Periodicity Flashcards
(33 cards)
Explain why xenon has a lower first ionisation energy than neon. (3)
- more shielding
- bigger atomic radius
- less nuclear attraction
Explain why the first ionisation energies show a general increase from Li to Ne. (3)
- atomic radii increases
- nuclear charge increases
- shielding remains the same
Explain the difference between the first ionisation energies of Li and Na. (3)
- atomic radii increases
- more shielding
- distance outweighs nuclear charge
Explain why first ionisation energies show a general increase across Period 3, Na - Ar. (3)
- atomic radii decreases
- number of protons in nucleus increases
- nuclear attraction increases
What is the equation, including state symbols, for the third ionisation energy of Sodium? (1)
Na2+ (g) –> Na3+ (g) + e-
Explain why less energy is needed to ionise gaseous atoms of rubidium than gaseous atoms of sodium. (3)
- Rb = more shells
^ more shielding
^ less attraction
Explain whether a Barium ion is larger, smaller or the same size as a Barium atom. (2)
- smaller
- less shielding
Define: first ionisation energy
the energy change when each atom in 1 mole of gaseous atoms loses an electron
Explain why the first ionisation energy of B is less than that of Be. (2)
- In B, the electron is being removed at a higher energy
- An s-orbital electron is lost in Be & a p-orbital electron is lost in B
Explain why less energy is needed to remove an electron from a radium atom than from a calcium atom. (3)
- atomic radii of radium is bigger
- radium has more shells
^ more shielding
Explain why a nitrogen atom is larger than an oxygen atom. (4)
- Nitrogen has less protons
- similar shielding
- Nitrogen has weaker nuclear attraction
- Shell drawn in less by nuclear charge in Nitrogen
What determines the order of elements in The Periodic Table?
the number of protons in the nucleus
Why does phosphorus have a larger first ionisation energy than sulfur? (1)
- P atoms have less repulsion between p-orbital electrons than S atoms
Define: periodicity
- trend in the properties of elements across a period repeated in the next row
Define: metallic bonding
electrostatic force of attraction between metal ions and delocalised electrons
State the element in Period 3 with the highest first ionisation energy
Silicon
Describe the structure and bonding present in silicon. (2)
- giant macromolecular structure
- covalent bonds between silicon atoms
Explain why aluminium has a higher melting point than sodium. (3)
- Aluminium has a higher charge
^ more delocalised electrons - stronger metallic bonding
Explain why aluminium does not follow the general trend in first ionisation energies across the period. (3)
- lower than Mg
- In Al, electron removed from p-subshell
- electron is of higher energy
Describe a similarity between copper and graphite. (1)
- both have layers of atoms that can slide over one another
Explain why copper has a high melting point. (2)
- contains positive metal ions
^ strong attraction b/w them
Identify the Period 3 element with the lowest melting point and explain why. (3)
- Argon
- simple molecular structure
- molecules held together by weak intermolecular forces which require very little energy to overcome
Explain why the second ionisation energy of Aluminium is larger than the first. (1)
- electron is being removed from an s-orbital rather than a p-orbital
Describe the structure & bonding of graphite + why it has such a high melting point. (5)
- arranged in layers
- Carbon atoms in each layer connected by 3 covalent bonds
- weak intermolecular forces b/w layers
high m.p: - millions of strong covalent bonds
^ require a large amount of energy to overcome
