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Covalent Bonding

Covalent Bonding. Bonding models for methane, CH 4 . Models are NOT reality. Each has its own strengths and limitations. CA Standards. Students know atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by exchanging electrons to form ionic bonds.

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Covalent Bonding

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  1. CovalentBonding Bonding models for methane, CH4. Models are NOT reality. Each has its own strengths and limitations.

  2. CA Standards • Students know atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by exchanging electrons to form ionic bonds. • Students know chemical bonds between atoms in molecules such as H2, CH4, NH3, H2CCH2, N2, Cl2, and many large biological molecules are covalent. • Students know how to draw Lewis dot structures.

  3. The Octet Rule and Covalent Compounds • Covalent compounds tend to form so that each atom, by sharing electrons, has an octet of electrons in its highest occupied energy level. • Covalent compounds involve atoms of nonmetals only. • The term “molecule” is used exclusively for covalent bonding

  4. The OctetRule: The Diatomic Fluorine Molecule F 1s 2s 2p Each has seven valence electrons F 1s 2s 2p F F

  5. The OctetRule: The Diatomic Oxygen Molecule O 1s 2s 2p Each has six valence electrons O 1s 2s 2p O O

  6. The OctetRule: The Diatomic Nitrogen Molecule N 1s 2s 2p Each has five valence electrons N 1s 2s 2p N N

  7. Lewis structures show how valence electrons are arranged among atoms in a molecule. Lewis structures Reflect the central idea that stability of a compound relates to noble gas electron configuration. Shared electrons pairs are covalent bonds and can be represented by two dots (:) or by a single line ( - ) Lewis Structures

  8. Hydrogen (and Halogens) form one covalent bond Oxygen (and sulfur) form two covalent bonds One double bond, or two single bonds Nitrogen (and phosphorus) form three covalent bonds One triple bond, or three single bonds, or one double bond and a single bond Carbon (and silicon) form four covalent bonds. Two double bonds, or four single bonds, or a triple and a single, or a double and two singles The HONCRule

  9. Completing a Lewis Structure -CH3Cl • Make carbon the central atom (it wants the most bonds, 4) • Add up available valence electrons: • C = 4, H = (3)(1), Cl = 7 Total = 14 • Join peripheral atoms • to the central atom with electron pairs. H .. .. .. .. .. H C Cl .. .. • Complete octets on • atoms other than hydrogen with remaining electrons H

  10. Bond Length and Bond Energy

  11. Occurs when more than one valid Lewis structure can be written for a particular molecule. Resonance • These are resonance structures. • The actual structure is an average of • the resonance structures.

  12. Resonance in Benzene, C6H6

  13. VSEPR Theory • Repulsion between electron pairs causes molecular shapes to adjust • Valence electron pairs stay as far apart as possible • Bonded/shared pairs & un-shared pairs used to predict shape

  14. Bond Polarity • Nonpolar Covalent Bond • Electron pairs pulled (relatively) equally • Electronegativity difference 0.0 – 0.4 H –H: 2.1 – 2.1 = 0.0 • Polar Covalent Bond • Electron pairs shared unequally • Electronegativity difference 0.4 – 2.0 • More electronegative atom attracts electrons & has a slightly negative charge H=2.1, Cl=3.0 δ+δ- + H – Cl H – Cl

  15. Bond Polarity • Ionic Bond • Electronegativity difference greater than or equal to 2.0 Na+Cl- 3.0 – 0.9 = 2.1

  16. Polar Molecules • Water (H2O) and Carbon Dioxide (CO2) O(3.5) – H(2.1) = 1.4 O(3.5) – C(2.5) = 1.0 O O C O H H δ- δ- δ+ δ- δ+ δ+ Water is polar b/c it is has bent geometry & polarities do not cancel; CO2 is linear and bond polarities cancel

  17. Intermolecular Forces • Intermolecular attractions are weak attractions; weaker than ionic or covalent bonds • Van der Waals Forces • Dipole Interactions -occur when polar molecules are attracted to each other -slightly negative region of polar molecule weakly attracted to slightly positive region of another polar molecule 2) Dispersion Forces -weakest of all molecular interactions -caused by motion of electrons -occur in nonpolar molecules -strength increases as number of electrons increase

  18. Intermolecular Forces • Hydrogen Bonds • A H covalently bonded to very electronegative atom is also weakly bonded to unshared electron pair of another electronegative atom • May be in same molecule or nearby molecule • Strongest of all intermolecular forces • 5% the strength of covalent bond

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