Chapter 9 : Ionic and Covalent Bonding

Transcription

1 Chapter 9 Ionic and Covalent Bonding Essentials of General Chemistry Ebbing Gammon Ragsdale 2nd Edition Dr. Azra Ghumman Memorial University of Newfoundland Chapter 9 Ionic and Covalent bonding Review Chemical bonding (9.4) 9.5 Polar covalent bonds Electronegativity 9.6 Writing Lewis Electron-Dot Formulas 9.7 Delocalized bonding Resonance 9.8 Exceptions to ctet rule 1

2 Chemical bond Valence electron play a fundamental role in chemical bonding. Electron transfer leads to ionic bonds. Monatomic - ions are formed by losing or gaining one or more electron from the highest energy principal level. Sharing of electrons leads to covalent bonds. e - are transferred or shared to give each atom a noble gas valence configuration Elements tend to have complete octet I.e. having 8 electrons in the valence shell. ydrogen prefer to have 2 electron in the valence shell (e e- configuration) Ionic Bonds Ionic Bond involves the transfer of electrons from one atom (usually a metal) to another (usually a nonmetal). A chemical bond formed by electrostatic attraction between positive and negative ions. ne extreme end of the bonding continuum Large number of oppositely charged ions gather together to form ionic solid 2

3 9.4 Covalent Bonds Covalent bond is a chemical bond that results from the sharing of electrons between nonmetal atoms and it is represented by a straight line. These atoms will share sufficient numbers of electrons in order to achieve a noble gas electron configuration (eight electrons and 2e- for 2 ) Covalent bond and ctet rule ctet rule. The tendency of atoms in a molecule to have eight electrons in their outer shell ( ns 2 p 6 except two e - for hydrogen) is called the octet rule. Figure 9.10 illustrates how two electrons can be shared by bonded hydrogen atoms. Figure 9.11 shows the potential energy of the atoms for various distances between nuclei. The decrease in energy is a reflection of the bonding of the atoms. 3

4 The electron probability distribution for 2 molecule Fig The electrons are attracted simultaneously by the positive charge of two nuclei. This attraction is the force holding both atoms together. This attraction is the force holding the atoms together. Potential Energy curve for 2 Bond length The distance between the nuclei at minimum potential energy is called bond length Bond dissociation energy Energy required to separate the two bonded atoms in a molecule is called bond dissociation energy Fig P.E of the atoms change The decrease in energy is a reflection of the bonding of the atoms 4

5 Lewis Structures You can represent the formation of the covalent bond in 2 as follows. +. This uses the Lewis dot symbols for the hydrogen atom and represents the covalent The shared electrons in 2 spend part of the time in the region around each atom. Lewis Formulas Lewis electron dot formula A formula using dots to represent valence electrons. Bonding electron pair An electron pair shared between two atoms. Non bonding or a lone pair An electron pair that remains on one atom and is not shared The formation of a bond between and Cl to give an Cl molecule can be represented as.. +. Cl Cl 5

6 Lewis Formulas Cl bonding pair lone pair The # of covalent bonds formed by an atom equals the number of unpaired electrons shown in Lewis formula e.g. formation of N 3 Coordinate Covalent Bonds When bonds form between atoms that both donate an electron, you have. +. A B A B It is, however, possible that both electrons are donated by one of the atoms. This is called a coordinate covalent bond. + A B A B 6

7 Coordinate Covalent Bonds Coordinate covalent bondone atom donates both the electrons (a lone pair) to another atom that has a vacant valence orbital Formation of N 4+ ion, in which N atom of N 3 donates a lone pair to + ion to form a bond that results in the formation of ammonium cation with a charge of plus one. N + + N + Multiple bonds Single bond A covalent bond formed by sharing a single pair of electrons by two atoms. Double bond A covalent bond formed by sharing two pairs of electrons by two atoms e.g. ethylene. Triple bond A covalent bond formed in which three pairs of electrons are shared by two atoms e.g. acetylene. 7

8 Multiple bonds Ethylene (double bond) C C Acetylene (triple bond) or C C C C or C C Double bonds form primarily with C, N,, and S atoms. Triple bonds forms mostly to C and N atoms. 9.5 Polar Covalent Bonds A polar covalent bond is one in which the bonding electrons spend more time near one (more electronegative) atom than the other. Non-polar covalent bond In which the bonding electrons are shared equally by bonded atoms. the bonded atoms are alike, as in the - bond of 2, A polar covalent bond is intermediate between two extremes; a nonpolar and an ionic bond - -Cl Na + Cl - Nonpolar covalent polar covalent ionic bond 8

9 Polar Covalent Bonds The bonding electrons are attracted somewhat more strongly by one atom in a bond e.g -Cl,,C 2 Electrons are not completely transferred Atom that attracts electrons more toward itself bears a partial negative charge (-δ Greek letter delta) Atom that has less attraction for the electrons bears a partial positive charge (+δ) e.g. C- bond in C 2 molecule δ δ + C δ Electronegativity Electronegitivity is a measure of the ability of an atom in a molecule to draw bonding electrons to itself. In general, Electronegativity increases from the lower-left corner to the upper-right corner of the periodic table. Electronegativity scale was developed by Linus Pauling, derived from bond energies assigns a value of 4.0 to fluorine and a value of 0.7 to cesium. The absolute difference in electronegativity values of two bonded atoms determines the polar character of the bond. 9

10 9_12 Electronegativities IA IIA 2.1 IIIA IVA VA VIA VIIA Li 1.0 Be 1.5 B 2.0 C 2.5 N F 4.0 Na 0.9 Mg 1.2 VIIIB IIIB IVB VB VIB VIIB IB IIB Al 1.5 Si 1.8 P 2.1 S 2.5 Cl 3.0 K 0.8 Ca 1.0 Sc 1.3 Ti 1.5 V 1.6 Cr 1.6 Mn 1.5 Fe 1.8 Co 1.8 Ni 1.8 Cu 1.9 Zn 1.6 Ga 1.6 Ge 1.8 As 2.0 Se 2.4 Br 2.8 Rb 0.8 Sr 1.0 Y 1.2 Zr 1.4 Nb 1.6 Mo 1.8 Tc 1.9 Ru 2.2 Rh 2.2 Pd 2.2 Ag 1.9 Cd 1.7 In 1.7 Sn 1.8 Sb 1.9 Te 2.1 I 2.5 Cs 0.7 Ba 0.9 La Lu f 1.3 Ta 1.5 W 1.7 Re 1.9 s 2.2 Ir 2.2 Pt 2.2 Au 2.4 g 1.9 Tl 1.8 Pb 1.8 Bi 1.9 Po 2.0 At 2.2 Fr 0.7 Ra 0.9 Ac No General trend in P.T Electronegativity increases from left to right and decreases from top to bottom Electronegativity and Polarity of the Bond Bond PolarityThe absolute value of the difference in electronegativity of two bonded atoms gives a rough measure of the polarity of the bond. When this difference is small (less than 0.5), the bond is nonpolar. When this difference is large (greater than 0.5), the bond is considered as polar. If the difference exceeds approximately 1.8, sharing of electrons is no longer possible and the bond becomes ionic. Electronegativity scale can be` used to predict the direction in which the electrons shift during bond formation. 10

11 Using electronegativities to obtain relative bond polarity Example Use the electronegativity values in Figure 9.5 to arrange the following bonds in order of increasing polarity C-Cl, P-, and As-Br. ( = 0.8) Solution P- C-Cl As-Br EN values P- EN = = 0 nonploar C-Cl EN = = 0.5 polar As-Br EN = = 0.8 polar The order of increasing bond polarity is P-, C-Cl, As-Br 9.6 Writing Lewis Electron-Dot Formulas Lewis electron-dot formula of a covalent compound is a simple two-dimensional representation of the positions of electrons in a molecule. Bonding electron pairs are indicated by either two dots or a dash. In addition, these formulas show the positions of lone pairs of electrons The given rules allow you to write electron-dot formulas even when the central atom does not follow the octet rule (p 267) 11

12 Steps to write Lewis structures Calculate the total number of valance electrons for the molecule by adding up the valence electrons for each element Add one electron for each negative charge and subtract one electron for each positive charge that an ion has. 2. Write the skeleton structure of the molecule or ion, connecting every bonded pair of atoms by a pair of dots or a dash subtract two electrons for each single bond formed to the central atom. 3. Distribute the remaining electrons to the terminal atoms surrounding the central atom to satisfy their octet rule. 4. Distribute the remaining electrons as pairs to the central atom (or atoms). Writing Lewis structures Note If there are fewer than 8 electrons on the central atom, it is likely that a multiple bond is present. To obtain a multiple bond, move one or two pair of electrons to the bond connecting to central atom. Usually C, N, and S form multiple bonds Note that group number indicates the number of valence electrons for main group elements (some exceptions). 12

13 Facts about central atom and terminal atoms Skeleton structure Must be found by experiment For simple molecules can be predicted by these facts ydrogen atoms are always terminal atoms. can only accommodate two electrons (a duet). Central atoms are generally those with the lowest electronegativity and higher valence except and. C atoms are always central atoms. Generally structures are compact and symmetrical. Most organic compounds are neither compact nor symmetrical. Writing Lewis Dot Formulas Write the Lewis electron-dot formulas (formula) of carbonyl chloride (phosgene gas, highly toxic) CCl 2, boron tetrafluoride ion BF 4-,and N 2 F 2, ClF

14 9.7 Delocalized bonding Resonance The structure of ozone, 3, can be represented by two different Lewis electron-dot formulas. orr Experiments show, however, that both bonds are identical (128 pm). 9.7 Delocalized bonding Delocalized bonding A type of bonding in which a bonding pair of electrons is spread over a number of atoms rather than localized between two. According to theory, one pair of bonding electrons is spread over the region of all three atoms. Symbolic description of the delocalized bonding in ozon Bond order is

15 Delocalized bonding Resonance Resonance description you describe the electron structure of a molecule having delocalized bonding by drawing all of the possible electron-dot formulas e.g. 3 r These are called the resonance formulas of the molecule Resonance structures Resonance can be expected when it is possible to draw more than one Lewis structure that follows octet rule. Actual structure is an average of various possible electron-dot structures for a molecule. Resonance is indicated by a straight double headed arrow ( ) The different resonance forms differ only in the placement of the valence electrons. total number of electrons and relative positions of the atoms remain the same. Resonance forms are not different kind of molecules; there is only one 3 molecule, its structure is average of two resonance forms drawn for it. 15

16 Problem Write all possible resonance formulas for the following species; C 3 2-, N 3-, S 2 and N Exceptions to the ctet Rule Although many molecules obey the octet rule, there are some exceptions; a. A molecule having odd # of electrons e.g. N b. where the central atom has fewer than eight electrons e.g. Be, B, and Al. c. where the central atom has more than eight electrons e.g. nonmetals in third period or great can accommodate more than eight electron in their valence e.g. P, S, Xe etc. These elements have empty nd orbitals that can be used for bonding. 16

17 Exceptions to the ctet Rule Incomplete octets. Some elements like Boron may be electron deficient, i.e. have less than eight electron in the valence shell. F F N Lewis base + F B BF 3 is strong Lewis acid and form a coordinate covalent bond with ammonia molecule N 3 (strong Lewis base) F Lewis acid N B F F Exceptions to the ctet Rule Expanded octets Nonmetals in third period or great can accommodate more than eight electron in their valence shell e.g. P, S, Xe etc Example Write the Lewis electron-dot formula for the following molecules; PF 5, XeF 4, SF 6. 17

18 Formal Charge and Lewis Structures In certain instances, more than one feasible Lewis structure can be illustrated for a molecule. For example, C N or N C The concept of formal charge can help to determine the best Lewis formula. write the skeleton structure of the molecule. e.g. Lewis structure for 2 2 Formal Charge and Lewis Structures The formal charge of an atom is simply the difference between the number of valence electrons on the free atom and the number it possesses within the structure. Rules to assign v.e - to individual atoms; Both e - s of a lone-pair belong to only one atom on which they are placed. An atom owns half of the bonding electrons. FC = (# of valance electrons) (# of unshared electrons on the atom) (½ of the shared electrons) r FC = (group number) (number of lone-pair electrons) ½(number of bonding pair) Sum of formal charges on the atoms should be equal to the charge on the molecular species (zero for neutral and charge on the ion) 18

19 Formal Charge and Lewis Structures Whenever there is a possibility of writing several Lewis formulas for a molecule use these rules to make the selection of the best possible structure 1. Choose the structure with the least number of atoms carrying formal charge. 2. If the structures have the same number of atoms carrying formal charge, choose the structure with the negative formal charge on the more electronegative atom. Example Write a Lewis formula for hydrogen peroxide, 2 2. Select the best possible structure according to rules of formal charge (will be solved in the class). 19

20 Examples 1. Write all possible Lewis structures for hydrogen cyanide, CN and assign the formal charges to all atoms in the molecule. Indicate the best possible structure according to the rules of formal charge (will be solved in the class). perational Skills Using electronegativities to obtain relative bond polarity. Writing Lewis formulas. Writing resonance structures and determining the best Lewis structure using formal charges. Exception to octet rule. 20