Unit 1 – Atomic Structure and the Periodic Table

The Atom
 Atoms are made from proton, neutrons and electrons.
 The protons and neutrons are in the centre of an atom, called the nucleus, which is a
dense structure, also where most of the atom’s mass is.
 The electrons orbit the nucleus in energy levels and have a relative mass of 0
compared to the protons and neutrons. (The mass is so small we can ignore it.)
o This means that the mass of an atom is calculated by the number of protons
and neutrons.
 Protons have a charge of +1, electrons have a charge of -1 while neutrons have no
relative charge.
o Therefore, we can conclude that the nucleus is a positively charged structure.
 Ions have different number of electrons.
o Positive ions have fewer electrons than protons.
o Negative ions have more electrons than protons.
 Isotopes are atoms of the same element, with the same number of protons and
electrons, however they vary in the number of neutrons.
o This means that isotopes can have different masses (since neutrons are in the
nucleus.)
o They also have slightly different physical properties, as they depend on the
mass of an atom.
 The relative atomic mass is the average mass of an atom compared to 1/12th of the
mass of a carbon-12 atom.
 The relative isotopic mass is the average mass of an isotope compared to 1/12th of the
mass of a carbon-12 atom.
 The relative molecular mass is the average mass of a molecule compared to 1/12th of
the mass of a carbon-12 atom.
 The relative mass of isotopes and the relative isotopic abundance can help us work out
the relative atomic mass of the element.
o Multiply each of the relative isotopic mass by the isotopic abundance and add
up the values.
o Divide by 100 (since abundance is calculated in percentage)
 The relative isotopic mass and the abundance of isotopes can be collectively presented
in a graph. The graph is called a mass spectrum.
  The relative atomic mass can be worked out by multiplying each of the m/z by it’s
corresponding relative abundance and doing it for all the other ones, then finally
dividing it by 100 since abundance is worked out in percentage.
 Once you work out the relative atomic mass, you can essentially work out what
element the spectrum is corresponding for.
 The mass spectra of a molecule can be given, and the same procedure can be used.
o Work out the total relative molecular mass and see which molecule it
makes us.
o Most of the times the type of molecule or the elements in the molecule are
given so it won’t be hard.
Electron Structure
 In atoms, the electrons orbit the positively charged nucleus in electron shells.
 Electron shells themselves are made from subshells and orbitals.
o The shells contain subshells, which can have different number of orbitals,
each holding up to a maximum of 2 electrons.
 Subshell is the name given to a collective group of orbitals.
 Each subshell can hold different amounts of orbitals.
 For example, electron shell 2 has subshells, S and P. The S subshell has a total of
1 s orbital and the P subshell has a total of 3 p-orbitals, x, y and z.
 Each orbital can only hold 2 electrons, so since electron shell 2 has a total of 4
orbitals, it can hold a total of 8 electrons.
o S subshell holds 1 S-orbital
o P subshell holds 3 P-orbitals
o D subshell holds 5 D-orbitals (shape not needed to know).
 This diagram below illustrates how the orbitals look like (again, the shape of a D-
orbital is not required.)
 It’s also required to know that the electrons inside those orbitals spin in opposite
directions, known as spin-pairing.

 Electron configuration can be shown using boxes and arrows.

 There are however rules to how electrons fill up energy shells in an atom.
o They always go from the lowest energy level to the highest, from the s-
subshell to the p-subshell to the d-subshell.
o When it is in the p and the d subshell, only half of the orbitals are filled and
once all the orbitals are half filled, it goes back to the start and starts
completing the orbitals.
o Chromium and Copper have an exception where the 4s subshell is only half
filled and the extra electron is donated to the 3d subshell.
o The 4s subshell has a lower energy level than the 3d subshell, meaning that
electrons will fill up the 4s subshell before moving on to the 3d subshell, even
though it has a greater quantum value.
Emission Spectra – Proof of shells
 When an atom is excited it means it has gained extra energy in a process called
excitation.
 This means that the electrons begin jumping between the energy levels and the greater
the energy level they jump to, the greater the energy required.
o This means that more energy is required to jump from level 1 to level 3 than
from level 1 to level 2 purely because the distance of travel is much greater.
 They also release energy in fixed amounts, meaning that they are discrete. This is
because there isn’t space between the levels for the electrons to situated themselves.
o This is also proof of existence for energy levels, or known as quantum shells.
 When an electron moves between shells, electromagnetic radiation is emitted.
However, the energy is fixed for each movement meaning that the radiation produced
has a fixed frequency.
 This means that when you look at a gas using emission spectra, you’ll get a set of
lines which represent energy radiation and they occur at greater frequencies, if the
energy released or absorbed from a movement is greater.
 Since the lines are dashed, it also means that the energy is discrete, and that there is
no space in between shells, so it means that there is proof for the existence for energy
levels.
Ionisation Energy
 Ionisation is the removal of electrons.
 The first ionisation energy if the energy required to remove 1 electron from each atom
in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions.
 There are factors however which affect ionisation energy.
 o The greater the positive charge of the nucleus, the greater the attraction within
the nucleus and the outer electrons (from which an electron is removed)
therefore more ionisation energy is required.
o The greater distance from the nucleus and the outer electron shell, the lesser
the attraction, meaning that less ionisation energy is required.
o The more number of electron shells means that there will more electrons
between the nucleus and the outer energy level. This means that there is
increased shielding, which means that there is repulsion for the outer electrons
away from the nucleus, so the ionisation energy is decreased with more shells.
 The first ionisation energy decreases down a group.
o This is because the electron shells increase, which decreases the attraction
between the nucleus and the outer shell.
o The number of levels also increase meaning that there is greater repulsion to
the outer electrons due to increase in the shielding electrons.
o These two things are both responsible for a lower ionisation energy therefore it
decreases down a group.
 Successive ionisation energy is the energy required to remove electrons continuously
from an atom.
 They can have represented in a graph and the shape of the line (jumps and drops) can
help us identify what element it is, since it essentially shows us the shell structure.

 From this graph we can see a jump after the first electron, meaning that after the first
electron is remove, the next is removed from a shell closer to the nucleus since the
ionisation energy increases. This indicated the electron is in group 1.
 Then we can count the number of electrons remove, which is 11 which tells us the
number of electrons, also the number of protons, hence the atomic number (assuming
it’s an atom, not an ion.)
 We can also see that there are 3 shells in total due to 3 jumps.
 Things like this can tell us a lot about an atom, including its electron structure.
 The periodic table has 4 sections. The sections are referred to as “blocks”.
o The section where group 1 and 2 elements are located is known as the P block,
since all their outer electrons are situated in a p-orbital.
o The section with the transition elements is known as the d block, since their
outer electrons are in a d-orbital.
 o The section with group 3,4,5,6,7 and 8 elements is known as the p-block, since
their outer electrons are in a p-orbital.
o You’re not required to know about the actinides and the acids, other than the
fact that there are known as the F-block elements.
 The atomic radius, the distance from the nucleus to the outermost energy level
decreases across a period.
o This is because the positive charge of the nucleus increases, meaning that
electrons are pulled closer to the nucleus, meaning that the distance decreases.
o The extra electrons added across a period don’t affect the increase since they
don’t affect the shielding.
 The ionisation energy increases across a period.
o The number of protons increases, meaning that there is stronger nuclear
attraction.
o The electrons in the outer energy level are at similar distance from the nucleus
and there is little change in repulsion.
 There are drops in ionisation energy however, between group 2 and 3, and group 5
and 6. This is due to the actual electron configuration of the atoms and the energy
levels.
o From group 2 and 3, ionisation energy drops because the outer electrons in
group 2 are in a S-subshell while the outer electrons in group 3 are in a P-
subshell.
o This means that the electron being removed from group 2 is closer to the
nucleus, hence more attraction, therefore a greater ionisation energy.
o The electron being from group 3 is further away from the nucleus than the
electron from group 2, meaning that there’s less attraction and more shielding
is provided by the s-subshell electrons; hence ionisation energy is much
smaller.
 The ionisation energy drops between group 5 and 6 is to do with the electrons inside
the orbitals, as they are in the same sub-shell.
o The electron being removed from the group 5 electron is from a p-orbital with
no other electron in it. This means there no repulsion from another electron.
o Whereas, in group 6, the electron is being removed from a p-orbital which is
preoccupied with an electron, so there’s repulsion to the electron being
removed.
o This means that that the ionisation energy required to remove the outer
electron from group 5 is much greater than the electron from group 6 due to
the orbitals the electrons are in.
Melting and Boiling points
 As you go across a period, the type of bond formed between the atoms of that element
changes. This changes the melting and boiling point of the element and alters their
chemical properties.
o For the metals (group 1-2), the melting and boiling points increase across a
period because their metallic bonds get stronger.
 o They get stronger as the charge of the nucleus gets stronger and so the overall
attraction between the nucleus and the delocalised electrons is stronger, since
the electron charge stays the same.
o For group 3 and 4 elements, they form giant covalent lattice structures. This
means that they simply have strong covalent bonds in groups linking all their
atoms together. This means that overall a lot of energy is required to break up
these bonds. This means that they will have a high boiling and melting points.
o The group 5-7 elements are the simple molecular structure. There melting and
boiling point generally depends on the strength of London forces between
their molecules (learnt in a future topic, however in-depth knowledge is not
required.)
o The London forces are weak and easy to overcome, so these elements have a
lower melting and boiling point. However, some of them form very large
molecules such as Sulphur which forms S8 which has a greater mass, hence
more electrons.
o If the number of electrons increase, then the strength of London forces
increases, therefore sulphur will be an odd one and will have a greater boiling
and melting point than the other simple molecular structures. However, it is
still not as strong as the other group bonds, so it will still be quite small
overall.
 Group 8 elements (noble gases) have the lowest melting and boiling points since they
can’t form bonds with each other as they have full outer shells, which means they are
found as individual atoms, hence monatomic. This means they have very weak
London forces, so they have low melting and boiling points.