structures and shapes of polyatomic molecules Flashcards
Lewis structures (the rules)
▪ Electrons are represented by dots (or crosses) and no distinction is made between different types of atomic orbitals (s, p, d, etc.).
▪ Electrons are (equally?) shared* between atoms to form two-electron(covalent) (single) chemical bonds, which are most clearly shown as lines.
▪ Many atoms ‘want’ to be associated with (have a share) of eight (valence)electrons (an ‘octet’) in their chemical compounds.
▪ Hydrogen is a notable exception, only ‘wanting’ to be associated with two electrons (a ‘duet’).
▪ Multiple (four-electron double and six-electron triple) bonds are possible and any ‘left over’ electrons form lone pairs localised on one particular atom.
Draw the exceptions to the Lewis structure rule because the molecules have too few electrons
Draw the exceptions to the Lewis structure rule because the molecules have too many electrons
How to work out a Lewis structure
1) Assemble your atoms and tot up the valence electrons(and account for charge, if any).
2) You may almost always assume that the least electronegative atom (not including H) is the central atom of the molecule.
3) Join the central atom to the peripheral atoms by single(two-electron) bonds.
4) For the ‘left over’ electrons (if any), assign them first as lone pairs on the peripheral atoms (except H!) as evenly as possible, and then (if necessary) as lone pairs on the central atom, even if this exceeds the octet for the central atom.
5) If the central atom ‘lacks its octet’ (has fewer than eight electrons), rearrange some of the lone pairs of the peripheral atoms to make multiple bonds to the central atom (for C, N, O, P and S but never for H)
Draw Lewis structures of hydrogenic triatomic molecules/ions
Lewis structures of Non-hydrogenic triatomic molecules/ions
Draw Lewis structures of tetra-atomic molecules and ion
define and explain formal charge
▪ The formal charge of an atom in a molecule is the difference between the number of valence electrons in the free atom and the number of electrons assigned to that atom in a Lewis structure.▪ It can be zero, positive or negative.▪ The sum of the formal charges of all the atoms must equal the overall charge (if any) of the molecule or ion in question.▪ If there are two or more possible Lewis structures, the structure with the lower formal charges is expected to be favoured.▪ For anions, the lower energy (favoured) structure is usually the one in which the more electronegative atom bears the negative charge. See Chemistry3 3rd edition, p. 222.▪ Formal charges can be reduced by resonance and hypervalence (see later)
Draw The VSEPR standard shapes
Analysis of bond-angle trends with VSEPR
▪ The classic sequence CH4 → NH3 → H2O can be expressed in VSEPR jargon as AX4 → AX3E → AX2E2; the experimentalX−A−X bond angles are 109.5º, 107.8º and 104.5º,respectively.
▪ This trend can be qualitatively explained by assuming that one pairs ‘need more space’ (or that they ‘push back’ the bonds more strongly).
▪ If we have two or more bonding pairs (BP) and two or more lone pairs (LP) on an atom, we can hand-wavingly suggest that the strength of repulsions between them go in the order
Failures of VSEPR model
- Most transition metal compounds (e.g., [CoCl4]2–, [Fe(H2O)6]2+…) because d electrons don’t form lone pairs!
- Gaseous group II halides, e.g., CaF2 (AX2, but bent rather than linear).
- Gaseous lithium oxide, Li2O (AX2E2, but linear rather than bent…).
- [C(CN)3]− is trigonal planar (flat) and not pyramidal despite the central C atom seeming to be an AX3E system.
- TeBr62− (AX6E?) is a regular octahedron
equation for dipole moment
pi=q x r
name two bonding models based on orbitals
valence bonds theory (VB), Molecular orbital theory (MO)
Compare valence bond and molecular orbital theory
VB theory essentially regards electrons as localised in bonds, whereas MO assumes that electrons are (can be) delocalised over the entire molecule.
draw the different types of overlaps
