Structure 2.2 Flashcards

Sum the valence electrons for all atoms and ADD one extra electron for each negative charge on the ion. For cations, subtract electrons.

No. Hydrogen forms only one bond and fluorine is highly electronegative with one bond. Both are always placed as terminal (outer) atoms, never as the central atom.

As multiplicity increases (single to double to triple): bond strength INCREASES and bond length DECREASES. More shared electrons create stronger electrostatic attraction, pulling nuclei closer.

A covalent bond where BOTH shared electrons originate from the same donor atom. The acceptor has an empty orbital. Shown with an arrow from donor to acceptor.

NO. Once formed, it is completely identical in length, strength, and reactivity. The term coordination only refers to how it was formed, not its final properties.

Electron domains (bonds and lone pairs) repel each other and arrange to maximise distance, minimising electrostatic repulsion. This determines molecular shape.

NO. The repulsion hierarchy is LP-LP > LP-BP > BP-BP. Lone pairs spread more broadly (attracted to only one nucleus) and compress bond angles below ideal values.

A difference in electronegativity between the two bonded atoms. The shared pair is drawn toward the more electronegative atom, creating partial charges.

Using partial charges or dipole vector arrows pointing toward the more electronegative atom. Use data booklet Pauling values to justify.

YES. If the geometry is highly symmetrical (e.g. CO2, CCl4, BF3), the individual bond dipoles cancel out perfectly, giving a net zero dipole moment.

Assuming polar bonds always make a polar molecule. Symmetrical molecules (no lone pairs, identical terminal atoms) have bond dipoles that cancel. Do not confuse bond polarity with molecular polarity.

Diamond (3D tetrahedral, sp3), graphite (2D layers, sp2), and graphene (single graphite layer). Note: fullerene C60 is MOLECULAR, not giant covalent.

Large numbers of atoms are linked by strong covalent bonds in continuous 3D networks. Breaking requires massive energy to overcome these strong intramolecular bonds.

London (dispersion) forces, dipole-dipole attractions, and dipole-induced dipole forces. Hydrogen bonding is NOT classified under van der Waals in the IB syllabus.

Hydrogen must be DIRECTLY covalently bonded to F, O, or N. The H then interacts with a lone pair on F/O/N of a neighbouring molecule. H-Cl does NOT form H-bonds.

They are held together by weak intermolecular forces (not covalent bonds). Little energy is needed to overcome these weak IMFs during changes of state.

NOTHING. Boiling only overcomes intermolecular forces BETWEEN molecules. The covalent bonds WITHIN each molecule remain completely intact. Never say covalent bonds break during boiling.

Like dissolves like. Polar/ionic solutes dissolve in polar solvents (water) via ion-dipole or dipole-dipole interactions. Non-polar solutes dissolve in non-polar solvents via London forces.

The differing intermolecular forces (affinities) each component has for the stationary phase vs the mobile phase. More affinity for mobile phase = higher Rf value = travels further.

When multiple valid Lewis formulas exist, the real molecule is a composite average of all resonance forms. Use double-headed arrows between structures. The molecule does NOT flip between forms.

Delocalised electrons make all three N-O bonds equivalent, with bond lengths and strengths intermediate between single and double bonds. Bond order = 1 and one-third.

All six C-C bonds are identical at 140 pm (between single 154 pm and double 134 pm). Hydrogenation enthalpy is less exothermic than predicted. Benzene favours substitution over addition.

Elements in Period 3 and beyond (e.g. P, S, Xe) because their larger atomic radius allows 5 or 6 electron domains (10 or 12 valence electrons in Lewis counting) around the central atom, without requiring significant d-orbital involvement.

Trigonal bipyramidal with 3 equatorial positions (120 degrees apart) and 2 axial positions (90 degrees from equatorial). Lone pairs preferentially occupy equatorial positions.

FC = V - L - B/2 where V = valence electrons, L = lone pair electrons, B = bonding electrons. The best Lewis formula minimises formal charges toward zero.

The structure where formal charges are closest to zero AND any negative formal charge sits on the most electronegative atom. Avoid adjacent same-sign charges.

Sigma bonds form by HEAD-ON orbital overlap with electron density along the bond axis. Pi bonds form by LATERAL (sideways) overlap of p-orbitals with density above and below the axis.

Double bond = 1 sigma + 1 pi. Triple bond = 1 sigma + 2 pi. Every single covalent bond is a sigma bond. Free rotation occurs around sigma bonds but NOT around pi bonds.

The mixing of different atomic orbitals (s and p) on the same atom to form new equivalent hybrid orbitals. Determines geometry: sp3 = tetrahedral, sp2 = trigonal planar, sp = linear.

Count electron domains (sigma bonds + lone pairs) around the central atom: 4 domains = sp3, 3 domains = sp2, 2 domains = sp. H2O and NH3 are both sp3 (lone pairs count!).