Periodicity Flashcards
Periodicity
Repeating trends of physical or chemical properties
Blocking of the Periodic table
The block of an element (s,d,p,f) corresponds to the sub level where the outermost (highest energy) electron is
Trends in atomic radius (Increasing)
Atomic Radius INCREASES DOWN a group :
- There are extra energy levels because there are more electrons
- More shielding
Attraction the the nucleus decreases
Trends in atomic radius (Decreasing)
Atomic radius DECREASES ACROSS a period:
-Number of protons increases
- Same energy levels
- Similar/Same shielding
- Attraction to the nucleus decreases
First Ionisation energy
The energy required to remove 1mol of electrons from 1mol of gaseous atoms
Trend in the 1st IE down group 2
-Ionisation Energy decreases
- Extra energy level because more electrons
- More shielding
- Weaker attraction between nucleus and outer electron
- Requires less energy to remove an electron
- The electron is removed from a higher principle energy level
Trend in the 1st IE across period 3 elements
The general trend across the period is an increase in ionisation energy
- Ionisation energy increases
- The number of protons increases
- Same shielding /Atomic radius decreases
- Same energy levels
- Stronger attraction between nucleus and outer electron
- Requires more energy to remove an electron
Deviations across the period (Group 3)
-Ionisation Energy decreases in each period at group 3 because u change from an S to a P sub level
- The electron is removed from a higher energy P sub level
- Weaker attraction between nucleus and outer electrons
Deviations across the period (group 6)
- Ionisation energy decreases within each period at group 6 because the electrons pair up in a P orbital
- There is a pair of electrons in a p-orbital
- Extra repulsion means less energy is required to remove an electron
Successive Ionisation Energies
Successive IEs will always increase because:
- The positive charge on the ion increases
- The ionic radius decreases
- Nuclear attraction on the outer electron increases
- Its more difficult to remove an electron from a more positive ion
e.g C+
Jumps in ionisation energies
Large jump = an electron is removed from a principal energy level that is much closer to the nucleus
e.g if the jump is between the 3rd and 4th IE then there must’ve been 3 outer electrons so the element is in group 3
Metallic Bonding Defintion
Strong electrostatic forces of attraction between positive ions and delocalised electrons
Metals
- Metal : A lattice of positively charged metal ions attracted to a sea of delocalised electrons
-Metals can’t transfer electrons because there’s no non metal present. - So, the atoms merge their outer shells and there electrons become shared
Delocalised Electrons
- Electrons have no fixed position
- Positive ions will repel but the sea of electrons is attracted to the positive ions and will make the lattice very strong
What factors affect the strength of a metallic bond?
- Higher ionic charge = stronger metallic bonds
- More delocalised electrons = stronger metallic bonds
- Smaller atomic radius = stronger metallic bonds
These make the positive ions more attracted to the delocalised sea of electrons
Properties of Metals
- Good conductors of electricity and heat as the delocalised electrons are free to move and flow.
- Have very high melting and boiling points as they have a a strong electrostatic attraction between the positive ions and delocalised electrons
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Malleable and ductile
Because there are layers of ions that can slide over each other
Giant Covalent Structures
In period 2 & 3
Diamond
Diamond (carbon)
Each carbon bonded to 4 other carbon atoms
Strong covalent binds - require lots of energy to break
Conduct electricity? No because there are no delocalised electrons
High melting/boiling point because strong covalent bonds
Giant Covalent Structures
Graphite (carbon)
Each carbon atom bonded to 3 other carbon atoms
- Forms a hexagonal layer
- Strong covalent bonds
Weak London forces between the layers
Conduct electricity? Yes. delocalised electrons are free to move and flow
Giant Covalent Structures
Graphene (carbon)
- 1 layer of graphite
- Each carbon atom bonded to 3 other carbon atom
- Forms a hexagonal layer
- Strong covalent bonds
- Conduct electricity? Yes delocalised electrons are free to move and flow
Giant Covalent Structures
Silicon
Each silicon atom is bonded to 4 other silicon atoms
Seeing covalent bonds
Conduct electricity?
No, no delocalised electrons
Bonding of period 2 elements
S block - Li & Be - Metallic bonding
P block - C,N,O,F - Covalent bonds
Neon - Monoatomic - Single atom
Structure of period 2 elements
Li and Be - Giant metallic lattice
Carbon - Giant Covalent lattice
N2, O2, F2, Neon - Simple molecular
Simple Molecular
Strong covalent bonds within the molecule
Weak IMFs between the molecules
All non-polar molecules
So weak London forces between the molecules that require less energy to break
Bonding of period 3 elements
Bonds increase as you go across
Na, Mg, Al - Metallic bonding
Si, P, S, Cl - Covalent bonds
Ar - Monoatomic , Single atom
Structure of period 3 elements
Na, Mg, Al - Giant metallic lattice
Si (silicon) - Giant covalent lattice
P, S , Cl , Ar - Simple molecular
P4, S8, Cl2, Ar
Melting point
Melting - Some of the attractive forces between the particles are broken
They now move freely around ravished but still close together
The stronger the forces, the more difficult it is to melt
Boiling point
All the attractive forces between the particles are broken
They are free to move randomly
Melting/Boiling point of Period 2&3 Metals
Giant metallic lattices
P2 - Li & Be
P3 - Na, Mg, Al
High MP/BP due to strong electrostatic attraction between positive ions and delocalised electrons
Melting/Boiling point of Period 2&3 elements
Giant covalent
P2 - C
P3 - Si
Very high MP & BP due to strong covalent bonds which require lots of energy to break
Melting/Boiling point of Period 2&3 elements
Simple Molecular
P2 - N2, O2, F2, Neon
P3 - S8, P4, Cl2, Ar
Low MP & BP because weak London forces between molecules that require less energy to break
