Periodicity and Main Group Chemistry Flashcards
Use knowledge about the general trends of the properties of elements (such as atomic radius, ionisation energy, electron affinity and electronegativity) to understand how chemical compounds form and explain their properties.
Atoms with the same outer shell electron configuration will have similar chemical properties (elements that lie in the same group on the PT).
Atomic radius: increases down a group as electrons add to new shells. Decreases across a row as the effective nuclear charge increases Zeff = Z - S (net positive charge experienced by valence electrons).
Ionisation energy: energy required to REMOVE an electron from a gaseous state e.g. for He:
First ionisation energy: He(g) -> He+(g) + e-
Second ionisation energy: He+(g) -> He2+(g) + e-
Increases from left to right of PT and decreases from top to bottom.
Electron affinity: similar to ionisation energy. Energy required to ADD an electron to a neutral atom ingas phase. e.g. Fe(g) + e- -> Fe-(g).
Decrease in magnitude (less negative) down a group. Increase in magnitude (more negative) across a row.
Electronegativity: describes the relative polarity of bonds and molecules which is a result of unequal electron distribution. Generally increases as atomic size decreases. Increases from left to right of PT. Decreases from top to bottom. Electronegativity ranges from 0 - 4. F = 4
Represent a dipole in a bond, and use electronegativity to identify the positive and negative ends.
Dipole: occurs when there is a separation of charge.
Representation: with delta negative or delta positive sign. Positive delta is on atom with least electronegativity and negative delta on atom with most electronegativity. For example, HCl = H(+ delta) - Cl(- delta).
Electron distribution in heteronuclear diatomic molecules (molecule containing two atoms of different elements) is uneven due to the higher nuclear charge and electron configuration of the two different atoms. The bonds between these atoms are polar.
Examples: NO favours O for the distribution of charge,
HF strongly favours F, HCl favours Cl.
Describe and explain the periodic trends in electronegativity.
Electronegativity generally increases as atomic size decreases. This is because the valence electrons are closer to the nucleus more tightly bound. So electronegativity increase left to right across a row and decreases down a group.
Understand and explain why electronegativities do not correctly predict the polarity of CO, given the MO (Molecular Orbital) diagram.
CO is polar but C is not the negative end! This is not what electronegativities predict. The HOMO (3 o with line on top) is closer in energy to C 2s than O 2s so electrons localised more on C.
Explain the origin of ionic bonding as a limiting case of MO theory.
MO theory = method for describing electronic structure of molecules (bonding) using quantum mechanics.
Molecular orbital meaning = orbitals around molecule.
Rule of MO theory:
1. The number of molecular orbitals produced is
always equal to the number of atomic orbitals
brought by the atoms that have combined.
2. Bonding molecular orbitals (sigma) are lower in
energy that the parent orbitals, and the antibonding
orbitals (sigma star) are higher in energy.
3. Electrons of the molecule are assigned to orbitals
from lowest to successively higher energy.
4. Atomic orbitals combine to form molecular orbitals
most effectively when the atomic orbitals are of
similar energy.
Explain why ionic interactions lead to crystals rather than small molecules.
Ionic bonding is the attraction between positively- and negatively-charged ions. These oppositely charged ions attract each other to form ionic networks, or lattices. Electrostatics explains why this happens: opposite charges attract and like charges repel. When many ions attract each other, they form large, ordered, crystal lattices in which each ion is surrounded by ions of the opposite charge.
