The periodic table isn't just a random arrangement of elements...
Chemistry301 views·Updated 8 Jul 2026·2 pagesUnderstanding Periodic Table Trends: Key Chemistry Insights
Hannah @hannah_studys1012
1
of 2
Atomic Radius and Ionisation Energy Trends
Atomic radius gets smaller as you move across a period because there's greater attraction between the outer electrons and the nucleus. Think of it like a magnet getting stronger - the electrons get pulled in closer.
Ionisation energy is the energy needed to remove one outer electron from a gaseous atom. As you go across a period, this generally increases because electrons are closer to the nucleus and harder to remove. It's like trying to pull something away from a stronger magnet.
There are some clever exceptions to remember. Group 3 elements have lower ionisation energy than Group 2 because their outer electron sits in a p-orbital, which is easier to lose. Group 6 elements have lower ionisation energy than Group 5 due to electron-electron repulsions making it easier to remove an electron.
Quick Tip: The biggest jump between successive ionisation energies tells you which group an element belongs to - this often comes up in exam questions!
Going down a group, ionisation energy decreases because there's more shielding from inner electrons, reducing the attraction between nucleus and outer electrons.
2
of 2
Melting Point Patterns
Melting points across periods show distinct patterns that make perfect sense once you understand the bonding. Across Period 2 and 3, you'll see metals on the left with increasing melting points, then a massive peak at carbon and silicon.
For metals (Li, Be, B and Na, Mg, Al), melting point increases because the charge on metal ions increases, creating stronger attraction between atoms. More charge equals more energy needed to break the structure apart.
Giant covalent structures like carbon and silicon have the highest melting points because you need to break many strong covalent bonds. Think of trying to demolish a brick wall versus knocking down a house of cards.
Remember: Simple molecular substances (like P₄, S₈, Cl₂) have low melting points because you're only breaking weak intermolecular forces, not actual bonds.
The reason sulfur (S₈) has a higher melting point than phosphorus (P₄) is simply that S₈ has more electrons, creating stronger London forces between molecules.
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Chemistry301 views·Updated 8 Jul 2026·2 pagesUnderstanding Periodic Table Trends: Key Chemistry Insights
Hannah @hannah_studys1012
The periodic table isn't just a random arrangement of elements - it reveals fascinating patterns that can help you predict how atoms behave. Understanding these trends will make chemistry much more logical and give you powerful tools for exams.
1
of 2
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Atomic Radius and Ionisation Energy Trends
Atomic radius gets smaller as you move across a period because there's greater attraction between the outer electrons and the nucleus. Think of it like a magnet getting stronger - the electrons get pulled in closer.
Ionisation energy is the energy needed to remove one outer electron from a gaseous atom. As you go across a period, this generally increases because electrons are closer to the nucleus and harder to remove. It's like trying to pull something away from a stronger magnet.
There are some clever exceptions to remember. Group 3 elements have lower ionisation energy than Group 2 because their outer electron sits in a p-orbital, which is easier to lose. Group 6 elements have lower ionisation energy than Group 5 due to electron-electron repulsions making it easier to remove an electron.
Quick Tip: The biggest jump between successive ionisation energies tells you which group an element belongs to - this often comes up in exam questions!
Going down a group, ionisation energy decreases because there's more shielding from inner electrons, reducing the attraction between nucleus and outer electrons.
2
of 2Sign up to see the content. It's free!
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Melting Point Patterns
Melting points across periods show distinct patterns that make perfect sense once you understand the bonding. Across Period 2 and 3, you'll see metals on the left with increasing melting points, then a massive peak at carbon and silicon.
For metals (Li, Be, B and Na, Mg, Al), melting point increases because the charge on metal ions increases, creating stronger attraction between atoms. More charge equals more energy needed to break the structure apart.
Giant covalent structures like carbon and silicon have the highest melting points because you need to break many strong covalent bonds. Think of trying to demolish a brick wall versus knocking down a house of cards.
Remember: Simple molecular substances (like P₄, S₈, Cl₂) have low melting points because you're only breaking weak intermolecular forces, not actual bonds.
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