Periodicity Flashcards
What is periodicity?
The study of repeating patterns or trends in physical and chemical properties across the periodic table
What determines an element’s block in the periodic table?
The subshell in which its outer electrons are located
What are the four blocks in the periodic table?
s - group 1 and 2
p - groups 3- 8
d - transition metals
f - lanthanides and actinides
What is atomic radius?
The distance from the nucleus to the outermost electrons in an atom
How do atomic radius change down a group?
- more electron shells
- greater shieling reduces attraction from the nucleus
What is electronegativity?
The ability of an atom to attract pairs of electrons in a covalent bond
Why are noble gases not included in the electronegativity series?
They don’t form covalent bonds readily so they don’t attract pairs of electrons
What happens to atomic radius across a period?
It decreases
Why does atomic radius decrease across a period?
- More protons in nucleus
- Same amount of shielding
- Stronger attraction pulls electrons closer
What happens to atomic radius down a group?
It increases
Why does atomic radius increase down a group?
- More energy levels
- Greater distance from nucleus
- Increased shielding
- Weaker attraction from nucleus
What happens to electronegativity across a period?
It increases
Why does electronegativity increase across a period?
- More protons in nucleus
- Smaller atomic radius
- Stronger attraction to bonding electrons
What happens to electronegativity down a group?
It decreases
Why does electronegativity decrease down a group?
- More electron shells (more shielding)
- Larger atomic radius
- Weaker attraction to bonding electrons
What happens to ionisation energy down a group?
It decreases
Why does ionization energy decrease down a group?
- more shells/ energy levels
- more shielding
- greater distance between nucleus and outer electron
- weaker attraction between nucleus and outer shell electrons
- so less energy required to remove one electron from outer shell
General trend of ionization energy across a period
It increases
Why does ionization energy increase across a period?
- More protons in the nucleus
- Same amount of shielding
- Smaller atomic radius
- Stronger attraction between nucleus and outer electron
Why is there a drop in ionization energy from Group 2 to Group 3?
- Group 2 loses an electron from an s orbital
- Group 3 loses an electron from a p orbital
p orbitals have higher energy than s orbitals, so the electron is easier to lose
Why is there a drop in ionization energy from Group 5 to Group 6?
- Group 6 loses an electron from a p4 orbital (two electrons in orbital)
- Extra electron-electron repulsion in p⁴ makes it easier to lose an electron
What determines melting and boiling points?
The strength of the forces between particles (metallic bonds, covalent bonds, or intermolecular forces)
What type of bonding do Na, Mg, and Al have?
Metallic bonding.
Why does Al have a higher melting point than Na and Mg?
- Higher charge on metal ion
- More delocalized electrons
- Smaller ions, leading to stronger metallic bonding
What type of bonding does Si have?
Giant covalent bonding
Why does Si have the highest melting point in Period 3?
It has a giant covalent structure, requiring many strong covalent bonds to be broken
What type of bonding do P₄, S₈, and Cl₂ have?
Simple molecular structures with van der Waals’ forces
Why does S₈ have a higher melting point than P₄ and Cl₂?
- S₈ has more electrons
- Bigger molecules create stronger van der Waals’ forces
What type of bonding does Ar have?
Monatomic structure with very weak van der Waals’ forces.
Why does Ar have the lowest melting point in Period 3?
It has only weak van der Waals’ forces between atoms
