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Chapter 8 Covalent Bonding.

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Presentation on theme: "Chapter 8 Covalent Bonding."— Presentation transcript:

1 Chapter 8 Covalent Bonding

2 Section 1 Molecular Compounds

3 Section 1 Learning Targets
8.1.1 – I can distinguish between the melting points and boiling points of molecular compounds and ionic compounds – I can describe the information provided by a molecular formula.

4 Molecules and Molecular Compounds
Covalent bond – atoms held together by sharing electrons. Molecule – neutral group of atoms joined together by covalent bonds.

5 Molecular compound – compound composed of molecules.
Diatomic molecule – molecule consisting of two atoms.

6 Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds.

7 Molecular Formulas Molecular formula – chemical formula of a molecular compound. A molecular formula shows how many atoms of each element a molecule contains.


9 Molecular formulas can not tell you about a molecules structure.

10 Section 2 The Nature of Covalent Bonding Part 1

11 Section 2 Learning Targets
8.2.1 – I can describe how electrons are shared to form covalent bonds and identify exceptions to the octet rule – I can demonstrate how electron dot structures represent shared electrons – I can describe how atoms form double or triple bonds.

12 Section 2 Learning Targets
8.2.4 – I can distinguish between a covalent bond and a coordinate covalent bond and describe how the strength of a covalent bond is related to its bond dissociation energy – I can describe how oxygen atoms are bonded in ozone.

13 The Octet Rule in Covalent Bonding
In covalent bonding, electron sharing usually occurs so that atoms can attain the electron configuration of the noble gases.

14 Single Covalent Bonds Structural formula – represents the covalent bonds by dashes and shows the arrangement of covalently bonded atoms.

15 An electron dot structure, such as H:H represents the shared pair of electrons of the covalent bond by two dots.

16 Single covalent bond – two atoms held together by sharing a pair of electrons.
Unshared pair – (lone pair) or nonbonding pair of electrons.

17 In the water molecule the two hydrogen atoms share electrons with the one oxygen to attain a noble gas electron configuration.

18 The halogens form single covalent bonds because they have seven valence electrons and need one more to attain the electron configuration of a noble gas.

19 Ammonia has three bonds and one unshared pair of electrons.

20 Carbon behaves differently than expected.
What you expect is: What happens is: This movement of electrons allows carbon to make four covalent bonds.


22 Double and Triple Covalent Bonds
Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two or three pairs of electrons. Double covalent bond – bond that involves two shared pairs of electrons. Triple covalent bond – bond formed by sharing three pairs of electrons.

23 Oxygen forms a double bond by sharing two pairs of electrons.

24 Nitrogen’s three 2p electrons allow it to form triple bonds.

25 Some elements exist as diatomic molecules.
There are seven of them.

26 Double and triple bonds can exist in molecules that are not diatomic.

27 Section 2 The Nature of Covalent Bonding Part 2

28 Coordinate Covalent Bonds
Coordinate covalent bond – covalent bond in which one atom contributes both bonding electrons. Represented by an arrow pointing to the atom that accepts the electrons.

29 In the CO molecule the oxygen is stable with the double bond but the carbon is not.

30 The problem is solved when oxygen donates a pair of electrons to the bond.

31 Polyatomic ions – a tightly bound group of atoms that has a positive or negative charge and behaves like a unit.

32 Polyatomic ions usually contain both covalent and coordinate bonds.


34 Bond Dissociation Energy
Bond dissociation energy – energy required to break the bond between two covalently bonded atoms. A large bond dissociation energy corresponds to a strong covalent bond. Type of Bond Dissociation Energy (kJ/mol) C-C Carbon single 347 C=C Carbon double 657 C≡C Carbon triple 908 H-H Hydrogen single 435

35 Resonance Resonance structures – structure that occurs when it is possible to draw two or more valid electron dot structures. Usually seen with double bonds – where they could shift around.

36 The actual bonding in ozone (O3) is a hybrid, mixture, of the extremes represented by the resonance forms.

37 Exceptions to the Octet Rule
The octet rule can not be satisfied in molecules whose total valence electrons in an odd number. There are also molecules in which an atom has fewer, or more, than a complete octet of valence electrons.

38 Some boron compounds have an unfilled octet.

39 Sulfur and phosphorus compounds expand to more than an octet.

40 Section 3 Bonding Theories

41 Section 3 Learning Targets
8.3.2 – I can describe how VSEPR theory helps predict the shapes of molecules.

42 VSEPR Theory VSEPR (Valence Shell Electron Pair Repulsion Theory
VSEPR video According to VSEPR the repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible.

43 Unshared pairs are just as important as bonds because they help determine the shapes of the molecules. Unshared pairs of electrons are held closer to the nucleus and push bonded atoms out of their way. Tetrahedral angle – a bond angle of 109.5° that results when a central atom forms four bonds.


45 I’m giving you a handout instead of this table for your notes
Linear Triatomic Trigonal Planar Bent Triatomic Pyramidal Tetrahedral Trigonal Bipyramidal Sawhorse or Seesaw Octahedral Square Pyramid Square Planar T-Shaped I’m giving you a handout instead of this table for your notes



48 How many bonds do each of these make?
Let’s Practice On the back of your yellow paper answer the following: How many bonds do each of these make? Carbon Nitrogen Oxygen Fluorine Phosphorus Sulfur Chlorine Bromine Iodine Hydrogen Group 5 Group 6 Group 7

49 Draw the molecules and tell their molecular shape
CO BF3 Shape = Shape =

50 Draw the molecules and tell their molecular shape
SO2 CH4 Shape = Shape =

51 Draw the molecules and tell their molecular shape
NH3 H2O Shape = Shape =

52 Draw the molecules and tell their molecular shape
PF5 SF4 Shape = Shape =

53 Draw the molecules and tell their molecular shape
BrF3 XeF2 Shape = Shape =

54 Draw the molecules and tell their molecular shape
SF6 BrF5 Shape = Shape =

55 Draw the molecules and tell their molecular shape
XeF4 Shape =

56 VSEPR Video Page 233 in your text book may help with the matching on your daily work if you’re struggling.

57 More Practice CO2 NO31- Shape = Shape =

58 More Practice NO2 NH41+ Shape = Shape =

59 More Practice ICl21- ICl41+ Shape = Shape =

60 Section 4 Polar Bonds and Molecules

61 Section 4 Learning Targets
8.4.1 – I can describe how electronegativity values determine the distribution of charge in a polar molecule – I can describe what happens to polar molecules when they are placed between oppositely charged metal plates.

62 Section 4 Learning Targets
8.4.3 – I can evaluate the strength of intermolecular attractions compared with the strength of ionic and covalent bonds – I can identify the reason why network solids have high melting points.

63 Bond Polarity Electrons are not always shared equally in compounds.
Nonpolar covalent bond – when atoms in the bond pull equally the bonding electrons are shared equally.

64 Polar covalent bond (polar bond) – covalent bond between atoms in which the electrons are shared unequally.

65 The more electronegative atom attracts electrons more strongly and gains a slightly negative charge.
Slight charges are represented by Greek delta symbols (δ+ δ-)

66 The electronegativity difference can tell you the kind of bond between two atoms.
Page 177 or Chapter 6 Section 3 has this information. What kind of bonds are these?

67 Polar Molecules Polar molecule – one end of the molecule is slightly negative and the other end is slightly positive. Dipole – a molecule that has two poles.

68 When polar molecules are placed between oppositely charged plates, they tend to become oriented with respect to the positive and negative plates.


70 Attractions Between Molecules
Intermolecular attractions are weaker than either ionic or covalent bonds.

71 Van der Waals Forces Van der Waals forces – the two weakest attractions between molecules. After Dutch chemist Johannes van der Waals (1837 – 1923). Two kinds of van der Waals forces: Dipole interactions Dispersions forces

72 Dipole interactions – polar molecules are attracted to one another.
Dispersion forces – weakest of all, caused by the motion of electrons. Neighbor electrons influence other neighbors momentarily. Grey dashes represent dipole interactions

73 Hydrogen Bonds Hydrogen bonds – attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared pair of another electronegative atom.

74 Hydrogen bonds are the strongest of the intermolecular forces.
HINT: hydrogen needs to already be in a compound first then attracted to an unshared pair. Red lines represent hydrogen bonding

75 Intermolecular Attractions and Molecular Properties
Because the kinds of covalent bonds vary so do the properties of covalent (molecular) compounds. Network solid – (network crystal) solid in which all of the atoms are covalently bonded to each other.

76 Melting a network solid would require breaking covalent bonds throughout the solid.
For example: the melting point of diamond is 3500°C and then it vaporizes instead of melting.



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