GROUPS 1&2: ATOMIC AND PHYSICAL PROPERTIES

Explain why the atomic radii of the Group 2 elements increase as you go down the group. As you go down the group, you are adding ​ (a)   which take up more space. The​ (b)   in the nucleus are exactly compensated for by the same number of extra shielding electrons, so there is no other effect to think about.

• extra layers of electrons
• additional protons
• additional neutrons
• valence electrons

Define the term first ionisation energy. The energy needed to remove ​ (a)   from each of ​ (b)   to make ​ (c)   .

• the most loosely held electron
• one mole of gaseous atoms
• one mole of singly charged gaseous ions
• one proton
• one mole of aqueous ions

Describe and explain how the first ionisation energy of Group 2 elements changes.

First ionisation energy ​ (a)   There are 3 considerations: ​ (b)   , the amount of shielding, and the distance of the electron from the nucleus.

As you go down the group, the extra charge on the nucleus is ​ (c)   by the inner electrons (shielding). The only factor affecting the ionisation energy is the distance from the nucleus. The ​ (d)   the outer electrons get, the less strongly they are attracted, and the lower the ionisation energy.

• decrease down the group.
• the charge on the nucleus
• exactly offset
• further away
• increase down the group.
• reinforced
• nearer

The electronegativity of the elements falls as you go down the group. Explain what is meant by electronegativity.

Electronegativity is a measure of the tendency of ​ (a)   to ​ (b)   a ​ (c)   .

• an atom
• attract
• bonding pair of electrons
• lone pair electrons
• repel

The most electronegative element in the group 2 is beryllium. Explain why beryllium is more electronegative than magnesium.

In both Be and Mg, the bonding pairs will feel a net pull of ​ (a)   – the number of protons minus the number of shielding electrons. But in the case of Be, the bonding pairs will be ​ (b)   to the nucleus than they are in Mg, because Be is a ​ (c)   atom than Mg. The ​ (d)   are more strongly attracted in Be and therefore the electronegativity of Be is greater than that of Mg.

• 2+ from the metal nucleus
• closer
• smaller
• bonding pairs
• 1+ from the metal nucleus
• further
• bigger
• lone pair

Answer explanation

When the atoms form bonds there will be bonding pairs of electrons – each pair made of one from the Group 2 element and one from whatever it is bonding to. This applies just as well to the formation of ionic bonds as to covalent ones. In the case of the ionic bond, the bonding pair ends up very close indeed to the more electronegative atom in the bond - so close that we consider that the electron from the Group 2 metal has been transferred entirely to the other element.

Explain why the single outer electron in both sodium (2,8,1) and potassium (2,8,8,1) feel the same net pull from the nucleus of +1. Sodium's nucleus has ​ (a)   , and there are ​ (b)   between the nucleus and the single outer electron. The outer electron therefore feels a net attraction of +1 from the nucleus. Potassium has ​ (c)   and ​ (d)   , also giving a net attraction of +1.

• 11 protons
• 19 protons
• 10 shielding electrons
• 18 shielding electrons

If both Na and K feel the same net pull from the nucleus, why is the potassium atom bigger than a sodium atom?

• Potassium has an extra layer of electrons, and these take up space.

• Sodium has an extra layer of electrons, and these take up space.

• Potassium is more electronegative.

• Potassium has bigger proton number.

Explain fully why first ionisation energy decreases as you go down Group 1. Going down Group 1, the nuclear charge increases, so the force of attraction for the electron being removed also ​ (a)   . As each quantum shell is added, energy of the outermost electrons​ (b)   . As the number of filled inner shells increases, force of repulsion on the electron being removed increases - IE decreases down the group. The combined effect of the ​ (c)   factors outweighs the effect of the​ (d)   factor.

• increases - IE increase down the group
• increases - IE decrease down the group
• second and third
• first
• first and second
• third

Explain how does electronegativity change as you go down Group 1. ​ (a)   down group 1.

When the atoms form bonds there will be a ​ (b)   of electrons – one from the Group 1 element and one from whatever it is bonding to. This applies just as well to the formation of ionic bonds as to covalent ones. In the case of the ionic bond, the bonding pair ends up very close indeed to the more electronegative atom - so close that we consider that the electron from the Group 1 metal has been transferred entirely to the other element. As the atoms get bigger, the bonding pair of electrons gets further from the nucleus of the Group 1 element, and so feels ​ (c)   Therefore the electronegativity​ (d)  

• Electronegativity decreases
• bonding pair
• less attraction.
• decreases.
• Electronegativity increases
• lone pair
• more attraction
• increases.

Explain why the melting and boiling points of the Group 1 metals decrease down the group? As the atoms get ​ (a)   , the delocalised electrons are ​ (b)   the metal nuclei, and so are ​ (c)   . That means that the attractions are ​ (d)   to break, and so ​ (e)   is needed to separate the atoms into a melt or a vapour.

• bigger
• less strongly attracted
• less heat
• smaller
• nearer to
• further away from
• more strongly attracted
• more heat
• easier
• more difficult