Saturday 8 April 2023

Quantum Phenomena

Excitation EnergyIonization Energy

Absorption and Emission Spectra

Energy Levels in the atomAs we considered, the electrons in the atom can only occupy quantized

orbits, i.e. energy levels. All electrons prefer to be in the lowest (unoccupied) energy level (lowest potential energy). That is why an electron at a higher energy level eventually falls in a lower level releasing a photon. We say there is a potential well in the atom where the top of the well is at zero potential and the other energy levels are at negative potential. The ground level will be at the lowest potential energy

Ground State

Higher Energy Levels

E = -13.61 eV

E = -3.40 eV

E = -1.51 eV

E = 0 eV

So, the energy levels are a bit like steps of a ladder that the electrons can climb, if provided with the right energy. What would happen if an electron is given enough energy to climb to the last energy level (E = 0)?

• The electron will escape from the atom, as it has escaped from the potential well that holds it in the atom

• The atom is left with one less electron in its electron cloud, so it becomes a positive ion

• The energy needed for this jump is called IONIZATION ENERGY

Ionization Energy

Ionization Energy is the energy required to cause an electron to escape from the atom

But what can give this energy to the electron to move up the ladder?

• An electric charge (discharge tubes)

• Heating the metal (thermionic emission)

• A photon with just the right energy can hit an electron giving its energy for the electron to jump to the next energy level

Excitation Energy

The energy of the photon must be exactly the same

as the energy gap between the two level

What would happen if a photon of E = 5.00 eV hits an electron in the ground state?

• There is not enough energy to cause the electron to jump to any higher energy level, so the electron stays in the ground state

Excitation Energy

E = -13.61 eV

E = -3.40 eV

E = -1.51 eV

E = 0 eV

Excitation EnergyWhat wavelength photon would you need to cause excitement to the 1st

energy level?• hf = E1 – E2 = 13.61 – 3.40 = 10.21 eV

E1 = -13.61 eV

E2 = -3.40 eV

E3 = -1.51 eV

E = 0 eV

HzJs

J

h

Ef 15

34

19

1046.21063.6

1060.121.10

nmms

ms

f

c1221022.1

1046.2

1000.3 7115

18

In what region of the EM spectrum is the wavelength from the previous question?

U.V.

Excitation Energy

E = -13.61 eV

E = -3.40 eV

E = -1.51 eV

E = 0 eV

What energy photon would you need to ionize the atom?

DE = 13.61 – 0 = 13.61 eV

Hydrogen is the simplest atom and so are its energy levels. When its only electron is excited to higher states, it can fall back to lower energy levels to emit different photons

Hydrogen Emission Spectrum

E = -13.61 eV

E = -3.40 eV

E = -1.51 eV

E = 0 eV

E = -0.85 eV

Lyman Series

Balmer Series

Ground State

1st excited state

What is the difference between the Lyman and the Balmer Series?

• In the Lyman Series the electron “relaxation” produces the emission of U.V. photons

• In the Balmer Series the “relaxation” of the electron produces emission of visible light photons

Hydrogen Emission Spectrum

Hydrogen Emission Spectrum

Emission SpectraWhen the electrons of an atom are exited they can jump to higher energy levels, if the energy provided is just right to make the jump. When they jump back down a photon of energy hf = E1 – E2 is emitted. That is why an atom will emit only light with wavelength (i.e. colour) characteristic of the energy levels in the atom.

Click on each energy level to reveal the emission spectrum

of hydrogen

We considered how excited electrons can fall from a higher energy level to a lower one and emit a photon. But, what gave the electrons the energy to get to those higher levels?

• One way is when an electron absorbs a photon

• The photon must have exactly the same energy as the energy between the two energy levels

Absorption Spectra

What wavelengths of visible light make up white light from, say, an incandescent filament lamp?

• All the wavelengths between 700 – 400 nm (all colours)

So, what would the spectrum of this “white” light be after passing through a gas of, say, hydrogen?

All the wavelengths will be present apart from the wavelengths of the emission spectrum of hydrogen, i.e. Balmer Series. In fact, these have been absorbed by the atoms in the hydrogen gas.

Absorption Spectra

Hydrogen Atoms

Hydrogen Absorption Spectrum

We can study absorption spectra from stars to understand their composition. By looking at the Sun’s absorption spectrum can you tell which element is most abundant?

Hydrogen is the most abundant element in the Sun. In fact, the darkest lines in its spectrum match the wavelengths of the absorption spectrum of hydrogen.

Absorption Spectra and Stars

Sun’s Absorption Spectrum

Hydrogen Absorption Spectrum

The atom and energy levelsClick on different areas to reveal their properties

Li7

3

E1

E2

Electron: atomic particle orbiting around the nucleus. –ve charge and mass 1/1800 of a proton

Neutron: neutral particle inside the nucleus. Different no of neutrons for the same element make different ISOTOPES of that element

Proton: equal and opposite charge as e-. Slightly less mass than a neutron.

Proton no: it tells the no of protons (and also electrons) in the atom. This no is a property of the element, so the atomic no tells what element the atom is.

Nucleon no: given by the sum of no of protons and no of neutrons. Using this no you can identify the different isotopes of the same element.

Ground level: lowest energy state possible for the electrons. Electrons tend to fill this level 1st and “fall” in it after being exited.

Higher energy levels: when not

completely filled with e-, they can

“host” exited e-.

hf = E1 – E2

Exited e - jumps

energy gap

e- falls

bac

k to

grou

nd le

vel

e- falls to a lower level and a photon with

energy hf = E1 – E2 is released.

A photon with energy hf

= E1 – E2 hits the e-, which is exited to the next energy level.

hf = E1 – E2