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Covalent Bonding Ch. 16. The Nature of Covalent Bonding 16-1 Skip pgs 444 - 451.

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Presentation on theme: "Covalent Bonding Ch. 16. The Nature of Covalent Bonding 16-1 Skip pgs 444 - 451."— Presentation transcript:

1 Covalent Bonding Ch. 16

2 The Nature of Covalent Bonding 16-1 Skip pgs

3

4 Covalent Bonds Covalent (molecular) bond = the attraction of two atoms for a shared pair of electrons –Neither atom will have an ionic charge –Usually between 2 nonmetals (some involve metalloids)! Covalent compound = a compound whose atoms are held together by covalent bonds Molecule = an uncharged group of two or more atoms held together by covalent bonds

5 Single Covalent Bonds Single Covalent Bond = 2 atoms share one pair of electrons. –H 2, F 2, H 2 O Structural Formula = chemical formula that shows the arrangement of atoms. –H + H  H H  H - H

6 – F + F  F F  F - F – H + O  O H  O - H H H H F2:F2: H 2 O:

7 H CH 4  C + 4H  HC H H H * HC H H NH 3  N + 3H  HNH H HNHHNH H

8 Double and Triple Covalent Bonds Double Covalent bonds = bonds that involve 2 shared pairs of electrons. –O 2  O O Triple Covalent bonds = bonds that involve 3 shared pairs of electrons. –N 2  N N

9 O 2  O + O  O O  O O

10 N 2  N + N  N N  N N

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13 Bonding Theories 16-2 Skip pgs. 452–

14 VSEPR Theory VSEPR Theory states that because electron pairs repel, molecular shape adjusts so that valence-electron pairs are as far apart as possible. –Ex: H 2 O bond is NOT linear! 2H + O  O H H

15 VSEPR Geometrics A = Central Atom, X = Attached Species, E = Lone Pair of e-’s on A Total # of Attached Species Species Type Molecular Geometry Example 2AX 2 LinearCO 2 4AX 4 AX 3 E AX 2 E 2 Tetrahedral Pyramidal Bent CH 4 NH 3 H 2 O

16 A A A A A A 107° 105° Triatomic 120° Triagonal PyramidalBent

17 Linear Example

18 Tetrahedral Example

19 Pyramidal Example

20 Bent Example

21 Polar Bonds + Molecules 16-3 Part I

22 Bond Polarity Bonding pairs of electrons are pulled, as in a tug-of-war, between nuclei of atoms sharing electrons. If bonding pairs are shared equally it is a nonpolar covalent bond. –Atoms will have equal electronegativities (pg. 405) –Ex: N 2, O 2, H 2, Cl 2, CO 2 If bonding pairs are shared unequally it is a polar covalent bond. –Atoms have unequal electronegativity. –H 2 O, HCl, CO

23 Polar Molecules Polar molecule = one end of molecule is slightly negative and other end is slightly positive. Ex: HCl Electronegativity: H = 2.1, Cl = 3.0 Difference = 0.9 Ex: H 2 O H = 2.1, O = 3.5 Difference = 1.4

24 Electronegativity Differences + Bond Types Electronegativity Difference Type of BondExample 0.0 – 0.3Nonpolar CovalentH – H (0.0) 0.4 – 1.0Moderate Polar Covalent ∂+ ∂- H – Cl (0.9) 1.1 – 2.0Very Polar Covalent ∂+ ∂- H – F (1.9) > 2.0IonicNa + Cl - (2.1)

25 The polarity of a molecule depends on the shape + orientation of the bonds. –Ex: CO 2 polarity cancels out since it is linear = nonpolar molecule –Ex: H 2 O poles add up due to its bent shape = polar molecule

26 Polar: F H C H H H N H H OH Nonpolar: H HCH H O C OHCH

27 Attractions Between Molecules 16-3 Part II

28 Attractions between Molecules van der Waals forces = weakest attractions (ionic + covalent are stronger); consist of dispersion forces, dipole interactions and hydrogen bonds. 1) Dispersion forces = weakest of all molecular interactions; caused by motion of electrons. Increases as the number of electrons increases. Halogen diatomic molecules (F 2, Cl 2, Br 2, I 2 ) Fluorine + Chlorine have weak dispersion forces (less electrons); thus are gases at STP. Bromine (more electrons) is a liquid at STP, and Iodine (most electrons) is a solid at STP.

29 2) Dipole Interactions = occurs when polar molecules are attracted to one another. –The slightly negative region of a polar molecule is attracted to the slightly positive region of another polar molecule

30 –When placed in an electric field, dipole molecules become oriented with respect to (-) and (+) charge

31 3) Hydrogen bonds = attractive forces in which a hydrogen covalently bonded to a very electronegative atoms is also weakly bonded to an unshared electron pair of another electronegative atom. -Strongest of intermolecular (van der Waals) forces

32 Van Der Waals Forces Summary

33 Comparing Ionic + Molecular Properties CharacteristicIonic CmpdCovalent Cmpd Representative UnitFormula UnitMolecule Bond FormationTransfer e-’sShare pairs e-’s Type of ElementMetal + nonmetalNonmetal (possible metalloid) Physical StateSolidS, L, or G Melting PointHigh (>300°C)Low (< 300°C) Solubility in WaterHighHigh to Low Electrical Conductivity as aqueous soln GoodPoor to none

34 Network Solids Most molecules are easy to break; however, a few molecular solids are very stable. Network Solids = solids in which all atoms are covalently bonded to each other. –Solid does not “melt” until 1000°C or higher, in which it vaporizes without melting at all. –Ex: Diamond; made of carbon, each carbon bonded to 4 other carbons

35 Diamond + Silicon carbide (SiC)


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