Atomic Structure And The Periodic Table Flashcards
mass of an electron
1/1840
mass number
total number of particles in the nucleus
atomic number
number of protons in the nucleus
number of neutrons
mass number - atomic number
isostopes
atoms of the same type with a different number of neutrons
relative atomic mass
the average mass of an atom of an element, relative to 1/12th the mass of as carbon-12 atom
relative isotopic mass
the mass of one mole of an isotope of the element expressed on a scale on which one mole of the atoms of the C-12 isotope has a mass of exactly 12 units.
relative molecular mass
Average mass of a molecule of a substance relative to 1/12th the mass of a carbon 12 atom
determining relative atomic mass from mass spectrometer
- the sample is vaporised
- the sample is ionised
- ion is accelerated by an electric field
- ions pass through the magnetic field and are deflected
- ions are detected and recorded
the size of the peaks in mass spectra
number/abundance of each type of ion
chlorine mass spectra
chlorine has two isotopes at mass 35 and mass 37, the height of 35 trace is three times that of 37 so its three times more abundant
first ionisation energy
energy required to remove one electron from one moles of gaseous atoms to produce one mole of 1+ gaseous ions
attraction of the electron to the nucleus
· The charge on the nucleus – the higher the charge, the greater the attraction
· The distance from the nucleus - the greater the distance, the less the attraction
· Shielding from inner electrons – the more inner electrons, the less the attraction.
reasons for first ionisation energy to increase across a period
- same number of shells
- increased nuclear charge
- greater attraction to the nucleus
reasons for first ionisation energy to decrease down a group
- increasing number of shells
- greater shielding
- decreased attraction to the nucleus
second/successive ionisation energy
The energy required to remove one electron from each ion in one mole of gaseous ions with a single positive charge producing one mole of gaseous ions with two positive charges
in shell one what are the sub shells
1s2
in shell two what are the sub shells
2s2 2p6
in shell three what are the sub shells
3s2 3p6 3d10
orbital
The area occupied by an electron wave
s orbital
2 electrons
p orbital
6 electrons in 3 orbitals
d orbital
10 electrons in 5 orbitals
how do electrons fill subshells
singly but then pair up and two electrons in the same orbitals have to have opposite spins
how does a flame test work
electrons in the sodium atoms gain energy from the heat and jump to a higher energy level. They then drop back down to a lower energy level giving out energy, this time in the form of visible light. This is known as emission.
group 4 bonding
tend to form giant covalent structures, although the elements towards the bottom of the group are metallic. Lead and tin are metallic because their large atoms have electrons a long way from the nucleus with a great deal of inner shielding which means they can be released quite easily
group 5, 6, 7 bonding
generally have simple covalent structures, although again at the bottom of the group the elements become more metallic.
group 8 bonding
noble gases, having complete electron shells are monatomic
periodicity
repeating pattern across different periods
periodicity in period 2, 3 metallic
metallic and boiling points increase as you go along
- the positive charge on the metal ion increase
- greater charge density
- The number of electrons in the delocalized region increases
- the attraction between the ion and the delocalized electron region increases.
periodicity in period 2, 3 giant molecular
melting and boiling points are at their highest as vast numbers of strong covalent bonds must be broken to release the atoms.
periodicity in period 2, 3 simple molecular
melting and boiling points are low
- intermolecular forces are weak
- roughly relative to the size of the molecules
- the noble gases have the lowest melting points as they are monatomic.
