Chapter 9 MODELS OF CHEMICAL BONDING

Transcription

1 Chapter 9 MODELS OF CHEMICAL BONDING 1 H H A + B H H A B

2 A comparison of metals and nonmetals 2

3 9.1 Atomic Properties & Chemical Bonds Chemical bond: A force that holds atoms together in a molecule or compound Types of Chemical Bonding Ionic bonding involves the transfer of electrons and is usually observed when a metal bonds to a nonmetal. Covalent bonding involves the sharing of electrons and is usually observed when a nonmetal bonds to a nonmetal. Metallic bonding involves electron pooling and occurs when a metal bonds to another metal. 3

4 Three models of chemical bonding 4

5 Gradations in bond type among Period 3 and Group 4A elements 5 Most bonds are somewhere in between ionic and covalent.

6 Lewis Symbols & the Octet Rules Lewis symbols for atoms show valence electrons only 6 Rules of the Game # of valence electrons of a main group atom = Group number

7 Lewis Symbols & The Octet Rule 7 Why do compounds form? -Elements with unfilled valence shells usually form compounds with other elements to gain stability. The OCTET RULE: Many atoms gain or lose electrons so as to end up with the same # of electrons as the noble gas closest to them on the periodic table Example: Phosphorus has 5 valence electrons and want to get 3 more electrons in its valence shell P Group 5A(15) 1s 2 2s 2 2p 6 3s 2 3p 3 Br Group 7A(17) [Ar] 4s 2 3d 10 4p 5 Core = [Ar] 3d 10, valence = 4s 2 4p 5 Bromine has 7 valence electrons and want to get 1 more electron in its valence shell

8 9.2 The Ionic Bonding Model Usually made up of a metal ion and one or more nonmetal ions»metals have low IE and tend to lose electrons (to form cations) e.g., Li becomes Li + ion.»nonmetals have large negative values of EA and tend to gain electrons (to form anions) e.g., F becomes F - ion. Li 1s 2 2s 1 + F 1s 2 2p 5 Li + 1s 2 + F - 1s 2 2s 2 2p 6 8 Li F Li + + F - - Why don t compounds such as Li 2 F, LiF 2 exist?

9 Why do ionic compounds form? The importance of Lattice Energy DH 0 lattice 9 Lattice energy is the energy required to separate 1 mol of an ionic solid into gaseous ions. Lattice energy is a measure of the strength of the ionic bond. The Born-Haber Cycle to determine Lattice Energy

10 Periodic Trends in Lattice Energy Coulomb s Law: 10 - As ionic size increases, lattice energy decreases. Lattice energy therefore decreases down a group on the periodic table.

11 Periodic Trends in Lattice Energy Coulomb s Law: 11 - As ionic charge increases, lattice energy increases. NaCl, Na + and Cl -, m.p. 804 o C MgO, Mg 2+ and O 2- m.p o C

12 Properties of Ionic Compounds 12 Ionic compounds tend to be hard, rigid, and brittle, with high melting points. Ionic compounds do not conduct electricity in the solid state. In the solid state, the ions are fixed in place in the lattice and do not move. Ionic compounds conduct electricity when melted or dissolved. In the liquid state or in solution, the ions are free to move and carry a current. ionic compounds are easily cracked

13 13 Interionic attractions are so strong that when an ionic compound is vaporized, ion pairs are formed.

14 9.3 The Covalent Bonding Model 14 Covalent bond arises from the mutual attraction of 2 nuclei for the same electrons. Electron sharing results. Covalent bond is a balance of attractive and repulsive forces. H H A + B H H A B

15 The Covalent Bonding Model 15 Formation of a covalent bond results in greater electron density between the nuclei. Covalent bond formation in H 2.

16 Bonding Pairs & Lone Pairs of electrons 16 Valence electrons are distributed as shared or BONDING PAIRS and unshared or LONE PAIRS. G. N. Lewis This is called a LEWIS ELECTRON DOT structure of a molecule.

17 Covalent Bond Properties: Order, Energy. & Length 17 What is the effect of bonding and structure on molecular properties? Free rotation around C C single bond No rotation around C=C double bond

18 Bond Order # of electron pairs being shared between a given pair of atoms 18 Double bond Single bond Acrylonitrile Triple bond

19 Bond Energy (Bond Strength, Bond Enthalpy) measured by the energy required to break a covalent bond. Bond Enthalpy has positive values. BOND STRENGTH (kj/mol) H H 436 C C 346 C=C 610 C C 835 N N 945 The GREATER the bond order the HIGHER the bond strength and the SHORTER the bond length. 19

20 Bond Strength 20 Bond Order Length Strength HO OH O=O pm 210 kj/mol kj/mol O O O ?

21 21 Bond Length The distance between the nuclei of two bonded atoms. H F Bond length depends on size of bonded atoms. H Cl H I Bond distances measured in Angstrom units where 1 A = 10-2 pm = 10-8 cm

22 22 Bond Length Bond length also depends on bond order. Bond distances measured in Angstrom units where 1 A = 10-2 pm.

23 Physical Properties of Covalent Substances 1) Poor electrical conductivity 2) Physical properties of - Molecular covalent substances: gas, liquid, lowmelting point solid due to Strong and localized intramolecular (bonding) forces Weak intermolecular forces - Network covalent solids: Covalent bonds throughout the sample 23 Quartz (SiO 2 ) mp C Diamond mp C

24 9.4 Bond Energy & Chemical Change Using Bond Energies to Calculate DH 0 rxn H 0 rxn = (Sum of energy required to break bonds) + (Sum of energy generated by newly formed bonds) 24 Guide: 1) Break ALL the reactant bonds to obtain individual atoms 2) Use the atoms to form ALL the product bonds 3) Add the bond energies with appropriate signs to obtain DH 0 rxn H 0 rxn > 0 for broken bonds H 0 rxn < 0 for formed bonds

25 Problem: Using bond energies to calculate DH rxn for HF formation. H 2 (g) + F 2 (g) 2 HF (g) 25 H 0 rxn = (BE H2 + BE F2 ) + 2 BE HF H 0 rxn = (432 kj/mol kj/mol) + ( 2x565 kj/mol) H 0 rxn = -539 kj/mol

26 Using bond energies to calculate DH rxn for the combustion of methane. CH4 (g) + 2 O 2 (g) CO 2 (g) + 2 H 2 O (g) 26 H 0 rxn = [4 BE CH4 + 2 BE O=O ] + [2 BE C=O + 4 BE O-H ] H 0 rxn = [(4 x 413) + (2 x 498)] + [(2 x -799) + (4 x -467)] H 0 rxn = -818 kj/mol

27 Bond Strengths & Heat Released from Fuels & Foods 27 Weaker bonds such as C-H bonds (less stable, more reactive) are easier to break than stronger bonds such as C-O bonds (more stable, less reactive) Relative bond strength and energy from fuels.

28 Bond Strengths & Heat Released from Fuels & Foods 28 Fuels with fewer bonds to O release more energy.

29 9.5 Electronegativity and Bond Polarity A covalent bond in which the shared electron pair is not shared equally, but remains closer to one atom than the other, is a polar covalent bond. The ability of an atom in a covalent bond to attract the shared electron pair is called its electronegativity. Unequal sharing of electrons causes the more electronegative atom of the bond to be partially negative and the less electronegative atom to be partially positive. Due to the bond polarity, the H Cl bond energy is GREATER than expected for a pure covalent bond H Cl BOND ENERGY pure bond BE HCl = 339 kj/mol real bond BE HCl = 432 kj/mol

30 Electronegativity (EN) A measure of the ability of a bonded atom to attract shared electrons 30 F has highest EN value Relative values of EN determine BOND POLARITY

31 Trends in Electronegativity 31 Electronegativity and atomic size. EN increases as atomic size decreases. Nonmetals have higher EN than metals. Figure 9.22

32 Electronegativity and Oxidation Number 32 Electronegativities can be used to assign oxidation numbers (ON): The more electronegative atom is assigned all the shared electrons. The less electronegative atom is assigned none of the shared electrons. Each atom in a bond is assigned all of its unshared electrons. O.N. = # of valence e - - (# of shared e - + # of unshared e - )

33 33 Example: Cl is more electronegative than H, so for Cl: valence e - = 7 shared e - = 2 unshared e - = 6 O.N. = 7 (2 + 6) = -1 H is less electronegative than Cl, so for H: valence e - = 1 shared e - = 0 (all shared e - assigned to Cl) unshared e - = 0 O.N. = 1 (0 + 0) = +1

34 Bond Polarity & Partial Ionic Character Depicting Polar Bonds The unequal sharing of electrons can be depicted by a polar arrow. The head of the arrow points to the more electronegative element. 34 A polar bond can also be marked using δ+ and δ- symbols. + - H Cl

35 The Importance of DEN 35 - Percent ionic character of a bond increases with DEN. - No bonds are purely ionic or covalent.

36 Determine Bond Polarity from DEN Which bond is more polar? O H O F DEN DEN OH bond is more polar than OF bond, and direction of polarity is opposite in these cases. - O + H + - O F

37 The Graduation in Bonding across a Period Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. 37 Properties of the Period 3 chlorides. As DEN decreases, melting point and electrical conductivity decrease because the bond type changes from ionic to polar covalent to nonpolar covalent.

38 9.6 Introduction to Metallic Bonding 38 The Electron Sea Model: All metal atoms in the sample contribute their valence electrons to form a delocalized electron sea. The metal ions (nuclei with core electrons) lie in an orderly array within this mobile sea. All the atoms in the sample share the electrons. The metal is held together by the attraction between the metal cations and the sea of valence electrons. Group 1A/1 Group 2A/2

39 Properties of Metals Metals are generally solids with moderate to high melting points and much higher boiling points. Melting points decrease down a group and increase across a period. 39

40 Properties of Metals 40 - Mechanical properties: Metals can be shaped without breaking. The electron sea allows the metal ions to slide past each other. Metals dent or bend rather than crack - Electric Conductivity: Metals are good conductors of electricity in both the solid and liquid states. The electron sea is mobile in both phases. - Heat Conductivity: Metals are good conductors of heat.