Electronic Structure and Periodicity
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What is the oxidation number of an element on its own?
This is always zero.
How do we assign oxidation numbers in compounds?
• Each atom within the compound is assigned an oxidation number, which is usually identical to its ionic charge. • The sum of these is always equal to the charge of the complex ion.
What are the oxidation numbers of the following exceptional cases of atoms in compounds? 1. H in hydrides 2. O in peroxides 3. O bonded to F
• This is -1, such as in NaH. • This is -1, such as in H2O2. • This is +2, such as in F2O.
What can we say about the oxidation number of an element bonded to a more electronegative element?
The oxidation number of the less electronegative atom will have a positive oxidation number.
What convention must we follow when writing oxidation numbers?
The sign always comes before the number itself, unlike with ionic charges where it comes afterwards.
How do we name atoms and compounds which can have different ionic charges? Give an example.
• We present the oxidation number of the atom or compound as a Roman numeral given after the name of an element. • Nitrate (III) i.e. nitrite, NO2- has a +3 oxidation number, and nitrate (V) i.e. NO3- has a +5 oxidation number.
How do we define an oxidation reaction?
This is the loss of electrons, and the gain in oxidation number. Remember the OILRIG mnemonic.
How do we define a reduction reaction?
This is the gain in electrons, but the loss in oxidation number. Think of it as, the oxidation number reduces.
What is a redox reaction defined as?
A reaction in which species are both oxidised and reduced - i.e. if one of the processes happens, so must the other. This could be for example a metal-acid reaction.
What is first ionisation energy defined as?
• The energy required to remove one electron from each atom In one mole of gaseous atoms of an element, in order to form one mole of gaseous 1+ ions. • Think of it is a measure of how easily an atom loses electrons to form ions.
Why must ionisation occur in the gaseous state?
This is because the removal of electrons is only achievable in high-energy states.
What is the general formula for 1st ionisation energy?
X(g) = X+(g) + e-
Why does first ionisation energy decrease down the group?
• The atomic radius increases due to the addition of new shells • This also gives rise to more shielding • Therefore the attraction between valence electrons and the nucleus is reduced
Why does first ionisation energy generally increase across the period?
• Atomic radius decreases due to the addition of valence electrons in the same shell, and shielding is the same • Nuclear charge increases due to the addition of protons • Increasing attraction between valence electrons and nucleus
How can we explain the group 2-3 dip in ionisation energies across the period?
• An electron is removed from a lower energy s-orbital in Group 2, whereas in Group 3 an electron is removed from a higher energy p-orbital. • This requires less energy.
How do we explain the group 5-6 dip in ionisation energies across the period?
• The Group 5 atom’s orbitals are singularly occupied, so the electron added in Group 6 must pair up, giving rise to repulsion as electrons have like charges. • This means that less energy is required to overcome forces of attraction.
What are successive ionisation energies defined as?
The energy required to successively remove electrons from each atom In one mole of gaseous atoms of an element.
What do the ‘jumps’ on a graph of successive ionisation energies represent?
These represent the loss of an energy level (shell). It requires more energy to remove an electron from a shell closer to the nucleus, as there is less shielding, and nuclear attraction is greater.
Why do successive ionisation energies increase logarithmically?
• As each electron is removed, the valence electrons are drawn closer to the nucleus since there are more protons than electrons. • This gives rise to an increase in nuclear attraction, due to reduced atomic radius.
How do we define periodicity?
A retreating trend of physical, chemical and atomic properties across the period, caused by regular and predictable variations in atomic structure.
How are elements classified in to ‘blocks’ of the periodic table?
The ‘block’ an element occupies is defined by which sub-shell its highest-energy electron occupies: s, p or d.
Explain the trend in atomic radius across period 3.
• This exhibits a decrease • Nuclear charge increases, and electrons are added to the same shell; shielding is unchanged • Greater attraction, electrons drawn closer to nucleus
Explain the trend in electronegativity across period 3.
• This exhibits an increase • Nuclear charge increases and atomic radius decreases • Giving rise to stronger attraction between the nucleus and the two bonded/shared electrons
Describe the trend in boiling point across period 3.
• Increases from sodium to magnesium • Peaks at silicon • Steep decrease to sulfur, then slow decrease to chlorine • Lowest at argon