Topic 2 : Atomic Structure

Question 1

History of the atom

Answer 1

1. Atom comes from the Greek word indivisible. They were thought for thousands of years to be the smallest building block of everything around us. Often they were represented as tiny spheres.
2. John Dalton recognised that different types of
   atoms must be responsible for the different
   chemical elements.
3. However thanks to J.J. Thomson, Ernest Rutherford and James Chadwick we now know that atoms are made up of sub-atomic particles called protons, neutrons and
   electrons.

Question 2

Developing models of an atom (in order)

Answer 2

1. Daltons model (solid sphere)
2. Thomsons model (plum-pudding)
3. Rutherford model (Nuclear model)
4. Bohr Model (planetary model)

Question 3

Solid Sphere model

Answer 3

1. Dalton thought that all matter was made of tiny particles called atoms, which he imagined as tiny solid balls.
2. It was believed that atoms cannot be broken down into anything simpler
3. the atoms of a given element are identical to each other
4. the atoms of different elements are different from one another
5. during chemical reactions atoms rearrange to make different substances

Question 4

Plum-pudding model

Answer 4

1. J.J. Thomson discovered the electron. Atoms are neutral overall, so in Thomson’s ‘plum pudding model’:
2. atoms are spheres of positive charge
3. electrons are dotted around inside
4. did not have a nucleus

Question 5

Nuclear Model

Answer 5

1. Geiger Marsden tested the plum pudding model.
2. They aimed beams of positively-charged particles at very thin gold foil.
3. These particles should have passed straight through, according to the plum pudding model. However, many of them changed direction.
4. Ernest Rutherford explained these results in his ‘planetary model’:
5. This model states that atoms have a central, positively charged nucleus with most of the mass (positive charge in the middle
6. It also explains that electrons orbit the nucleus, like planets around a star

Question 6

Atoms

Answer 6

1. Atoms contain a positively charged dense nucleus composed of protons and neutrons
2. Negatively charged electrons occupy the space outside the nucleus.

Question 7

Protons

Answer 7

1. located inside of the nucleus
2. charge of +1
3. Relative Mass of 1

Question 8

Neutrons

Answer 8

1. located inside of the nucleus
2. charge of 0
3. Relative Mass of 1

Question 9

Electrons

Answer 9

1. orbiting the nucleus
2. charge of -1
3. Relative Mass of -0.005

Question 10

Proton Number

Answer 10

1. Proton number (commonly referred to as atomic number) indicates how many protons there are in the nucleus. Each element has a unique number of protons.
2. Uranium for example has 92 protons in its nucleus.
3. Located at the bottom of the box

Question 11

Nucleus Number

Answer 11

1. Located at the top of the box
2. Nucleon number (more commonly referred to as mass number) indicates how many nucleons there are in the nucleus.
3. A nucleon is a particle in the nucleus, so more simply it is just the number of protons AND neutrons.

Question 12

Isotopes

Answer 12

1. Atoms of the same element with the same atomic number/Z/same number of protons, but different mass number/A/different number of neutrons.

Question 13

Mass Spectrometer

Answer 13

1. Used to determine the relative atomic mass of an element from its isotopic composition
2. Vaporised sample (gas) is brought in the mass spectrometer
3. The mass spectrometer sorts out the ions based on their mass/charge ratio.
4. The relative height of the peaks tells you the relative abundance of each isotope.
5. Steps in the mass spectrometer are as followed in order : ionisation, acceleration, deflection, detector

Question 14

Ionisation in the mass spectrometer

Answer 14

1. The element is fired with electrons at high speed, causing it to be ionized → like negative charges of firing electrons + negative electrons which causes it to repel, removing the electrons from the atom and making the ion positive
2. e.g. X + e- → X+ + 2e-

Question 15

Atomic Mass Units

Answer 15

1. AMU
2. 1/12th of the mass of a carbon – 12 atom in its ground state. This is used to express masses of atomic particles.
3. 1 AMU = 1.6605402 x 10-27 kg

Question 16

Relative atomic mass

Answer 16

1. Ratio of the average mass of the atom to the unified atomic mass unit
2. “the weighted average of the isotopes of the atoms of an element relative to the carbon-12 isotope”.

Question 17

Acceleration in the mass spectrometer

Answer 17

A negatively charged plate accelerates the ions

Question 18

Deflection in the mass spectrometer

Answer 18

1. An electromagnet deflects the ions - heavier ions deflect less than light ions
2. Some ions are too heavy that they don’t deflect enough, while too light ions deflect too much → need to find the perfect mass
3. mass to charge ratio=M/Z
   M=mass, Z=charge
   The larger the mass to charge ratio, the less it deflects in a mass spectrometer

Question 19

Detection in the mass spectrometer

Answer 19

1. A detector is used to find the percentage of isotopes of different masses
2. This allows the relative atomic mass to be determined

Question 20

Calculations to find relative atomic mass

Answer 20

1. A =(isotope abundance percentage * atomic mass)+(isotope abundance percentage * atomic mass)…
2. you add the products of the abundance and their isotopic mass for each isotope

Question 21

Strong Nuclear Force

Answer 21

1. Like charges repel and opposite charges attract so it doesn’t make sense how atoms can exist without the nucleus repelling and breaking apart
2. The strong nuclear force is what allows a nucleus to remain stable despite the repulsion between the protons.
3. You can think of it being like the glue that holds a nucleus together.
4. A very small percentage of a nucleons mass is converted into energy which helps to ‘glue’ the nucleus together.
5. Neutrons also help with the process by reducing the repulsion between protons.
6. This is why no stable isotopes of helium can exist without neutrons, the repulsion between the protons would be greater than the strong nuclear force and the nucleus would break apart.

Question 22

quantum model of an atom

Answer 22

1. States that electrons aren’t fixed particles with a definite location but are probability waves.
2. Theoretically an electron in an atom in your hand could be on the other side of the room.
3. According to this model it’s not impossible that electrons that make up the atoms in your body could be affecting electrons in atoms on the other side of the universe.

Question 23

Electromagnetic Spectrum (EMS)

Answer 23

1. Visible light, radio waves, infrared waves (IR), ultraviolet (UV), x-rays and gamma rays are forms of electromagnetic radiation
2. c=vλ where λ is wavelength and v is frequency and c is the speed of light (3.00 x 108)

Question 24

Emission Spectra

Answer 24

1. When element in gaseous state is subjected to high voltage under reduced pressure, the gas will emit a certain light.
2. When passed through a prism, the spectrum is not continuous but a black background with certain line spectrums
3. Emission spectra are produced when photons are emitted from atoms as excited electrons return to a lower energy level

Question 25

maximum electrons on main energy level

Answer 25

2n^2

Question 26

Orbitals

Answer 26

1. Schrodinger and others found that the electrons behaved in strange ways. The electron couldn’t be found in one place at one time.
2. Instead they could only calculate the probability of where an electron would be. This is often imagined as a cloud taking a specific shape.
3. These areas of electron density are called orbitals. They are made up of two electrons with opposite ‘spin’.
4. “regions of space where there is a high probability of finding an electron” → one orbital can hold a pair of electrons
5. The boundaries of an orbital are often defined in probabilities

Question 27

Principle quantum number

Answer 27

1. represented as n
2. represents energy levels
3. Lines get smaller and smaller → less and less energy, at infinity the lines converge
4. Each energy level can be divided into s, p, d and f sublevels of successively higher energies
5. Sublevels contain a fixed number of orbitals

Question 28

Nodes

Answer 28

areas without electron density

Question 29

Pauli Exclusion Principal

Answer 29

1. The Pauli exclusion principle states that each orbital may contain no more than two electrons of opposite spin.
2. It also introduces a property of electrons called spin, which has two states: ‘up’ and ‘down’. The spins of electrons in the same orbital must be opposite, i.e. one ‘up’ and one ‘down’.
3. Spin does not literally mean the electrons are spinning. It is a way of representing complicated mathematical properties of the electron with an easily understood concept.

Question 30

Hund’s rule

Answer 30

1. When filling “degenerate” (meaning the same energy) orbitals, we follow Hund’s rule
2. every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.

Question 31

max electrons in each sub-level

Answer 31

1. s - 2
2. p - 6
3. d - 10
4. f - 14

Question 32

sub-levels in each main energy level

Answer 32

1. n = 1 - 1s
2. n = 2 - 1s, 2s, 2p
3. n = 3 - 1s, 2s, 2p, 3s, 3p, 3d
4. n = 4 - 1s, 2s, 2p, 3s, 3p, 3d, 4s, 4p, 4d, 4f

Question 33

spin diagram

Answer 33

A spin diagram shows how the orbitals are filled. Orbitals are represented by squares, and electrons by arrows pointing up or down.

Question 34

Aufbau principle

Answer 34

1. states that electrons fill lower-energy atomic orbitals before filling higher-energy ones
2. Exception to Aufbau principle → 4s has lower energy level than 3d - parallel spins in the d-orbitals help stabilize the electrons and reduce the energy, there is also less repulsion in the 4s orbital. This exception occurs in chromium and copper

Question 35

Quantization of Energy

Answer 35

1. The line spectrums in the line emission have specific wave lengths λ.
2. Each wavelength corresponds to discrete amount of energy. Quantisation is based on this idea that ER comes in discrete packets or “quanta”

Question 36

Calculating Energy evolved

Answer 36

1. The energy evolved (absorbed or emitted) from an electrons transition is called a photon (discrete packet of energy).
2. ΔE = hν = (hc) / λ where
   h = Planck’s Constant = 6.63 x 10^-34 J s
   v = frequency of radiation
   c = speed of light = 3.00 x 10^8

Question 37

Transition metals form coloured compounds

Answer 37

1. When transition metals form compounds, the d-orbital splits into a higher and lower orbital
2. Electrons absorb certain frequencies of electromagnetic radiation to get promoted to higher energy orbitals
3. This make transition metal compounds appear coloured
4. Since Zinc has a complete d-orbital, no electron can jump between the orbital meaning it cannot not form coloured compounds.

Question 38

The hydrogen spectrum

Answer 38

1. Hydrogen atom contains one electron
2. When a hydrogen atom in the ground state receives energy in the form of heat or electricity (ionisation energy), the hydrogen atom is promoted to a higher energy level
3. However it cannot remain at a higher level (excited state) for very long, and falls back to a lower level
4. When the electron falls back down it must lose the energy difference between the two energy levels
5. This loss of energy is performed by releasing electromagnetic energy in the form of infrared, visible light or ultraviolet radiation
6. The movement of electrons between the shells is called an electron transition
7. When electron transitions take place the energy emitted can be detected and its wavelength measured → this provides information about the relative energies of the shells

Question 39

Absorption spectrum

Answer 39

1. coloured background
2. the black lines represent energy with specific frequency absorbed by the electron which corresponds to that electron being promoted from a lower discrete energy level to a higher energy level

Question 40

Emission spectrum

Answer 40

1. black background
2. the coloured lines represent the photon of energy given out when an electron moves from a higher energy level to a lower energy level

Question 41

The difference between a continuous spectrum and line spectrum

Answer 41

1. A continuous spectrum contains all wavelengths
2. A line spectrum contains only selected or certain wavelengths - the analysis of line spectra reveals electronic arrangement