Atoms, Structure And Bonding

• Relative atomic mass is the weighted mean mass of an atom. • Relative isotopic mass relates to the weighted mean mass of an isotope of the element in question. • Both of these are given in comparison to 1/12th the mass of a carbon-12 atom.

• The transfer of electrons between the highest energy levels • Production of ions with opposite charges and stable electronic configurations isoelectric to noble gases • The ionic bond itself is the electrostatic attraction between cations and anions

• Atoms share electrons from their highest energy levels • Covalent bonds are the electrostatic attractions between nuclei and shared electrons • Multiple pairs of electrons can be shared • Sometimes involves expanding the octet

Max no. electrons in a given shell = 2n^2, where n is the principal quantum number.

• Subshells are labelled s, p and d in relation to the orbitals they contain. • The order of energy of electrons increases s<p<d. • The order of subshell energies overlaps as energy increases, hence explaining why 4s precedes 3d.

These are areas about nuclei is which up to two electrons can be present, with opposite spin arrangements.

• S-orbitals have a spherical shape. • P-orbitals have a dumbbell-like shape.

• Electrons undergo spin-pair repulsion. When they spin, a magnetic field is induced causing repulsion. • To prevent this electrons initially occupy separate orbitals as they are added. • They then begin to pair up.

• Generally, electrons are gained and lost to/from the highest-energy outermost subshell. • The transition metals have an exception to the rule - the 4s subshell is filled AND emptied first. • Also note that Cr and Cu have unusual configurations - 3d5 and 3d10 are more stable.

We write from the lowest-energy subshell upwards, with a superscript number describing the number of electrons which fill the sub shell. This is based on the number of orbitals within an s, p or d sub shell - 2, 3 and 5 respectively.

• Molecules with lone electron pairs donate these to form a bond with an electron-deficient ion. • So both electrons in the covalent bond come from one ion.

• The outer shell electrons of a metal atom are donated to form a pool of delocalised electrons. • This leads to strong electrostatic attractions between the oppositely-charged ions. • The donated electrons are now delocalised so can move freely through the lattice structure, carrying a charge.

• We use valence shell electron pair repulsion theory. This essentially means that, pairs of electrons repel to move as far away as possible. • Lone pairs have a greater repulsive force than bonded pairs.

• The ability of an atom to attract a bonded pair of electrons in a covalent bond. • This exhibits an increase across the period and a decrease down the group.

• Polar bonds are formed with atoms of varying electronegativities. Where one atom is more electronegative, electrons are attracted more - inducing a dipole. • Non-polar bonds form when atoms of equal electronegativity interact, so the electron pair is equally distributed.

These are non-polar, because the electronegativities of the atoms involved are the same. That means that the electrons in their overlapping orbitals are shared evenly and equally between them; so the bond between them is not polarised and and therefore has no dipole.

• Two molecules with permanent dipoles attract one another where their constituent atoms with opposite charges attract one another. • Therefore an intermolecular force occurs - weakly, since no electrons are actually transferred.

Random electron motion in all molecules causes the formation of an instantaneous dipole moment. This induces a dipole moment in another neighbouring molecule, causing a slight attraction between the molecules.

• Molecules with a permanent dipole (polar molecules) can induce a dipole moment in another neighbouring molecule supplementary to the pre-existing London forces. • This causes a weak induced dipole, and therefore a weak attraction between the two molecules. • The resulting induced dipole then induces dipoles on neighbouring molecules - hence the attraction spreads throughout the system.

• As the number of electrons increases, the greater the fluctuation in the movements of these electrons. • This causes the difference in the electronegative and electropositive ends to increase. • Therefore the London forces become stronger.

• This occurs between a highly electronegative atom with a lone pair of electrons and the the lone pair of a hydrogen atom in another molecule. • This is the strongest type of intermolecular bonding.

• Solid ice is less dense than liquid water. This occurs because the hydrogen bonds arrange the water molecules in a more open lattice structure, so the water molecules in ice are further apart than in water. • Water has a high melting/boiling point. This is because hydrogen bonds require extra forces to overcome as they are stronger.

• Bonds between atoms continues indefinitely, there are no molecules. • Strong covalent bond between all atoms. This gives rise to a high melting point and boiling point. • For example, the allotropes of carbon.

• Within molecules there are strong covalent bonds. • Between these molecules we find much weaker intermolecular bonds, which require less energy to overcome. • This gives rise to a much lower melting/boiling point.