Chemical Bonding Chemical Bonding I: The Covalent Bond I. Types of bonds a) Ionic b) Covalent II. Lewis Dot Structures a) ctet Rule b) Multiple Bonds c) Resonance d) Polyatomic Ions e) ormal Charge on Atoms Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chemical Bonding III. Exceptions to ctet Rule a) Incomplete ctet b) Expanded ctet c) dd # of Electrons IV. Electronegativity V. Polar vs. NonPolar Covalent Bonds VI. Use of Lewis Dot Structures a) Bond Strength/Length b) Estimate ΔH Bonding Generalities Unlike Charges Attract Electrons will Be in Pairs nly Valence Electrons are Involved 1
Two Main Types of Bonds Lewis Dot Symbols 1. Ionic Bonds- Electrostatic Attraction Between a + and Ion That Holds Together Ionic Compounds 2. Covalent Bond Sharing of an Valence Electron Pair Between two Nonmetal Atoms to orm a Molecule The Ionic Bond Li + Li + - ig. 9.6 Li Li + + e - e - + - Li + + - Li + - 9.2 2
ΔH latt Periodic Trends in Lattice Energy ig. 9.7 Energy of electrostatic force is proportional to: Cation charge X anion charge Cation radius + anion radius Proportional to Lattice energy, ΔH latt A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Why should two atoms share electrons? + 7e - 7e - 8e - 8e - Lewis structure of 2 single covalent bond lone pairs lone pairs lone pairs lone pairs single covalent bond 9.4 3
ig. 9.12 ctet Rule When Nonmetal Atoms Share Valence Electrons to orm a Covalent Bond- It Will Be to Have the Same Number of Valence Electrons as the Closest Noble Gas»H 2 electrons» Everything Else 8 Electrons Comparison of Ionic and Covalent Compounds Writing Lewis Structures for Molecular Compounds 1. Draw skeletal structure of compound showing what atoms are bonded to each other. 2. Count total number of valence e -. Add 1 for each negative charge. Subtract 1 for each positive charge. 3. Connect surrounding atoms to central atom with single bonds. Add remaining electrons (2 at a time) such that surrounding atoms follow octet rule (H -2 electrons). Stop adding electrons once the number exceeds value calculated in step 2. 9.4 9.6 4
Writing Lewis Structures (cont) 4. Make sure that every atom satisfies octet rule and the total number of valence electrons in Lewis Structure is correct. 5. If a surrounding atom does not have an octet; move lone pair to bonding position from an adjacent atom to form a double or triple bond. Group IV C 4 bonds 0 Lone Pairs # Bonds = 8 Group# # LonePairs = 4 # Bonds Group V N, P 3 bonds 1 Lone Pair Group VI,S 2 bonds 2 Lone Pairs Group VII,Cl, Br, I 1 bond 3 Lone Pairs Double bond two atoms share two pairs of electrons C or C double 8e - 8ebonds - - double bonds Triple bond two atoms share three pairs of electrons N N or N N triple 8e - 8e bond - triple bond Exception to Generality of # Bonds and Lone Pairs for Atom 1. Resonance Structures 2. Polyatomic Ions 3. Exceptions to the ctet Rule Incomplete ctet Expanded ctet dd # Electrons 9.4 5
Write the Lewis structure of nitrogen trifluoride (N 3 ). Step 1 N is less electronegative than, put N in center Step 2 Count valence electrons N - 5 and - 7 5 + (3 x 7) = 26 valence electrons Step 3 Draw single bonds between N and atoms and complete octets on N and atoms. Step 4 - Check, are # of e - in structure equal to number of valence e -? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons N Resonance Structures 1. RS are Imaginary; the Real Structure is Intermediate All the RS. 2. nly Electrons Move Between RS; Atoms Positions Never Change 3. At Least ne Atom in RS Will Have a Non-zero ormal Charge and Won t ollow Generality About # Bonds and Lone Pairs. 9.6 Write the Lewis structure of the carbonate ion (C 3 2- ). An atom s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure. formal charge on an atom in a Lewis structure = total number of valence electrons in - the free atom total number of nonbonding - electrons 1 2 ( ) total number of bonding electrons The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion. 9.7 Step 1 C is less electronegative than, put C in center Step 2 Count valence electrons C - 4 (2s 2 2p 2 ) and - 6 (2s 2 2p 4 ) -2 charge 2e - 4 + (3 x 6) + 2 = 24 valence electrons Step 3 Draw single bonds between C and atoms and complete octet on C and atoms. Step 4 - Check, are # of e - in structure equal to number of valence e -? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons Step 5 - Too many electrons, form double bond and re-check # of e - C 2 single bonds (2x2) = 4 1 double bond = 4 8 lone pairs (8x2) = 16 Total = 24 9.6 6
What are the resonance structures of the carbonate (C 32 -) ion? Check Lewis Dot Structures - - C C - - - C - 9.8 1. Correct Number of Valence Electrons in Structure 2. Every Atom ollows the ctet Rule 3. Every Atom ollows Generality About # Bonds and Lone Pairs UNLESS it ollows ne of the Exceptions Exceptions to the ctet Rule (on Central Atom) The Incomplete ctet Incomplete ctet! 4 not 8 e- BeH 2 BH 3 B 3e - 3H 3x1e - 6e - Be 2e - 2H 2x1e - H B H 4e - H Be Incomplete ctet! 6 not 8 e- H H Coordinate Covalent Bonds In a coordinate covalent bonds, both electrons in the bond come from one of the atoms the bond is between. Usually, coordinate covalent bond is between An atom with incomplete octet and An atom with a lone pair 9.9 7
Exceptions to the ctet Rule (on Central Atom) The Expanded ctet (central atom with principal quantum number n 3 S 6 S 6e - 6 42e - 48e - S 6 single bonds (6x2) = 12 18 lone pairs (18x2) = 36 Total = 48 9.9 dd-electron Molecules N Exceptions to the ctet Rule N 5e - 6e - 11e - * When there is an odd # of valence electrons; ne atom will have 7 rather than 8 electrons * Use the electronegativity of atoms to determine which atom has 7 rather than 8 electrons Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond. Electronegativity - relative, is highest 9.5 9.5 8
Lewis Dot Structure for N (cont) is more electronegative than N The atom with a lower electronegativity has a lower attraction for the electrons in the bond and will be the atom with 7 rather than 8 electrons Polar covalent bond or polar bond is a covalent bond with between 2 nonmetal atoms of different electronegativities. The electron pair is closer to the atom with a higher electronegativity electron poor region electron rich region e - poor H H δ + δ - e - rich ree radical unpaired electron N 9.5 Bond Type Comparison NonPolar covalent bond or nonpolar bond is a covalent bond between atoms of the about the same electronegativity; the electrons pair is equally shared between atoms Bond Type Type of Atoms Involved Ionic Metal & Nonmetal Polar Covalent 2 Nonmetal Atoms of Different EN NonPolar Covalent 2 Nonmetal Atoms with Similar EN Example NaCl ICl Cl 2 H-H 9
Classification of bonds by difference in electronegativity Electronegativity Difference Bond Type 0 NonPolar Covalent 2 Ionic 0 < and <2 Polar Covalent Increasing difference in electronegativity NonPolar Covalent Polar Covalent Ionic Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H 2 S; and the NN bond in H 2 NNH 2. Cs 0.7 Cl 3.0 3.0 0.7 = 2.3 Ionic H 2.1 S 2.5 2.5 2.1 = 0.4 Polar Covalent N 3.0 N 3.0 3.0 3.0 = 0 NonPolar Covalent share e - partial transfer of e - or unequal sharing transfer e - 9.5 9.5 The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy. Bond Bond Type Energy (kj/mole) C-C 347 C=C 620 C C 812 Bond Energies Single bond < Double bond < Triple bond Lengths of Covalent Bonds Bond Lengths Bond Type Bond Length (pm) C-C 154 C=C 133 C C 120 C-N 143 C=N 138 C N 116 9.10 Triple bond < Double Bond < Single Bond 9.4 10
Table 9.2 Bond energy reflects electronegativity difference of bonded atoms Bond length reflects atomic or ionic radius Table 9.3 Bond length decreases with increasing energy Estimating ΔH for a Reaction from Bond Energies 1. Balance Chemical Equation 2. Write Lewis Dot Structures for all reactants and products. 3. Calculate energy needed to break reactant bonds and make product bonds. Breaking bonds is endothermic (+) Making bonds is exothermic (-) 11
Use bond energies to calculate the enthalpy change for: H 2 (g) + 2 (g) 2H (g) Type of bonds broken Number of bonds broken Bond energy (kj/mol) Energy change (kj) H H 1 436.4 436.4 1 156.9 156.9 Type of bonds formed Number of bonds formed Bond energy (kj/mol) Energy change (kj) H 2 568.2 1136.4 ΔH 0 = 436.4 + 156.9 2 x 568.2 = -543.1 kj 9.10 12