Chapter 8. Homework. Valence Electrons. Molecular Structure & Bonding. Example of Lewis Dot Symbols

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1 Homework Chapter 8 Bonding and Molecular Shapes: Fundamental Concepts Chapter 8 21, 23, 31, 35, 39, 47, 51, 57, 61, 65, 71, 73, 81, 83, 89, 105, 109, 113 Molecular Structure & Bonding Structure Refers to the way atoms are arranged in space Bonding Defines the forces that hold adjacent atoms together The bonding and structural characteristics for individual atoms are very similar from molecule to molecule. Knowing a compound s structure and understanding the bonding in that compound are prerequisites to explaining its chemical properties. Valence Electrons The concept of valence electrons was first introduced by the American chemist G. N. Lewis. Valence Electrons Those electrons in the outermost shell of an atom that determines the chemical properties of that element. Core Electrons The remaining electrons (the inner electrons) that are not involved in chemical behavior. For the main group elements the valence electrons are the s and p electrons. For the transition elements the valence electrons are the s and d electrons. Lewis Symbols for Atoms (Lewis Electron Dot Symbols) First introduced by G. N. Lewis, this is a useful way of representing the core and valence electrons of an element. 1. The element s symbol is chosen to represent the atomic nucleus together with the core electrons. 2. Up to four valence electrons, represented by dots, are placed around the symbol one at a time. 3. If any electrons remain, they are paired with the ones already there. Example of Lewis Dot Symbols Ca Se Ca Se

2 Chemical Bond The valence electrons are reorganized so that a new attractive force occurs between atoms. Two basic types of bonds Ionic bond Forms when one or more valence electrons is transferred from on atom to another, creating positive and negative ions. Covalent bonding Involves sharing of valence electrons between atoms. Bond Polarity Pure covalent bonding only occurs when two identical atoms are bonded. When two dissimilar atoms form a covalent bond, the electron pair will be unequally shared. Three types of bonds Covalent (nonpolar covalent) bond Electrons are completely shared Polar Covalent bond A partial charge is formed on each atom used in the bond Ionic (polar) bond A full charge separation occurs (+ and poles form) Figure 8.3: The Pauling electronegativity values. Electronegativity generally increases across a period and decreases down a group. EN 0.4 = Non-polar covalent 0.5 EN 1.6 = Polar covalent EN 1.7 = Ionic Examples LiF EN = = 3.0 Therefore it is IONIC. SiO EN = = 1.7 Therefore it is on the border of IONIC and POLAR. We consider this VERY POLAR. CO EN = = 1.0 Therefore it is (moderately) POLAR. Example Ionic Bonding (NaCl) Na + Cl [Na Cl ] [Na + Cl - ] Covalent Bonding (Cl 2 ) Cl + Cl Cl Cl Ionic Bonds Formed from electrostatic attractions of closely packed, oppositely charged ions. Formed when an atom that easily loses electrons reacts with one that has a high electron affinity.

3 Lattice Energy The change in energy when separated gaseous ions are packed together to form an ionic solid. M + (g) + X (g) MX(s) Lattice energy is negative (exothermic) from the point of view of the system. Lattice Energy = k( QQ / r) Q 1, Q 2 = charges on the ions 1 2 r = shortest distance between centers of the cations and anions Covalent Bonding For the rest of the chapter, we will be concerning ourselves with only covalent bonding. Remember that covalent bonding means that the electrons are shared between to atoms. Notation Take our Cl 2 example Cl + Cl Cl Cl The electrons around the Cl s that are not used in bonding are called lone pairs or nonbonding electrons. The electrons used in bonding is the chemical bond. Cl + Cl Cl Cl Cl Cl This is an example of a single bond, there are also double and triple bonds. Octet Rule With regards to the main group elements, each atom has to have eight electrons associated with it. There are exceptions! Obvious ones are H and He Fewer than eight electrons Boron has three valence electrons and so is expected to form three covalent bonds with other nonmetallic elements. BF 3 Greater than eight electrons Elements in the third or higher periods often form compounds and ions in which the central element is surrounded by more than four valence electron pairs. [ClF 4 ] - Drawing Lewis Dot Structures 1. Determine the total number of valence electrons available. Set this number equal to T. 2. Draw the skeleton of the compound. The central atom is usually the one with the lowest electron affinity. 3. Count up the number of bonding electrons used and subtract that from T. Set the number you get equal to R. 4. Count up the number of electrons needed to fill the octet rule for all of the atoms. Set this number equal to N. 5. If N = R then fill in the remaining electrons 6. If N > R then you need at add multiple bonds. (N-R) 2 = the number of bonds to add. 7. If N < R then you need to add electrons to the central atom (make sure the central atom is a 3 rd period or greater atom).

4 Example Draw the Lewis dot structure for the following: CCl 4 [CO 3 ] 2- [ClF 4 ] - [CN] - Isoelectronic Species Are molecules and ions having the same number of valence electrons and the same Lewis structure. Example (draw the following Lewis structures) NO + N 2 CO CN - Resonance Draw the Lewis structure for ozone (O 3 ) O=O O But there are alternate way of drawing the same structure. These are called resonance structures. The actual structure is a composite of all of the resonance structures. O=O O Resonance Occurs when more than one valid Lewis structure can be written for a particular molecule. These are resonance structures. The actual structure is an average of the resonance structures. Bond Order The number of bonding electrons pairs shared by two atoms in a molecule. It tells us what type of bond it is. 1 = single bond 2 = double bond 3 = triple bond Fractional = resonance Bond Length The distance between the nuclei of two bonded atoms. Largely determined by the sizes of the atoms. Table 8.5 lists average bond distances. Remember this is an AVERAGE! It should be observed that the higher the bond order the shorted the bond length. Bond Order = number of sharded pairs linking X and Y number of X Y links

5 Bond Energy The bond dissociation energy, D, is the enthalpy change for breaking a bond in a molecule. In order for chemical reactions to occur bonds in the reactants must be broken and the new bonds in the products must be formed. Breaking bonds is always endothermic Formation of bonds is always exothermic Bond Energy (cont d) See p. 372, Table 8.4 for some average bond energies. We can use these bond energies to calculate the H rxn. H rxn = ΣD(bondsbroken) ΣD(bondsformed) Example H H H H H H H C C C H + H H H C C C H H H H H 1. Break a C=C bond and a H-H bond 2. Form a C-C bond and two C-H bonds H rxn = (602kJ+ 436kJ) (346kJ kJ) = -134kJ Formal Charge Formal Charges on Atoms O C O (-1) (0) (+1) Not as good O C O (0) (0) (0) Better The charge calculated for that atom based on the Lewis structure of the molecule or ion using the following equation Formal Charge = # of valence electrons [# lone pair electrons + 1/2(# bonding electrons)] Calculate the following formal charges on all the atoms OH - NO - 3

6 Valence Shell Electron-pair Repulsion (VSEPR) model Devised by Ronald J. Gillespie and Ronald S. Nyholm. Reliable method for predicting the shapes of covalent molecules. The basic idea is that bond and lone electron pairs in the valence shell of an element repel each other and seek to be as far as apart as possible. Central Atoms Surrounded Only by Bond Pairs Predict the shape of SiCl 4 CH 2 Cl 2 Central Atom with Bond Pairs and Lone Pairs Lone pairs of electrons on the central atom occupy spatial positions even though their location is not included in the verbal description of the shape of the molecule or ion. Example NH 3 The lone pairs not only effect the shape of the molecule, but also the angles between the bonds. Central Atoms with More Than Four Valence Electron Pairs

7 Multiple Bonds and Molecular Geometry Double and triple bonds involve more electron pairs than single bonds, but this does not affect the overall molecular shape. Examples: CO 2 CO 3 2- NO 2 - Molecular Polarity Because most molecules have at least some polar bonds, molecules can also be polar. In a polar molecule, electron density accumulates toward one side of the molecule, giving that side a negative charge and leaving the other side with a positive charge. First, we consider each individual bond s polarity. Next, we look at the over all shape of the molecule and try to determine if all of the individual bond polarities would cause a molecular polarity (dipole moment). Examples: H 2 O CO 2 BF 3 Cl 2 CO