S1.3

Question 1

white light

Answer 1

= light from the Sun

Question 2

physical behaviour of light and subatomic particles

Answer 2

behave as particles (light - photon) and obey the laws of waves

Question 3

energy in relation to wavelength

Answer 3

energy of electromagnetic radiation is indirectly proportional to the wavelength of electromagnetic radiation (E = hγ, c = γλ)

Question 4

regions of the electromagnetic spectrum in terms of increasing energy

Answer 4

radio waves → microwaves → infrared (IR) → visible → ultraviolet (UV) → X-rays → gamma rays

Question 5

types of spectrums

Answer 5

continuos
emission line
absorption

Question 6

continuous spectrum

Answer 6

= spectrum that contains all the frequencies/wavelengths across a range of electromagnetic radiation without any visible lines
shows all wavelengths of visible light

Question 7

emission line spectrum

Answer 7

• = 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)

Question 8

absorption spectrum

Answer 8

= a spectrum of electromagnetic radiation transmitted through a substance, showing dark lines or bands due to absorption at specific wavelengths
opposite of emission spectrum

Question 9

Bohr model of the atom

Answer 9

• 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)

Question 10

transitions of electrons in between energy levels

Answer 10

• 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)

Question 11

hydrogen line emission spectrum

Answer 11

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

Question 12

limitations of the Bohr theory

Answer 12

• 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

Question 13

electron energy levels

Answer 13

1. levels (n … principal quantum number) => principal energy level (higher n => higher orbital energy => greater distance from the nucleus)
2. sublevels (l … second quantum number) => define the shape of the orbitals, four types of sublevels (s, p, d, f)
3. orbitals (ml … third quantum number) => each orbital can contain 2 electrons of opposite spin

Question 14

electron spin

Answer 14

ms … fourth quantum number
indicates orientation
+ ½ ⇒ “spin up”
- ½ ⇒ “spin down”

Question 15

writing electron configuration

Answer 15

Aufbau principle
Pauli exclusion principle
Hund’s rule of maximum multiplicity

Question 16

Aufbau principle

Answer 16

electrons first fill the lowest energy orbitals
ns has lower energy than (n-1)d

Question 17

Pauli exclusion principle

Answer 17

any orbital can hold a maximum of 2 electrons of opposite spin

Question 18

Hund’s rule of maximum multiplicity

Answer 18

electrons fill the orbitals alone, before occupying them in pairs

Question 19

electron configuration of ions

Answer 19

gained/lost electrons should be added/removed from the highest energy level

Question 20

valence electrons

Answer 20

= electrons in the outermost energy level (n)

Question 21

ionization

Answer 21

= the process of removing electrons from an atom in its ground state

Question 22

ionization energy

def, reaction, calculation

Answer 22

= 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

Question 23

first, second, third ionization energy

Answer 23

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]

Question 24

determinants of IE

Answer 24

greater nuclear charge (number of protons) ⇒ greater IE
greater distance distance from nucleus to outer electrons ⇒ smaller IE
shielding effect

Question 25

tendencies of elements with low and high IE?

Answer 25

Low IE ⇒ tendency to lose electrons, i.e. becoming oxidized
High IE ⇒ high value for electronegativity ⇒ tendency to gain electrons, i.e. becoming reduced