1. Chapter 8 Basic Concepts of Chemical Bonding

2. Chemical Bonds Three basic types of bonds Ionic Covalent Metallic
Electrostatic attraction between ions
Covalent
Sharing of electrons
Metallic
Free electron hold metal atoms together
Whenever two atoms or ions are strongly held together, we say there is a chemical bond between them. There are three general types of chemical bonds: ionic, covalent, and metallic.
Table salt is sodium chloride, NaCl, which consists of sodium ions, Na+, and chloride ions, Cl−. The structure is held together by ionic bonds, which are due to the electrostatic attractions between oppositely charged ions. The water consists mainly of H2O. H2O molecules. The hydrogen and oxygen atoms are bonded to one another through covalent bonds, in which molecules are formed by the sharing of electrons between atoms. The spoon consists mainly of iron metal, in which Fe atoms are connected to one another by metallic bonds, which are formed by electrons that are relatively free to move from one atom to another.

3. Lewis Symbols
G. N. Lewis developed a method to denote potential bonding electrons by using one dot for every valence electron around the element symbol.
When forming compounds, atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons (the octet rule).
The electrons involved in chemical bonding are the valence electrons, which, for most atoms, are those in the outermost occupied shell. The American chemist G. N. Lewis suggested a simple way of showing the valence electrons in an atom and tracking them during bond formation, using what are now known as either Lewis electron-dot symbols or simply Lewis symbols.
The Lewis symbol for an element consists of the element’s chemical symbol plus a dot for each valence electron. Sulfur, for example, has the electron configuration [Ne]3s23p4[Ne]3s23p4 and therefore six valence electrons. Its Lewis symbol consists of a symbol for the element with the six valence electrons positioned around it.
The dots are placed on the four sides of the symbol—top, bottom, left, and right—and each side can accommodate up to two electrons. All four sides are equivalent, which means that the choice of sides for placement of two electrons rather than one electron is arbitrary. In general, we spread out the dots as much as possible.
Electrons involved in chemical bonding are valence electrons

4. Ionic Bonding Metals and nonmetals (except group 8A) Electron transfer
Very exothermic
Na(s) + ½ Cl2(g)  NaCl(s) DHf = -411 kJ

5. Ionic Bonding
One element readily gives up an electron (has a LOW ionization energy).
Another element readily gains an electron (has a HIGH electron affinity).
Arrow(s) indicate the transfer of the electron(s).
Each ion has an octet of electrons, the Na+ octet being the 2s22p62s22p6 electrons that lie below the single 3s valence electron of the Na atom. We put a bracket around the chloride ion to emphasize that all eight electrons are located on it.

6. Properties of Ionic Substances
Evidence of well-defined 3-D structures:
Brittle
High melting points
Crystalline
Cleave along smooth lines
Ionic substances possess several characteristic properties. They are usually brittle substances with high melting points. They are usually crystalline. Furthermore, ionic crystals often can be cleaved; that is, they break apart along smooth, flat surfaces. These characteristics result from electrostatic forces that maintain the ions in a rigid, well-defined, three-dimensional arrangement.

7. Energetics of Ionic Bonding— Born–Haber Cycle
Many factors affect the energy of ionic bonding.
Start with the metal and nonmetal elements: Na(s) and Cl2(g).
Make gaseous atoms: Na(g) and Cl(g).
Make ions: Na+(g) and Cl–(g).
Combine the ions: NaCl(s).
What factors make the formation of ionic compounds so exothermic?
We represent the formation of NaCl as the transfer of an electron from Na to Cl. Recall that the loss of electrons from an atom is always an endothermic process. Removing an electron from Na(g) to form Na+(g) for instance, requires 496kJ/mol. Recall that when a nonmetal gains an electron, the process is generally exothermic, as seen from the negative electron affinities of the elements. Adding an electron to Cl(g), for example, releases 349kJ/mol. From the magnitudes of these energies, we can see that the transfer of an electron from an Na atom to a Cl atom would not be exothermic—the overall process would be an endothermic process that requires 496−349=147kJ/mol. This endothermic process corresponds to the formation of sodium and chloride ions that are infinitely far apart—in other words, the positive energy change assumes that the ions do not interact with each other, which is quite different from the situation in ionic solids.
The principal reason ionic compounds are stable is the attraction between ions of opposite charge. This attraction draws the ions together, releasing energy and causing many ions to form a solid array, or lattice. A measure of how much stabilization results from arranging oppositely charged ions in an ionic solid is given by the lattice energy, which is the energy required to completely separate one mole of a solid ionic compound into its gaseous ions.
The coming together of Na+(g) and Cl−(g) to form NaCl(s) is therefore highly exothermic (ΔH=−788kJ/mol).

8. Energetics of Ionic Bonding
We already discussed making ions (ionization energy and electron affinity).
It takes energy to convert the elements to atoms (endothermic).
It takes energy to create a cation (endothermic).
Energy is released by making the anion (exothermic).
The formation of the solid releases a huge amount of energy (exothermic).
This makes the formation of salts from the elements exothermic.

9. Lattice Energy
Energy required to completely separate one mole of a solid ionic compound into its gaseous ions.
That amount of energy is RELEASED to MAKE the ionic compound (in the Born–Haber cycle).

10. Lattice Energy
What are trends here related to charge and size of ions?

11. Trends in Lattice Energy
Lattice energy increases with:
Increasing charge on the ions
Decreasing size of ions
Coulomb’s Law
Eel - Lattice energy
Q - Charge on particle
d – distance between nuclei
k - constant

12. Electron Configuration of Ions
Main group metals lose electrons, resulting in the electron configuration of the previous noble gas.
Nonmetals gain electrons, resulting in the electron configuration of the nearest noble gas.
Transition metals do NOT follow the Octet rule.
Transition metals lose the VALENCE electrons FIRST, THEN lose the d- electrons necessary for the given ion charge.
The octet rule, although useful, is clearly limited in scope.

13. Practice Exercise
Which of the following orderings of lattice energy is correct for these ionic compounds?
a) NaCl > MgO > CsI > ScN
b) ScN > MgO > NaCl > CsI
c) NaCl > CsI > ScN > MgO
d) MgO > NaCl > ScN > CsI
e) ScN > CsI > NaCl > MgO

14. Practice Exercise
Which substance do you expect to have the greatest lattice energy?
a) MgF2
b) CaF2
c) ZrO2

15. Covalent Bonding In covalent bonds, atoms share electrons.
There are several electrostatic interactions in these bonds:
Attractions between electrons and nuclei
Repulsions between electrons
Repulsions between nuclei
For a bond to form, the attractions must be greater than the repulsions.
The vast majority of chemical substances do not have the characteristics of ionic materials.

16. Covalent Bonding
In a covalent “single” bond, 2 electrons are “shared” between 2 atoms
Covalently bound species are different than ionic
exist as individual, discrete species (vs. 3-D crystal lattice structure for ionic)
tend to exhibit much lower melting and boiling points (vs. ionic)

17. Covalent Bonding
For the very large class of substances that do not behave like ionic substances, we need a different model to describe the bonding between atoms. G. N. Lewis reasoned that atoms might acquire a noble-gas electron configuration by sharing electrons with other atoms. A chemical bond formed by sharing a pair of electrons is a covalent bond.

18. Lewis Structures
Sharing electrons to make covalent bonds can be demonstrated using Lewis structures.
We start by trying to give each atom the same number of electrons as the nearest noble gas by sharing electrons.
The simplest examples are for hydrogen, H2, and chlorine, Cl2, shown below.

19. Number of Bonds for Nonmetals
The group number is the number of valence electrons.
To get an octet, like the nearest noble gas, in the simplest covalent molecules for nonmetals, the number of bonds needed will be the same as the electrons needed to complete the octet.

20. Electrons on Lewis Structures
Lone pairs: electrons located on only one atom in a Lewis structure
Bonding pairs: shared electrons in a Lewis structure; they can be represented by two dots or one line, NOT both!

21. Multiple Bonds
Some atoms share only one pair of electrons. These bonds are called single bonds.
Sometimes, two pairs need to be shared. These are called double bonds.
There are even cases where three bonds are shared between two atoms. These are called triple bonds.

22. a double bond is shorter and stronger than a single bond
As a general rule, the length of the bond between two atoms decreases as the number of shared electron pairs increases.
a double bond is shorter and stronger than a single bond
a triple bond is shorter and stronger than a double bond

23. Polarity of Bonds
The electrons in a covalent bond are not always shared equally.
Bond polarity is a measure of how equally or unequally the electrons in a covalent bond are shared.
In a nonpolar covalent bond, the electrons are shared equally.
In a polar covalent bond, one of the atoms attracts electrons to itself with a greater force.

24. Polar or Nonpolar Covalent Bonds
In elemental fluorine, the atoms pull electrons equally. The bond is a nonpolar covalent bond.
Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does. Therefore, the fluorine end of the molecule has more electron density than the hydrogen end, making it a polar covalent bond.

25. Electronegativity
Electronegativity is the ability of an atom in a molecule to attract electrons to itself.
On the periodic table, electronegativity generally increases as you go:
from left to right across a period.
from the bottom to the top of a group.

26. Arrange the following in order of increasing electronegativity: Na, F, O, K, Al, Si, Mg

27. Electronegativity and Polar Covalent Bonds
When two atoms share electrons unequally, a polar covalent bond results.
Electrons tend to spend more time around the more electronegative atom. The result is a partial negative charge (not a complete transfer of charge). It is represented by δ–.
The other atom is “more positive,” or δ+.

28. Polar Covalent Bonds
The greater the difference in electronegativity, the more polar is the bond.

29. Dipoles
When two equal, but opposite, charges are separated by a distance, a dipole forms.
A dipole moment, , produced by two equal but opposite charges separated by a distance, r, can be calculated:
 = Qr
It is measured in debyes (D).

30. Is a Compound Ionic or Covalent?
Simplest approach: Metal + nonmetal is ionic; nonmetal + nonmetal is covalent.
There are many exceptions: It doesn’t take into account oxidation number of a metal (higher oxidation numbers can give covalent bonding).
Electronegativity difference can be used; the table still doesn’t take into account oxidation number.
Properties of compounds are often best: Lower melting points mean covalent bonding, for example.

31. Rank the following in order of increasing bond polarity:. H─F. H─Br
Rank the following in order of increasing bond polarity: H─F H─Br F─F Na─Cl
H─F < H─Br < F─F < Na─Cl
F─F < H─F < H─Br < Na─Cl
H─Br < H─F < F─F < Na─Cl
F─F < H─Br < H─F < Na─Cl
Na─Cl < H─F < H─Br < F─F

32. Metallic Bonds
The relatively low ionization energy of metals allows them to lose electrons easily.
The simplest theory of metallic bonding involves the metal atoms releasing their valence electrons to be shared as a pool by all the atoms/ions in the metal.
An organization of metal cation islands in a sea of electrons
Electrons delocalized throughout the metal structure
Bonding results from attraction of cation for the delocalized electrons.

33. Metallic Bonding

34. Which of the following compounds has ionic bonding?
NO2 Al
CaCO3

35. Writing Lewis Structures (Covalent Molecules)
Sum the valence electrons from all atoms, taking into account overall charge.
If it is an anion, add one electron for each negative charge.
If it is a cation, subtract one electron for each positive charge.
PCl3
Keep track of the electrons:
(7) = 26

36. Writing Lewis Structures
Write the symbols for the atoms, show which atoms are attached to which, and connect them with a single bond (a line representing two electrons).
Chemical formulas are often written in the order in which the atoms are connected in the molecule or ion. The formula HCN, for example, tells you that the carbon atom is bonded to the H and to the N. In many polyatomic molecules and ions, the central atom is usually written first, as in CO32− and SF4. Remember that the central atom is generally less electronegative than the atoms surrounding it. In other cases, you may need more information before you can draw the Lewis structure.
Keep track of the electrons:
26 − 6 = 20

37. Writing Lewis Structures
Complete the octets around all atoms bonded to the central atom.
Keep track of the electrons:
26 − 6 = 20; 20 − 18 = 2

38. Writing Lewis Structures
Place any remaining electrons on the central atom.
If there are not enough electrons to give the central atom an octet, try multiple bonds.
Keep track of the electrons:
26 − 6 = 20; 20 − 18 = 2; 2 − 2 = 0
(multiple bonds unnecessary here)

39. Writing Lewis Structures
6. If there are not enough electrons to give the
central atom an octet, try multiple bonds.
Do NOT (under any circumstance…..ever) form a multiple bond to a halogen or hydrogen

40. Writing Lewis Structures
Then assign formal charges.
Formal charge is the charge an atom would have if all of the electrons in a covalent bond were shared equally.
Formal charge = (valence electrons) – ½ (bonding electrons) – (all nonbonding electrons).
This can be a method to determine structure.
In some instances, we can draw two or more valid Lewis structures for a molecule that all obey the octet rule. All of these structures can be thought of as contributing to the actual arrangement of the electrons in the molecule, but not all of them will contribute to the same extent. How do we decide which one of several Lewis structures is the most important? One approach is to do some “bookkeeping” of the valence electrons to determine the formal charge of each atom in each Lewis structure. The formal charge of any atom in a molecule is the charge the atom would have if each bonding electron pair in the molecule were shared equally between its two atoms.

41. Writing Lewis Structures
The dominant Lewis structure:
is the one in which atoms have formal charges closest to zero.
puts a negative formal charge on the most electronegative atom.
As such, it can be used to decide which structure is best.

42. Draw the Lewis structure for F2

43. Draw the Lewis structure for H2O

44. Draw the Lewis structure for ethene, C2H4
In drawing Lewis structures, always put halogens and hydrogens on the outside. (They only form one bond)

45. Draw the Lewis structure for NO+
Place electrons on the electronegative atom first.
Triple bond between N and O.

46. The Best Lewis Structure?
Following our rules, this is the Lewis structure we would draw for ozone, O3.
However, it doesn’t agree with what is observed in nature: Both O-to-O connections are the same.

47. Resonance
One Lewis structure cannot accurately depict a molecule like ozone.
We use multiple structures, resonance structures, to describe the molecule.
To describe the structure of ozone properly, we write both resonance structures and use a double-headed arrow to indicate that the real molecule is described by an average of the two.
To understand why certain molecules require more than one resonance structure, we can draw an analogy to mixing paint. Blue and yellow are both primary colors of paint pigment. An equal blend of blue and yellow pigments produces green pigment. We cannot describe green paint in terms of a single primary color, yet it still has its own identity. Green paint does not oscillate between its two primary colors: It is not blue part of the time and yellow the rest of the time. Similarly, molecules such as ozone cannot be described as oscillating between the two individual Lewis structures shown previously—there are two equivalent dominant Lewis structures that contribute equally to the actual structure of the molecule.
The actual arrangement of the electrons in molecules such as O3 must be considered as a blend of two (or more) Lewis structures. By analogy to the green paint, the molecule has its own identity separate from the individual resonance structures.

48. Resonance
In truth, the electrons that form the second C—O bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon.
They are not localized; they are delocalized.

49. Resonance
The organic compound benzene, C6H6, has two resonance structures.
It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.
Localized electrons are specifically on one atom or shared between two atoms; delocalized electrons are shared by multiple atoms.

50. Determine the Lewis structure for NO2-
What are your bond expectations for nitrite ?

51. Go to lab and measure the actual bond lengths in a real nitrite anion.
The N-O bonds in nitrite are identical (in every sense; same length; same strength)
Figure: 08-10
A single Lewis structure can NOT be drawn to describe the “real” nitrite species

52. The Real molecule is somewhere in between these two extremes
Figure: 08-10

53. Draw the Lewis structure for HNO3
Calculation of formal charges is important here

54. Exceptions to the Octet Rule
There are three types of ions or molecules that do not follow the octet rule:
Ions or molecules with an odd number of electrons
Ions or molecules with less than an octet
Ions or molecules with more than eight valence electrons (an expanded octet)

55. Odd Number of Electrons
Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons.

56. Fewer Than Eight Electrons
Consider BF3:
Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine.
This would not be an accurate picture of the distribution of electrons in BF3.

57. Fewer Than Eight Electrons
If filling the octet of the central atom results in a negative formal charge on the central atom and a positive formal charge on the more electronegative outer atom, don’t fill the octet of the central atom.

58. Draw the Lewis structure for PF5
octet expansion – some atoms can exceed 8 valence electrons (usually P & S)

59. More Than Eight Electrons
The only way PCl5 can exist is if phosphorus has 10 electrons around it.
It is allowed to expand the octet of atoms on the third row or below.
Presumably d orbitals in these atoms participate in bonding.

60. MoreThan Eight Electrons
The only way PF5 can exist is if phosphorus has 10 electrons around it.
It is allowed to expand the octet of atoms on the third row or below.
Presumably d orbitals in these atoms participate in bonding.
(Note: Phosphate will actually have four resonance structures with five bonds on the P atom!)