Higher Chemistry: Unit 1- Inorganic And Physical Chemistry > Quantum Theory > Flashcards

Flashcards in Quantum Theory Deck (43)
Loading flashcards...
1
Q

What did Niels Bohr predict?

A

electrons exist in orbits with specific energy

2
Q

What does atomic emission spectra provide evidence f?

A

subshells within each energy level

3
Q

What does classical theory assume electrons exist as?

A

a particle (with mass and charge)

4
Q

What does experimental evidence show that electrons can have?

A

wave-like properties (show diffraction and interference)

5
Q

What does the wave-like nature of electrons make it impossible to do? What is this known as?

A

specify exactly the position and momentum of the electron - Heisenburg Uncertainty Principle

6
Q

How can the discrete lines seen in the emission spectra be explained?

A

if the electron is regarded as having the properties of both particles and waves

7
Q

What does quantum theory suggest? (3)

A

Electrons within atoms behave as waves of different shapes and sizes

The waves are known as orbitals

Each orbital can hold a maximum of two electrons

8
Q

What are orbitals defined by?

A

Quantum numbers

9
Q

What is an orbital?

A

a region of space where the probability of finding the electron is >95% (i.e. there is uncertainty)

10
Q

What are the shapes of orbitals?

A

S
P
D

11
Q

What are S orbitals said to be?

A

spherically symmetrical

12
Q

How many types of S orbitals are there?

A

1

13
Q

What is the maximum number of electrons S orbitals hold?

A

2

14
Q

What is the maximum number of electrons P orbitals hold?

A

6

15
Q

What are the 3 types of P orbitals?

A

Px
Py
Pz

16
Q

What are the 3 P orbitals described to be?

A

degenerate, i.e. have the same energy

17
Q

What is the maximum number of electrons D orbitals hold?

A

10

18
Q

What are the 5 types of D orbitals?

A
dxy
dxz
dyz
dx2-y2
dz2
19
Q

What are the 5 D orbitals described to be?

A

degenerate, i.e. have the same energy

20
Q

Describe the principal quantum number

A

n
Tells us the main energy level
n has values from 1 to infinity
first shell n=1

21
Q

Describe the angular momentum quantum number

A

This describes the shape of the orbital and tells us which subshells are present in the principal shell
l=0 - S
l=1 - p
l=2 - d

22
Q

Describe the magnetic quantum number

A

Gives u the number of each type of orbital
m has values from -l to +l
l= 0 = one S orbital
l=1 = -1,0,1 = 3 p orbitals

23
Q

What do electrons appear to do?

A

spin on their axis

24
Q

What do spinning charges create?

A

a magnetic field

25
Q

Describe the spin quantum number

A

gives the direction of electron spin and has values of s= +1/2 or -1/2

26
Q

What must each electron in the same orbital have?

A

opposite spins

27
Q

How can electron arrangements be determined?

A

by placing electrons in ‘boxes’, where each box represents an atomic orbital

28
Q

What are the 3 principles must be considered when determining the order of filling of orbitals?

A
  1. Aufbau principle
  2. Hund’s Rule
  3. Pauli’s Exclusion Principle
29
Q

What is Aufbau principle?

A

Electrons occupy the lowest energy level available i.e. those closest to the nucleus

30
Q

What is Hund’s Rule?

A

Electrons occupy degenerate orbitals singly before any orbital gets a second electron

31
Q

What is Pauli’s Exclusion Principle?

A

the maximum number of electrons in any atomic orbital is two and if there are two electrons in an orbital they must have opposite spins

32
Q

What is the order of filling orbitals?

A

1s 2s 2p 3s 3p 4s 3d 4p

33
Q

What does shortened electronic configuration use?

A

the notation of the preceding noble gas in place of filled inner shells

34
Q

What are the chemical properties of an element dictated by?

A

the electrons in the outer shell

35
Q

The periodic table is divided into blocks depending on what?

A

the type of orbital which holds the outer electrons

36
Q

What is the first ionisation energy?

A

the energy required to remove one mole of electrons from one mole of gaseous atoms

37
Q

What two factors affect I.E?

A

atomic radius and nucleur charge

38
Q

Describe I.E for across a period

A

nucleur charge increases/ same number of shells

39
Q

Describe I.E for down a group

A

increasing atomic radius

40
Q

Across period 2 what are the 2 anomalies?

A

1st I.E. shows a dip from

  • . Be to B (group 2 to group 3)
  • N to O (group 5 to group 6)
41
Q

How can the variations in ionisation energies can be explained?

A

in terms of the relative stabilities of the electronic configuratons of the element

42
Q

Explain the anomalies in I.E. for group 2 to group 3 (Be to B)?

A

In Be, the electron is removed from the 2s orbital, which is a FULL sub-shell. This is a STABLE ARRANGEMENT, therefore more energy is required to remove the electron.
In B, the electron is removed from a 2p orbital which is further from the nucleus, therefore electron requires less energy to be removed

43
Q

Explain the anomalies in I.E. for group 5 to group 6 (N to O)?

A

In N, all p orbitals are HALF-FILLED. This is a relatively STABLE ARRANGEMENT, therefore more energy is required to remove an electron
In O, there are 2 paired electrons in one of the 2p orbitals. Repulsion between the paired electrons makes the electron easier to remove. Also, the stable half-filled arrangement is lost.