We will not be doing these type of calculations however, if interested then can read on your own

Transcription

1 Chemical Bond Lattice Energies and Types of Ions Na (s) + 1/2Cl 2 (g) NaCl (s) ΔH= -411 kj/mol Energetically favored: lower energy Like a car rolling down a hill We will not be doing these type of calculations however, if interested then can read on your own Question why Cl - and not Cl 2- Na + and not Na 2+ Mg 2+ and not Mg + Answer energetically most favored always have cation and anion Get out more energy than you put in Born Haber cycle is formal method of keeping track of energy changes Periodic Table (older Roman numerals) IA IIA IIIA s 2 and s 2 p 6 ions VA VIA VIIA s 2 and s 2 p 6 ions IB IIB d 10 ions IIIA IVA VA d 10 s 2 ions Transition metals frequently form more than one type of cation Most will form +2 by losing 4s 2 electrons and some others are more complex structures Example: Pour ammonia solution in bottle and add concentrated HCl to produce a white smoke NH 4 Cl which stays in bottle. Add water and shake and white smoke disappears because NH 4 Cl is soluble in water NH 3 + HCl NH 4 + Cl - (ammonium chloride)

2 Bonding General Concepts Chemical Bonds atoms combine, Electron distribution changes Types of Bonds: Ionic, Covalent and Metallic Ionic Bonds is the transfer of electrons (ex. salts, crystals) and is composed of a positive (cation) and negative (anion) Covalent is the sharing of electrons and the compound is either molecular (CH 4 ) or networking (diamond) Metallic Bond is the outer electrons move throughout solid and holds nuclei together and the compound is only composed of metals Conduction of Electricity Type of Bond Compound Conduct Explanation Electricity Ionic Solid NaCl No Na+Cl- ions held rigid Ionic Liquid NaCl Yes Ions free to move Covalent (molecules) Solid Cl 2 No Cl Cl all electrons tied up Covalent Liquid Cl 2 No Cl Cl all electrons tied up Metallic Solid Na Yes Outer electrons move throughout metal Metallic Liquid Na Yes Electrons free to move throughout More unpaired than would expect from electron configuration because of hybrid orbitals

3 Valence Electrons By sharing or transferring valence electrons the electronic configuration of noble gases frequently obtained from A family elements Element Inner electrons Outer electrons He 1s 2 Ne (He) 2s 2 2p 6 Ar (Ne) 3s 2 3p 6 Kr (Ar)3d 10 4s 2 4p 6 Xe (Kr)4d 10 5s 2 5p 6 Metal and Nonmetal Electrons Lost Electrons Gained Na Cl [Ne] 1s 1 [Ne] 3s 2 3p 5 Octet Rule Rules of Eight and Two works for IA, IIA, Nonmetals H, Li, Be H -, Li +, Be 2+ Isoelectronic with the same electronic structure Eight (Octet Rule) Na +, F -1, O -2 Isoelectronic with (Ne) Exceptions especially in d orbital Lose electrons in different order than build up Transition element loses electron from outer s subshell first Zn [Ar] 3d 10 4s 2 Zn 2+ [Ar] 3d e - Pseudonoble gas configuration ns 2 np 6 nd 10 Ionic Bond Electron dot structure Sodium metal explosive metal when added to water Chlorine gas is a poisonous gas Sodium chloride white crystal essential for life transfer e-

4 Bonds: Ionic ( Covalent Cl + Cl Cl 2 The Cl Cl bond shares electrons to complete the octet Element Valence e- s p p p Al 3s 2 3p 1 Na 3s 1 O 2s 2 2p 4 Cl 3s 2 3p 5 Show valence electrons The valence electrons are the ones beyond the noble gas or beyond noble gas and d electrons (The number at the top of the periodic table) Water H 2 O ( ( Lewis Structure is a powerful technique to predict the structure of water

5 In transfer the electrons lost must be equal to electrons gained Mg + F + F Mg F - Ionic Compound (cation) (anion) = lost 2 + 2(gain 1) All halogens form diatomic molecules (single bond) F F Cl Cl Br Br I I What about N 2 and O 2? Lewis dot structures of valence bond structures Number of valence electrons + group number for nonmetals Group Example # of Bonds 14 (4A or IV) C, Si 4 15 (5A or V) N, P 3 16 (6A or VI) O, S 2 17 (7A or VII) F, Cl, Br, I 1 18 (8A) 0 (normally) Once a particular element in the table bonds to their respective number then they have a noble gas like structure Notice the number of bonds directly correlates to how far they are from the noble gas on the periodic table ( 4 nonbonding electrons and 6 bonding electrons

6 ( 8 nonbonding electrons and 4 bonding electrons Bonds = (electrons need electrons have)/2 Electrons Need refers to the rule of eight and rule of two Electrons Have refer to valence electrons N 2 (16-10)/2 = 3 O 2 (16-12)/2 = 2 HF (10-8)/2 = 1 Covalent Bond and Lewis Structures Share electrons a pair of electrons ~ 1 bond H 2 simplest molecule ( Positive nuclei are attracted toward each other by negative charge in between them Each Hydrogen atom satisfied the noble gas structure Think of 2 electrons being around each Hydrogen

7 1s 2 structure for each H 2 discrete unit NaCl ionic compound, simplest ratio, not separate entity or species F [He] 2s 2 2p 5 ( Each F 2s 2 2p 6 12 nonbonding electrons 2 bonding electrons Family A elements (not transition elements) satisfy rule of 8 or rule of 2 (H) Covalent Examples of Covalent Bonds: ( ( ( (

8 Lewis Structural Formula 1916 Unshared electrons are referred to as lone pairs ( Whenever possible the Rule of Eight is followed Exceptions to Rule of Eight (Octet Rule) Examples: Boron trifluoride ( Phosphorus pentachloride (

9 Multiple Bonds Single Bonds Double Bonds Triple Bonds All can be used to satisfy the Rule of Eight Ethane, Ethene, Ethyne ( Ethane ( Bonds = (need have)/2 (( ) ( ))/2 (28-14)/2 = 14/2 = 7 (

10 Ethene ( Bonds = (need have)/2 (( ) ( ))/2 (24-12)/2 = 12/2 = 6 Chemist use lines when drawing Lewis structures One line equals two electrons shared between two atoms and these two electrons compose one bond Ethyne (Acetylene) ( index/multiplebonds/multiplebonds.htm) Bonds = (need have)/2 (( ) ( ))/2 (20 10)/2 = 10/2 = 5 Resonance Hybrid Lewis Structure for SO 2 (sulfur dioxide) ( These are resonance forms with equivalent structures just the bonds are in different places This molecule is a resonance hybrid Really doesn t mean that the structure flips back and forth rather there is a hybrid (intermediate) that includes both of theses structures The hybrid has two 1 ½ bonds in which 2 e- are spread between 3 atoms Bonds = (need have)/2 (24 18)/2 = 6/2 = 3 bonds Normal Lewis Structures don t always work

11 Limitation of the theory We can have fractional bond order Lewis Structure for N 2 H 4 (hydrazine) ( Bonds = (need have)/2 (( ) ( ))/2 (24 14)/2 = 10/2 = 5 bonds Coordinate Covalent Bond both electrons donated by one atom Lewis Structure of HNO 2 (hydrogen nitrite) O=N N H To form HNO 3 (hydrogen nitrate) in solution with nitric acid ( Bonds = (need have)/2 (34-24)/2 = 10/2 = 5 bonds Formal Charge = + (group #) ( # of bonds + # unshared electrons) +5 (4) = +1 for N +6 (7) = -1 for O with single bond +6 (6) = 0 for O with double bond In solution break the O H bond

12 ( (Note: negative adds an electron and positive removes an electron) Bonds = (need have)/2 (32 24)/2 = 8/2 = 4 bonds Formal Approach to Drawing Lewis Structures 1. Total number of valence electrons positive ion decrease by charge negative ion increase by charge 2. Number of needed electrons to give each atom 8 e- and each Hydrogen 2 e- 3. Number of bonding e- + ( needed e-) (valence e-) 4. Bonds = Bonding e-/ 2 5. Indicate position of atoms 6. Draw dash for each bond 7. Complete octet (other than H) for each atom filling in unshared e- # unshared e- = (total e-) ( # bonding e-) Example: N 2 Formal charge = +(group valence #) ( # bonds + # unshared electrons) Formal charge is a way of saying if an atom has fewer or more electrons than it needs to satisfy octet rule. In N 2 below each N has 0 (zero) formal charge because each has 5 valence electrons to start with and ( 3 bonds + 2 electrons in lone pair= 5 ) so 5 5 = 0 zero formal charge 1. Valence given = Needed + 16 ( 2 x8) 3. need 6 bonding electrons 4. or 3 bonds (6/2) 5. so N 2 has triple bond 6.

13 ( 7. ( # unshared e- = (total e-) ( # bonding e-) 10 6 = 4 unshared e- Measure of Charge Separation in Molecules (dipole moment) Dipole moment = (charge)(distance) HCl observed dipole moment Polar molecules affect charge that pair of electrically charged plates can hold ( HCl = 1.03 D debye (Coloumb)(m) ~ 3.34 x C x m If charge completely transferred = (1.60 x C)( 1.27 x m )(D/3.34 x C x m) = 6.08 D

14 Observed/ predicted by transfer = 1.03/6.08 = 17% ionic Left (metal) with right (nonmetal) is a ionic bond Nonmetal with Nonmetal is polar covalent Nonmetal with Nonmetal can also be pure covalent NaCl ionic MgCl 2 ionic PCl 3 covalent