Announcements. Chem 7 Final Exam Wednesday, Oct 10 1:30-3:30AM Chapter or 75 multiple choice questions

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1 Exam III (Chapter 7-0) Wednesday, ctober 3, 202 Time 600PM - 730PM SEC A 24A and 25A SKIPPING THIS STUFF Announcements Chem 7 Final Exam Wednesday, ct AM Chapter or 75 multiple choice questions Classical distinction between energy and matter (p 27) 2 Numerical problems involving the Rydberg equation (equations 73/4) 3 Spectral analysis in the laboratory (pp ) 4 Numerical problems involving the Heisenberg uncertainty principle (p 23) 5 Trends among the transition elements (p 26) 6 Trends in electron affinity (pp ) 7 Pseudo-noble gas configuration (p 269) 8 Lattice energy (pp ) 9 IR spectroscopy (p 292) 0 Numerical problems involving electronegativity (p 296) Electronegativity with oxidation number (p 297) 2 Section 3 M theory and electron delocalization 3 All sections in chapter 2 except 23 (types of intermolecular forces) An ionic compound (salt) results from the chemical reaction between a metal and non-metal Cadmium Metals compounds result from chemical reactions between non-metals Non-Metals Bromine Ionic and covalent compounds have different physical properties Ionic Compounds Solid high melting and boiling points (>600 C) 2 Soluble in polar solvents (like water) 3 Insoluble in non-polar solvents (like organic solvents) 4 Molten compounds conduct electricity very well 5 Aqueous solutions are electrolytes 6 Formed between elements with very different electronegativity Compounds Gas, liquid or solid with low melting points (<300 C) 2 Insoluble in polar solvents 3 Soluble in organic solvents (not water soluble) 4 Non-electrolytes as liquids 5 No charged particles and (electrolytes) in water 6 Formed between elements of similar electronegativity Valence electrons are the electrons in the outer shell (highest value of n quantum number) of the atom that dictate chemistry & in chemical bonding Group e - configuration # of valence e - A ns 2A ns 2 2 3A ns 2 np 3 4A ns 2 np 2 4 5A ns 2 np 3 5 6A ns 2 np 4 6 7A ns 2 np 5 7 Given by the Group Number for Group A For B group elements, the valence electrons are in the highest value ns orbital and the (n-)d orbitals Lewis dot structures are used to depict valence electrons and bonding between atoms Lewis structures are used to depict how ionic bonds are formed by chemists Formation of sodium chloride Mg Cl Cl noble gas electronic configuration Mg Cl Formation of magnesium chloride A chemical symbol represents the nucleus and all core e - A single dot around the symbol represents one valence e - Ba + Ba Formation of barium oxide

2 Chemists use different ways to conceptualize and draw ionic bonding Know both! Using spdf electron configurations Li (s 2 2s ) + F (s 2 2s 2 2p 5 ) Li + s 2 + F - (s 2 2s 2 2p 6 ) The for main group or representative elements the group number in the periodic table can tell us the charge of the metal cation and the charge of the non-metal anion that will be formed in an ionic reaction Non-metals Charge of Anion Metals Charge of Cation = = Group Number - 8 Group Number (A, 2A, 3A) NN-METALS 2 rbital box diagrams Li s 2s 2p + F s 2s 2p Li + s 2s 2p + F - s 2s 2p Transition Metals starts with filling of d-orbitals Ni 2+ 3 Lewis electron-dot symbols Li + F Li + + Ḟ Use partial orbital diagrams and Lewis symbols to depict the formation of + and 2- ions from the atoms, and determine the formula of the compound bonding results from sharing of one or more electrons between non-metals atom SLUTIN 3s 3p 2-2s 2p ccurs generally between non-metal elements forming molecules or molecular compounds share electrons 3s 3p 2s 2p The formation of a covalent bond releases energy and reduces the potential energy (ie minimizes energy) Atoms to Close Together Atoms Far Apart We also use Lewis structures and the rule to show covalent bonds and bonding between atoms F + F F F 7e - 7e - single covalent bond 8e - 8e - 8 e - or an around each atom Every covalent bond has a characteristic bond length that leads to maximum stability Released When Bond Formed lone pairs F F F F single covalent bond lone pairs Lewis Dot Structure Lewis Skeletal Structure Note the outer electrons are not shown!

3 What is the Lewis dot structure and dash formula for water, H2 and for carbon dioxide? Bond order is the number of covalent bonds between two atoms H H H H H H C C Lewis Dot Dash Formula What is the Lewis dot structure and the dash formula for carbon dioxide? C Bonding Pairs (double bond) C Non-bonding Pairs Bond rder is the number of bonds between two atoms Single bond, order = Double bond, order = 2 Triple bond, order = 3 H H Higher bond orders give rise to Shorter bond lengths between bonded atoms Stronger bond energies (more energy needed to break) C Higher bond orders give shorter bond lengths and require more energy to break a bond Comparing Bond Length and Bond Strength The trends in atomic can be extended to the bond length in molecules in a family of compounds Using the periodic table, but not Tables 92 and 93, rank the bonds in each set in order of increasing bond length and bond strength (a) S - F, S - Br, S - Cl (b) C =, C -, C Bond Lengths Triple bond < Double Bond < Single Bond (c) C-Br, C-I, C-Cl Comparing Bond Length and Bond Strength Using the periodic table, but not Tables 92 and 93, rank the bonds in each set in order of decreasing bond length and bond strength Bond length depends on the size of the bonded atoms PLAN (a) S - F, S - Br, S - Cl (b) C =, C -, C (a) The bond order is one for all and sulfur is bonded to halogens; bond length should increase and bond strength should decrease with increasing atomic (b) The same two atoms are bonded but the bond order changes; bond length decreases as bond order increases while bond strength increases as bond order increases (a) Atomic size increases going down a group Bond length S - Br > S - Cl > S - F (b) Using bond orders we get Bond length C - > C = > C F2 Cl2 72 pm 72pm 00 pm 00pm 4 pm Br2 4pm 33 pm I2 33pm Bond strength S - F > S - Cl > S - Br Bond strength C > C = > C -

4 VSEPRT explains the geometry of molecules but NT how covalent bonds are formed with that geometry Write the Lewis dot and skeletals structure of nitrogen trifluoride (NF 3 ) Molecular formula Lewis structure VSEPRT Geometry Hybrid orbitals Write the Lewis dot and skeletal structures of the carbonate ion (C 3 2- ) Write the Lewis dot and skeletal structures structure of the carbonate ion (Br 3- ) Lewis Structure VSEPRT Valence Bond Theory Write the Lewis dot and skeletal structures structure of the carbonate ion HCN? Write the Lewis structure of nitrogen trifluoride (NF 3 ) Write the Lewis structure of the carbonate ion (Br 3- ) Step N is less electronegative than F --> N is central atom! Step 2 - Count valence electrons = A; Nitrogen = 5, Fluorine = 3 X 7 = 2 A = = 26 valence electrons Step 3 - Write structure with N central and three bonds and rest nonbonding electrons around the central atom Look at the formula sometimes it gives clues to the central atom [Br3] Br Valence e- = 7 + 3(6) + = 26 Step 4 - Write structure with N central and three bonds and rest nonbonding electrons F N F HCN Valence e- = = 0 Carbon is central atom, watch for hydrogen-- bond F H C N Bond Enthalpies and Reactions Hess s Law allows us to analyze the energetics of a reaction This provides insight into the relative stability of products and reactants The bond energy, BE is the positive amount of energy (enthalpy) required to break or make one mole of bonds in a gaseous covalent compound to form products in the gaseous state at constant T and P H 2 (g) Example H (g) + H (g) Bond Dissociation!H 0 = kj/mol Cl 2 (g) Cl (g) + Cl (g)!h 0 = kj/mol Endothermic Exothermic-energy released HCl (g) H (g) + Cl (g)!h 0 = +439 kj/mol

5 All chemical reactions give off a net heat of reaction that can be either (!H rxn ) absorbed or gained The source of heat is the breaking and making of chemical bonds! H 2 (g) + F 2 (g) 2HF(g)!H rxn =? H rxn = ΣBE reactant ΣBE product H H F F H F !H rxn = ( ( (2 x 5682) = kj Calculating Bond Enthalpies In A Reaction Calculate the!hrxn for the chlorination of methane to form chloromethane gas using bond energies CH 4 (g) + Cl 2 (g) " CH 3 Cl(g) + HCl(g)!H rxn =? Write the balanced chemical equation and set up table below 2 Write Lewis structures break all bonds in products and reform all bonds in products 3 Use the simple equation below to calculate!h rxn H rxn = ΣBE reactant ΣBE product Calculate the!hrxn for the chlorination of methane to form chloromethane gas using bond energies CH 4 (g) + Cl 2 (g) " CH 3 Cl(g) + HCl(g)!H rxn =? + + H rxn = ΣBE reactant ΣBE product C-H Cl-Cl C-H C-Cl H-Cl 43 43!H rxn = [656kJ/mol kj/mol] - [232kJ/mol kj/mol + 43 kj/mol]!h rxn = -4 kj/mol (exothermic) Calculate the standard enthalpy change H in kj for the hydrogenation of ethyne (acetylene) to ethane + 2H 2 (g) H 3 C CH 3 (g) Calculate (in kj) the standard enthalpy change H for the hydrogenation of ethyne (acetylene) to ethane + 2H 2 (g) H 3 C CH 3 (g) H rxn = ΣBE reactant ΣBE product C-H H-H = -296 kj/mol C-C C-H !H rxn = [826kJ + 839kJ kj] - [347 kj/mol kj/mol]

6 Calculating Enthalpy Changes from Bond Energies Use Table 92 to calculate!h 0 rxn for the following reaction CH 4 (g) + 3Cl 2 (g) CHCl 3 (g) + 3HCl(g) Write the balanced chemical equation 2 Write Lewis structures to see and 3 Use the given Bond energies in the problem and a table account for /formed to calculate!h rxn Calculating Enthalpy Changes from Bond Energies Use Table 92 to calculate!h 0 rxn for the following reaction CH 4 (g) + 3Cl 2 (g) CHCl 3 (g) + 3HCl(g) Calculate the number of and bond formed using the bond energies found in Table 92 H rxn = ΣBE reactant ΣBE product 4C-H = 4 mol(43 kj/mol) = 652 kj 3Cl-Cl = 3 mol(243 kj/mol) = 729 kj #!H o = 238 kj 3 C-Cl = 3 mol(-339 kj/mol) = -07 kj C-H = mol(-43 kj/mol) = -43 kj 3H-Cl = 3 mol(-427 kj/mol) = -28 kj #!H 0 = -27 kj!h o reaction = #!H o + #!H o = 238 kj + (-27 kj) Electronegativity is an element s inherent ability to draw electrons to itself when chemically bonded to another atom in a molecule (relative to Li) Notice F,, N, Cl, Br, C are highly electronegative with F being the most electronegative Group I and II metals are the least electronegative Differences in electronegativity between bonding atoms blur the distinction between covalent, polar covalent and ionic bonding types Bonding Polar Bonding Ionic Bonding!+!- F2 HF LiF The delta s are used to show a partial charge on atoms The arrow is used to denote a polar covalent bond Arrow points negative F F H F Li F For a polyatomic molecule we must consider the vector sum of polar bonds in the molecule to see if there is a net dipole moment polar Electronegativity values can be used to judge the extent of ionic or covalent bond character in a chemical bond absolute value of electronegativity difference between 2-bond atoms Mostly Ionic 30 No Net χ = EN X - EN M Polar! 20 No Net Mostly 00

7 How ionic or covalent a bond is between two atoms can be judged by looking at the the absolute value of the difference in electronegativity χ = EN X - EN M Classify the following bonds as ionic, polar covalent, or covalent The bond in CsCl; the bond in H 2 S; and the NN bond in H 2 NNH 2 Cs 07 Cl = 23 Ionic H 2 S = 04 Polar N 30 N = 0 Ionic compounds are typically solids of high melting and boiling points Across a given period compounds become more covalent with lower melting points than their ionic counterparts Melting point of compounds increase across a period as compounds become less ionic and more covalent Across Period 3 chloride compounds go from high melting solids to very low boiling point gases (as it gets more covalent)