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Ch 6 Covalent bonds

Ch 6 Covalent bonds. Share electrons (between nonmetal/metaliods): to reach a stable octet of atoms in a chemical bond. Forming molecular orbital. Simplest: diatomic molecules. Attraction & repulsion between atoms balanced and at lowest energy.

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Ch 6 Covalent bonds

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  1. Ch 6Covalent bonds Share electrons (between nonmetal/metaliods): to reach a stable octet of atoms in a chemical bond

  2. Forming molecular orbital • Simplest: diatomic molecules. Attraction & repulsion between atoms balanced and at lowest energy. • Shared pair of electrons to form a covalent bond. • Covalent bond: atoms share one or more pairs of electrons. • Molecular orbital: overlap of two atomic orbits • Energy and stability: • Unbonded atoms have high potential energy • Bonded atoms have lower PE & high stability

  3. P.E. & bond length (pg 192 figure 4) • When P.E. between two atoms are at a minimum --> bond length • Covalent bonds are flexible, nuclei vibrate back and forth --> bond length is an average. • Bond energy: energy released when bonds are formed & the energy required to break bonds • Bond E & length (pg 193, table 1) units kJ/mol & pm. • Trend: increase bond strength decrease bond length.

  4. Electro negativity & covalent bonding • E.N. useful to predict type of bond(pg194, figure 6) either ionic or covalent. • Covalent bonds: • Electrons shared equally or unequally • Nonpolar covalent bond: • Bonded electrons are equally attached to both atoms. Electro negativity value the same • Polar covalent bond: • Bonded electrons is held more closely by one atom more. Electro negativity value higher than the other atom.

  5. Electro negativity and bond type • E.N. • 0-.5 nonpolar bond, .5-2.1 polar, 2.1-3.3 ionic bond • Polar molecules: • Opposite charges between atoms in a molecule • + more positive, lower E.N. • - more negative, higher E.N. • Pair of electrons attracted to the atom with higher E.N. • ==> form dipole, molecules contains both positive and negative charged region. • Greater E.N., greater polarity and bond strength. (pg 196, table 2) - O H H Water molecule +

  6. Bond character • Greater the EN the greater (higher %) of ionic character • Properties depended on bond type (pg 197 table 3) • Metallic bond: • Good electrical/heat conductor, solid rm temp. valence e- attracted to all atoms, e- move freely • Ionic substance: • Attraction between cation & anion. Conductor as liquid or aqueous solution. • Molecular: • Shared electrons, non conductor.

  7. Drawing and naming molecules • Lewis electron dot structure: • Valance electrons (outer electrons) used to form chemical bonds & determine chemical properties. • Lewis structure (electron dot diagram) • Structural formula, electrons represented with dots between two atomic symbols representing pairs of electrons in a bond. • Only show valance electrons, symbol of atom represents kernel electrons (inner electrons) • Mirror s&p orbital notation. s p X p X = symbol of element p

  8. Chemical bonds/Lewis struct. • Unshared electrons: • Nonbonding pair of electrons in the valance shell (lone pair) • Single bond: • Covalent bond in which two atoms share one pair of electrons. • Drawing Lewis structure: • 1. Determine total valance electrons. • 2. Determine total electrons need for octet • 3. Determine # of shared pair of electrons (bonds) • #in step 2 - # in step 1 divided by 2 (1 bond = 2 electrons) • 4 place bonds & lone pairs in structure. • Draw the following: SCl2, AsF3, SiH4, CHF3

  9. Lewis struct./polyatomic ions • Same as before with the addition of adding or subtracting valence electrons due to charge • Place brackets, [ ], around ion with charge on the outside. • Draw the following: HS-, NF4+, IO-, PCl4+, CH3+ • Multi bonds: • More than one shared pair of electrons between 2 atoms. • Double bond: • Covalent bond, 2 atoms share 2 pair of electrons. • Triple bond: • Share 3 pair of electron between 2 atoms (covalent) • Draw the following: CO2, HCN, HC2Cl, N2F2, NO2-

  10. Resonance structure • Molecules that can not be represented by a single Lewis structure. • Compounds have identical geometry but different arrangement of of electrons/bonds. • Compounds with multi bonds. • Draw all possible Lewis structures with double headed arrows between each. • Draw the following: NO2, O3, SO2

  11. Naming covalent cpds • Binary cpds (pg 207 table 5) • Just two nonmetal/metaliods covalently bonded • 1. Use prefixes to indicate # of atoms of each element (no mono for first element) • 1-mono, 2-di, 3-tri, 4-tetra, 5-penta, 6-hexa, 7-hepta, 8-octa, 9-nona, 10-deca. • 2. Name 1st element as is 2nd element end in ide • 3. Drop vowel in prefix if element begins with a vowel (not always) • Name or write the following: CO2, N2O5, diphosphorus tetroxide, sulfur hexafluoride

  12. Molecular shape • 3-d shape of a molecule • Help determine properties of the molecule. • Lewis structure predicts molecular shape. • VSEPR theory: • Valence shell electron pair repulsion theory • Valence electrons surrounding an atom repel each other --> spread out to maximize space between electron pairs. • Lone (unbonded) electrons contort molecular shape.

  13. Molecular geometry • Formula shape bond angle(o) • AB2 linear 180 • AB3 trigonal planar 120 • AB2E1 bent >120 • AB4 tetrahedral 109.5 • AB3E1 pyramidal >109.5 • AB2E2 bent >109.5 • AB5 trigonal bipyramidal 90 & 120 • ** symmetrical molecules are nonpolar**

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