Topic 2/12: Atomic Structure Flashcards
Describe Rutherford’s experiment.
- Shot alpha particles straight towards a sheet of gold foil
- A small % of particles were deflected through angles larger than 90 degrees/scattered back
- Conclusion: atom was mostly empty space, atoms have a nucleus
Define mass number.
Sum of the number of protons and neutrons in a nucleus
Define atomic number.
The number of protons in the nucleus. Equal to the number of electrons.
Define isotopes.
Atoms of the same element with the same number of protons but different number of neutrons.
Describe the properties of isotopes.
Chemical properties: Remains the same as the number of electrons is the same
Physical properties: Differs as the number of neutrons changes. E.g. density, MP, BP, rate of diffusion
Describe the use of radioisotopes.
Carbon-14
- Radiocarbon dating: estimate age of organisms
- Treat cancerous cells
Cobalt-60
- emits gamma rays
- treat cancerous cells
- sterilize foods
Iodine-131
- Treat thyroid cancer/prostate cancer/brain tumours
Describe the operation of a mass spectrometer.
Use: Separates individual isotopes from a sample of atoms and determines the mass of each isotope
- Vaporization
- Ionization: atoms bombarded by high energy electrons, resulting in ions of 1+ charge
- Acceleration: attraction of 1+ ions to negatively charged plates
- Deflection: pass through strong magnetic field
- Detection: measures location and number of particles
Factors determining deflection of ions in a mass spectrometer.
- Absolute mass
- Charge
- Strength of magnetic field
- Velocity (speed)
Describe the electromagnetic spectrum from low to high energy.
Radio waves
Microwave infrared
Visible
Ultraviolet
X-ray
Gamma-ray
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Define first ionization energy.
The minimum amount of energy required to remove one mole of electrons from one mole of gaseous atoms.
Describe the trend of ionization energy.
Across a period
- Increases due to greater nuclear charge + similar shielding effect ➜ valence electrons experience greater pull
Down a group
- Decreases due to valence electrons being further away from the nucleus ➜ less tightly held
Why does nitrogen (period 5) have a higher IE than oxygen (period 6)?
Nitrogen has a half-filled 2p, making it more stable than oxygen. The increased electron-electron repulsion in oxygen makes it less stable than nitrogen.
Why does beryllium (period 2) have a higher IE than boron (period 3)?
The electron in BE is removed from the s-orbital while the electron in B is removed from the p-orbital (higher energy level), which counteracts the stronger nuclear charge of B
Factors that influence ionization energy.
Size of nuclear charge
- proton number increases ➜ greater nuclear charge ➜ valence electrons experience greater pull ➜ IE increases
Distance of valence electrons from the nucleus
- further from nucleus ➜ less tightly held ➜ IE decreases
Shielding effects
- electrons in full inner shells repel electrons in outer shells
- greater shielding by full inner shells ➜ lower electrostatic attractive forces ➜ IE decreases
Explain Bohr’s model and its limitations.
- electrons orbit the nucleus in ring-like paths at fixed energy levels ➜ The higher the energy level, the further from the nucleus
- electrons can only occur in one energy level or another but nothing in between
- electrons can move from one orbital to another at one time
- electrons absorb energy ➜ enter the excited state ➜ move up a higher energy level
Weakness
- could not explain why only certain energy levels were allowed
- only applies to atoms with one electron/hydrogen
- does not represent sub-levels/orbitals
- does not take into account the interactions between atoms
