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Covalent Bonding. Molecular Compounds 8.1.1 – I can distinguish between the melting points and boiling points of molecular compounds and ionic compounds.

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Presentation on theme: "Covalent Bonding. Molecular Compounds 8.1.1 – I can distinguish between the melting points and boiling points of molecular compounds and ionic compounds."— Presentation transcript:

1 Covalent Bonding

2 Molecular Compounds

3 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  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 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 The Nature of Covalent Bonding Part 1

11 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 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  In covalent bonding, electron sharing usually occurs so that atoms can attain the electron configuration of the noble gases.

14  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  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 The Nature of Covalent Bonding Part 2

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

35  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 (O 3 ) is a hybrid, mixture, of the extremes represented by the resonance forms.

37  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. video

40 Bonding Theories

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

42  VSEPR (Valence Shell Electron Pair Repulsion Theory  VSEPR video 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 Linear TriatomicTrigonal PlanarBent Triatomic PyramidalTetrahedralTrigonal Bipyramidal Sawhorse or SeesawOctahedralSquare Pyramid Square PlanarT-Shaped I’m giving you a handout instead of this table for your notes



48  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 Shape =  CO  BF 3

50 Shape =  SO 2  CH 4

51 Shape =  NH 3 H2OH2O

52 Shape =  PF 5  SF 4

53 Shape =  BrF 3  XeF 2

54 Shape =  SF 6  BrF 5

55 Shape =  XeF 4

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

57 Shape =  CO 2  NO 3 1-

58 Shape =  NO 2  NH 4 1+

59 Shape =  ICl 2 1-  ICl 4 1+

60 Polar Bonds and Molecules

61 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 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  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. What kind of bonds are these? Page 177 or Chapter 6 Section 3 has this information.

67  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  Intermolecular attractions are weaker than either ionic or covalent bonds.

71  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 – 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  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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