9 1 Chapter 9 Chemical Bonding I Lewis Theory Introduction to Ionic and Covalent Bonding 1 Three types of chemical bonding 1 Ionic bonding 2 Covalent bonding 3 Metallic Bonding 2 Na 3 Lewis electron dot symbols Mg S Cl C N O The Octet Rule Ionic and Covalent Bonding A Continuum Ionic Polar Covalent Nonpolar Covalent F 9 2 Ionic Bonding and Lattice Energy Consider a binary ionic bond between a metal atom and a non metal atom There are three things to consider in terms of energy 1 Metals have low 2 Non metals have high 3 Energy is when gaseous ions form the solid crystal Lattice energy How does Lattice Energy relate to some of the physical properties of ionic substances Important Bond breaking is always and bond formation is always So you can think of lattice energy this way Coulomb s law Compare these lattice energies NaCl MgCl2 KI CaI2 9 3 To predict these apply Coulomb s law 1 In general the the ionic charge the the lattice energy 2 In general the the ion the the lattice energy And remember is usually more important than Example Problems 1 Arrange the following ionic compounds in order of increasing lattice energy that is from least exothermic to most exothermic NaF CsI CaO 2 Which compound has the greatest i e the most exothermic lattice energy AgCl 3 Based upon a consideration of lattice energy which compound is expected to have the higher melting point CuO CrN NaCl or NaBr Covalent Bonding 1 A covalent bond is formed when 2 Covalent bonds are typically seen between 3 Two measures of the strength of a bond 4 The the bond the the bond 5 Example 6 Triple bonds are and than double bonds and C C C C C C double bonds are and than single bonds 7 Bond breaking is always bond formation is always 9 4 Lewis Structures Steps for drawing Lewis Structures 1 Sum the from all atoms take into account any positive or negative charge 2 Determine which atoms Usually the is listed first In some cases the formula describes the order in which the atoms are bonded such as In oxyacids the hydrogens are attached to and the oxygen atoms are attached to HNO2 H2SO3 The less electronegative atom is usually the 3 Distribute all of the valence electrons in this order First Second 4 Third Check to see if the central atom has an If not use There are exceptions AlCl3 H2SO4 5 Assign 6 Draw important Examples Draw the correct Lewis structures for the following molecules or ions PCl5 SF4 H2SO4 SO2 XeF4 9 5 Exceptions to the Octet Rule 1 Three atoms routinely have 2 Molecules with an odd number of electrons 3 3 row elements and below can have rd Formal Charge Formal charge on an atom is defined as the number of valence electrons on the atom minus the number of electrons assigned to the atom The number assigned corresponds to all the unshared electrons plus one half of the shared electrons Shortcut Formal Charge Example NO3 Resonance O3 1 Resonance is used when 2 To indicate resonance you must use 3 Resonance is not referring to or 4 Resonance means that the true molecule is a of the resonance structures 5 Be sure to distinguish between these terms vs 9 6 6 Resonance structures are said to 7 Resonance structures are not necessarily Some resonance structures 8 To evaluate resonance structures consider Three Points 1 The the better 2 It is usually better to have the negative formal charge on the 3 Avoid on Example 9 In general the more resonance structures that can be formed the more This is due to 10 When drawing resonance structures move only Never move Example Draw the resonance structures for N2O nitrogen is the central atom and evaluate which contributes most and which contributes least to the resonance hybrid Example Draw resonance structures for the sulfate ion and evaluate which contributes most to the resonance hybrid Electronegativity Of the bonds shown below which is the most polar C C C N C O C F 9 7 Question Why is it best to draw the Lewis Structures for molecular BF3 AlCl3 and BeCl2 with less than an octet Drawing Lewis structures for Organic Molecules A Shortcut Second period elements never never never never have Does this mean that nitrogen always has three bonds HCHO C2H6 C2H4 C2H6O CH3CH2CH3 Bond Energy Bond Energy Of the bonds listed below which is expected to have the highest bond energy Why C C C C C O C O C2H2 CH3CHO 9 8 Using Bond Energies to estimate H rxn Three keys to doing this kind of problem 1 Rewrite the balanced equation with of each reactant and product shown 2 Identify which bonds are and which bonds are 3 Remember bond breaking is always and bond formation is always Bond Energies kJ mol H H 436 H Cl 431 C H 414 C C 347 C O 360 C O in CO2 799 C Cl 339 O O in O2 498 O H 464 Cl Cl 243 Examples Using the bond dissociation energies listed in the table above calculate the approximate H in kJ for the following reaction CH4 g 2Cl2 g CH2Cl2 g 2HCl g H Determine the molar heat of combustion of methane CH4 in kJ mol The standard enthalpy of formation of H2O2 g is 136 kJ mol Use this information along with the bond energies in the table above to estimate the bond energy of an O O bond