Chapter 8 Covalent bonding

Valence Electrons Elements with similar chemical behavior have the same number of valence electrons. For the representative elements (1A, 2A, 3A, 4A, 5A, 6A, 7A, 8A) the group number corresponds to the number of valence electron in each group (with the exception of He) When examining electron configurations, the electrons that are present in the highest principle energy level represent the valence electrons of those atoms. Br: [Ar] 4s2 3d10 4p5 Bromine has 7 valence electrons

Valence Electrons and Electron Dot Structures Valence electrons are the electrons that participate in chemical bonds Electron dot structures consist of the atom symbol and its valence electrons represented as dots. Br: [Ar]4s2 3d10 4p5

Covalent Bonding Covalent bonds occur between two or more non-metals Unlike ionic bonds where electrons are transferred from one atom to another, electrons are shared between atoms in a covalent bond. Atoms joined together by covalent bonds are called molecules A compound composed of molecules is called a molecular compound

Molecular and Structural Formulas A molecular formula indicates the types and numbers of each atom in a molecule The structural formula indicate the arrangement of the atoms in the molecule H2O

Covalent Bonds and the Octet Rule Atoms share electrons in a covalent bond so that each atom has enough electrons to satisfy the octet rule

Varieties of Covalent Bonds Single bonds (sigma bonds) One pair of electrons is shared between two atoms Overlap of orbitals occurs between the two nuclei Lone pair

Varieties of Covalent Bonds Double Bonds (1sigma bond, 1 pi bond) Atoms share two pairs of electrons Overlap of orbitals in the pi bond occur off to the side of adjoining nuclei Triple Bonds (1sigma bond, 2 pi bonds) Atoms share three pairs of electrons

Coordinate Covalent Bonds A covalent bond in which one atom contributes both bonding electrons. Carbon

Resonance Structures A condition when more than one valid Lewis structure can be written for a molecule or ion.

Exceptions to the Octet Rule Too few electrons surrounding the central atom (ex: BH3) Boron will not have a full octet, only 6 electrons. It can only achieve a full octet when another atom shares an entire pair of electrons with it (Coordinate covalent bonding) Too many electrons surrounding the central atom (ex: PCl5) An odd number of electrons

How to Draw a Lewis Structure for Molecules Predict the location of atoms If there are more than two atoms, place the least electronegative atom in the center and surround it by the remaining atoms. Hydrogen is always terminal (outside) because it can only make one bond Determine the total number of electrons if each atom had a full set of valence electrons (2 for H, 8 for all others) Add up the number of valence electron that you have to work with Subtract total valence electrons from total electrons and divide by two. This is the number of bonding pairs that are needed to put together the molecule. Connect the atoms with the number of bonds that you calculated above Add lone pairs where needed so that each atom has a full octet (except for hydrogen which can only have two electrons) Molecule Total Electrons Valence Electrons Bonding Pairs HCN

Polyatomic Ions Polyatomic ions are a cluster of non-metals that carry a charge. To draw the structure of a polyatomic ion, follow the procedure for drawing ordinary molecules but add or subtract the number of electrons gained or lost to the total number of valence electrons in your structure as indicated by the charge on the ion. Molecule Total Electrons Valence Electrons Bonding Pairs IO3-

Molecular Shape (VSEPR) Valence Shell Electron Pair Repulsion – minimizes the repulsion of shared and unshared pairs of electrons around the central atom. The shape of a molecule determines many of its physical and chemical properties. The VSEPR is based on the arrangement of bonding and lone electrons around a central atom to minimize repulsion. The repulsion of electrons creates a specific bond angle between a central atom and two terminal atoms. Lone pairs of electrons occupy more space than bonding pairs of electrons

Molecular Geometry

Electronegativity and Polarity Recall: Electronegativity is the ability of an atom to attract an electron.

Chemical bonding is like “Tug-o-War” Electronegativity Difference Bond Type Non-polar Covalent 0-0.4 Polar Covalent 0.5-2.0 Ionic >2.0

Molecular Polarity Linear Trigonal Planar Tetrahedral Molecules are either polar or non-polar Both polar and non-polar molecules may contain polar bonds. What determines whether a molecule is polar or non-polar is the symmetry of the molecule PolarBonds Present Symmetry Polar/ Non-Polar Examples No NO2 Yes Non-polar SiH4 Polar NH3 CO2 VSEPR shapes that can demonstrate symmetry are: Linear Trigonal Planar Tetrahedral

Naming Binary Covalent Compounds If there is more than one of the 1st atom, precede the atom name by the appropriate prefix (di, tri, tetra, penta, hexa, hepta, octa, nona, deca) Example: C6O2 hexacarbon dioxide If there is only one of the first atom, do not precede the atom name by mono. CO2 = monocarbon dioxide CO2 = carbon dioxide Precede the second atom name by the appropriate prefix, including mono if there is only one of that atom. Drop the last syllable (or 2) and add –ide to the element name C2O Dicarbon monoxide 2nd Element Name 2nd Element C Carbide S Sulfide N Nitride Cl Chloride O Oxide Se Selenide F Fluoride Br Bromide P Phosphide I Iodide