8.2 and 8.3 Properties of period 3 Flashcards
What structures do the first 3 (Na,Mg,Al) have
Electron arrangement
Giant metallic structures, they lose their outer electrons to form ionic compounds
Ne(3s1)
Ne(3s2)
Ne(3s2,3p1)
Structure of silicon, what type of bonding in it
Electron arrangement (shorthand)
Has 4 electrons on its outer shell which forms 4 covalent bonds
It has some metallic properties and is classed as a semi-metal
Ne(3s2, 3p2)
Structures of (p, S, and Cl)
Electronic configuration
Non metals, they accept electrons to form ionic compounds, or share outer electrons to form covalent compounds.
Ne(3s2, 3p3)
Ne(3s2, 3p4)
Ne(3s2, 3p5)
Structure of argon
Group 0 so has full outer shell and is unreactive
Ne(3s2, 3p6)
Trend in melting and boiling point across period 3
On the left, melting point increases up to Silicon , then it decreases up to sulfur where it increases slightly, then decreases to argon.
Why does melting and boiling point increase up to aluminum ( metals)
They’re giant structures: the strength of metallic bonding increases.
Because as you go from left to right, charge of ion increases so more electrons join sea of delocalised electrons, holding the giant metallic lattice together more tightly so the forces of attraction require more energy to break
Why does silicon have the highest melting and boiling point in period 3
It exists as a giant macromolecular structure with covalent bonds
So these covalent bonds are stronger than the metallic bonds in the first 3 elements
So loads of energy is required to break them
Trend in melting point of the after silicon, non metals in period 3
Why is it not a constant order
. Depends on size of van der walls forces between the molecules.
. These depend on the number of electrons in the molecule, and how closely the electrons can pack together.
Sulfur exists as S8 molecules, phosphorus exists as P4 molecules, and chlorine is Cl2.
So because sulfur has biggest molecules, it has the most electrons so the most Van Der Waals.
. This means melting points are ordered s8, p4, cl2
What is atomic radii and why is it used
You can’t measure radius of an isolated atom because there’s not a clear point of when electron cloud density drops to 0.
Half the distance between the centres of a pair of atoms is used
Trend in atomic radii across period
. As you move from sodium to chlorine across group, you add protons to the nucleus, and electrons to the outer level (the third shell.)
. The charge of the nucleus increases going across, which pulls the electrons closer to the nucleus
. There are no additional electron shells to provide shielding so the size of atom decreases across period.
Radii of atoms down a group
Increases
. As you go down group, atoms of each element have another complete level of electrons
. This increases the distance of the outer shell from the nucleus, and increases shielding
What is first ionisation energy
Energy required to convert a mole of isolated gaseous atoms into a mole of singly positively charged gaseous ions
Eg remove one electron from each atom
Half equation for first ionisation energy
E stands for element
E(g) —> E+(g) + e-(g)
Pattern of ionisation energy across period 3
Generally increase across the period.
Eg sodium and lithium have the lowest values, whilst helium etc have the highest values
Cause for the pattern of ionisation energy across period
-As you go from left to right, number of protons increases but the electrons occupy the same energy level.
- This increased charge of the nucleus means the outer electrons are more attracted to nucleus so that it gets increasingly difficult to remove an electron
First ionisation energy down a group
Decreases down a group
Cause for the pattern of ionisation energy down a group
It decreases
. the number of filled shells increase down a group, which results in an increase in shielding.
. Also, the distance between outer electrons and nucleus increases so the electron is held less strongly/is less attracted to nucleus
. Thus outer electrons get easier to remove going down a group
Why is there a large drop in ionisation energy as you move from one period to the next
. It decreases
. Eg neon period 0 with electron arrangement 2,8 to sodium group 1 with electron arrangement 2,8,1.
Because at sodium a new energy level starts so there is an increase in atomic radius , the outer electron is further from nucleus so it takes less energy to remove
Explain the reason for the drop in ionsiation energy in period 3 between
Magnesium to aluminium:
Ne(3s2) Ne(3s23p1)
Although ionisation energy overall increases across the period, there are two drops in it.
Mg to Al: It decreases as magnesium has electron arrangement Ne(3s2) whilst Al has an added subshell making it Ne(3s23p1).
This new subshell is further from positive nucleus and has more shielding so an electron is removed more easily than Mg which has a full subshell closer to the nuceus.
Explain the reason for the drop in ionisation energy in period 3
From phosphurus to sulfur
Ne(3s2,3p3) Ne(3s2,3p4)
Phosphurus has outer 3p3 subshell so each electron is in its own orbital
Whereas sulfur has 3p4 which means there are two electrons in one orbital
This pair of electrons in one orbital repel due to having the same charge, so it takes less energy to remove an electron from sulfur as opposed to removing an electron from phosphurus
