Periodicity Flashcards
How does the radius change across the period?
The radius decreases
Because the increase in protons therefore increase in nuclear attraction
So electrons feel stronger attraction and are pulled closer
Does the increasing electrons affect the atomic radius?
The increasing electron-electron repulsion by adding electrons does NOT affect the radius as its being added to the same shell
How does the atomic radius change with the group?
Increases as you go down
Because the number of shells increases by 1 as there are more electrons to occupy these
So the subsequent electrons are at a further distance form the nucleus
Does the number of protons affect the radius of an atom as you go down the group?
There is more nuclear charge (which would pull the shells in) but the increases number of electrons that occupy more shells outweighs this and so the radius increases
Valence electrons
Outer shell electrons
Melting point in metals
Greater number of valence electrons = high melting point
So the more electrons in an atoms outer shell, the more energy required to melt it
Shorter atomic radius = high melting point
Why does the melting point increase in metals if there are more valence electrons?
Because in metallic bonding the valence electrons are DONATED in the sea or delocalised electrons
Electrostatic attraction between these electrons and nucleus is greater so therefore more energy required to break the bond
Outer electron configuration of group 1
ns¹
n = the period that specific element is in
Outer electron configuration of group 2
ns²
n = the period that specific element is in
Outer electron configuration of group 3
ns² np¹
n = the period that specific element is in
Outer electron configuration of group 4
ns² np²
n = the period that specific element is in
Outer electron configuration of group 5
ns² np³
n = the period that specific element is in
Outer electron configuration of group 6
ns² np⁴
n = the period that specific element is in
Outer electron configuration of group 7
ns² np⁵
n = the period that specific element is in
Outer electron configuration of group 8
ns² np⁶
n = the period that specific element is in
Groups
Vertical columns in periodic table that all contain the same outer electron configuration
Periods
Horizontal rows in periodic table that all contain the same quantum levels that have electrons
Blocks
Dividing the periodic table into 3 by sorting each group into a different block
S block
Groups 1 and 2
The highest energy electron (outer shell electrons) are in an s orbital
D block
All elements between Sc —> Zn
Y —> Cd
As you go across the period, the number of electrons in the d su shell increases
P block
Groups 3, 4,5, 6, 7, 8
Highest outer electron is in the p orbital
Overall ionisation trend across periods
Increases as you go from left to right
Then sudden drop when moved to the next row which is further down the group
What are the anomalies in ionisation trends across the period?
It DECREASES between Be and B
It DECREASES between Mg and Al
It DECREASES between N and O
When following the ionisation trend it should increase
Why is there an anomaly in ionisation energy between Be and B?
Although nuclear charge of the B atom is greater than Be atom,
Outer B electron has more energy (since it is in a 2p orbital unlike in the 2s orbital for Be)
So the energy required to remove this electron is less than that required to remove a 2s electron from a beryllium atom so IE decreases
Also the 2p electron in boron experiences greater electron–electron repulsion (so greater shielding) because there are two inner electron sub-shells as opposed to only one in the beryllium atom
Anomaly in ionisation between Mg and Al
Aluminium has lower IE than Mg
Because the outer shell electron in Al is in a 3p orbital which is a higher energy level than 3s, the outer electron for Mg.
So the energy put in to ionise it is higher for Mg than Al giving Mg a higher IE
Also greater shielding for Al means lower ionisation energy
Anomaly in ionisation between N and O
O has a lower IE than N
Because in the p subshell in O there is a pair of electrons in 1 p orbital whereas in N all the P subshell electrons are in diff p orbitals
This increases electron-electron impulsion in O so more shielding = less electron is required to remove the electron therefore the IE in O is lower than in N
Periodicity
A regularly repeating pattern of atomic, physical and chemical properties with increasing atomic number.
Trend in melting points across group 3
Increase from Na, Mg to Al
Large increase to Si
Big decrease down to P, S, Cl then Ar
Why do melting points of metals increase with increasing atomic number?
Because there are more valence electrons per atom and a shorter radius
So greater electrostatic attraction between positive ions and electrons = more energy required to melt
Why does si have such a high melting point?
Because silicon is a giant covalent structure it contains many strong covalent bonds which require lots of energy to break these
Why does the melting point from Si to the simple molecules like P, S and Cl decrease massively ?
Because these are simple molecules rather than giant covalent structure
Held together by van Der waals forces instead
Why is the melting point from P to S increase but decrease to Cl?
Because van der waals forces are stronger with stronger induced dipole-dipole forces
So the more electrons per molecule (which is in sulfur) the higher the melting point
Why is the melting point in Ar very low?
Because it exists as a monoatomic element so has very weak van der waals forces so little energy is needed to overcome it
