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Chapter 7 Section 2 Notes

Chapter 7 Section 2 Notes. Covalent (=molecular) Bonding. Formed when nonmetal atoms share electrons (Ex) HF, CO 2 , NH 4 Cl, NO 3 - Molecules: the smallest building block of molecular compounds (Ex) HF: a molecular compound Polyatomic ions also have covalent bonds

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Chapter 7 Section 2 Notes

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  1. Chapter 7 Section 2 Notes

  2. Covalent (=molecular) Bonding • Formed when nonmetal atoms share electrons (Ex) HF, CO2, NH4Cl, NO3- • Molecules: the smallest building block of molecular compounds (Ex) HF: a molecular compound • Polyatomic ions also have covalent bonds (Ex) NO3- , CO32- • Diatomic molecules have covalent bonds • Molecules made of two atoms of the same nonmetal element • Technically, they are elements because they have one kind of atoms (Ex) N2, O2, F2, Cl2, Br2, I2

  3. Molecule Representations *Draw the best you can * We will be using molecular formula and structural formula

  4. Ionic Bond Formation (Ex) Na + Cl (Ex) Ca + F

  5. Covalent Bond Formation • Exclusively shows by structural formulas (=Lewis dot structures) • Must follow the octet rule • Move the lone electrons to pair up until each atom except for H atom has 8 electrons around it • Two atoms except for H can share 2, 4, or 6 electrons (Ex) Formation of H2O molecule

  6. Examples • Show the covalent bond formation of: • HCl (2) NH3

  7. (3) CO2 (4) NO3-

  8. 5) N2

  9. Useful Hints to Drawing Lewis Dot Structure • Follow the octet rule for all elements except for hydrogen (Ex) 2) A hydrogen atom 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)

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

  11. More about Covalent • Coordinate covalent bond: bond formed when an atom donates two electrons to form a bond (Ex) formation of NH4+

  12. 2) single bond: 2 electrons shared by two atoms 3) double bond: 4 electrons shared by two atoms 4) triple bond: 6 electrons shared by two atoms ** No such thing as quadruple bond 5) The strength of bond is measured with bond dissociation energy ‒ Energy required to break a bond ‒ Triple bond is the shortest in length and strongest, therefore has the highest bond dissociation energy

  13. 6) Resonance: molecular structures that differ only in the location of the double bond (Ex) Resonance of SO3 7) polyatomic ions: two or more nonmetal atoms covalently bonded and either has gained or lost (an) electron(s) (Ex) NO3‒, SO42‒, NH4+ 8) Exception to the octet rule: happens often with B, P, S, and Xe

  14. (Ex) BF3 PF5 SF6 XeF4

  15. 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 • Has dipole moment (δ+, δ−) • Ionic bonds • Consider extremely polar bonds • △EN > 2

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

  17. Properties of Ionic vs. Covalent Compounds

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