1. Chemistry 068, Chapter 10

2. Bonding, Lewis Structures, and Molecular Geometry
Chemical compounds can be broadly split into two different kinds of compounds – ionic and molecular (or covalent).
They are distinguished from one another by the type of chemical bonding they undergo.
Chemical bonds are attractive forces that hold two or more atoms together.
To represent the shape of compounds, chemists often use Lewis structures, which are a short hand way to represent the arrangement of atoms within a molecule.
Lewis structures are only two dimensional representations; three dimensional geometry is handled by a variety of theories.

3. Ionic and Covalent Bonding
An ionic bond is a chemical bond formed by electrostatic attraction between positive and negative ions.
A covalent bond is a chemical bond formed by two or more atoms sharing electrons.

4. Bonding and Valence Electrons
Not all of the electrons in an atom participate in bonding.
Only the outermost electrons, called valence electrons, participate in bonding.
Only the outermost shell can bond.
Technically, only s and p electrons are valence electrons.

5. Lewis Dot Structures
Lewis dot structures represent the molecular geometry and valence shell electrons of a molecule.
Sometimes called Lewis formulas, diagrams or symbols.
Most useful for period 2 and 3 atoms and molecules.
Only includes valence electrons.

6. Lewis Dot Structures (Cont’d)
Electrons are represented as dots and dashes, showing the positions of both bonding and nonbonding electrons.
Elements are represented by their atomic symbols.
The positions of the atoms (without electrons) is referred to as the skeleton structure.

7. Lewis Dot Structures of Elements
Determine the number of valence electrons from the periodic table.
The skeleton structure consists only of the symbol for the element.
Write dots/dashes to represent the electrons.
By convention, electrons are first written singly, then paired.

8. Lewis Dot Structures of Elements Problems
Write Lewis dot structures for the following atoms. C Na O S Ar

9. The Octet Rule
The octet rule is the tendency of atoms within molecules to try to have 8 valence electrons.
This corresponds to a noble gas configuration.
Atoms will gain/loose electrons via bonding to try to reach this configuration.
Both bonding and nonbonding electrons count toward the octet rule.
Most atoms follow the octet rule.
Lewis structures that violate the octet rule are generally less stable than those that obey it.
Unless they are an exception to the rule.
The maximum number of electrons on an atom is NOT eight.

10. Exceptions to the Octet Rule
Atoms with more or fewer than 8 valence electrons will violate the octet rule.
Hydrogen, for example, has 2 valence electrons.
This is sometimes called the duet rule.
Atoms from the fourth and lower period can have more than 8 electrons because of the d orbital.
Phosphorous can also make additional bonds by filling the 3d orbital.
Metals and group IIIA also violate the octet rule.
They instead make one bond per valence electron (thus 1, 2, or 3 bonds).

11. Ionic Bonding
The ionic bond forms when electrons are transferred from one valence shell to another.
The cation looses electrons.
The anion gains electrons.
The electron transfer occurs so that each atom has a noble gas configuration.
Each atom will have a full valence shell.

12. Ionic Bonding (Cont’d)
Most ionic compounds are solids at room temperature.
This is due to the strong electrostatic attractions between atoms.
It is also due to the lattic energy.
Most ionic solids have very high melting points.
It takes energy to overcome electrostatics and the lattice energy and move molecules far enough apart to melt.

13. Isoelectronic Species
Isoelectronic species have the same number of electrons.
This does NOT mean they have the same properties, they are only alike in the number of electrons they posses.

14. Lewis Structures of Ionic Compounds
Typically, the cation will be the central atom with the anions surrounding it.
Metal atoms (cations) make one bond for each valence electron that they have.
These can be single, double, or triple bonds.
In the case of polyatomic ions, they will be bonded together as units around the central atom.

15. Writing Lewis Structures of Ionic Compounds
Determine the number of valence electrons on each atom.
Write the skeleton structure of the molecule and add bonds between atoms. The metal atom will make one bond for each valence atom.
These can be single, double, or triple bonds.
Distribute electrons to the atoms surrounding the central atom.
Ensure that every atom which can have an octet does.
The metal atoms will not have octets.

16. Drawing Lewis Structures of Ionic Compounds Problems
Draw Lewis structures for each of the following:
NaCl
CaO
K2O
BF3

17. Covalent Bonding
The vast majority of compounds are formed by covalent bonds.
Covalent bonds form when the valence orbitals of atoms overlap.
This is how atoms “share” electrons.
The bond length is the distance between nuclei which is most stable.
This is the distance which minimizes the potential energy of the bond.
A coordinate covalent bond is formed when both electrons of a bonding pair are donated by a single atom.

18. Covalent Bonding (Cont’d)
Bonds are formed when valence electrons are shared between atoms.
Bonds are always made up of pairs of electrons – this corresponds to filled orbitals.
Non-bonding pairs of electrons are called lone pairs – again, pairs of electrons to be filled orbitals.

19. Covalent Bonding (Cont’d)
Covalent bonds can be single, double, or triple bonds. These correspond to one, two, or three pairs of electrons.
So 2, 4, or 6 electrons shared between two atoms.
Quadruple bonds (4 pairs, 8 electrons) are not geometrically possible.

20. Lewis Structures of Covalent Compounds
While only experiment can determine the molecular geometry, the central atom is usually the least electronegative atom (other than hydrogen.)
Atoms will make a number of bonds equal to the number of electrons they need to achieve a noble gas configuration.
These can be single, double, or triple bonds.
Most atoms will obey the octet rule.

21. Writing Covalent Lewis Structures
Calculate the total number of valence electrons.
Write the skeleton structure of the molecule and add bonds between atoms.
Distribute electrons to the atoms surrounding the central atom.
Distribute any remaining electrons to the central atom.
Rearrange electrons as needed, so that each atom has an octet.

22. Calculating Valence Electrons
Add up the number of valence electrons for each atom in the molecule.
Only count valence electrons.
Usually you do not include d or f orbital electrons.
Count along the periodic table to get the number of valence electrons.

23. Writing the Skeleton Structure
How the individual atoms are arranged has to be determined experimentally.
Usually, the central atom will be surrounded by more electronegative atoms.
Hydrogen is a notable exception – it is never a central atom.
Indicate bonds with dash marks between atoms. Use one dash per bond.

24. Distributing Electrons to the Surrounding Atoms
Add lone pairs to the atoms surrounding the central atom.
Add enough electrons so that each atom has an octet.
Represent lone pair electrons with dots, using two dots per pair (one per electron).

25. Distributing Electrons to the Central Atom
Add lone pairs to the central atom.
Add enough electrons so that the central atom has an octet.
Represent lone pair electrons with dots, using two dots per pair (one per electron).

26. Rearrange Electrons Add up the number of electrons you’ve drawn.
If it is greater than the number of valence electrons, replace lone pairs with bonds between atoms.
As needed, replace two lone pairs (one on each of two atoms) with one extra bond.

27. Different Ways to Write Lewis Structures
Some text books use a different way to draw Lewis structures.
Some books use all dots or all dashes to write structures.
In organic chemistry, hydrogens are often left out of a structure, and are written as a dash only.

28. Drawing Lewis Structure Problems
Draw Lewis structures for each of the following:
NH4+
PF3
H2O
CH4

29. Bond Order, Strength, and Length
Bond order is the number of electron pairs involved in a bond.
This is the same as a bond being single (bond order 1), double (bond order 2), or triple (bond order 3).
As bond order increases, the bond strength also increases.

30. Bond Order, Strength, and Length (Cont’d)
Bond length is the distance between bonded nuclei.
As bond order increases, bond length will decrease.
This is because the stronger bond will pull the nuclei closer together.

31. Bond Order, Strength, and Length (Cont’d)
Bond energy is the average enthalpy change to break a bond in the gas phase.
Bond energy increases with bond order, as is expected for a stronger bond.
Bond energies can also be used to estimate the enthalpy of a reaction by comparing bonds formed/broken during the reaction.

32. Resonance Resonance structures show delocalized bonding.
Delocalized bonding is bonding spread over multiple atoms rather than just two.
Resonance structures show the different possible arrangements of electrons.
Electrons move, atoms do not.
Some resonance structures will be more stable than others. The electrons will spend most of their time in that arrangement.

33. Isomers
Isomers are molecules that have the same formula but a different structure.
The atoms will be arranged differently.
Unlike resonance structures, the atoms are arranged differently, rather than the electrons.

34. Isomer and Resonance Problems
Find all possible structures for each of the following, ignoring structures that violet the octet rule:
C2H2Cl2 O3
HCN
CO2

35. The VSEPR Model Valence Shell Electron Pair Repulsion.
Predicts molecular shape by assuming that the valence shell pairs are arranged to minimize electron pair repulsions by staying the maximum possible distance apart from one another.
Electrostatic repulsion is thus minimized.
Repulsion of electron pairs is a source of instability.

36. The VSEPR Model (Cont’d)
Molecules are three dimensional.
Both bonds and nonbonding (lone) pairs take up space.
Atoms and electron pairs are arranged to be as far apart as possible – to minimize the interactions.

37. Electron Groups VSEPR uses the idea of electron groups.
Electron groups are valence electron pairs present on an atom.
Bonds (including multiple bonds) count as one group.
Nonbonding pairs also count as one group per nonbonding pair.

38. Effect of Lone Electron Pairs on Shape
Lone pairs not only take up space but also push other atoms away.
Lone pairs take up more “space” than atoms do because they repel the bonding pairs more than other atoms do.
They have no positive component (no protons).
The most stable geometry minimizes interactions between bonding and nonbonding pairs. This means keeping them as far apart as possible.
Lone pairs usually shift bond angles by several degrees.

39. Effect of Multiple Bonds on Shape
Atoms with multiple bonds take up slightly more space than single bonds.
Shifts are much smaller than those do to lone pairs.
The effects of multiple bonds are usually ignored in regards to shape as far as naming goes.

40. Determining Shape Using VSEPR
First, write the Lewis structure.
Next, determine the number of electron groups around the central atom. Count bonds as one pair.
Determine the general shape from the number of electron groups.
Determine the specific shape from the number of nonbonding (lone) electron pairs.

41. One Electron Group Shapes
Molecules of the sort AX (1 total pair) are linear.
The only possibility for two atoms is a linear shape, regardless of nonbonding pairs.
Hence it not being included in the text.

42. Two Electron Group Shapes
Molecules of the sort AX2 (2 total pairs) are also linear.
No nonbonding pairs on the central atom, and two bonding pairs.
The only possibility is a linear shape with 180o angles.

43. Three Electron Group Shapes
Molecules of the sort AX3 or AX2E (3 total pairs) are called trigonal planar.
All atoms and nonbonding pairs on the central atom will be in the same plane.
Angles are nominally 120o.
There are two specific geometries.
Trigonal planar for AX3.
Bent (or angular) for AX2E.

44. Four Electron Group Shapes
Molecules of the sort AX4, AX3E, or AX2E2 (4 total pairs) are called tetrahedral.
No longer all in the same plane – only three atoms are in one plane.
Angles are nominally 109.5o.
There are three specific geometries.
Tetrahedral for AX4.
Trigonal pyramidal for AX3E.
Bent (or angular) for AX2E2.

45. VSEPR Problems
Use VSEPR to assign a molecular geometry to each of the following molecules:
CH4
KCl
CO2
NH3
BCl3

46. Electronegativity
Electronegativity is a measure of an atoms ability to draw electrons to itself, specifically the electrons in a bond.
It increases from left to right and bottom to top on the periodic table.
With the exception of the noble gasses.
Fluorine (F) is the most electronegative element.

47. Electronegativity (Cont’d)
A comparison of the difference in electronegativity of two atoms gives a rough qualitative indication of the ionic or covalent character of bonds formed between them.
A large difference, such as between a metal and a nonmetal indicates an ionic bond.
Conversely, a small difference indicates a covalent bond.

48. Polar Bonds
A polar covalent bond is one in which the bonding electrons spend an unequal amount of time between the two nuclei.
They spend most of their time near the more electronegative atom.
Thus, there is uneven sharing of the electrons.
Polar covalent bonds are an intermediate between non-polar covalent bonds and ionic bonds.
Shared electrons like covalent bonds.
Unequal electron distribution like an ionic bond.

49. Polar Bonds (Cont’d)
Molecules will be polar if there is any asymmetry in the three dimensional structure.
This especially includes unpaired electrons. Partial charges shown by δ+ and δ-.
Direction of dipole is also shown by arrows pointing toward the δ-.
How polar a molecule is depends on the difference in electronegativity of the individual atoms.
The larger the difference, the greater the polarity of the bond.
The larger the electronegativity difference, the more ionic the bond will be (more polar).
% ionic character is a measure of how ionic the bond in compared to a true ionic bond.

50. Polar Bond Problems Which of the following molecules are polar? NaCl
CO2 O2