period 3 elements Flashcards
structure and boning in each element
Na - metallic
Mg- metallic
Al - metallic
Si - giant covalent / macromolecular
P4- simple molecular / VdW
S8 - simple molecular / VdW
Cl2 - simple molecular / VdW
Ar - atomic / VdW
Describe the structure of magnesium. Use this to explain why magnesium
has a high melting point
-Mg2+ ions in a sea of delocalised electrons
-strong electrostatic forces between ions and delocalised electrons
-require lots of energy to overcome
Explain why magnesium is a good conductor of electricity
delocalised electrons are able to flow through the metallic lattice and carry charge
Describe the structure and bonding in phosphorus. Explain how the forces in
phosphorus arise
-small P4 molecules bonded by covalent bonds
-weak intermolecular (VdW) forces between P4 molecules
-VdW arise due to random movement of electrons inducing dipoles in neighbouring molecules
Put the following elements in order of their melting points. Explain your
answer in terms of structure and bonding
chlorine, sulfur, phosphorus, argon
-S8 > P4 > Cl2 > Ar
-larger molecules have more electrons and a greater surface area
-stronger VdW forces between molecules (or atoms)
-more energy is required to overcome stronger VdW forces
Explain why silicon has a very high melting point
-giant covalent / macromolecular lattice
-lots of strong covalent bonds between silicon atoms
-each Si is covalently bonded to 4 others
-lots of energy is needed to break the forces
State and explain the trend in atomic radius across period 3. In your answer
you should refer to nuclear charge and electron shielding
-atomic radius decreases across the period as:
-nuclear charge increases across the period
-nuclear shielding is constant
-electrostatic attraction between nucleus and outer electron shell increases
State and explain the trend in ionic radius from Na-Al and P-Cl. In your
answer you should first give the charge of the most likely ion formed by that
element
-Na+, Mg2+ and Al3+ have lost an electron shell compared to non-metal ions
-therefore smaller than non-metal ions
-Al3+ has smallest ionic radius because it has a greater nuclear charge
(electron shielding for Na+, Mg2+ and Al3+ is the same)
-Ionic radius decreases P3-, S2-, Cl-
-Nuclear charge increases from P to Cl with constant electron shielding
-electrostatic attraction between nucleus and outer electron shell increases
Write an equation for the first ionisation energy for magnesium
Mg(g) ⎯→ Mg+(g) + e-
Explain why Mg has a more endothermic first ionisation energy than Na
-Mg has a greater nuclear charge
-electron shielding is constant
-electrostatic attraction between nucleus and outer electron is greater,
-therefore more energy is required to remove the electron
Explain why the first ionisation energy for Al is less endothermic than that of
Mg despite Al have a greater nuclear charge
-Electron removed from Al is in the 3p sub-shell compared to 3s for Mg
-3p sub shell is further from the nucleus with weaker electrostatic attraction,
therefore easier to remove
Explain why the first ionisation energy for S is less than that for P
-The electron removed for S is the first time a 3p sub-shell electron is paired
in an orbital
-The repulsion between electrons in the orbital means the energy needed to
remove the electron is less
explain why atomic radii of the elements across period 3 decrease from sodium to chlorine
- nuclear charge increases (no. of protons) across a group
-so attraction between nucleus and e- increases
explain why the melting point of sulfur (S8) is greater than of phosphorus (P4)
-S8 Molecules are bigger than of P4 molecules
-therefore VDW are stronger in S8
identify element from period 3 from sodium to argon that has the highest second ionisation energy
give equation using state symbols to show processes involved when second ionisation energy of this element is measured
-Na
-Na+ (g) —> Na2+ (g) + e-
