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Lewis Structure Shows how valence electrons are arranged among atoms in a molecule. Reflects central idea that stability of a compound relates to noble.

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Presentation on theme: "Lewis Structure Shows how valence electrons are arranged among atoms in a molecule. Reflects central idea that stability of a compound relates to noble."— Presentation transcript:

1 Lewis Structure Shows how valence electrons are arranged among atoms in a molecule. Reflects central idea that stability of a compound relates to noble gas electron configuration. Copyright © Cengage Learning. All rights reserved

2 Duet Rule Hydrogen forms stable molecules where it shares two electrons.

3 Octet Rule Elements form stable molecules when surrounded by eight electrons. Copyright © Cengage Learning. All rights reserved

4 H–H Single Covalent Bond
A covalent bond in which two atoms share one pair of electrons. H–H Copyright © Cengage Learning. All rights reserved

5 O=C=O Double Covalent Bond
A covalent bond in which two atoms share two pairs of electrons. O=C=O Copyright © Cengage Learning. All rights reserved

6 Triple Covalent Bond A covalent bond in which two atoms share three pairs of electrons. Copyright © Cengage Learning. All rights reserved

7 Steps for Writing Lewis Structures
Sum the valence electrons from all the atoms. Use a pair of electrons to form a bond between each pair of bound atoms. Atoms usually have noble gas configurations. Arrange the remaining electrons to satisfy the octet rule (or duet rule for hydrogen). Copyright © Cengage Learning. All rights reserved

8 Steps for Writing Lewis Structures
Sum the valence electrons from all the atoms. (Use the periodic table.) Example: H2O 2 (1 e–) + 6 e– = 8 e– total Copyright © Cengage Learning. All rights reserved

9 Steps for Writing Lewis Structures
Use a pair of electrons to form a bond between each pair of bound atoms. Example: H2O Copyright © Cengage Learning. All rights reserved

10 Steps for Writing Lewis Structures
Atoms usually have noble gas configurations. Arrange the remaining electrons to satisfy the octet rule (or duet rule for hydrogen). Examples: H2O, PBr3, and HCN Copyright © Cengage Learning. All rights reserved

11 Draw a Lewis structure for each of the following molecules: H2 F2 HF
CONCEPT CHECK! Draw a Lewis structure for each of the following molecules: H2 F2 HF Copyright © Cengage Learning. All rights reserved

12 Draw a Lewis structure for each of the following molecules: NH3 CO2
CONCEPT CHECK! Draw a Lewis structure for each of the following molecules: NH3 CO2 CCl4 Copyright © Cengage Learning. All rights reserved

13 Boron tends to form compounds in which the boron atom has fewer than eight electrons around it (it does not have a complete octet). BH3 = 6e– Copyright © Cengage Learning. All rights reserved

14 When it is necessary to exceed the octet rule for one of several third-row (or higher) elements, place the extra electrons on the central atom. SF4 = 34e– AsBr5 = 40e– Copyright © Cengage Learning. All rights reserved

15 CONCEPT CHECK! Draw a Lewis structure for each of the following molecules: BF3 PCl5 SF6 Copyright © Cengage Learning. All rights reserved

16 Let’s Review C, N, O, and F should always be assumed to obey the octet rule. B and Be often have fewer than 8 electrons around them in their compounds. Second-row elements never exceed the octet rule. Third-row and heavier elements often satisfy the octet rule but can exceed the octet rule by using their empty valence d orbitals. Copyright © Cengage Learning. All rights reserved

17 Let’s Review When writing the Lewis structure for a molecule, satisfy the octet rule for the atoms first. If electrons remain after the octet rule has been satisfied, then place them on the elements having available d orbitals (elements in Period 3 or beyond). Copyright © Cengage Learning. All rights reserved

18 More than one valid Lewis structure can be written for a particular molecule.
NO3– = 24e– Copyright © Cengage Learning. All rights reserved

19 Actual structure is an average of the resonance structures.
Electrons are really delocalized – they can move around the entire molecule. Copyright © Cengage Learning. All rights reserved

20 CONCEPT CHECK! Draw a Lewis structure for each of the following molecules: CO CO2 CH3OH OCN– Copyright © Cengage Learning. All rights reserved

21 Formal Charge Used to evaluate nonequivalent Lewis structures.
Atoms in molecules try to achieve formal charges as close to zero as possible. Any negative formal charges are expected to reside on the most electronegative atoms. Copyright © Cengage Learning. All rights reserved

22 Formal Charge Formal charge = (# valence e– on free neutral atom) – (# valence e– assigned to the atom in the molecule). Assume: Lone pair electrons belong entirely to the atom in question. Shared electrons are divided equally between the two sharing atoms. Copyright © Cengage Learning. All rights reserved

23 Rules Governing Formal Charge
To calculate the formal charge on an atom: Take the sum of the lone pair electrons and one-half the shared electrons. Subtract the number of assigned electrons from the number of valence electrons on the free, neutral atom. Copyright © Cengage Learning. All rights reserved

24 CONCEPT CHECK! Consider the Lewis structure for POCl3. Assign the formal charge for each atom in the molecule. P: 5 – 4 = +1 O: 6 – 7 = –1 Cl: 7 – 7 = 0 P: 5 – 4 = +1 O: 6 – 7 = -1 Cl: 7 – 7 = 0 Copyright © Cengage Learning. All rights reserved

25 Rules Governing Formal Charge
The sum of the formal charges of all atoms in a given molecule or ion must equal the overall charge on that species. Copyright © Cengage Learning. All rights reserved

26 Rules Governing Formal Charge
If nonequivalent Lewis structures exist for a species, those with formal charges closest to zero and with any negative formal charges on the most electronegative atoms are considered to best describe the bonding in the molecule or ion. Copyright © Cengage Learning. All rights reserved

27 VSEPR Model VSEPR: Valence Shell Electron-Pair Repulsion.
The structure around a given atom is determined principally by minimizing electron pair repulsions. Copyright © Cengage Learning. All rights reserved

28 Steps to Apply the VSEPR Model
Draw the Lewis structure for the molecule. Count the electron pairs and arrange them in the way that minimizes repulsion (put the pairs as far apart as possible. Determine the positions of the atoms from the way electron pairs are shared (how electrons are shared between the central atom and surrounding atoms). Determine the name of the molecular structure from positions of the atoms. Copyright © Cengage Learning. All rights reserved

29 VSEPR To play movie you must be in Slide Show Mode
PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE Copyright © Cengage Learning. All rights reserved

30 VSEPR: Two Electron Pairs
To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE Copyright © Cengage Learning. All rights reserved

31 VSEPR: Three Electron Pairs
To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE Copyright © Cengage Learning. All rights reserved

32 VSEPR: Four Electron Pairs
To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE Copyright © Cengage Learning. All rights reserved

33 VSEPR: Iodine Pentafluoride
To play movie you must be in Slide Show Mode PC Users: Please wait for content to load, then click to play Mac Users: CLICK HERE Copyright © Cengage Learning. All rights reserved

34 CONCEPT CHECK! Determine the shape for each of the following molecules, and include bond angles: HCN PH3 SF4 HCN – linear, 180o PH3 – trigonal pyramid, 109.5o (107o) SF4 – see saw, 90o, 120o HCN – linear, 180o PH3 – trigonal pyramid, 107o SF4 – see saw, 90o, 120o Copyright © Cengage Learning. All rights reserved

35 O3 – bent, 120o KrF4 – square planar, 90o, 180o
CONCEPT CHECK! Determine the shape for each of the following molecules, and include bond angles: O3 KrF4 O3 – bent, 120o KrF4 – square planar, 90o, 180o O3 – bent, 120o KrF4 – square planar, 90o, 180o Copyright © Cengage Learning. All rights reserved

36 CONCEPT CHECK! True or false: A molecule that has polar bonds will always be polar. -If true, explain why. -If false, provide a counter-example. False, a molecule may have polar bonds (like CO2) but the individual dipoles might cancel out so that the net dipole moment is zero. Copyright © Cengage Learning. All rights reserved

37 Let’s Think About It Draw the Lewis structure for CO2.
Does CO2 contain polar bonds? Is the molecule polar or nonpolar overall? Why? Copyright © Cengage Learning. All rights reserved

38 CONCEPT CHECK! True or false: Lone pairs make a molecule polar. -If true, explain why. -If false, provide a counter-example. False, lone pairs do not always make a molecule polar. They might be arranged so that they are symmetrically distributed to minimize repulsions, such as XeF4. Copyright © Cengage Learning. All rights reserved

39 Let’s Think About It Draw the Lewis structure for XeF4.
Does XeF4 contain lone pairs? Is the molecule polar or nonpolar overall? Why? Copyright © Cengage Learning. All rights reserved

40 Arrangements of Electron Pairs Around an Atom Yielding Minimum Repulsion
Copyright © Cengage Learning. All rights reserved

41 Structures of Molecules That Have Four Electron Pairs Around the Central Atom

42 Structures of Molecules with Five Electron Pairs Around the Central Atom
Copyright © Cengage Learning. All rights reserved

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