Quantum Theory Flashcards Preview

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

Flashcards in Quantum Theory Deck (43):
1

What did Niels Bohr predict?

electrons exist in orbits with specific energy

2

What does atomic emission spectra provide evidence f?

subshells within each energy level

3

What does classical theory assume electrons exist as?

a particle (with mass and charge)

4

What does experimental evidence show that electrons can have?

wave-like properties (show diffraction and interference)

5

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

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

6

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

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

7

What does quantum theory suggest? (3)

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

What are orbitals defined by?

Quantum numbers

9

What is an orbital?

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

10

What are the shapes of orbitals?

S
P
D

11

What are S orbitals said to be?

spherically symmetrical

12

How many types of S orbitals are there?

1

13

What is the maximum number of electrons S orbitals hold?

2

14

What is the maximum number of electrons P orbitals hold?

6

15

What are the 3 types of P orbitals?

Px
Py
Pz

16

What are the 3 P orbitals described to be?

degenerate, i.e. have the same energy

17

What is the maximum number of electrons D orbitals hold?

10

18

What are the 5 types of D orbitals?

dxy
dxz
dyz
dx2-y2
dz2

19

What are the 5 D orbitals described to be?

degenerate, i.e. have the same energy

20

Describe the principal quantum number

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

21

Describe the angular momentum quantum number

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

Describe the magnetic quantum number

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

What do electrons appear to do?

spin on their axis

24

What do spinning charges create?

a magnetic field

25

Describe the spin quantum number

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

26

What must each electron in the same orbital have?

opposite spins

27

How can electron arrangements be determined?

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

28

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

1. Aufbau principle
2. Hund's Rule
3. Pauli's Exclusion Principle

29

What is Aufbau principle?

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

30

What is Hund's Rule?

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

31

What is Pauli's Exclusion Principle?

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

What is the order of filling orbitals?

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

33

What does shortened electronic configuration use?

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

34

What are the chemical properties of an element dictated by?

the electrons in the outer shell

35

The periodic table is divided into blocks depending on what?

the type of orbital which holds the outer electrons

36

What is the first ionisation energy?

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

37

What two factors affect I.E?

atomic radius and nucleur charge

38

Describe I.E for across a period

nucleur charge increases/ same number of shells

39

Describe I.E for down a group

increasing atomic radius

40

Across period 2 what are the 2 anomalies?

1st I.E. shows a dip from
-. Be to B (group 2 to group 3)
- N to O (group 5 to group 6)

41

How can the variations in ionisation energies can be explained?

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

42

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

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

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

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.