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Chapter 6.2 and 6.5 Covalent Compounds.

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Presentation on theme: "Chapter 6.2 and 6.5 Covalent Compounds."— Presentation transcript:

1 Chapter 6.2 and 6.5 Covalent Compounds

2 Covalent Bonds Sharing Electrons
Covalent bonds form when atoms share one or more pairs of electrons nucleus of each atom is attracted to electron cloud of other atom neither atom removes an electron from the other

3 Covalent Bonding

4 Covalent Bonds Sharing Electrons Covalent bonds
space where electrons move is called molecular orbital made when atomic orbitals overlap

5 Molecules

6 Covalent Bonds Energy and Stability
Noble gases are stable (full octet) (low P.E.) Other elements are not stable (high P.E.) covalent bonding decreases potential energy because each atom achieves electron configuration like noble gas

7 Covalent Bonds Energy and Stability
because P.E. decreases when atoms bond, energy is released i.e., atoms lose P.E. when they bond loss of P.E. implies higher stability


9 Covalent Bonds Energy and Stability
potential energy determines bond length at minimum P.E., distance between two bonded atoms is called bond length bonded atoms vibrate therefore, bond length is an average length

10 Covalent Bonds Energy and Stability bonds vary in strength
bond energy is the amount of energy required to break the bonds in 1 mol of a chemical compound bond energy predicts reactivity bond energy is equal to loss of P.E. during formation

11 Bond Energies and Lengths

12 Bonds and Energy Single bonds are the longest bonds with the least bond energy Double bonds are shorter, stronger and have intermediate bond energy Triple bonds are the shortest, strongest, and have the highest bond energy

13 Covalent Bonds Electronegativity
Atoms share electrons equally or unequally nonpolar covalent bond: bonding electrons shared equally polar covalent bond: shared electrons more likely to be found around more electronegative atom

14 Covalent Bonds Electronegativity
Atoms share electrons equally or unequally difference in electronegativity can be used to predict type of bond (but boundaries are arbitrary)


16 Bond Types 0.3 1.7

17 Practice: Calculate the bond type
N and H F and F Ca and Cl C and O Polar Non-polar Ionic

18 Covalent Bonds Electronegativity
Polar molecules have positive and negative ends such molecules called dipoles δ means partial in math and science positive end—δ+ negative end—δ- example: Hδ+Fδ-

19 Electronegativity Difference for Hydrogen Halides

20 Covalent Bonds Electronegativity Polarity is related to bond strength
greater electronegativity means greater polarity greater bond strength

21 Covalent Bonds Electronegativity
Bond type determines properties of substances metallic bonds: electrons can move from one atom to another—good conductors ionic bonds: hard and difficult to break apart covalent bonds: low melting/boiling points

22 Properties of Substances with Different Types of Bonds

23 Polarity of Molecules Two atoms: bond polarity is the molecular polarity More than 2 atoms: the geometry of the molecule must be considered If the bonds are non-polar, the molecular is non-polar Some molecules with polar bonds can be non-polar

24 More Sometimes the partial charges cancel each other out because they are directly opposite each other Consider CO2 and CCl4 The symmetrical distribution of the bonds leads to cancellation of the charges

25 Polar and Non-polar Molecules

26 Drawing and Naming Lewis Electron-Dot Structures
Lewis structures represent valence electrons with dots position of electrons is symbolic (not literal) shows only the valence electrons of an atom dots around atomic symbol represent electrons

27 Lewis Structures of Second-Period Elements

28 Drawing and Naming Lewis Electron-Dot Structures Cl2 HCl

29 Drawing and Naming Lewis Electron-Dot Structures 1. Gather information
draw Lewis structure for each atom in compound; place one electron on each side before pairing determine total number of valence electrons

30 Drawing and Naming Lewis Electron-Dot Structures Drawing
2. Arrange atoms arrange structure to show bonding halogens and hydrogen usually make one bond at end of molecule carbon usually in center

31 Drawing and Naming Lewis Electron-Dot Structures Drawing
3. Distribute the dots so that each atom satisfies octet rule (except H, Be, B) 4. Draw the bonds as long dashes 5. Verify the structure by counting number of valence electrons

32 Drawing and Naming Lewis Electron-Dot Structures Polyatomic Ions
use brackets [] to show overall charge example:

33 Drawing and Naming Lewis Electron-Dot Structures Multiple Bonds
sharing two pairs of electrons is a double bond sharing three pairs of electrons makes triple bonds example:

34 Drawing and Naming Lewis Electron-Dot Structures Resonance Structures
sometimes, multiple structures are possible show all possibilities example:

35 Drawing and Naming Naming Covalent Compounds
First name: name of first element in formula usually least electronegative requires a prefix if more than one of them Second name: ends in –ide

36 Naming Covalent Compounds

37 Practice Naming 1. antimony tribromide 2. hexaboron silicide
3. chlorine dioxide 4. iodine pentafluoride 5. dinitrogen trioxide 6. ammonia

38 More Naming Practice P4S5 2. PI3 3. NO 4. N2F4 5. CO2 6. H2O

39 Molecular Shapes Determining Molecular Shapes
Three-dimensional shape helps determine physical and chemical properties valence shell electron pair repulsion (VSEPR) theory predicts molecular shapes based on idea that electrons repel one another

40 Molecular Shapes Lewis structures show which atoms are
connected where, and by how many bonds, but they don't properly show 3-D shapes of molecules. To find the actual shape of a molecule, first draw the Lewis structure, and then use VSEPR Theory.

41 VSEPR Valence Shell Electron Pair Repulsion theory. MOLECULAR GEOMETRY
Most important factor in determining geometry is relative repulsion between electron pairs.

42 MOLECULAR GEOMETRY Molecule adopts the shape that minimizes the electron pair repulsions.

43 VSEPR Rules To apply VSEPR theory:
1: Draw the Lewis structure of the molecule and identify the central atom 2: Count the number of electron charge clouds (lone and bonding pairs) surrounding the central atom. 3: Predict molecular shape by assuming that clouds orient so they are as far away from one another as possible

44 VSEPR Shape Predictor Table -

45 VSEPR Shape Predictor Table -

46 Bond Angles Lone-pairs of electrons behave as if they are slightly bigger than bonded electron pairs and act to distort the geometry about the atomic center so that bond angles are slightly smaller than expected:

47 Bond Angles Methane, CH4, has a perfect tetrahedral bond angle of 109.5°, while the H-N-H bond angle of ammonia, NH3, is slightly less at 107°, trigonal pyramidal

48 Bond Angles                                                                              The oxygen of water has two bonded electron pairs and two nonbonded "lone" electron pairs giving a total VSEPR coordination number of 4. But the geometry is defined by the relationship between the H-O-H atoms and water is said to be "bent" or "angular" shape of 105°. 105°

49 Molecular Shapes Determining Molecular Shapes and angles.
Let’s try some. CO CO2 BF3 CH4 CBr4 PCl3

50 Simple Shapes Bond angle for linear is 180°

51 Trigonal Planar Bond angle of 120°

52 Tetrahedral Bond angle of 109.5°

53 Bent Bond angle of 105°

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