- Electron - negative (-1) charge, orbit nucleus in orbitals, take up large volume, little mass.
- Proton - positive (1) charge, set in nucleus, mass (1).
- Neutron - no charge (0), mass (1), set in nucleus.
- A - represents mass number (protons plus neutrons).
- Z - represents atomic (proton) number.
- X - represents element symbol.
- Different number of electrons compared to protons.
- Negative - more electrons (than normal).
- Positive - less electrons (than normal).
- Isotopes of an element are atoms with the same number of protons but different numbers of electrons.
- Number of electrons and electron configuration determines chemical propeties, therefore isotope (have the same number and configuration) chemically react the same way.
- But they do change physically because that usually dependant on mass and isotopes have different masses.
Atomic models have changed throughout history as new evidence and new therories are made.
John Dalton - 19th century: Atoms are solid spheres.
J J Thompson - 1897: Weren't solid. Experiments showed they had charge and mass, negative particles (electrons). New model provided "plum pudding" model.
Rutherford - 1909: Gold foil experiement, alpha particles fired at thin gold sheet. They expected most to be deflected slightly by positive sphere.But most passed through and very few were deflected backwards. Hypothesis, Plum Pudding model was wrong. New model needed:
Nuclear model of atom:
- Tiny positive charge in centre (nucleus).
- Surrounded by cloud of negative electrons.
- Mostly empty space.
Niels Bohr: Electrons in a cloud would quickly spiral down into the nucleus causing collapse. So they gave a new model with 4 basic principals:
- Electrons exist in fixed orbits (shells).
- Each shells had a fixed energy.
- When electrons move in between shells electromagnetic radiation is emited or absorbed.
- Electromagnetic radiation has a fixed energy because electron shells have a fixed energy.
Bohr's model fits observations from already known experiments. But they redifined to include sub shells.
- Bohr's model explains inert gasses.
- Shells can only hold a fixed number of electrons, full shells means stable atom.
More than one model in use today for different observations and experiments.
Relative masses are masses of atoms compared to Carbon 12. The acctual mass of an atom is tiny, too small to weigh, so masses are compared. This is relative mass.
- Relative atomic mass (Ar) average mass on scale where carbon 12 is 12.
- Relative isotopic mass mass of an isotopewhere carbon 12 is 12.
- Relative molecular mass (relative formula mass, Mr) average mass of a molecule or formula unit on a scale where carbon 12 is 12.
Mass Spectrometer: Tells you the relative isotopic mass, relative molecular mass, relative isotopic abundance, molecular structure & relative atomic mass.
- Vaporisation (gas)
- Ionisation (bombarded with high-energy electrons)
- Acceleration (electric field)
- Deflection (magnetic field, lighter deflected more, higher charge deflected more)
- Detection (slowly increased, mass spectrum produced)
A mass sprectrum:
- Y axis = abundance.
- Peak height = isotopic abundance.
- X axis = mass:charge ratio (relative isotopic mass).
Using mass spectrometry:
- Relative atomic mass % (calculating):
1. find % relative isotopic abundance for each peak (y-axis) and relative isotopic mass (x-axis). Multiply for each peak.
2. Add up totals.
3. Divide by 100.
- Relative atomic mass (calculating):
1. & 2. the same.
3. Divide by total relative abundance.
- Sub shells.
- Fixed energy shells/ energy levels.
- Principal quantum number (each shell given a number).
- Further away, higher the energy, larger the quantum number.
- Shells have sub levels, which have orbitals which hold up to 2 electrons:
Sub-shell Orbitals Max electrons
s 1 1 x 2 = 2
p 3 3 x 2 = 6
d 5 5 x 2 = 10
f 7 7 x 2 = 14
- the two electrons have opposite spins
Electronic structure decides chemical properties.
What is ionisation: The removal of one or more electrons. Removing the first electron is the first ionisation energy.
1. First ionisation energy is the energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions:
2. Equation - O(g) -> O+(g) + e- (1st ionisation energy = +1314 kJ mol-1)
Important points about ionisation energy:
- Must use state symbol (g).
- Always refer to 1 mole.
- Lower ionisation energy, easier to form ion.
Factors that effect ionisation energy:
- Nuclear charge: More protons, stronger the attraction.
- Distance from nucleus: attraction falls off rapidly with distance.
- Shielding: nimber of electrons between outer electrons and nucleus increases less attraction felt by outer electrons, but they can also cause screening (or sheilding).
High ionisation energy, high attraction between electrons and the nucleus.
Ionisation energy decreases down group 2:
- Provides evidence electron shell exists.
- Each element down group 2 has an extra electron shell, inner shells shield attraction of nucleus.
- Extra shells also mean electron is further away.
- All above show its easier to remove electrons that are further away from nucleus, lowering ionisation energy.
Ionisation energy increases across the period:
- General trend is it becomes harder to remove electrons.
- This is because number of protons increases, mean stronger nuclear attraction.
- all extra electrons have roughly the same energy level.
- Therefore little extra shielding or extra distance effect.
- There are small drops between group 2 & 3 & 5 & 6.
Drops between group 2 & 3 shows sub-shell structure:
- 3p orbital has slightly higher energy than 3s, so on average is found further away from nucleus.
- 3p orbitals has extra sheilding.
- override effect of increased nuclear charge ie. ionisation energy drops.
- Evidence for theory of electron sub-shells.
Drop between 5 & 6 due to electron repulsion:
- Sheilding identical.
- Electron removed from orbital containing 2 electrons.
- Repulsion means electrons easier to remove.
- More evidence for electronic structure models.