2. What is meant by the term “chemical bond”?
Questions to Consider
What is meant by the term “chemical bond”?
Why do atoms bond with each other to form compounds?
How do atoms bond with each other to form compounds?
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3. A Chemical Bond No simple, and yet complete, way to define this.
Forces that hold groups of atoms together and make them function as a unit.
A bond will form if the energy of the aggregate is lower than that of the separated atoms.
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4. The Interaction of Two Hydrogen Atoms
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5. The Interaction of Two Hydrogen Atoms
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6. Key Ideas in Bonding Ionic Bonding – electrons are transferred
Covalent Bonding – electrons are shared equally by nuclei
What about intermediate cases?
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7. Polar Covalent Bond
Unequal sharing of electrons between atoms in a molecule.
Results in a charge separation in the bond (partial positive and partial negative charge).
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8. The Effect of an Electric Field on Hydrogen Fluoride Molecules
indicates a positive or negative fractional charge.
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9. What is meant by the term “chemical bond?”
CONCEPT CHECK!
What is meant by the term “chemical bond?”
Why do atoms bond with each other to form molecules?
How do atoms bond with each other to form molecules?
A chemical bond is difficult to define but in general it can be described as forces that hold groups of atoms together and make them function as a unit.
Atoms will bond with each other to form molecules if the energy of the aggregate is lower than that of the separated atoms.
Atoms bond with each other by either sharing electrons or transferring electrons to form oppositely charged ions which then attract each other.
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10. The ability of an atom in a molecule to attract shared electrons to itself.
For a molecule HX, the relative electronegativities of the H and X atoms are determined by comparing the measured H–X bond energy with the “expected” H–X bond energy.
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11. On the periodic table, electronegativity generally increases across a period and decreases down a group.
The range of electronegativity values is from 4.0 for fluorine (the most electronegative) to 0.7 for cesium (the least electronegative).
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12. The Pauling Electronegativity Values
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14. Arrange the following bonds from most to least polar:
EXERCISE!
Arrange the following bonds from most to least polar: 
a) N–F O–F C–F
b) C–F N–O Si–F
c) Cl–Cl B–Cl S–Cl
a) C–F, N–F, O–F
b) Si–F, C–F, N–O
c) B–Cl, S–Cl, Cl–Cl
The greater the electronegativity difference between the atoms, the more polar the bond.
C-F, N-F, O-F
Si-F, C-F, N-O
B-Cl, S-Cl, Cl-Cl
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15. CONCEPT CHECK!
Which of the following bonds would be the most polar without being considered ionic? Mg–O C–O O–O Si–O N–O
The correct answer is Si-O. To not be considered ionic, generally the bond needs to be between two nonmetals. The most polar bond between the nonmetals occurs with the bond that has the greatest difference in electronegativity.
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16. Dipole Moment
Property of a molecule whose charge distribution can be represented by a center of positive charge and a center of negative charge.
Use an arrow to represent a dipole moment.
Point to the negative charge center with the tail of the arrow indicating the positive center of charge.
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17. Dipole Moment

18. No Net Dipole Moment (Dipoles Cancel)
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19. Stable Compounds
Atoms in stable compounds usually have a noble gas electron configuration.
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20. Electron Configurations in Stable Compounds
When two nonmetals react to form a covalent bond, they share electrons in a way that completes the valence electron configurations of both atoms.
When a nonmetal and a representative-group metal react to form a binary ionic compound, the ions form so that the valence electron configuration of the nonmetal achieves the electron configuration of the next noble gas atom. The valence orbitals of the metal are emptied.
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21. Isoelectronic Series
A series of ions/atoms containing the same number of electrons.
O2-, F-, Ne, Na+, Mg2+, and Al3+
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22. CONCEPT CHECK!
Choose an alkali metal, an alkaline earth metal, a noble gas, and a halogen so that they constitute an isoelectronic series when the metals and halogen are written as their most stable ions.
What is the electron configuration for each species?
Determine the number of electrons for each species.
Determine the number of protons for each species.
One example could be:
K+, Ca2+, Ar, and Cl-
The electron configuration for each species is 1s22s22p63s23p6.
The number of electrons for each species is 18.
K+ has 19 protons, Ca2+ has 20 protons, Ar has 18 protons, and Cl- has 17 protons.
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23. Periodic Table Allows Us to Predict Many Properties
Trends for:
Atomic size, ion radius, ionization energy,
electronegativity
Electron configurations
Formula prediction for ionic compounds
Covalent bond polarity ranking
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24. What are the factors that influence the stability and the structures of solid binary ionic compounds?
How strongly the ions attract each other in the solid state is indicated by the lattice energy.
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25. Lattice Energy
The change in energy that takes place when separated gaseous ions are packed together to form an ionic solid.
k = proportionality constant
Q1 and Q2 = charges on the ions
r = shortest distance between the centers of the cations and anions
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26. Operational Definition of Ionic Compound
Any compound that conducts an electric current when melted will be classified as ionic.
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27. Bond Energies
To break bonds, energy must be added to the system (endothermic, energy term carries a positive sign).
To form bonds, energy is released (exothermic, energy term carries a negative sign).
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28. Bond Energies
ΔH = Σn×D(bonds broken) – Σn×D(bonds formed) D represents the bond energy per mole of bonds (always has a positive sign).
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29. Given the following information: Bond Energy (kJ/mol) C–H 413 C–N 305
CONCEPT CHECK!
Predict ΔH for the following reaction:
Given the following information:
Bond Energy (kJ/mol)
C–H
C–N
C–C
891
ΔH = –42 kJ
[3(413) ] – [3(413) ] = –42 kJ
ΔH = –42 kJ
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30. Localized Electron Model
A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.
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31. Localized Electron Model
Electron pairs are assumed to be localized on a particular atom or in the space between two atoms:
Lone pairs – pairs of electrons localized on an atom
Bonding pairs – pairs of electrons found in the space between the atoms
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32. Localized Electron Model
Description of valence electron arrangement (Lewis structure).
Prediction of geometry (VSEPR model).
Description of atomic orbital types used by atoms to share electrons or hold lone pairs.
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33. 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.
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34. Duet Rule
Hydrogen forms stable molecules where it shares two electrons.

35. Octet Rule
Elements form stable molecules when surrounded by eight electrons.
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36. H–H Single Covalent Bond
A covalent bond in which two atoms share one pair of electrons.
H–H
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37. O=C=O Double Covalent Bond
A covalent bond in which two atoms share two pairs of electrons.
O=C=O
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38. Triple Covalent Bond
A covalent bond in which two atoms share three pairs of electrons.
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39. Steps for Writing Lewis Structures
1. Sum the valence electrons from all the atoms. For a polyatomic anion, add as many electrons as the charge. For a polyatomic cation subtract as many electrons as the charge. 2. Identify the central atom. F, H are never central atoms. The central atom is generally the first atom (A) in formula. 3. Connect all atoms to the central atom using only single bonds. Atoms connected to a central atom are end atoms (X). 4. Use any remaining electrons to complete the octets of all end atoms except H. Do not add more than 1 line or put dots on H in a Lewis structure. 5. If any electrons remain, place them as dots on the central atom even if doing so exceeds octet. A pair of dots on the central atom is denoted E.
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40. Steps in Drawing Lewis Structure Continued
6. If the central atom does not still have an octet, make 1 or more lone pairs on the end atom a bond pair till octet is achieved. H, F, Cl, Br, I will not form double & triple bonds. So, if your end atom is one of these, this rule won’t apply to it.
The octet rule cannot always be obeyed. Exceptions include:
A. The central atom having less than an octet of electrons
B. The central atom having an odd number of electrons
C. The central atom having more than an octet of electrons.
Your Lewis structure is complete when you have used up all the valence electrons & as many as possible atoms have been satisfied

41. In a lewis structure, the central atom is denoted A, the end atoms X and lone pairs on the central atom E
We combine these to obtain a general molecular formula.
For example, if in a Lewis structure there are 3 end atoms attached to the central atom and 2 lone pairs on the central atom, the general molecular formula is AX3E2.
When we follow rule 6, if multiple structures can be drawn, then they are called resonance structures. The actual structure is a hybrid (resonance hybrid) of these resonance structures.
If the central atom is in periods 3, 4, 5 etc (S, P, Cl, Br, I, As, Xe), it can accommodate more than 8 electrons because of the availability of d orbitals. In exceptions to the octet rule where octet on the central atom is exceeded, the central atom is therefore in periods 3 & below.

42. 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–
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43. 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–
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44. 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.
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45. More than one valid Lewis structure can be written for a particular molecule.
NO3– = 24e–
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46. Actual structure is an average of the resonance structures.
Electrons are really delocalized – they can move around the entire molecule.
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47. CONCEPT CHECK!
Draw a Lewis structure for each of the following molecules: CO CO2 CH3OH OCN–
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48. 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.
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49. 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.
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50. 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.
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51. 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
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52. 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.
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53. 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.
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54. VSEPR Model VSEPR: Valence Shell Electron-Pair Repulsion.
The structure around a given atom is determined principally by minimizing electron pair repulsions.
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55. 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.
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56. 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
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57. 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
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58. 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.
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59. Let’s Think About It Draw the Lewis structure for CO2.
Does CO2 contain polar bonds?
Is the molecule polar or nonpolar overall? Why?
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60. 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.
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61. Let’s Think About It Draw the Lewis structure for XeF4.
Does XeF4 contain lone pairs?
Is the molecule polar or nonpolar overall? Why?
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62. Arrangements of Electron Pairs Around an Atom Yielding Minimum Repulsion
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63. Structures of Molecules That Have Four Electron Pairs Around the Central Atom

64. Structures of Molecules with Five Electron Pairs Around the Central Atom
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