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Chapter 12:

Chapter 12: . Chemical Bonding. Expect to learn about. Types of chemical bonding Forming chemical bonds ionic covalent Lewis dot structures Bond polarity Molecular polarity VSEPR theory Geometries of molecules. Types of Chemical Bonding.

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Chapter 12:

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  1. Chapter 12: Chemical Bonding

  2. Expect to learn about.. • Types of chemical bonding • Forming chemical bonds • ionic • covalent • Lewis dot structures • Bond polarity • Molecular polarity • VSEPR theory • Geometries of molecules

  3. Types of Chemical Bonding *Chemical bonding: electrostatic attraction that holds two or more atoms together 1) Metallic bonding a) holds together metal atoms b) forms by moving electrons in the sea of electrons 2) Ionic bonding • forms between metal and nonmetal atoms • forms by transferring electrons from metal to nonmetal atoms 3) Covalent bonding (or molecular bonding) a) forms between nonmetal atoms b) forms by sharing electrons

  4. Ions? 1. form when an atom gains or loses one or more electrons 2. cations a) metal atoms become by losing valence electrons ** valence electrons = the electrons in the highest energy level b) each electron lost = 1+ c) the charge written as a positive number upper right of the atomic symbol (Ex) Ca → 2e- + Ca2+ 3. anions a) nonmetal atoms become by gaining electrons b) each electron gained = 1‒ c) the charge written as a negative number (Ex) P + 3e-→P3-

  5. Predicting Valence Electrons • *Why gain or lose electrons? • To achieve 8 valence electrons (called “octet rule”) and become pseudo noble gas • 8 valence electrons = s and p orbitals filled up

  6. Practices Draw the (Lewis) dot structure for each atom. • Ca (2) K (3) Ar (4) Al (5) Br (6) C (7) He (8) O (9) P (10) H

  7. Electron Configurations of Ions 1. In cations, electrons in the highest energy level (valence electrons) are removed first (Ex) Ga       [Ar] 4s23d104p1 Ga3+  [Ar] 3d10 2. In anions, additional electrons are added to the existing electrons (Ex) S     [Ne] 3s2 3p4   S2-   [Ne] 3s2 3p6 = [Ar]

  8. Lewis Dot Structure • Show how valence electrons are arranged among the atoms in a formula unit or molecule (Ex)

  9. Ionic Bond Formation • Cations(+) and anions(‒) attract each other by electrostatic force • # of electrons given = # of electrons accepted • Total positive charge = Total negative charge • Total charge = 0 (or neutral) (Ex) Form the ionic bond between: • Na & Cl • Ca & Br • Al & S

  10. Properties of Ionic Compounds • Crystalline solid at room temperature • Ions are arranged in repeating 3 dimensional pattern = lattice structure **lattice = rectangular or square structure like windows or doors • High melting point due to strong electrostatic attraction • Conduct electricity in solutions or molten liquid, but not in solid, of ionic compounds Write “See Figure 7.8 on Pg 195”

  11. Covalent (=molecular) Bonding • Formed when nonmetal atoms share electrons (Ex) HF, CO2, NH4Cl, NO3-

  12. Molecule Representations • We will be using molecular formula and structural formula.

  13. Useful Hints to Drawing Lewis Dot Structure • Follow the octet rule for all elements except for hydrogen (Ex) 2) A hydrogen atom (duet rule) can form only one single bond (Ex) (What’s wrong?) 3) The total number of valence electrons is conserved (Ex) :N≡O: (What’s wrong?) 4) Avoid making closed structures (Ex) CO32- (24 valence electrons)

  14. 5) If necessary, an electron can be moved to another atom within a molecule (coordinate covalent bond) (Ex) NO3‒

  15. Examples • Show the covalent bond formation of: • F2 How many bond pairs and lone (unshared) pairs of electrons? (2) CO2 (3) NO3-

  16. Multiple Bonds • Double bond = sharing 2 bonds or 4 electrons between two atoms • Triple bond = sharing 3 bonds or 6 electrons between two atoms *Bond energy (the amount of energy needed to break a bond) is higher and the bond distance is shorter for triple bonds

  17. Exception to Octet Rule • B (3 bonds), P (5 bonds), Xe (6 bonds), S (6 bonds)

  18. Bond Polarity • Shifting of the bond electrons • Bond electrons shift toward the element of higher electronegativity • Electronegativity : the ability of an atom to pull the bond electrons (Electronegativity increases across a period and decreases down a group. Leave out the noble gases) • Nonpolar (covalent) bond • 0 < △EN < 0.4 *△EN means the difference in electronegativity • Polar (covalent) bond • 0.4 < △EN < 2 • Consider ionic bonds extremely polar bonds • △EN > 2

  19. Example • Which has more polar bond? (1) HF and HCl (2) H2Se and H2O

  20. Molecular Structure • Linear • bent • trigonal planar • square planar • trigonal pyramid • tetrahedral • pyramid • bipyramid (= octahedral)

  21. VSEPR • VSEPR: Valence Shell Electron-Pair Repulsion. • The structure around a given atom is determined principally by minimizing electron pair repulsions. • Lone pairs of electrons and bigger atoms occupy greater space • Bond angle gets smaller

  22. Example Determine the molecular geometry and bond angle. • CH4 • CH3Cl • CH2Cl2 • H2O

  23. Molecular Polarity • the vector sum of all bond polarities *Don’t confuse bond polarity with molecular polarity.

  24. ∑ EN = 0 (nonpolar molecule) • Even symmetrical shape • No lone-pair electrons on the central atom • Peripheral atoms of the same element • ∑ EN ≠ 0 (polar molecule) • Uneven asymmetrical shape • Forms dipole (= two poles) *Show the dipole with:

  25. Examples

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