Unit 2: Compound Structure and Properties

21 terms. Say what each one means, then open the card to check.

• Ionic bond

  The attraction between oppositely charged ions, usually formed when a metal gives electrons to a nonmetal.

  Example. In calcium fluoride, CaF₂, each Ca²⁺ ion is held by its attraction to the F⁻ ions around it.

• Covalent bond

  A bond in which two atoms, usually nonmetals, share one or more pairs of electrons.

  Example. In hydrogen chloride, HCl, the H and Cl atoms share one pair of electrons.

• Metallic bonding

  Metal cations held together in a sea of valence electrons that are free to move through the whole solid.

  Example. A copper wire conducts electricity because its valence electrons can move freely through the metal.

• Polar covalent bond

  A covalent bond between atoms of different electronegativity, so the shared electrons sit closer to one atom, which gains a partial negative charge.

  Example. In an H–Br bond, bromine pulls the shared pair closer, so Br is partly negative and H is partly positive.

• Lattice energy

  The energy needed to pull apart one mole of an ionic solid into gaseous ions. It is larger for ions with higher charges and smaller radii.

  Example. MgS has a larger lattice energy than NaCl, because its ions carry charges of 2+ and 2− instead of 1+ and 1−.

• Bond order

  The number of shared electron pairs between two atoms: 1 for a single bond, 2 for a double and 3 for a triple. Higher order means shorter and stronger.

  Example. The O=O bond in an oxygen molecule, O₂, has a bond order of 2.

• Bond length

  The distance between two bonded nuclei where the potential energy of the pair is lowest.

  Example. The N≡N triple bond in N₂ is shorter than an N–N single bond, because three shared pairs pull the nuclei closer.

• Bond energy

  The energy needed to break one mole of a bond in the gas phase. Breaking a bond always takes energy; forming one releases it.

  Example. Breaking the bonds in a mole of N₂ molecules takes far more energy than breaking a mole of N–N single bonds.

• Alloy

  A mixture of a metal with other elements. Atoms of similar size can replace metal atoms; much smaller atoms fit into the gaps between them.

  Example. In brass, zinc atoms take the places of some copper atoms, because the two atoms are close in size.

• Lewis structure

  A diagram of a molecule or ion that shows every valence electron as a bonding pair (a line) or a lone pair (two dots).

  Example. The Lewis structure of methanol, CH₃OH, shows five single bonds and two lone pairs on the oxygen atom.

• Octet rule

  Atoms in many molecules end up with eight valence electrons around them, counting both shared and lone pairs. Hydrogen needs only two.

  Example. In Cl₂, each chlorine has one shared pair and three lone pairs, for eight electrons around it.

• Lone pair

  A pair of valence electrons that belongs to one atom and is not shared in a bond.

  Example. In hydrogen sulfide, H₂S, the sulfur atom has two bonds and two lone pairs.

• Formal charge

  Valence electrons minus lone-pair electrons minus half the bonding electrons, for one atom in a Lewis structure. The best structure keeps these near zero.

  Example. In ozone, O₃, the central oxygen has 6 − 2 − 3 = +1, and the singly bonded end oxygen has 6 − 6 − 1 = −1.

• Resonance

  When one Lewis structure cannot show the bonding, the real particle is an average (a hybrid) of two or more structures that differ only in where electrons are.

  Example. In benzene, C₆H₆, all six carbon-carbon bonds are the same length, between a single and a double bond.

• Expanded octet

  More than eight electrons around a central atom from period 3 or below. Period 2 atoms never exceed eight.

  Example. In PCl₅, the phosphorus atom has five bonds, so ten electrons surround it.

• VSEPR model

  Electron domains around a central atom repel one another and spread out as far apart as they can, which sets the shape.

  Example. BeCl₂ has two electron domains around Be, so they point in opposite directions and the molecule is linear.

• Electron domain

  A region of electron density around a central atom: one lone pair, or one bond of any order (single, double or triple).

  Example. The carbon in formaldehyde, H₂CO, has three domains: two C–H bonds and one C=O double bond.

• Molecular geometry

  The arrangement of the atoms only, ignoring lone pairs, even though the lone pairs still help set it.

  Example. H₂S has four domains arranged in a tetrahedron, but with two of them lone pairs its molecular geometry is bent.

• Hybridization

  The mixing of an atom's orbitals to match its number of electron domains: 2 is sp, 3 is sp² and 4 is sp³.

  Example. The carbon in formaldehyde, H₂CO, has three electron domains, so it is sp² hybridized.

• Sigma and pi bonds

  Every single bond is one sigma bond. A double bond is one sigma and one pi bond, and a triple bond is one sigma and two pi bonds.

  Example. The triple bond in N₂ is made of one sigma bond and two pi bonds.

• Polar molecule

  A molecule with a net dipole: it has polar bonds arranged so that their pulls do not cancel.

  Example. Sulfur dichloride, SCl₂, is bent, so its two polar S–Cl bonds do not point in opposite directions and their dipoles do not cancel.