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Chapter 8: Covalent Bonding. The bonds of nature. Covalent Bond. Shared valence electrons Complete outer energy levels Molecule – neutral group of atoms joined by covalent bonds - water, Carbohydrates, proteins, fats, DNA Molecular Compound – compound consisting of molecules

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Chapter 8: Covalent Bonding


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    1. Chapter 8: Covalent Bonding The bonds of nature.

    2. Covalent Bond • Shared valence electrons • Complete outer energy levels • Molecule – neutral group of atoms joined by covalent bonds - water, Carbohydrates, proteins, fats, DNA • Molecular Compound – compound consisting of molecules • Diatomic Molecule – O2, N2, Cl2 etc… P. 247 - Q 39, 42 P. 216 Q 3

    3. How do they form? Distance is to great E - E – repulsive P – E - attractive P – P – repulsive E – E - repulsive Repel = Attract Attractive forces balance the repulsive forces P. 247 - Q 44

    4. 1 valence e 1 s1 1 valence e 1 s1 1 s2 Stable Helium configuration P. 220 - Q - 7

    5. Structural Formula – represents covalent bonds by dashes and shows the arrangement of the bonded atoms Single Covalent bond – shared electrons (also known as a sigma bond σ) 1s1 1s1 1s2 2s2 2p4 H –1s2 - He O - 1s2 2s2 2p6 - Ne Lone pairs – unshared elecrons The molecules are more stable because they have complete outer energy levels – Octet Rule

    6. Drawing Lewis Dot Structures • Predict the location of the atoms • Hydrogen is a terminal atom • The central atom has the smallest electronegativity. • Count the valence electrons. • Draw a single covalent bond between the central atom and the surrounding atoms. • Subtract the number of electrons in the single covalent bonds from the total number of electrons in 2. • Use the remaining electrons to complete the octets of each atom. • If the central atom does not have a complete octet then try double or triple bonds.

    7. Drawing Lewis Dot Structures Draw Lewis Dot Structures for: PH3 H2S HCl CCl4 SiH4 CH2Cl2 P. 220 - Q 8, P. 247 - Q - 46

    8. Draw Lewis Dot Structures Cl2 NF3 CS2 BH3 CH4 SCl2 C2H6 BF3

    9. Multiple Covalent Bonds – Double and Triple Bonds 6 valence electrons 6 valence electrons 12 valence electrons Octet satisfied More stable and stronger 2 shared pairs of electrons P. 247 - Q - 45

    10. 5 valence electrons 5 valence electrons 10 valence electrons Octet satisfied More stable and stronger 3 shared pairs of electrons

    11. Strength of Bonds • Based on proximity – bond length • Influenced by atom size and number of shared electrons • F2 – 1.43 x 10-10 m - single • O2 – 1.21 x 10-10 m - double • N2 – 1.10 x 10-10 m - triple

    12. Lewis Dot Structures – Polyatomic Ions • Count the total number of electrons in the ion. • Add or subtract the extra electrons in the ion. • NH4+ - subtract an electron • CO32- - add 2 electrons

    13. Polyatomic Lewis Dot Structures • Nitrate • Sulfite • Chlorate • Hydroxide • Ammonium P. 225 - Q - 10,11,12

    14. Resonance Structures – occurs when there is a single and double bond • Structures that have more than one possible Lewis Dot Structures • atoms remain stationary • electrons move • Resonance structures behave as the same molecule • Bond lengths are shorter than single, longer than double

    15. Representing Resonance Structures Dash line represent the possible double bond P. 247 - Q - 49

    16. Resonance Practice • Draw Lewis Resonance Structures for: • Sulfur Trioxide • Sulfur Dioxide • O3 – ozone

    17. Electronegativity and Bond Type • Ability to attract electrons in a chemical bond • Ionic Bonds > 1.7 • Covalent Bonds < 1.7

    18. What types of bonds are they? MgO, water, Calcium Carbide, Potassium Oxide, Nitrogen trihydride

    19. Dash Line Represents the Cut Off

    20. Polar Covalent Bonds • Electrons are shared, but not equally • Non polar bonds – equal sharing, H2, O2, Cl2 P. 247 - Q - 58

    21. Molecular Polarity and the VSEPR Theory P. 232-233 • Symmetry – free electron pair = polar • Non-polar molecule – equal pull from the same atoms • Valence Shell Electron Pair Repulsion Theory – the repulsion between electron pairs causes molecular shapes to adjust so valence-electron pairs stay as far apart as possible

    22. Exceptions to the Octet Rule P. 228 • An odd number of electrons N = 5, O = 12 • Some form with fewer than 8 electrons (rare) Very reactive • Some atoms contain more than 8 electrons (expanded octet) • D orbitals are utilized P. 247 - Q - 50

    23. Draw Lewis Structures for the following compounds – Expanded Octets SF6 PCl5 ClF3

    24. Molecular Shape • Linear • Bent • Trigonal Planar • Trigonal Pyramidal • Tetrahedral • Trigonal Bipyramidal • Octahedral P. 236 - Q – 29, P. 247 Q - 54

    25. Coordinate Covalent Bonds P. 226 • An unshared pair of electrons are donated to another atom to achieve stability. P. 229 - Q - 16, P. 247 - Q 48, P. 236 – Q - 29

    26. Bond dissociation energy – P. 226 • energy needed to break a bond apart • F2 – 159 kJ/mol – single • O2 – 498 kJ/mol - double • N2 – 945 kJ/mol - triple Potential energy of a molecule is the sum of the bond dissociation energies P. 229 - Q - 22, P. 247 - Q - 51, 52

    27. Sigma Bonds - σ- P. 230 • Overlapping electron orbitals create a bonding orbital – the likely hood of finding an electron • S overlaps with S • S overlaps with P • P overlaps with S

    28. Hybridization P. 234-236 • The mixing of atomic orbitals in an atom to generate a new set of bonding orbitals (hybrids) • Different shapes than atomic orbitals • Requires energy but the energy is returned during bond formation sp3 P. 236 - Q - 23

    29. sp2 sp

    30. sp3d

    31. sp3d2 P. 248 - Q - 62

    32. Acids • any substance that produces H+ when dissolved in water • 2 types • Binary – hydrogen and one other element • Oxyacids – contains hydrogen and oxygen

    33. Follow Up Problems – 248 and 249 • Questions 63, 64, 65, 66, 67, 70, 72, 73, 75, 76

    34. Naming Acids Binary • Use the prefix hydro for hydrogen • Use the root of the 2nd element and add -ic • Write the word acid HF, HBr, HI, HCl Note: If a polyatomic ion does not contain oxygen these rules apply. Use the root of the polyatomic. HCN – hydrocyanic acid

    35. Naming Acids - continued • Oxyacids • Depends on the oxyanion (polyatomic ion containing oxygen) • Use the root of the oxyanion • Add the suffix • Write the word acid ate = ic ite = ous

    36. Examples • NO3 – nitrate - HNO3 -Nitric Acid • NO2 – nitrite - HNO2 -Nitrous Acid • H2SO4 • H2SO3 • HClO3 • HClO2 • H2CO3 • H3PO3