Basics + Relative Mass
A - Mass (total number of protons + neutrons)
Z - Atomic (number of electrons)
- Isotopes are atoms of the same element with different numbers of neutrons.
- The Bohr model explains why noble gases are inert; The shells of an atom can only hold fixed numbers of electrons.
- When an atom has full shells of electrons, it is stable and doesn't react.
- Masses of atoms are compared to Carbon-12.
- Relative atomic mass : The average mass of an atom of an element on a scale where an atom of Carbon-12 is 12.
- Relative isotopic mass - The mass of an atom of an isotope of an element on a scale where an atom of Carbon-12 is 12.
- Relative molecular mass - Average mass of a molecule or formula unit on a scale where an atom of Carbon-12 is 12.
Mass Spectrometer: 1) Vaporisation - The sample is turned into a gas using an electrical heater. 2) Ionisation - The gas particles are bombarded with high-energy electrons to knock off electrons and ionise them. (M(g) > M+(g) + e-)
3) Acceleration - The positive ions are accelerated using an electric field.
4) Deflection - The positive ions paths are altered with a magnetic field. Lighter ions have less momentum so are deflected more than heavier ions. (Only ions with a particular mass/charge ratio will make it to the detector)
5) Detection - Magnetic field strength is slowly increased and different ions (with a lower mass/charge ratio) can reach the detector, producing a mass spectrum.
Calculating Ar from mass spectrum: 1) Read the % relative isotopic abundance and relative isotopic mass for each peak. Multiply together to get the total mass for each isotope. (% abundance x Mass = z)
2) Add up the totals for all the peaks (z + z + z = Y)
3) Divide by 100 since % is used (Y/100 = Ar) or if abundance is not %, divide by total relative abundance (Y/abundance for peak 1 plus abundance for peak 2)
Calculating Mr: The mass of M+ which is formed during ionisation (M+1)
- Electron shells are given a principle quantum number. The further a shell is from the nucleus, the higher its energy and larger its principle quantum number.
- Not all electrons in a shell have exactly the same energy, so they're divided into sub-shells with different numbers of orbitals.
Subshell s; p; d; r
Number of Orbitals 1; 3; 5; 7
Max. number of Electrons (1x2) = 2; (3x2) = 6; (5x2) = 10; (7x2) = 14
- The 2 electrons in each orbital spin in opposite directions as they're both negatively charged so repel eachother; Electrons fill the lowest energy sub-shells first and also singularly first before they start sharing an orbital.
- Cr and Cu donate one 4s electron to the 3d sub-shell because a full or half-full sub-shell is more stable; Electronic structure decides the chemical properties of an element; Groups 4-7 can share electrons when they form covalent bonds.
First ionisation energy - The energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms. (M(g) > M+(g) + e-)
Factors affecting ionisation energy:
1) Nuclear charge - The more protons there are in the nucleus, the more positively charged it is so the stronger the attraction for the electrons, therefore ionisation energy will be higher;
2) Distance from Nucleus (atomic radius) - An electron close to the nucleus will be more strongly attracted to it, making it harder to lose;
3) Shielding - As the number of electron shells between the outer electrons and the nucleus increases, the outer electrons are less attracted to the nucleus.
High ionisation energy means there is a high attraction between the electrons and the nucleus.
Trends in ionisation energy:
1) ionisation energy decreases down group 2 due to the increase in electron shells, so there is increased shielding and atomic radius.
2) ionisation energy increases across a period due to the increasing number of protons, so an increase in nuclear charge. Little extra sheilding/distance.
3) The drop in ionisation energy between Mg and Al - Aluminiums outer electron is in a 3p-orbital, rather than 3s, which has a higher energy than 3s so the electron is found further from the nucleus. Also 3p has extra shielding provided by the 3s2 electrons. Therefore ionisation energy is lower.
4) The drop in ionisation energy between P and S - Shielding is identical and electrons are removed from the same 3p-orbital, but in Phosphorus the electron is being removed from a singly-occupied orbital, whereas in Sulfur the electron is being removed from an orbital containing 2 electrons. The electron repulsion means the electrons are easier to remove from shared orbitals.