OCTET RULE Generally atoms prefer electron configurations with 8 valence electrons. - Filled s and p subshells

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1 TYPES EMIAL BDIG 1 Ionic Bonding - Bond between ions whose charges attract each other - ne atom gives electrons and one atom takes electrons. Example a + l - ionic bond ovalent Bonding - two atoms each sharing electrons within a molecular orbital covalent bond - both fluorine atoms own electrons within bond TET RULE Generally atoms prefer electron configurations with 8 valence electrons. - illed s and p subshells Atoms bond with each other so that every atom has 8 electrons in its outer shell. - Atoms may take electrons from each other or they may share electrons. ctet rule is able to explain a lot of chemistry. I. e., it is able to explain why certain elements combine together in specific proportions. Exceptions to the octet rule are plentiful. We will consider these later.

2 LEWIS STRUTURES 2 nly valence electrons are important in bonding. Lewis dot structures show valence electrons surrounding atom. We visualize the four valence orbitals of an atom as the sides of a box. Electrons are put into orbitals according to und s rule. Examples Be has 2 valence electrons. Therefore Lewis structure is Be has 5 valence electrons. Therefore Lewis structure is Br has 7 valence electrons. Therefore Lewis structure is Br has 4 valence electrons. Therefore Lewis structure is II BDIG (again) - ppositely charged ions attract each other. - Metal atoms lose e - and nonmetal atoms gain e -. - Ions attract each other to form ionic lattice (crystal). Lewis structures can be used to illustrate ionic bonding. onsider Potassium and Bromine K + Br K + Br - onsider alcium and luorine a + - a 2+ - Lewis structures are much more illuminating when we consider the sharing of electrons (covalent bonding).

3 Lattice Energy energy of released when positive and negative ions form crystal lattice due to their attraction for each other. reation of ionic compound can be decomposed into many small steps. I. e., lattice energy can be decomposed into many smaller energy steps. ULMB S LAW - fundamental law of physics (hugely important in chemistry) 3 qq E = k r E energy of interaction q 1 charge of ion 1 q 2 charge of ion 2 k proportionality constant r 12 distance between ion 1 and ion 2 Three items of note about oulomb s law 1. If the energy is negative, ions must be oppositely charged. 2. The higher the charge, the greater the energy. 3. The higher the distance, the smaller the energy. oulomb s Law and Lattice Energy - oulomb s law can be used to make relative assessments of lattice energy. Example: Which compound has greater lattice energy: LiBr or a? Li + Br - a higher charges means greater lattice energy - Distance between ions is also important but to a much lesser degree. - Smaller distance means higher lattice energy. - ote that since ions are oppositely charged, the lattice energy is a negative value. - Greater lattice energy means a more negative value. Example: Which compound has higher lattice energy: al or sl? harges are identical; however, cesium ion is much larger than sodium ion. a l - s l - greater distance between charges means smaller lattice energy Therefore, al has higher lattice energy.

4 VALET BDIG - Atoms that are bonded share valence electrons. - Sharing is what creates covalent bond. - When atoms bond covalently, new entity termed molecule is formed. 4 ovalent Bonding in Diatomic Molecules ydrogen each atom contributes e - to the bond - e - in bond belongs to both atoms (e - are shared) - when sharing e -, each atom has full shell (like e) - note hydrogen is an exception to octet rule luorine Since both atoms share electrons in bond, both atoms have 8 valence electrons, octet rule is satisfied. ydrogen fluoride (hydrofluoric acid) + xygen + - note that a molecule can have more than one covalent bond itrogen +

5 Bonding in Polyatomic Molecules 5 Water + Ammonia arbon Dioxide + Useful reminders in polyatomic bonding 1. arbon will form 4 covalent bonds per atom. 2. xygen will form 2 covalent bonds per atom. 3. itrogen will form 3 covalent bonds per atom. 4. ydrogen will form 1 covalent bond per atom. 5. alogens will form 1 covalent bond per atom. DRAWIG LEWIS DT STRUTURES R MLEULES Do the following steps in order. 1. Identify whether compound is molecule or polyatomic ionic compound. -I. e., Identify the charge of the molecule or ion. 2. Sum the number of valence electrons from all atoms in the molecule. 3. Add or subtract appropriate number of electrons for polyatomic ion. 4. Identify the central atom and decide how other atoms are bonded to it. - least electronegative atom is usually central atom - hydrogen is never central atom - oxygen is rarely central atom 5. Draw bonds between atoms. (Subtract 2 e - for each bond from total number of valence e - ) 6. omplete octets of peripheral (outer) atoms. (Subtract each electron used from total number of valence e - ) 7. Put remaining electrons in pairs on central atom. 8. If central atom has too few electrons to complete octet, change lone pairs on peripheral atoms into a double bond between central atom and peripheral atom.

6 Example: Draw the Lewis structure for Pl 3. 6 # of valence e - of P = 5 # of valence e - of l = 7 # of e - = 5 + 3(7) = 26 e - entral atom is P. l P l Used in the manufacture of pesticides and flame retardants. l Example: Draw the Lewis structure for a. - is the polyatomic ion # of valence e - of = 6 # of valence e - of = 1 # of valence e - added = 1 # of e - = = 8 Including the sodium ion we could write + - a ote: In a, there is ionic and covalent bonding. a is also known as lye and is the essential ingredient in drain cleaners. Example: Draw the Lewis structure of S # of e - = 6 + 3(6) = 24 e - 3. entral atom is S. S 7. entral atom has too few electrons, use multiple bond. S S 3 is a component of pollution from burning coal. When mixed with water it makes a component of acid rain. - doesn t matter which oxygen atom makes double bond.

7 Example: Draw the Lewis Structure of e 3 (P 4 ) 2. 7 P 4 3- is the polyatomic ion # of e - = 5 + 4(6) + 3 = 32 e - ote: three electrons are added. entral atom is P. P Iron (II) phosphate is used by organic farmers to control snails and slugs e 2+ P e 2+ P e 2+ Example: Draw the Lewis Structure of 4 l. # of e - = 5 + 4(1) 1 = 8 e - ote: one electron is subtracted. entral atom is. + Ammonium chloride can be used as an expectorant, an agent to loosen mucus in the respiratory tracts. RESAE - Sometimes more than one correct Lewis structure can be drawn. - In that case, actual structure is a blend of correct structures. Example: Draw a Lewis structure for the nitrate ion, 3 - # of e - = 5 + 3(6) + 1 = 24 e - is the central atom rule = 26 double bond is expected

8 hange lone pair into double bond to satisfy octet rule. 8 - The oxygen atom on the left was chosen arbitrarily. Any of the three atoms could have been used to form the double bond. hoosing the oxygen atom on the right would yield the structure: hoosing the oxygen atom on the bottom would yield the structure: The actual structure is a blend of resonance structures. - Double-headed arrows indicate resonance structures. - Double bond is not confined to a single pair. - Double bond is distributed over all three pairs. - Resonance is good, resonance makes a molecule more stable.

9 BD LEGTS 9 The length of the covalent bond depends on two items. 1. The atomic radii of the atoms involved in the bond. Example: d( ) = 109 pm r() = 37 pm d( ) = 133 pm r() = 72 pm d( l) = 177 pm r(l) = 100 pm d( Br) = 194 pm r(br) = 114 pm d( I) = 213 pm r(i) = 133 pm r() = 77 pm 2. The number of bonds between the atoms. Example: d( ) = 146 pm d( ) = 154 pm d( = ) = 122 pm d( = ) = 134 pm d( ) = 110 pm d( ) = 121 pm Generally, the shorter the bond length, the higher the bond dissociation energy. d( ) = 154 pm d( = ) = 134 pm d( ) = 121 pm E dis = 348 kj/mol E dis = 614 kj/mol E dis = 839 kj/mol SAPES MLEULES (VSEPR MDEL) Valence Shell Electron-Pair Repulsion model - Electron pairs surrounding atom spread out as to minimize repulsion. - Electron pairs can be bonding pairs or nonbonding pairs (multiple bonds not included). - arrangement of all the atoms surrounding central atom depends on electron pairs surrounding central atom. Two similar, but different geometries 1. Electron set geometry - arrangement of e - sets around central atom - remember: e - set from multiple bonds is only one set. 2. Molecular geometry - arrangement of atoms around central atom **A molecular geometry is decided only after an electron set geometry has been determined.** *- need to write Lewis structure to determine number of electron sets.*

10 Geometries with two e - sets about central atom. 1. electron set geometry linear 10 A A generic atom - angle between e - pairs is possible molecular geometries a) Linear - only linear geometry is possible with two electron pairs Example: Bel 2 l Be l Example: 2 - ote: nly two electron sets around central atom since double bonds only count as one. Geometries with three e - sets about central atom. 1. electron sets geometry trigonal planar A A generic atom - angle between e - pairs is possible molecular geometries a) Trigonal Planar - all three electron sets are bonding pairs Examples: B 3 and 3 - B b) Bent (V-shaped) - two bonding pairs and one nonbonding pair Example: S 2 S *onbonding e- pairs take up more room than bonding pairs. Therefore bond angle between oxygen atoms is slightly less than 120.*

11 Geometries with four e - sets about central atom. 1. electron set geometry tetrahedral - tetrahedron is three dimensional object - angle between electron pairs is A aution! Representation is 2-D, not 3-D, bond angle between e - pairs is not possible molecular geometries a.) Tetrahedral - all four electron sets are bonding pairs Example: 4 2-D picture Example: P D picture P P b) Trigonal Pyramidal - 3 bonding pairs and 1 nonbonding pair Example: 3 - Bond angle is Bond angle is less than because nonbonding pair takes more room than bonding pair.

12 Example: l 3-12 l l c) Bent (V-shaped) - two bonding pairs and two nonbonding pairs Example: 2 - Bond angle is Redraw 2 to show tetrahedral angle. Example: S 2 S

13 ELETREGATIVITY 13 In covalent bonding, sharing of electrons is rarely perfect. ne atom will draw electrons closer to itself than the other atom. onsider Methanol = - in the bond, oxygen draws e - to itself perfect sharing imperfect sharing Electronegativity ability of an atom in a molecule to attract electrons to itself. - symbol is χ (Greek chi) Periodicity of Electronegativity increasing PLAR VALET BDS - bond between two atoms where electrons are imperfectly shared - more electronegative atom has a greater electron density about it - bond has poles, regions that are have more positive or negative charge - more electronegative atom is slightly negative relative to less electronegative atom δ + δ - δ small change in charge density (lower case Greek delta)

14 ELETREGATIVITY AD BDIG 14 When the electronegativities of the two atoms involved in a bond are known, the type of bonding can be predicted. onsider two atoms, A and B. Atom A Atom B Bonding Less than 2.20 Less than 2.20 metallic Greater or = 2.20 Less than 1.70 ionic Greater or = 2.20 Greater or = 1.70 covalent Is one of the atoms greater or equal to 2.20? o Bonding is metallic Yes Is the other atom greater or equal to 1.70? o Bonding is ionic Yes Examples Li 1.5 Be Bonding is covalent 2.0 B haracterize the bonding of is used in Teflon for non-stick cookware. 4 χ = 2.5 χ = 4.0 Bonding is covalent - Both nonmetal atoms want the electrons, so the compromise is a covalent bond Li 2 2 is used in subs to convert 2 to haracterize the bonding of Li 2 2 χ Li = 1.0 χ = 3.5 Bonding is ionic - Metal atom doesn t want electrons and nonmetal atom want electron, so ionic bonding is a perfect result Be 2 forms polymer [long chain] rather than molecule. 3. haracterize the bonding of Be 2 χ Be = 1.5 χ = 2.1 Bonding is metallic - Both metals atoms don t want the electrons, so the valence electrons become delocalized in a metallic bond

15 PLAR MLEULES 15 A polar molecule has one side slightly positive and the other slightly negative. Two conditions must be met in a polar molecule. 1.) Polar covalent bonds 2.) orrect geometry To emphasize necessity of correct geometry, compare two examples. Ge 2 2 χ Ge = 2.0 χ = 2.1 χ l = 3.5 χ = 3.5 χ Ge χ = 1.5 χ χ = 1.4 δ + δ - δ + δ - Ge indicates positive end of bond Ge bond is more polar than bond; however, Ge 2 is nonpolar molecule and 2 is a polar molecule. Question: ow can this be? Answer: Ge 2 has a linear geometry and 2 has a bent geometry. Ge δ + total polarity adds to zero δ - total polarity is nonzero Example: Is either ammonia or methane a polar molecule? δ - δ + Answer: Ammonia is a polar molecule, but methane is a nonpolar molecule. Scheme: hemical formula Lewis structure e - set geometry molecular geometry polarity