Chapter 7 Chemical Bonding and Molecular Structure

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1 Chapter 7 Chemical Bonding and Molecular Structure Three Types of Chemical Bonding (1) Ionic: formed by electron transfer (2) Covalent: formed by electron sharing (3) Metallic: attraction between metal ions and their delocalized electrons 1

2 Ionic Bonding Metal to Nonmetal; Metal to Polyatomic Ions Examples: NaCl, MgCl 2, NaNO 3, NH 4 NO 3 Binary Compounds: Metals lose electrons and nonmetals gain electrons Metal ions are called cations (positively charged) Nonmetal ions are called anions (negatively charged) Cations are attracted to anions and form crystal lattices called ionic compounds Example: An atom of Na loses one electron to form Na + An atom of Cl gains one electron to form Cl - Na + is attracted to Cl - to form neutral NaCl Remember all compounds are neutral Ionic Bonding Question Describe how calcium fluoride would form from atoms of calcium and fluorine. 2

3 Lewis Electron-dot Symbols The Lewis dot symbol of an atom depicts the number of s and p electrons in the outer energy level. The number of dots around the symbol of an element ranges from 1-8 and reflects the valence electrons for that particle. Generally dots are placed around the element s symbol one at a time starting at the 9 o clock position and continuing clockwise with no more than two dots shown at o clock. Practicing Electron-Dot Symbols Write the Lewis-dot symbols for: Li Be B C N O F Ne What is the Lewis-dot symbol for: K S Al The Octet Rule When atoms form bonds they lose, gain, or share electrons to attain a filled outer shell (either two or eight electrons). 3

4 Use of Orbital Diagrams and Electron Configurations in Bonding Example: Lithium Fluoride Another Example of an Ionic Compound Use electron-dot symbols to show how aluminum oxide forms. 4

5 The Ionic Bonding Model Energy considerations in ionic bonding Overall, energy is released when ions come together in the formation of a compound However, the process of the formation of an ionic compound involves a number of steps Consider the simplified formation of the ionic compound NaCl. Let s start with atoms of sodium and chlorine. A sodium atom has to lose one electron and a chlorine atom has to gain one electron. To remove an electron from a sodium atom it requires energy (ionization energy) Na(atom) + energy Na + (ion) + e - When an atom of chlorine gains an electron it releases energy (electron affinity) Cl(atom) + e - Cl - (ion) + energy 5

6 When a Na + and a Cl - ion come together into a solid crystal a great amount of energy is released (lattice energy). Na + + Cl - NaCl + energy Lattice Energy Lattice Energy is the energy absorbed that occurs when an ionic solid is separated into isolated ions in the gas phase. (Delta H is positive) Also, the energy released when gaseous ions come together to form the crystal. (Delta H is negative) Lattice energies for alkali metal-halogen compounds 6

7 Summary Note that in comparing compounds of like charges the smaller the ion, the greater the lattice energy. BUT Lattice Energy is proportional to: charge M + X charge NM - distance between nuclei Which has the larger lattice energy? KF or LiF KF or CaF 2 AlCl 3 or AlBr 3 or Al 2 S 3 Note: Magnitude of charge dominates over size. Properties of Ionic Compounds Rigid, fixed positions of ions in solid state Hard, brittle Conduct electricity when melted or dissolved in water High melting (mp) and boiling points (bp). Note: The larger the lattice energy, the higher the mp and bp 7

8 Covalent Bonding Nonmetal to Nonmetal Attraction: Shared electrons Example H.. H H-H A single covalent bond is formed by 2 atoms sharing 2 electrons. Each atom has ownership of the 2 electrons. Covalent Bond Formation 8

9 Problem Solving Using Lewis Dot Structures, explain how F 2 forms. F F F-F How many bonding pairs are present? How many lone pairs or nonbonding pairs are present? Sharing of Bonding Pairs Each atom in a covalent bond counts the shared electrons as though they belong entirely to that atom. Let s return to H 2 and F 2 and discuss the meaning of this. Another example: hydrogen fluoride Bond Energy and Bond Length Bond Energy: The amount of energy required to break a bond (endothermic). It is also the amount of energy released in bond formation (exothermic) Bond Length: The distance between the nuclei of the two bonded atoms. 9

10 Examples of Bond Strength/Bond Energy C O C O C O 1070 kj/mol 745 kj/mol 358 kj/mol Note: Larger atoms result in longer, thus weaker bonds. Bond Energy and Chemical Change 10

11 Electronegativity (EN) Electronegativity is the ability of a bonded atom to attract the shared electrons. Change in Electronegativity ( EN) is the difference in electronegativity of the two bonded atoms. Ionic, Polar Covalent, and Nonpolar Covalent Bonding 11

12 Linus Pauling s Work Problem Solving Consider the C-O bond, what is the EN for this bond? Consider the Br-Br bond, what is the EN for this bond? Which of the bonds above is/are polar covalent? nonpolar covalent? 12

13 Illustrating Bond Polarity Conventional Methods for Illustrating the Polarity of a bond. C-O N-H C-H S-O H-F Linus Pauling s Work Lewis Dot Structures Review Draw Lewis Dot Structures of: SiO 2 SO 3 2- HNO 2 Place your answer(s) on the board. 13

14 Exceptions to the Octet Rule Examples: Deficient Octet BF 3 Expanded Octet H 2 SO 4 Radicals NO Resonance Structures Consider the following molecules: CO 3 2- C 6 H 6 CH 3 COO - Note: Atoms do not change position! Only pi and lone pair electrons. Formal Charge Determination of formal charge for each atom within a molecule: V.E. (NB.E. + ½ B.E.) = formal charge What are the formal charges of each atom in HCN? What are the formal charges of each atom in NO 3 -? What are the formal charges of each atom in SO 4 2-? 14

15 Selecting the Preferred Lewis Dot Structure Use the concept of formal charge to select the preferred structure of the CON - ion. Valence Bond Theory Central Themes: A covalent bond forms when orbitals of two atoms overlap and the overlap region is occupied by two electrons. The greater the overlap the stronger the bond. The stronger the bond the more stable the bond. Orbitals must become oriented so as to obtain the greatest overlap possible. Types of orbital overlap: sigma (end-to-end) pi (side-by-side) Note: sigma bond side-by-side overlap is not permitted. 15

16 Let s Consider CH 4 How can carbon form four bonds? What is the shape of methane? What are the bond angles? Hybrid Orbital Theory Theory depicts a mix of the atomic orbitals (hybridize) to form the necessary number of orbitals needed for bonding. The number of hybrid orbitals formed equals the number of atomic orbitals mixed. The type of hybrid orbitals obtained varies with the type of atomic orbitals mixed. 16

17 Types of Hybrid Orbitals sp sp 2 sp 3 sp 3 d sp 3 d 2 sp Hybrids Consider BeH 2 Central Be needs two hybrid orbitals to accommodate the two bonded H atoms. To aid the visualization of sp hybrid formation use an orbital diagram, shapes of hybrids, and bond angles. Bond angles for sp hybrids = 180 o 17

18 Note: The bonding between the sp hybrid orbital of beryllium and the s orbital of the hydrogen atom is considered a sigma bond overlap. sp 2 Hybrids Consider BF 3 Central boron atom needs three hybrid orbitals to accommodate three bonded fluorine atoms. Visual of hybrid orbital formation theory. Bond angles for sp 2 hybrids = 120 o 18

19 Note: The bond involving the sp 2 hybridized orbital of boron is an end-to-end overlap (sigma bond) with the p orbital of the fluorine atom. sp 3 Hybrids Back to CH 4 The central carbon atom needs four hybrid orbitals. Bond angles between sp 3 hybrids = o 19

20 Other atoms using sp 3 Hybrids If the central atom has only two or three (rather than four) bonded atoms, hybrid orbitals may also contain the lone pairs of electrons. Examples: NH 3, H 2 O sp 3 d Hybrids Consider PCl 5 The central phosphorus atom needs five hybrid orbitals to accommodate the five bonded chlorine atoms. Note that since there is only one s orbital and only three p orbitals available per energy level, a d orbital must be used in the hybrid. 20

21 Bond angles between sp 3 d hybrids = 90 o ( axial) and 120 o (equatorial) sp 3 d 2 Hybrids Consider SF 6 The six bonded atoms require six hybrids. Bond angles between sp 3 d 2 hybrids = 90 o 21

22 Summary of Hybrid Theory Hybrid Quiz Predict the type of hybrid orbital you would expect in the central atom of the following molecules: SiH 4 BH 3 AsF 5 AlCl 3 SF 4 Multiple Bonds Double Bond: A=B One bond is a sigma bond The other is a pi bond Sigma bonds are formed by hybrid orbitals overlapping. Pi bonds are formed by unhybridized p orbitals overlapping. 22

23 Double Bond Triple Bond: A B One bond is a sigma bond formed from overlapping hybrid orbitals. Two bonds are pi bonds formed from overlapping p orbitals. Triple Bond 23

24 Orbital Overlap and Molecular Rotation Sigma bonds allow free rotation of bonded atoms. Pi bond overlap restricts rotation of bonded atoms. Double bonds lead to cis- and transisomers. Example: C 2 H 2 Cl 2 Quiz Describe the type of hybrid orbitals used by each carbon and oxygen atom in the following molecule: O H C C O H H H Valence Shell Electron Pair Repulsion Theory (VSEPR) Each group of valence electrons around a central atom is located as far away from the other atoms as possible. Strength of electron pair repulsions: L.P.-L.P. > L.P.-B.P. > B.P.-B.P. A group of electrons is defined as any number of electrons that occupies a space around an atom and may consist of a single, double, triple bond or lone pair of electrons. The 3-dimensional arrangement of these groups determines the molecular arrangement (shape). 24

25 Molecular Shape The arrangement of the atoms around a central atom determines the shape of the molecule or portion of the molecule. Possible Shapes of Molecules Linear Bent Trigonal Planar Trigonal Pyramidal Square Planar Square Pyramidal Octahedral T-Shaped Seesaw Trigonal Bipyramidal Different Shapes of Molecules (or portions of molecules) Two electron groups around the central atom: Linear shape Bond angle = 180 o 25

26 Three electron groups Trigonal planar arrangement Bond angles = 120 o 26

27 Four electron groups Tetrahedral arrangement Bond angles = o Five electron groups Bipyramidal arrangement Consists of three equatorial groups And two axial groups 27

28 Six electron groups -Octahedral arrangement -Bond angles = 90 o 28

29 Polar vs. Nonpolar Molecules Generally speaking, a molecule will be nonpolar if (1) All of the bonded atoms (or groups of atoms) to the central atom are the same and equidistant from each other. i.e. BH 3 vs. BF 3 (2) There are no lone pairs of electrons on the central atom(s). (3) It is a hydrocarbon. 29

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