1. BONDING
Let’s get together…
Barbara A. Gage PGCC CHM 1010

2. Why Do Atoms Bond?
Chemical bonds form because they lower the potential energy between the charged particles that compose atoms
A chemical bond forms when the potential energy of the bonded atoms is less than the potential energy of the separate atoms
Barbara A. Gage PGCC CHM 1010

3. Types of Bonds
We can classify bonds based on the kinds of atoms that are bonded together
Types of Atoms
Type of Bond
Bond Characteristic
metals to
nonmetals
Ionic
electrons
transferred
nonmetals to
Covalent
shared
metals
Metallic
pooled
Barbara A. Gage PGCC CHM 1010

4. Types of Bonding
Barbara A. Gage PGCC CHM 1010

5. Ionic Bonds When a metal atom loses electrons it becomes a cation
metals have low ionization energy, making it relatively easy to remove electrons from them
When a nonmetal atom gains electrons it becomes an anion
nonmetals have high electron affinities, making it advantageous to add electrons to these atoms
The oppositely charged ions are then attracted to each other, resulting in an ionic bond
Barbara A. Gage PGCC CHM 1010

6. Covalent Bonds
Nonmetal atoms have relatively high ionization energies, so it is difficult to remove electrons from them
When nonmetals bond together, it is better in terms of potential energy for the atoms to share valence electrons
potential energy lowest when the electrons are between the nuclei
Shared electrons hold the atoms together by attracting nuclei of both atoms
Barbara A. Gage PGCC CHM 1010

7. an organization of metal cation islands in a sea of electrons
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
Barbara A. Gage PGCC CHM 1010

8. Lewis Electron-Dot Symbols
For main group elements -
The A group number gives the number of valence electrons.
Place one dot per valence electron on each of the four sides of the element symbol.
Pair the dots (electrons) until all of the valence electrons are used.
Example:
Nitrogen, N, is in Group 5A and therefore has 5 valence electrons. N : . . N : : N . : N .
Barbara A. Gage PGCC CHM 1010

9. Determining the Number of Valence Electrons in an Atom
The column number on the Periodic Table will tell you how many valence electrons a main group atom has
Transition Elements all have two valence electrons. Why?
Barbara A. Gage PGCC CHM 1010

10. Lewis Structures of Ions
Cations have Lewis symbols without valence electrons
lost in the cation formation
Anions have Lewis symbols with eight valence electrons
electrons gained in the formation of the anion
Barbara A. Gage PGCC CHM 1010

11. Ionic Bonding & the Crystal Lattice
Ionically bonded substances form a structure in which every cation is surrounded by anions, and vice versa
This structure is called a crystal lattice
The crystal lattice is held together by the electrostatic attraction of the cations for all the surrounding anions
The crystal lattice maximizes the attractions between cations and anions, leading to the most stable arrangement
Barbara A. Gage PGCC CHM 1010

12. Crystal Lattice Electrostatic attraction is nondirectional!!
no direct anion–cation pair
Therefore, there is no ionic molecule
the chemical formula is an empirical formula, simply giving the ratio of ions based on charge balance
One unit of the
empirical formula is
called a formula unit
Barbara A. Gage PGCC CHM 1010

13. Ionic compound dissolved in water
Electrical conductance and ion mobility.
Solid ionic compound
Molten ionic compound
Ionic compound dissolved in water
Barbara A. Gage PGCC CHM 1010

14. Lewis Theory of Covalent Bonding
Lewis theory implies that another way atoms can achieve an octet of valence electrons is to share their valence electrons with other atoms
The shared electrons would then count toward each atom’s octet
The sharing of valence electrons is called covalent bonding
Barbara A. Gage PGCC CHM 1010

15. Covalent Bonding
Atoms with incomplete octets can share rather than transfer electrons.
Each pair of shared electrons = 1 bond
Shared electrons move around the nuclei of both atoms in the bond so both atoms have possession of the shared electrons.
Barbara A. Gage PGCC CHM 1010

16. Lewis Dot Structures for Covalent Compounds
Sum the valence electrons of all atoms.
Determine the central atom.
Position the central atom and place the additional atoms equally around it.
Place the required number of electrons around the outside atoms first and then around the central atom to be each one meets the octet rule (or the number needed if it is an exception).
Barbara A. Gage PGCC CHM 1010

17. Lewis Dot Structures for Covalent Compounds
CCl4 Total electrons = 1(4) + 4(7) = 32 Cl Cl C Cl
Barbara A. Gage PGCC CHM 1010

18. Lewis Dot Structures for Covalent Compounds
SO3 Total electrons = 1(6) + 3(6) = 24 O O S O O O O S O O S O
Barbara A. Gage PGCC CHM 1010

19. Covalent Bonding: Bonding and Lone Pair Electrons
Electrons that are shared by atoms are called bonding pairs
Electrons that are not shared by atoms but belong to a particular atom are called lone pairs
aka nonbonding pairs . O S
Lone pairs
Bonding pairs
Barbara A. Gage PGCC CHM 1010

20. Single Covalent Bonds •• •• • • •• •• O H F O H F
When two atoms share one pair of electrons it is called a single covalent bond
2 electrons
One atom may use more than one single bond to fulfill its octet
to different atoms
H only duet H • O •• F •• • H O •• F •• F
Barbara A. Gage PGCC CHM 1010

21. Double Covalent Bond •• • •• • ••
When two atoms share two pairs of electrons the result is called a double covalent bond
four electrons O •• • O •• • O ••
Barbara A. Gage PGCC CHM 1010

22. Octet Rule Exceptions
Some elements are stable with fewer or more than 8 e-.
H 2e- Be 4e- B 6e-
P, Cl, Br (and more) 10e-
S, Se, Xe (and more) 12e-
Barbara A. Gage PGCC CHM 1010

23. Resonance: Delocalized Electron-Pair Bonding
O3 can be drawn in 2 ways -
Neither structure is actually correct but can be drawn to represent a structure which is a hybrid of the two - a resonance structure.
Resonance structures have the same relative atom placement but a difference in the locations of bonding and nonbonding electron pairs.
is used to indicate that resonance occurs.
Barbara A. Gage PGCC CHM 1010

24. Writing Resonance Structures
PROBLEM:
Write resonance structures for the nitrate ion, NO3-.
SOLUTION:
Nitrate has 1(5) + 3(6) + 1 = 24 valence e-
N does not have an octet; a pair of e- will move in to form a double bond.
Barbara A. Gage PGCC CHM 1010

25. Bond Lengths
The distance between the nuclei of bonded atoms is called the bond length
Because the actual bond length depends on the other atoms around the bond we often use the average bond length
averaged for similar bonds from many compounds
Barbara A. Gage PGCC CHM 1010

26. Bond Energies
Chemical reactions involve breaking bonds in reactant molecules and making new bonds to create the products
The DH°reaction can be estimated by comparing the cost of breaking old bonds to the income from making new bonds
The amount of energy it takes to break one mole of a bond in a compound is called the bond energy
in the gas state
homolytically – each atom gets ½ bonding electrons
Barbara A. Gage PGCC CHM 1010

27. Barbara A. Gage PGCC CHM 1010

28. What is the relationship between bond order and bond length for bonds
Silberberg, Principles of Chemistry
What is the relationship between bond order and bond length for bonds
between the same two elements?
What is the relationship between bond length and bond energy for bonds
Barbara A. Gage PGCC CHM 1010

29. Comparing Bond Length and Bond Strength
PROBLEM:
Using the periodic table, rank the bonds in each set in order of decreasing bond length and bond strength:
(a) S - F, S - Br, S - Cl
(b) C = O, C - O, C O
(a) The bond order is one for all and sulfur is bonded to halogens; bond length should increase and bond strength should decrease with increasing atomic radius. (b) The same two atoms are bonded but the bond order changes; bond length decreases as bond order increases while bond strength increases as bond order increases.
SOLUTION:
(a) Atomic size increases going down a group.
(b) Using bond orders we get
Bond length: S - Br > S - Cl > S - F
Bond length: C - O > C = O > C O
Bond strength: S - F > S - Cl > S - Br
Bond strength: C O > C = O > C - O
Barbara A. Gage PGCC CHM 1010

30. Break 1 mol C─H +414 kJ Make 1 mol Cl─Cl +243 kJ 1 mol C─Cl −339 kJ
1 mol H─Cl −431 kJ
Barbara A. Gage PGCC CHM 1010

31. Electron Distribution in a Covalent Bond
Are electrons shared equally in a covalent bond?
If not, why not?
Distance of electrons from nucleus and number of protons in the nucleus
Electronegativity – attraction of one atom in a bond for the electrons in that bond
Barbara A. Gage PGCC CHM 1010

32. Polar Covalent Bonding
Covalent bonding between unlike atoms results in unequal sharing of the electrons
one atom pulls the electrons in the bond closer to its side
one end of the bond has larger electron density than the other
The result is a polar covalent bond
bond polarity
the end with the larger electron density gets a partial negative charge
the end that is electron deficient gets a partial positive charge
Barbara A. Gage PGCC CHM 1010

33. The Pauling electronegativity (EN) scale.
Barbara A. Gage PGCC CHM 1010

34. Polarity
When atoms in a bond have different electronegativities, the electron sharing is unequal.
As the ΔEN increases, the electron distribution becomes more uneven and the molecule becomes polar.
Barbara A. Gage PGCC CHM 1010

35. Polarity HCl ENH = 2.1 ENCl = 3.0 ΔEN = 0.9
The end with the higher EN will be slightly negative and the other will be slightly positive
δ+H – Clδ- H – Cl
Barbara A. Gage PGCC CHM 1010

36. Electronegativity Difference and Bond Type
If difference in electronegativity between bonded atoms is 0, the bond is pure covalent
equal sharing
If difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent
If difference in electronegativity between bonded atoms is 0.5 to 1.9, the bond is polar covalent
If difference in electronegativity between bonded atoms is larger than or equal to 2.0, the bond is ionic
0.4
2.0
4.0 4%
51%
Percent Ionic Character
Electronegativity Difference
“100%”
Barbara A. Gage PGCC CHM 1010

37. Bond Dipole Moments Dipole moment, m, is a measure of bond polarity
a dipole is a material with a + and − end
it is directly proportional to the size of the partial charges and directly proportional to the distance between them
m = (q)(r)
not Coulomb’s Law
measured in Debyes, D
Generally, the more electrons two atoms share and the larger the atoms are, the larger the dipole moment
Barbara A. Gage PGCC CHM 1010

38. Determining Bond Polarity from EN Values
PROBLEM:
(a) Use a polar arrow to indicate the polarity of each bond: N-H, F-N, I-Cl.
(b) Rank the following bonds in order of increasing polarity: H-N, H-O, H-C.
(a) Find EN values; the arrow should point toward the negative end.
(b) Polarity increases across a period.
SOLUTION:
(a) The EN of N = 3.0, H = 2.1; F = 4.0; I = 2.5, Cl = 3.0
N - H
F - N
I - Cl
(b) The order of increasing EN is C < N < O; all have an EN larger than that of H.
H-C < H-N < H-O
Barbara A. Gage PGCC CHM 1010

39. Percent Ionic Character
The percent ionic character is the percentage of a bond’s measured dipole moment compared to what it would be if the electrons were completely transferred
The percent ionic character indicates the degree to which the electron is transferred
Barbara A. Gage PGCC CHM 1010