Ch. 8 Covalent Bonding

8.1 Molecular Compounds

Sharing Electrons Atoms held together by covalent bonds (sharing electrons) Molecules are neutral groups of atoms Diatomic molecules contain only two atoms

Representing Molecules Molecular Formula: chemical formula of a molecule, shows the number of each type of element in the molecule Structural Formula: shows the bonds between atoms Space-filling Model: shows the arrangement between atoms Perspective Drawing: shows the arrangement between atoms Ball and Stick Molecular Model: shows the arrangement between atoms

Representing Molecules What is molecular formula? Structural formula?

Comparing Molecular and Ionic Compounds Representative units Molecules (I.e. C6H12O6 or C3H6O) Number of atoms of each element Composed of nonmetals Formula units (I.e. NaCl or MgCl2) Ratio of ions of each element Composed of a metal cation and a nonmetal anion

8.2 The Nature of Covalent Bonding

Octet Rule Electrons are shared (not lost or gained) to achieve the electron configuration of a noble gas (full octet unless H or He) Single bond: one pair of electrons shared Double bond: two pairs of electrons shared Triple bond: three pairs of electrons shared In a structural formula, shared pairs of electrons are signified by dashes

Octet Rule Unshared pairs of electrons are left as dots Also called “lone pairs”

Drawing Dot Structures for Molecules H2O H2 HCl F2 O2

Coordinate Covalent Bonds Coordinate covalent bond: one atom contributes both bonding electrons NH3  NH4+ Often form polyatomic ions Polyatomic ions: group of atoms with a net positive or negative charge SO32- SO42-

Exceptions to the Octet Rule Octet rule cannot be satisfied if a molecule has an odd number of valence electrons NO2, NO, ClO2 Phosphorus and sulfur expand the octet to ten or twelve electrons

Bond Dissociation Energies Energy required to break a bond Large bond dissociation energies denote strong covalent bonds

Resonance Resonance structures are multiple possible dot structures for the same molecule or polyatomic ion Actual bonding is a hybrid of both, used when neither is a sufficient description

8.3 Bonding Theories

Molecular Orbitals Formed when atoms bond Belong to the molecule as a whole (not to one atom) If occupied by electrons of a covalent bond it is called a bonding orbital

Bonding Orbitals Sigma bond: combine end-to-end Pi bond: combine side-to-side

VSEPR Theory Valence-Shell Electron-Pair Repulsion Theory Explains three-dimensional shapes of molecules Lone pairs of electrons strongly repel bonding pairs of electrons 3D shapes of molecules

Hybrid Orbitals Hybridization provides information about molecular bonding and shape Describes double and triple covalent bonds

8.4 Polar Bonds and Molecules

Bond Polarity When atoms share electrons equally, a nonpolar covalent bond is formed When atoms do not share electrons equally, a polar covalent bond is formed The more electronegative atom is partially negative The less electronegative atom is partially positive Polar molecules are formed when one end of the molecule is partially positive and the other end is partially negative A molecule with two poles is called a dipole

Attractions Between Molecules Intermolecular attractions are weaker than ionic or covalent bonds Hydrogen bonds- hydrogen bonded to very electronegative atom is also weakly bonded to another highly electronegative atom (F, O, N) Van der Waals forces Dipole interactions- based on partial charges Similar to, but weaker than ionic bonds Dispersion forces- based on movement of electrons Increase with number of electrons in a molecule

Intermolecular Attractions and Molecular Properties Varying intermolecular attractions explain the variety of physical properties among covalent compounds Network solids- all atoms are covalently bonded to each other