white light
= light from the Sun
physical behaviour of light and subatomic particles
behave as particles (light - photon) and obey the laws of waves
energy in relation to wavelength
energy of electromagnetic radiation is indirectly proportional to the wavelength of electromagnetic radiation (E = hγ, c = γλ)
regions of the electromagnetic spectrum in terms of increasing energy
radio waves → microwaves → infrared (IR) → visible → ultraviolet (UV) → X-rays → gamma rays
types of spectrums
continuos
emission line
absorption
continuous spectrum
= spectrum that contains all the frequencies/wavelengths across a range of electromagnetic radiation without any visible lines
shows all wavelengths of visible light
emission line spectrum
- = range of frequencies/wavelengths of electromagnetic radiation emitted during an electron transition from a higher to a lower energy state
- energy is applied to a sample of hydrogen gas, which is then viewed through a spectroscope ⇒ a distinctive pattern (emission line spectrum of hydrogen)
- elements can be identified by their unique emission spectra
- shows only specific wavelengths of light (coloured lines on a black background)
- coloured lines get closer (converge) at high energy (short wavelength)
absorption spectrum
= a spectrum of electromagnetic radiation transmitted through a substance, showing dark lines or bands due to absorption at specific wavelengths
opposite of emission spectrum
Bohr model of the atom
- electrons are located in energy levels at fixed distances from the nucleus
- each level is assigned the letter n (principal quantum number)
- ground state ⇒ n = 1, closest to the nucleus
- n increases, so does the distance from the nucleus and its energy
- these levels begin to converge (border of an atom) at higher energy levels (higher n) because main energy levels also begin to converge
- an electron has a larger energy level ⇒ may transition to n = ∞ (become free, ionising the atom)
transitions of electrons in between energy levels
- electrons transition between energy levels by either absorbing or emitting energy (photons)
- absorption of energy ⇒ transition from a lower energy level to a higher energy level ⇒ the excited state of the electron (unstable relative to the ground state)
- emission of energy ⇒ transition from a higher energy level to a lower energy level (energy emitted corresponds to the wavelengths of visible light (E = hγ)
hydrogen line emission spectrum
Bohr ⇒ the electron in a hydrogen atom moves in a circular path (orbit) and electrons can only exist within these orbits, not between them
each orbit has a definite energy: E = -RH/n^2 (RH … Rydberg constant)
electron transitions to ground-state (n=1) ⇒ emission of UV radiation
electron transitions to n=2 ⇒ emission of visible light
electron transitions to n=3 ⇒ emission of infrared (IR) radiation
limitations of the Bohr theory
- limitation to one electron per atom ⇒ only the hydrogen atom
- assumes the electron is a subatomic particle in a fixed orbit
- does not account for the effect of electric and magnetic fields on the spectral lines of atoms and lines
- cannot not explain molecular bonding and geometry
electron energy levels
- levels (n … principal quantum number) => principal energy level (higher n => higher orbital energy => greater distance from the nucleus)
- sublevels (l … second quantum number) => define the shape of the orbitals, four types of sublevels (s, p, d, f)
- orbitals (ml … third quantum number) => each orbital can contain 2 electrons of opposite spin
electron spin
ms … fourth quantum number
indicates orientation
+ ½ ⇒ “spin up”
- ½ ⇒ “spin down”
writing electron configuration
Aufbau principle
Pauli exclusion principle
Hund’s rule of maximum multiplicity
Aufbau principle
electrons first fill the lowest energy orbitals
ns has lower energy than (n-1)d
Pauli exclusion principle
any orbital can hold a maximum of 2 electrons of opposite spin
Hund’s rule of maximum multiplicity
electrons fill the orbitals alone, before occupying them in pairs
electron configuration of ions
gained/lost electrons should be added/removed from the highest energy level
valence electrons
= electrons in the outermost energy level (n)
ionization
= the process of removing electrons from an atom in its ground state
ionization energy
def, reaction, calculation
= the minimum energy required to remove one mole of electrons from one mole of gaseous atoms in the ground state
X(g) → X(g)+ + e– (endothermic reaction)
calculated using the formula: E = hγ
first, second, third ionization energy
first ionization energy = the energy required to remove 1 mole of the first electron
second ionization energy = removal of the second mole of electrons
third ionization energy = removal of the third mole of electrons
1st IE < 2nd IE < 3rd IE < …
[kJ/mol]
determinants of IE
greater nuclear charge (number of protons) ⇒ greater IE
greater distance distance from nucleus to outer electrons ⇒ smaller IE
shielding effect
