topic 1 Flashcards
relative mass and charge of protons, neutrons & electrosn
proton : 1, +1
neutron : 1, 0
electron : 0.005, -1
atomic number
number of protons in an atom
mass number
total number of protons and neutrons in an atom
isotopes
atoms with the same number of protons but different number of neutrons
relative isotopic mass
mass of one atom of an isotope compared to one twelfth of the mass of one atom of carbon-12
relative atomic mass
the avg mass of an atom compared to 1/12th of the mass of a carbon-12 atom
relative molecular mass
the average mass of a molecule compared to one twelfth of the mass of one atom of carbon-12
relative atomic mass calculation from mass spectrometry
R.A.M =
(isotopic mass x % abundance) / 100
if relative abundance is used then:
R.A.M = (isotopic mass x relative abundance) / total relative abundance
predicting mass spectra for diatomic molecules including Cl
mass spectrometry of diatomic molecules produces molecules, not individual atoms.
fragmentation occurs. This is when Cl is put through the ionisation chamber, an electron is knocked off. This produces a molecular ion, Cl2+. This ion is unstable so eventually fall apart and give a chlorine atom and a Cl+ ion.
how can mass spectrometry be used to determine the relative molecular mass of a molecule
the heaviest ion - the one with the greatest m/z value - is likely to be the molecular ion.
Finding the relative molecular mass from a mass spectrum you look for the peak with the highest value for m/z, and that value is the relative formula mass of the compound
e.g. if highest m/z value is 72 then the relative molecular mass of that compound will be 72
first ionisation energy
the energy required for one mole of gaseous atoms to form one mole of gaseous ions with a single positive charge
successive ionisation energies
the energy that is required to remove the electron one after the other
factors that affect ionisation energies
- attraction of the nucleus (more protons in the nucleus the greater the attraction - large IE)
- distance of the electrons from the nucleus (weaker attraction for bigger atoms - small IE)
- shielding of the attraction of the nucleus (bigger the atom, more shielding, outer electrons repelled by inner shells - lower IE)
why is there a general increase in first IE across a period?
because the proton number increases which makes the attraction to the nucleus greater. Electrons are being added to the same shell which therefore has the same shielding effect and electrons are pulled in closer to the nucleus
why do first ionisation energies decrease down a group?
as you go down the group, the outer electrons are further away from the nucleus so the attraction from the nucleus decreases
why are successive ionisation energies always larger
second IE is always larger than the first. this is because when the first electron is removed, a positive ion is formed. this increases attraction on the remaining electrons, so the energy required to remove the next electron is larger
why does helium have the largest first ionisation energy
The first electron is in the shell closest to the nucleus, so has no shielding effects from the inner shells. He has a bigger first IE than H as it has one more proton
why does Na have a much lower first ionisation energy then Neon?
Na will have electrons in a 3s sub-shell further from the nucleus and is more shielded. So, Na’s outer electron is easier to remove and has a lower IE
Why is there a small drop from Mg to Al?
Al starts to fill 3p sub shell whereas Mg has its outer electrons in the 3s sub shell. The 3p electrons are slightly easier to remove because they are higher in energy and also less shielded than 3s electrons
Why is there a small drop form P to S?
Sulphur has 4 electrons in the 3p sub-shell and the 4th is starting to doubly fill the first 3p orbital. When a second electron is added to 3p orbital, there is slight repulsion between the two electrons, hence making it easier to remove the second electron
how did ideas about electronic configuration develop
- atomic emission spectra provide evidence for the existence of quantum shells
- successive ionisation energies provide evidence for the existence of quantum shells and the group to which the element belongs
- first ionisation energy of successive elements provides evidence for electron sub-shells
orbital
a region within an atom that can hold up to two electrons with opposite spins
s, p, d & f blocks
s - 2 electrons/ orbital
p - 6 electrons/ orbital
d - 10 electrons/ orbital
f - 14 electrons/ orbital
periodicity
a repeating pattern across a period
