Covalent Bonds Chapter 8

Review Define: Valence electrons Cation Anion Octet rule Ionic bond Ionic compound Draw Lewis structures for Carbon, Hydrogen, and Chlorine Write compound formed when boron and sulfur combine. Name it. Write compound formed when calcium and iodine combine. Name it. Write compound formed when barium and a phosphate ion combine. Name it.

Covalent bonds Sharing electrons to achieve octet (noble gases) Molecule: atoms joined together with covalent bond Diatomic molecule: molecule of 2 atoms H2, N2, O2, F2, Cl2, Br2, I2 Usually the same atom, but CO is also considered diatomic molecule Compound: two or more DIFFERENT atoms held together by covalent OR ionic bond

Molecular compounds Lower melting and boiling points than ionic compounds Many are gases or liquids at room temperature 2 or more nonmetals/metalloids (Groups 4A-7A) Molecular formula: shows how many atoms of each element a molecule contains H2O is 2 hydrogen atoms and 1 oxygen atom Whereas, NaCl is a 1:1 ratio of sodium atoms to chloride atoms

Types of covalent bonds Single covalent bond: sharing a single pair of electrons (H2) Double covalent bond: sharing 2 pairs of electrons (O2) Triple covalent bond: sharing 3 pairs of electrons (N2)

Coordinate covalent bonds One atom contributes both bonding electrons The shared electron pair comes from one of the bonding atoms Carbon monoxide and polyatomic ions

Naming covalent bonds Name elements in order listed in formula Use prefix to indicate number of each kind of atom Omit mono- when formula contains one atom in the first position Suffix of second element is –ide H2O: Dihydrogen monoxide CO: Carbon monoxide Table 9.4 on page 269

Bond dissociation energies Energy required to break covalent bond Large bond dissociation energy = strong covalent bond C-C 347KJ/mol C=C 657KJ/mol C≡C 908KJ/mol Strong carbon-carbon bonds help explain stability of carbon compounds Compounds with C-C or C-H single covalent bonds (methane) are unreactive because dissociation energy is so high

Resonance Resonance structure: occurs when it is possible to draw 2 or more valid electron dot structures that have same # of electron pairs for a molecule or ion Electrons do not change back and forth Ozone (O3)

Exceptions Total # valence electrons is an ODD number. There are also molecules in which atoms have fewer, or more, than 8 NO2, ClO2, and NO Several with EVEN # of valence electrons Atom acquires less than octet Some atoms, S and P, sometimes expand octet to include 10 to 12 electrons PCl5 and SF6

Molecular orbitals When 2 atoms combine, atomic orbitals overlap to produce molecular orbitals Atomic orbital belongs to an atom and molecular orbital belongs to the molecule 2 electrons in each atomic orbital and 2 electrons therefore in molecular orbital

Valence-shell electron-pair repulsion theory (VSEPR) The repulsion between electron pairs causes molecular shapes to adjust so that electrons stay as far apart as possible Unshared pairs are important in predicting shapes of molecules Tetrahedral: Methane (CH4) Pyramidal: Ammonia (NH3) Bent: Water (H2O) Linear: Carbon dioxide (CO2)

Review Define electronegativity and describe trend Table 6.2 p177 Electronegativity difference tells you what kind of bond is likely to form

Polar bonds and molecules Nonpolar covalent bond: atoms pull equally and electrons are equally shared H2, O2, Cl2 Polar covalent bond: electrons are shared unequally More electronegative atom attracts electrons more strongly and gains a slight negative charge. The less electronegative atom has a slightly positive charge. Increase electronegativity, increase polarity of bond Electronegativity greater than 2.0=ionic bond

Polar molecules One end of the molecule is slightly - and other end is slightly + Molecule with 2 poles is called a dipole Effect of polar bonds on the polarity of entire molecule depends on shape and orientation of polar bonds Bent=polar Linear=nonpolar Tetrahedral=nonpolar Pyramidal=polar

Attractions between molecules Intermolecular attractions are weaker than either ionic or covalent bonds Van der Waals forces (weakest) Dipole interactions and dispersion forces Hydrogen bonds

Dipole interactions Occur when polar molecules are attracted to one another Attractions occurs between oppositely charged regions of polar molecules

Dispersion forces Occurs between polar or non-polar molecules Caused by movement of electrons Weakest of all molecular interactions

Hydrogen bonds Attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom Must involve H because it is only chemically reactive element with valence electrons that are not shielded from nucleus by other electrons Bonded with oxygen, nitrogen, or fluorine Strongest intermolecular force Surface tension Cohesion/adhesion

Hydrogen bonding

Surface tension “skin” on surface of water Water molecules at surface cannot form H-bonds with air molecules so they are drawing into body of the liquid producing surface tension Water is pulled together creating the smallest surface area possible

Cohesion and adhesion Cohesion: water attracted to other water molecules Adhesion: water attracted to other materials *Capillary action*