Basic Concepts of Chemical Bonding. In this Chapter. Chemical Bonds. Chapter 8

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1 Basic Concepts of Chemical Bonding Chapter 8 In this Chapter Types of chemical bonds. The Lewis symbol. Electronegativity and its periodic variation. Lewis structures of molecules. Resonance structures. Exceptions to the octet rule. Thermodynamics of bonding. 2 Chemical Bonds Chemical bond: attractive force holding two or more atoms together. Covalent bond: results from sharing electrons between the atoms. 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. 3 1

2 Lewis Symbols Useful as a pictorial understanding of where the valence electrons are in an atom. The valence electrons are represented as dots around the symbol for the element. The number of electrons available for bonding are indicated by unpaired dots. These symbols are called Lewis symbols. The electrons are generally placed on four sides of a square around the element s symbol. 4 Lewis Symbols 5 The ctet Rule All noble gases except He have a s 2 p 6 configuration. ctet rule: atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons (4 electron pairs). Caution: there are many exceptions to the octet rule. 6 2

3 Ionic Bonding Consider the reaction between sodium and chlorine: Na(s) + ½Cl 2 (g) NaCl(s) ΔHº f = kj 7 Ionic Bonding The reaction is violently exothermic, producing NaCl which is more stable than Na and Cl 2. Why? Na loses an electron to become Na + and the Cl gains the electron to become Cl -. Both Na + and Cl - have an octet of electrons surrounding the central ion. 8 Ionic Bonding NaCl forms a very regular structure in which each Na + ion is surrounded by 6 Cl - ions. Similarly, each Cl - ion is surrounded by six Na + ions. This results in a regular arrangement of Na + and Cl - in 3D which is known as an ionic lattice or a crystal lattice. The ions are packed as closely as possible. It is hard to find a molecular formula to describe the ionic lattice. 9 3

4 Covalent Bonding When two similar atoms bond, none of them wants to lose or gain an electron to form an octet. When similar atoms bond, they share pairs of electrons to each obtain an octet. Each pair of shared electrons constitutes one chemical bond. Example: H + H H 2 has electrons on a line connecting the two H nuclei. 10 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: 11 Multiple Bonds It is possible for more than one pair of electrons to be shared between two atoms (multiple bonds): ne shared pair of electrons = single bond (e.g. H 2 ); Two shared pairs of electrons = double bond (e.g. 2 ); Three shared pairs of electrons = triple bond (e.g. N 2 ). Generally, bond distances decrease as we move from single through double to triple bonds. 12 4

5 Bond Polarity & Electronegativity In a covalent bond, electrons are shared. Sharing of electrons to form a covalent bond does not imply equal sharing of those electrons. There are some covalent bonds in which the electrons are located closer to one atom than the other. Unequal sharing of electrons results in polar bonds. 13 Bond Polarity & Electronegativity Electronegativity: The ability of an atom in a molecule to attract electrons to itself. Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F). Electronegativity 14 Bond Polarity & Electronegativity 15 5

6 Bond Polarity & Electronegativity Difference in electronegativity is a gauge of bond polarity: electronegativity differences around 0 result in non-polar covalent bonds (equal or almost equal sharing of electrons); electronegativity differences around 2 result in polar covalent bonds (unequal sharing of electrons); electronegativity differences around 3 result in ionic bonds (transfer of electrons). 16 Drawing Lewis Structures HANDUT IN LECTURE 17 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, formal charges are used. Formal charge is calculated by the formula: Formal charge of an atom in a molecule/ion (FC) = Group number of the atom [# of lone pair electrons + ½ (# of bonding electrons) 18 6

7 Formal Charge Consider the hydroxide ion - H 19 Resonance Structures Resonance is a way of explaining the structure if a single Lewis structure fails to give an accurate picture. The atoms have the same arrangement in space in all resonance structures. Resonance structures only differ in the assignment of the electrons, never the assignment of atoms. The # of bond pairs between in a given pair of atoms can be differ. The actual structure of the molecule is considered to be a composite of all resonance structures. 20 Resonance Structures The carbonate ion can show the following resonance structures C C C Pauling s Electroneutrality principle: Electrons in a molecule are distributed so that the charges on the atoms are as close to zero as possible. 21 7

8 Electroneutrality Consider the possible resonance structures for C 2 : FC (for the C) = 4-[0+1/2(8)] = FC = 6-[4+1/2(4)] = 0 C C FC = 6-[4+1/2(4)] = 0 The first structure is the favored one and is what is observed experimentally. 22 Exceptions to the ctet Rule There are three classes of exceptions to the octet rule: Molecules with an odd number of electrons; Molecules in which one atom has less than an octet; Molecules in which one atom has more than an octet. dd Number of Electrons Few examples. Generally molecules such as Cl 2, N, and N 2 have an odd number of electrons. N N 23 Exceptions to the ctet Rule Less than an ctet Relatively rare. Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A. Most typical example is BF 3. Formal charges indicate that the Lewis structure with an incomplete octet is more important than the ones with double bonds. 24 8

9 Exceptions to the ctet Rule More than an ctet This is the largest class of exceptions. Atoms from the 3 rd period onwards can accommodate more than an octet. Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density. 25 Strengths of Covalent Bonds The energy required to break a covalent bond is called the bond dissociation enthalpy, D. That is, for the Cl 2 molecule, D(Cl-Cl) is given by ΔH for the reaction: Cl 2 (g) 2Cl(g). When more than one bond is broken: CH 4 (g) C(g) + 4H(g)ΔH = 1660 kj the bond enthalpy is a fraction of ΔH for the atomization reaction: D(C-H) = ¼ΔH = ¼(1660 kj) = 415 kj. Bond enthalpies can either be positive or negative

10 Strengths of Covalent Bonds We can use bond enthalpies to calculate the enthalpy for a chemical reaction. We recognize that in any chemical reaction bonds need to be broken and then new bonds get formed. The enthalpy of the reaction is given by the sum of bond enthalpies for bonds broken less the sum of bond enthalpies for bonds formed. 28 Enthalpies of Reactions Mathematically, if ΔH rxn is the enthalpy for a reaction, then ΔH rxn D bonds broken ( ) D( bonds ) = formed Consider the reaction between methane, CH 4, and chlorine: CH 4 (g) + Cl 2 (g) CH 3 Cl(g) + HCl(g) ΔHrxn =? 29 Strengths of Covalent Bonds 30 10

11 Enthalpies of Reactions In this reaction one C-H bond and one Cl-Cl bond gets broken while one C-Cl bond and one H-Cl bond gets formed. {[ ( ) + D( Cl - Cl) ] [ D( C - Cl) + D( H- Cl) ]} ΔH rxn = C -H D = 104 kj The overall reaction is exothermic meaning the bonds formed are stronger than the bonds broken