Periodicity Flashcards
Describe the trend of atomic radius across Period 3.
The atomic radius decreases as you move across Period 3 due to the increasing number of protons, which leads to a stronger attraction between the nucleus and the electrons, pulling the electrons inwards.
Explain the relationship between atomic radius and melting point in metals across Period 3.
As the atomic radius decreases across Period 3, the number of delocalised electrons increases, resulting in stronger metallic bonds and higher melting points.
Define the term ‘macromolecular structure’ in the context of Silicon.
A macromolecular structure consists of strong covalent bonds that hold the atoms together, as seen in Silicon.
How does the number of delocalised electrons affect the melting point of metals in Period 3?
The melting point increases with the number of delocalised electrons because stronger electrostatic attractions between positive ions and delocalised electrons lead to stronger metallic bonds.
Do the melting points of Sodium, Magnesium, and Aluminium increase or decrease across Period 3?
The melting points of Sodium, Magnesium, and Aluminium increase as you move across Period 3 due to stronger metallic bonding.
Explain why Silicon has a different melting point compared to Sodium, Magnesium, and Aluminium.
Silicon has a macromolecular structure with strong covalent bonds, which results in a different melting point compared to the metallic bonding in Sodium, Magnesium, and Aluminium.
How does the atomic radius trend relate to the strength of metallic bonds in Period 3 elements?
The decreasing atomic radius leads to stronger metallic bonds due to increased electrostatic attraction between the positive ions and delocalised electrons.
Describe the relationship between covalent bonds and melting points in silicon.
Silicon has strong covalent bonds that require a large amount of energy to break, resulting in a high melting point.
Explain the trend in melting points from Phosphorous to Chlorine.
The melting points of Phosphorous (P4), Sulfur (S8), and Chlorine (Cl2) are low due to weak van der Waals forces, with sulfur having the highest melting point due to its larger size and stronger induced dipole-dipole forces.
Define van der Waals forces and their impact on melting points.
Van der Waals forces are weak intermolecular forces that require a small amount of energy to break, leading to low melting points in substances like Phosphorous, Sulfur, and Chlorine.
How does the size of a molecule affect its melting point?
A larger molecule contains more electrons and larger electron clouds, resulting in stronger induced dipole-dipole forces, which increases the melting point.
Do larger molecules have stronger or weaker van der Waals forces?
Larger molecules have stronger van der Waals forces due to the greater number of electrons and larger electron clouds.
Explain why sulfur has a higher melting point than phosphorous and chlorine.
Sulfur is the largest molecule among the three, containing the most electrons and the strongest van der Waals forces, resulting in a higher melting point.
Describe the melting point of Argon and its significance.
Argon has a very low melting point due to its monoatomic nature and weak van der Waals forces, requiring only a small amount of energy to break these forces.
Explain the trend of electrical conductivity across Period 3.
Electrical conductivity increases across metals in Period 3 due to the presence of giant metallic lattices and an increase in mobile electrons, but drops significantly when reaching non-metals like silicon.
Define first ionisation energy and its relevance in Period 3.
First ionisation energy is the energy needed to remove one electron from the outermost shell of an atom in its gaseous state, and it generally increases across Period 3.
How do van der Waals forces affect the melting point of Argon?
The weak van der Waals forces in Argon result in a low melting point, as minimal energy is required to overcome these forces.
Do metals in Period 3 have higher or lower electrical conductivity compared to non-metals?
Metals in Period 3 have higher electrical conductivity compared to non-metals, as they possess giant metallic lattices and more mobile electrons.
Explain the change in electrical conductivity from aluminum to silicon in Period 3.
Electrical conductivity drops dramatically from aluminum to silicon because silicon has a giant covalent structure with no free ions or electrons to conduct electricity.
How does the first ionisation energy change across Period 3?
The first ionisation energy increases across Period 3 as the nuclear charge increases, making it more difficult to remove an electron.
Describe the relationship between mobile electrons and electrical conductivity in metals.
In metals, an increase in the number of mobile electrons contributes to better electrical conductivity as these electrons can move freely and carry electric current.
Describe the trend in first ionisation energy across a period.
As you move across a period, the number of protons increases, leading to a stronger attraction between the nucleus and the outermost electrons. This results in an increase in the first ionisation energy.
Explain the exception in first ionisation energy between Magnesium and Aluminium.
The first ionisation energy of aluminium is slightly lower than that of magnesium because the outermost electron in aluminium is in a p orbital, which is further from the nucleus and experiences more shielding compared to the s orbital electron in magnesium.
How does the first ionisation energy change from Phosphorous to Sulphur?
The first ionisation energy of sulphur is slightly lower than that of phosphorous due to the outermost electron in sulphur being in a complete atomic orbital, which experiences spin pair repulsion, making it easier to remove.
