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Understanding Covalent Bonding: Formation, Examples, and Properties

Learn about covalent bonds formed by nonmetal atoms sharing electrons, molecular compounds, diatomic molecules, Lewis dot structures, and bond polarity in this educational guide. Explore examples like HCl, NH3, and more.

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Understanding Covalent Bonding: Formation, Examples, and Properties

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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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