IAL Chemistry Topic 2 Flashcards
Isotope
Isotopes are different atoms of the same element that have the same number of protons but a different number of neutrons
Atomic number
Atomic number is the number of protons in the nucleus of an atom.
Mass number
Mass number is the sum of the number of protons and the number of neutrons in the nucleus of an atom.
Relative isotopic mass
Relative isotopic mass is the mass of an atom of a particular isotope relative to 1/12 mass of a carbon-12 atom.
How does a mass spectrometer work?
Vaporise sample into gas → passes through ionisation chamber to make positive ions → pass through electric field which are accelerated → pass through the magnetic field which are deflected (different paths) → detected by the ion detector
Electronic configuration
The electronic configuration of an atom of an element is the distribution of electrons among atomic orbitals.
Ionisation energy
Ionisation energy is the energy required to remove one electron from each atom in one mole of gaseous atoms of the element.
Energy of electrons when removed - energy of electrons in the orbital
Trend of ionisation energy
The more electrons removed, the less electron-electron repulsion so the ionisation energy increases.
- across a period there is an increase in the first ionisation energy (more protons)
- down a group the first ionisation energy decreases
Electron-electron repulsion effect
raises the energy in the shell of the electrons to the value they would have if there was no repulsion between them –> lower ionisation energy
Factors that affect the energy of an electron
- the orbital in which the electron exists
- nuclear charge of the atom (number of protons)
- repulsion experienced by the electron from the other electrons in the atoms (high energy orbital increases repulsion / doubly occupied orbital increases repulsion)
- increase repulsion causes decrease in ionisation energy
Which group has the highest energy in s,p,d orbital?
s orbital: group 1 and 2
p orbital: group 2-8
d orbital: transition metals
How to measure the atomic radius?
→ distance from the nucleus to the boundary of an electron cloud. (orbitals → definite edge)
→ distance between 2 nuclei and divide it by two
Type of radius and their type of atoms
Covalent radius (two bonded atoms)
Van Der Waals’ radius (two non-bonded atoms)
Metallic radius (metal atoms)
Trend of atomic radii across a group
Atomic radii decrease across a period.
Electrons are being added to the same energy level and protons are being added to the nucleus.
This resulted in an increased nuclear charge (as the number of protons increases) and therefore the attractive force between the nucleus and electrons increases.
This force counterbalances the increased electron-electron repulsion that would occur due to the additional electrons in the quantum shells.
Trend of atomic radii down a group
Atomic radii increase down a group
The number of occupied quantum shells increases making atoms bigger
Melting and boiling points of different structures
Elements with giant lattice structures have high melting and boeing points.
Elements with simple molecular structures have low melting and boiling points.
Metallic → high
Covalent lattice → very high
Covalent molecular → low
Trend of first ionisation energy across a period
Increase in the first ionisation energy.
More protons are being added to the nuclei of the atoms. This results in an increase in nuclear charge.
The electrons in the outer energy levels will be more tightly held (decrease energy IN SHELL) and more difficult to remove.
The increase in nuclear charge is stronger than the increase in electron-electron repulsion. Thus, increased attraction, higher ionisation energy.
Trend of first ionisation energy down a group
First ionisation energy decreases.
Electrons removed from the outer energy level are increasingly distant from the nucleus as a result of the additional quantum shell. The attraction of the positive nucleus is decreased and it becomes easier to remove.
The outer electrons experience increased repulsion from the inner electrons which leads to a decreased ionisation energy.
First ionisation anomalies
- First ionisation of boron is lower than that of beryllium. Although the nuclear charge of the boron atom is greater than that of the beryllium atom, the outer electron of boron has more energy as it is in a 2p orbital as opposed to the 2s orbital for beryllium.
- First ionisation of oxygen is lower than that of nitrogen. Although the nuclear charge of the oxygen atom is greater than that of the nitrogen atom, the electron-electron repulsion decreases the first ionisation energy of oxygen atom.
