Quantum Theory Flashcards
What did Niels Bohr predict?
electrons exist in orbits with specific energy
What does atomic emission spectra provide evidence f?
subshells within each energy level
What does classical theory assume electrons exist as?
a particle (with mass and charge)
What does experimental evidence show that electrons can have?
wave-like properties (show diffraction and interference)
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
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
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
What are orbitals defined by?
Quantum numbers
What is an orbital?
a region of space where the probability of finding the electron is >95% (i.e. there is uncertainty)
What are the shapes of orbitals?
S
P
D
What are S orbitals said to be?
spherically symmetrical
How many types of S orbitals are there?
1
What is the maximum number of electrons S orbitals hold?
2
What is the maximum number of electrons P orbitals hold?
6
What are the 3 types of P orbitals?
Px
Py
Pz
What are the 3 P orbitals described to be?
degenerate, i.e. have the same energy
What is the maximum number of electrons D orbitals hold?
10
What are the 5 types of D orbitals?
dxy dxz dyz dx2-y2 dz2
What are the 5 D orbitals described to be?
degenerate, i.e. have the same energy
Describe the principal quantum number
n
Tells us the main energy level
n has values from 1 to infinity
first shell n=1
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
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
What do electrons appear to do?
spin on their axis
What do spinning charges create?
a magnetic field
Describe the spin quantum number
gives the direction of electron spin and has values of s= +1/2 or -1/2
What must each electron in the same orbital have?
opposite spins
How can electron arrangements be determined?
by placing electrons in ‘boxes’, where each box represents an atomic orbital
What are the 3 principles must be considered when determining the order of filling of orbitals?
- Aufbau principle
- Hund’s Rule
- Pauli’s Exclusion Principle
What is Aufbau principle?
Electrons occupy the lowest energy level available i.e. those closest to the nucleus
What is Hund’s Rule?
Electrons occupy degenerate orbitals singly before any orbital gets a second electron
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
What is the order of filling orbitals?
1s 2s 2p 3s 3p 4s 3d 4p
What does shortened electronic configuration use?
the notation of the preceding noble gas in place of filled inner shells
What are the chemical properties of an element dictated by?
the electrons in the outer shell
The periodic table is divided into blocks depending on what?
the type of orbital which holds the outer electrons
What is the first ionisation energy?
the energy required to remove one mole of electrons from one mole of gaseous atoms
What two factors affect I.E?
atomic radius and nucleur charge
Describe I.E for across a period
nucleur charge increases/ same number of shells
Describe I.E for down a group
increasing atomic radius
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)
How can the variations in ionisation energies can be explained?
in terms of the relative stabilities of the electronic configuratons of the element
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
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.
