Chapter 8 Basic Concepts of Chemical Bonding CHEMISTRY The Central Science 9th Edition Chapter 8 Basic Concepts of Chemical Bonding Chapter 8

8.1: Chemical Bonds, Lewis Symbols, and the Octet Rule Chemical bond: attractive force holding two or more atoms together Covalent bond: electrons are shared Usually found between nonmetals Ionic bond results from the transfer of electrons from a metal to a nonmetal Metallic bond: attractive force holding pure metals together Chapter 8

Valence electrons are represented as dots around the symbol for the element Electrons available for bonding are indicated by unpaired dots Text, P. 300 Chapter 8

All noble gases except He have an ns2np6 configuration The Octet Rule All noble gases except He have an ns2np6 configuration Octet rule: atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs) Chapter 8

Na(s) + ½Cl2(g)  NaCl(s) DHºf = -410.9 kJ 8.2: Ionic Bonding Consider the reaction between sodium and chlorine: Na(s) + ½Cl2(g)  NaCl(s) DHºf = -410.9 kJ Chapter 8

NaCl forms a very regular structure Regular arrangement of Na+ and Cl- in 3D Ions are packed as closely as possible Chapter 8

Energetics of Ionic Bond Formation Lattice energy: the energy required to completely separate an ionic solid into its gaseous ions NaCl(s)  Na+(g) + Cl-(g) is endothermic (H = +788 kJ/mol) Lattice energy depends on the charges on the ions and the sizes of the ions Chapter 8

Lattice energy (and thus stability) increases as The charges on the ions increase The distance between the ions decreases High lattice energies make ionic compounds hard and brittle with high melting points Chapter 8

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8.3: Covalent Bonding When two similar atoms bond, neither of them wants to lose or gain an electron to form an octet Similar electron affinities They share pairs of electrons to each obtain an octet Example: H + H  H2 Chapter 8

Lewis Structures Covalent bonds can be represented by the Lewis symbols of the elements: In Lewis structures, each pair of electrons in a bond is represented by a single line: Chapter 8

Bond distances decrease from single  triple bonds Multiple Bonds It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds): One shared pair of electrons = single bond Two shared pairs = double bond Three shared pairs = triple bond Bond distances decrease from single  triple bonds Chapter 8

8.4: Bond Polarity and Electronegativity Sharing of electrons to form a covalent bond does not imply equal sharing of those electrons Unequal sharing of electrons results in polar bonds Chapter 8

Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F) Electronegativity Electronegativity: The ability of one atom in a molecule to attract electrons to itself Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F) Electronegativity increases across a period decreases down a group Chapter 8

Electronegativity Text, P. 310

Electronegativity and Bond Polarity Difference in electronegativity is a gauge of bond polarity: (NO NOTES) differences around 0: non-polar covalent bonds (equal or almost equal sharing of electrons) differences around 2: polar covalent bonds (unequal sharing of electrons) differences around 3: ionic bonds (transfer of electrons) Chapter 8

There is no sharp distinction between bonding types The positive end (or pole) in a polar bond is represented + and the negative pole - Chapter 8

There is more electron density on F than on H Dipole Moments Consider HF: Polar bond There is more electron density on F than on H Since there are two different “ends” of the molecule, we call HF a dipole Dipole moment, m, is the magnitude of the dipole Chapter 8

Bond Types and Nomenclature Ionic Molecular MgH2 Magnesium hydride H2S Hydrogen sulfide FeF2 Iron(II) fluoride OF2 Oxygen difluoride Mn2O3 Manganese(III) oxide Cl2O3 Dichlorine trioxide The least EN element is first The more EN element ends in –ide Use prefixes for molecular compounds Chapter 8

8.5: Drawing Lewis Structures Calculate valence electrons needed to achieve noble gas configuration for each element in molecule (N) Calculate the actual # of valence electrons for each element in molecule (A) (add e-/subtract e- for polyatomic ions) Calculate the number of shared electrons in molecule (S): S = N – A Develop a skeletal structure (least EN element central; symmetry) Place S electrons in the structure as shared pairs Place remaining A electrons on atoms to make octets (H and He) Chapter 8

To determine which structure is most reasonable, use formal charge It is possible to draw more than one Lewis structure with the octet rule obeyed for all the atoms To determine which structure is most reasonable, use formal charge Formal charge is the charge on an atom that it would have if all the atoms had the same electronegativity Chapter 8

To calculate formal charge: FC = valence electrons - number of bonds - nonbonding electrons (For a polyatomic ion, sum equals the charge of the ion) Chapter 8

There are 4 valence electrons Consider: For C: There are 4 valence electrons In the Lewis structure there are 2 nonbonding electrons and 3 bonds Formal charge: 4 – 3 - 2 = -1 Chapter 8

There are 5 valence electrons Consider: For N: There are 5 valence electrons In the Lewis structure there are 2 nonbonding electrons and 3 bonds Formal charge = 5 – 3 - 2 = 0 The sum of the formal charges on the ion equals the charge of the ion (-1 + 0 = -1) We write: Chapter 8

The most stable structure has: the lowest formal charge on each atom (close to zero) the most negative formal charge on the most electronegative atom no adjacent atoms with FC of the same sign Chapter 8

Lewis structures: N2, SiF4, CO2, CS2, C3H8, (SO4)-2, (NH4)+, H2SO4, SOCl2, NH2OH Sample Problems # 47, 51 Chapter 8

8.6: Resonance Structures Some molecules are not well described by a single Lewis Structure Resonance structures are attempts to represent a real structure that is a mix between several extreme possibilities Chapter 8

Common examples: O3, SO3, NO3-, NO2, and benzene Chapter 8

Resonance in Benzene In the resonance structures for benzene there are single bonds between each pair of C atoms and the 6 additional electrons are delocalized over the entire ring: Benzene belongs to a category of organic molecules called aromatic compounds (due to their odor) Chapter 8

8.7: Exceptions to the Octet Rule There are three classes of exceptions to the octet rule: 1. Odd Number of Electrons Few examples. Generally molecules such as ClO2, NO, and NO2 have an odd number of electrons dimerization occurs Chapter 8

2. Less than an Octet Relatively rare Molecules with less than an octet are typical for compounds of Groups 2 and 13 Be: N=4 B: N=6 Chapter 8

This is the largest class of exceptions 3. More than an Octet This is the largest class of exceptions Atoms from the 3rd period onward can accommodate more than an octet the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density Lewis Structures: If everything has an octet and “A” electrons remain, put them on the central atom Chapter 8

Lewis structures for: NO, BeCl2, BCl3, SF6, KrF2, PF5, (I3)- Sample Problems # 63, 65 Chapter 8

8.8: Strengths of Covalent Bonds The energy required to break a covalent bond is called the bond dissociation enthalpy, D Bond enthalpies can either be positive or negative Chapter 8

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Bond Enthalpies and the Enthalpies of Reactions Bond enthalpies can be used to calculate the enthalpy for a chemical reaction In any chemical reaction bonds need to be broken and then new bonds get formed Chapter 8

Mathematically, if Hrxn is the enthalpy for a reaction, then Sample Problem # 69 Chapter 8

Bond Enthalpy and Bond Length Multiple bonds are shorter than single bonds Multiple bonds are stronger than single bonds As the number of bonds between atoms increases, the atoms are held closer and more tightly together (bond enthalpy increases) Chapter 8

End of Chapter 8: Basic Concepts of Chemical Bonding