EL4 - Periodicity Flashcards
Organising the elements of life
On the periodic table, what is a period?
A row.
What is periodicity?
The occurrence of a regular pattern in a property as you go across a period
The regular pattern is also repeated in other periods
How are elements in the periodic table arranged?
How did they used to be arranged?
Arranged by atomic number (no. protons)
Used to be arranged by Ar
What trend do melting/boiling points follow across a period?
e.g. period 3
Melting point increases then decreases across the period
This is because the metals on the left-hand side of a period
are metallically bonded,
and have an increasing positive charge,
increasing the number of delocalised electrons
and smaller ionic radius.
This means a stronger metallic bond,
so higher melting point.
The further across the period, the more electrons and the more positive the nucleus becomes, so the stronger the bonds. So higher melting point.
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Silicone has a high melting point because it is a giant covalent structure which requires a lot of energy to break
The remaining non-metals are simple molecules. They are only held together by weak intermolecular forces (e.g. id-id). To melt these molecules you don’t need to break the strong covalent bonds, only the weak intermolecular bonds.
What is first ionisation enthalpy?
The energy needed to remove one electron from each atom in one mole of isolated gaseous atoms of an element.
What is the general equation for first ionisation enthalpy?
X(g) –> X+(g) + e-
Remember state symbols!
What is the trend in first ionisation enthalpies as you go across the period?
Why?
The first ionisation enthalpy will increase.
This is because as you go along the period the atomic number increases and this causes a stronger electrostatic attraction between the outer electrons and the nucleus.
There are 2 ‘dips’ in this trend:
There is a drop in energy from Magnesium to Aluminium.
This is because in magnesium the outer electron is in a 3s orbital. The 3p orbital is further from the nucleus, and hence there is more electron shielding and distance from the nucleus so electrostatic attraction decreases.
The second drop is from Phosphorus to Sulphur.
This is because it is the first time electrons are paired up in 3p orbitals, so there is added electron-electron repulsion.
What is the trend in first ionisation enthalpy down a group?
Why?
The first ionisation enthalpy decreases.
There is more shielding between the nucleus and the outer electrons and the distance between the nucleus and the outer electron increases and therefore the force of attraction between the nucleus and outermost electrons is reduced.
What is the trend for atomic radii across a period?
Atomic radius decreases from left to right within a period.
This is caused by the increase in the number of protons and electrons across a period.
One proton has a greater effect than one electron; thus, electrons are pulled towards the nucleus, resulting in a smaller radius.
NOTE:
THE ISLAND OF STABILITY
The protons are positively charged and repel each other, with the repulsion increasing as the protons get closer.
However, the strong nuclear force is stronger than this repulsion of charge, and holds the nucleus together.
Neutrons don’t have an overall charge, but they have some magnetic properties.
Neutrons act as a buffer between protons. Usually, the more protons you have, the more neutrons you need…
…Beyond element 92, uranium, the strong nuclear force is insufficient to hold the nucleus of the element together, and all isotopes are unstable, breaking down by radioactive decay.
Compare the electronic arrangement of the noble gases:
He 1s2
Ne 1s2 2s2 2p6
Ar 1s2 2s2 2p6 3s2 3p6
Kr 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6
All have subshells fully occupied by electrons.
Such arrangements are called closed shell arrangements.
These are particularly stable arrangements.
NOTE:
The elements in Group 1 all have one electron in the outermost s subshell and therefore show similar chemical properties.
Group 2 elements all have two electrons in the outermost s subshell. Groups 1 and 2 are known as s-block elements.
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How does this compare to the other groups? And transition metals?
In groups 3, 4, 5, 6, 7 and 0,
the outermost p subshell is being filled. These elements are known as p-block elements.
The elements where a d-subshell is being filled (transition metals) are called d-block elements,
and those where an f-subshell is being filled are called f-block elements (two lines at the bottom of the periodic table).
NOTE:
Electronic configuration of s- and p-block ions:
- Group 1 elements form +1 ions
- Group 2 elements form +2 ions
- Group 7 elements form -1 ions
- Group 6 elements form -2 ions
- Aluminium forms +3 ions
Knowing this, you can write the electronic configurations of ions formed by the s- and p-block elements in the first four periods.
Worked example: Electronic configuration of an anion
What is the electronic configuration of S^2-?
Step 1:
Write out the electronic configuration of a sulfur atom.
1s2 2s2 2p6 3s2 3p4
Step 2:
Use the charge on the ion to decide how many electrons to add to give the required charge.
S^2- has a charge of -2 so need to add two electrons.
Step 3:
Add the required number of electrons to the electronic configuration.
1s2 2s2 2p6 3s2 3p6
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Worked example: Electronic configuration of a cation
Draw out the electronic configuration of Na+, showing the subshells and atomic orbitals.
Step 1: Write out the electronic configuration of a sodium atom.
1s2[^v] 2s2[^v] 2p6[^v][^v][^v] 3s1[^]
Step 2: Use the charge on the ion to decide how many electrons to take away to give the required charge.
Na+ has a charge of +1 so need to take away one electron.
Step 3: Add / take away the required number of electrons to the electronic configuration.
1s2[^v] 2s2[^v] 2p6[^v][^v][^v]
Why are s-block elements more reactive than p-block elements?
S-block metals have lower ionisation enthalpies compared to p-block metals.
Due to their low ionisation enthalpies, s-block metals tend to form cations since they are more willing to give their valence electrons away. This means they are more reactive.
Explain the difference between the melting points of silicon and phosphorus.
Silicon has the highest melting point in period 3.
It has a giant covalent structure.
Many strong covalent bonds hold it together.
A large amount of energy is needed to overcome these strong covalent bonds.
Phosphorus has the structural formula P4.
It has a lower melting point than silicon
due to a weaker simple molecular structure.
Melting point is determined by weaker intermolecular forces.
