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Chemical Bonding Unit 5 Introduction Chemical bonding refers to the attractive forces that hold atoms together in a compound Two major classes of bonding: Ionic Results from electrostatic interactions among ions Involves transfer of electrons Takes place b/w metals & non-metals Covalent Results from sharing one or more electron pairs b/w atoms Takes place b/w non-metals Ionic Compounds 1. 2. 3. 4. 5. They are solids with high m.p. (typically > 4000C) Many are soluble in polar solvents e.g. water Most are insoluble in non-polar solvents Molten compounds conduct electricity well because they contain mobile ions Aqueous solutions conduct electricity well because they contain mobile ions Covalent Compounds 1. 2. 3. 4. 5. They are gases, liquids or solids with low m.p. (typically < 3000C) Many are insoluble in polar substances Most are soluble in non-polar solvents e.g. hexane, C6H12 & carbon tetrachloride (CCl4) Liquid and molten compounds do not conduct electricity Aqueous solutions are usually poor conductors of electricity because most do not contain charged particles Fig. 7-CO, p. 250 Lewis Dot Formulas of Atoms The number and arrangement of valence electrons determine: Physical properties Chemical properties Chemical bonding Lewis dot formulas or Lewis structures show these chemically important valence electrons Shows only electrons in outermost s and p orbitals as dots Electron pairs are represented as a pair of dots & an unpaired paired electron as a single dot Not useful for representing transition and inner transition metals Table 7-1, p. 252 Formation of Ionic Compounds Ion atom or group of atoms that carries an electrical charge Cation- positively charged; more protons than electrons Anion- negatively charged: more electrons than protons Monoatomic ions – consists of only one atoms e.g. Cl-, Mg2+ Polyatomic ions - consists of a group of covalently bonded atoms e.g. NH4+, SO42- Formation of Ionic Compounds Ionic bonding can occur easily when elements that low EN and low IE (i.e. metals) react with elements that have high EN and very negative electron affinities (i.e. non-metals) Many metals are easily oxidized i.e. they lose electrons to form cations Many non-metals are easily reduced i.e. they gain electrons to form anions The farther apart across the periodic table two Group A elements are, the more ionic their bonding will be Formation of Ionic Compounds Lewis dot formula for ionic compounds e.g. Na+ is isoelectric with Ne In contrast, Cl- is isoelectric with Ar p. 254 Formation of Ionic Compounds General representation of the reaction of 1A metals with 7A elements (halogens): p. 255 Formation of Ionic Compounds Like other simple ionic compounds, NaCl exists in a regular, extended array of +ve and –ve ions. Distinct molecules of solid ionic substances do not exist, therefore referred to as formula units Fig. 7-1, p. 255 Formation of Ionic Compounds Reaction of 1A metals with 6A elements e.g. very small in size of Li+ gives it a much higher charge density (ratio of charge to size) than that of Na+ Similarly, the O2- is smaller than Cl- because of its smaller size and double charge The more concentrated charges & smaller sizes bring the Li+ and O2- closer together in Li2O than the ions in NaCl stronger ionic bond This is consistent with higher mp of Li2O (>1700oC) than NaCl (801oC) Formation of Ionic Compounds Reaction of 2A metals with 6A elements e.g. Ca2+ is about the same size as Na+ but carries twice the charge, so its charge density is higher Attraction b/w two small, highly charged ions is high very strong ionic bond M.p. of CaO is 2580oC p. 256 † group 1A and 2A can also form peroxides and superoxides Table 7-2, p. 257 Energy relationships in Ionic Compounds Why does ionic bonding occur? Why is solid NaCl more stable than a mixture of individual Na and Cl atoms? Consider a gaseous mixture of Na and Cl atoms: Step 1: 1st IE of Na atoms is a positive value less stable than original mixture of atoms Step 2: energy change for the gain of 1 mole e-s by one mole of Cl atoms is given by the electron affinity of Cl This –ve value lowers the energy of the mixture, but the mixture of separated ions Na+ and Cl- ions is still higher in energy ( less stable) than original mixture of atoms Fig. 7-2, p. 258 Energy relationships in Ionic Compounds Thus, the formation of ions does not explain why the process occurs The strong attractive forces b/w ions of opposite charges draw the ions together in a regular array The energy associated with this attraction (step 3) is the crystal lattice energy This further lowers the energy to (147789)kJ/mol = -642kJ/mol Formation of Ionic Compounds d-transition metals Have s electrons in outermost shell and one d electrons one energy level lower (e.g. 3d4s in 4th period transition elements) Outer s electrons are always lost before d electrons d- and f-transition elements form compounds that are essentially ionic 3d 21Sc Ar 21 Sc Ar 4p 4s Configurat ion Ar 4s 2 3d1 3d 3 4s 4p Configurat ion Ar p. 256 Formation of Ionic Compounds d-transition metals 3d 4s 4p 30 Zn Ar 3d 4s Configurat ion Ar 4s 2 3d10 4p 2 Ar Zn 30 Many transition metals ions are highly coloured Configurat ion Ar 4s0 3d10 Covalent Bonding Covalent bonding occurs when the electronegativity difference b/w elements (atoms) is zero or relatively small. In covalent compounds the bonds b/w atoms within a molecule (intramolecular bonding) are relatively strong BUT the attractive forces between molecules (intermolecular forces) are relatively weak. Hence covalent compounds have lower mp and bp than ionic compounds Covalent Bonding e.g. b/w hydrogen atoms a) Two H atoms are separated by a large distance b) As the atoms approach, the e- cloud of each atom is attracted by the +vely charged nucleus of the other atom (blue arrow). At the same time the e- clouds repel one another, as do the two nuclei c) The 2 e-s can both be in the region where the two 1s orbitals overlap; the e- density is highest b/w the nuclei of the two atoms Fig. 7-3, p. 259 The potential energy of the H2 molecule as a function of the distance b/w the to nuclei • The bonded atoms are lower in energy (more stable) than the separate atoms • The result of sharing is that each atoms gains an electron configuration of the nearest noble gas Fig. 7-4, p. 259 Polar and Non-Polar Covalent Bonds Nonpolar bond- the electrons are shared equally b/w atoms e.g. H2 or H- H Both H atoms have the same electronegativity the shared e-s are equally attracted to both H nuclei and spend equal amts of time near each nucleus The covalent bond in all homonuclear diatomic molecules must be nonpolar Polar and Non-Polar Covalent Bonds Polar bonds – the e- pairs are shared unequally e.g. H-F There is a large difference in electronegativity b/w H (2.1) and F (4.0) e-s spend more time close to the F nucleus p. 276 The separation of charge in a polar covalent bond creates an electric dipole p. 276 Polar and Non-Polar Covalent Bonds Each halogen can form a single bond to another halogen to form an interhalogen Bond polarities decreases as the electronegativity differences b/w atoms decreases: p. 276 Dipole Moments The polarity of a molecule is indicated by its dipole moment, which is given by: μ = d x q where d = distance separating opposite charges of equal magnitude q = magnitude of charge Table 7-5, p. 277 Dipole Moments Measured by placing a sample of the molecule in an electric field Polar molecules e.g. H-F tend to line up slightly in a direction opposite to the field. Non-polar molecules are not oriented in by an electric field Fig. 7-5, p. 278 Dipole Moments NOTE: Dipole moments of individual bonds can only be measured for diatomic molecules Dipole moments reflect overall polarities of molecules For polyatomic molecules overall dipole is affected by molecular geometry and the presence of lone pairs of e-s. Bond lengths & Bond Energies The internuclear distance at which the attractive & repulsive forces balance and the bond is most stable is called the bond length Bond dissociation energy is the energy needed to separate the atoms, breaking the covalent bond p. 260 Comparison of Carbon-carbon bond lengths and energies Two nuclei are more strongly attracted to 2 shared e- pairs than to 1 pair Atoms are pulled closer together & more difficult to pull apart Multiple bonds are shorter & are stronger than single bonds. p. 260 The greater the difference in electronegativity between bonded atoms, the stronger the polar bond, and the greater the "extra" bond energy. Table 7-3, p. 261 Trends Stronger bonds tend to be shorter Bonds b/w H and 2nd row elements are very strong Bonds become longer & weaker as atomic number increases e.g. C-halogen bonds Bonds b/w C and 2nd row elements are reasonably strong N-O & O-O bonds weaker than C-C bonds because of repulsion b/w unshared e-s on N and O Lewis Formulas A covalent bond is represented by writing each shared electron pair as either: A pair of two dots b/w atoms As a dash connecting atoms OR p. 261 Lewis Formulas Showing double bonds… p. 262 Lewis Formulas Polyatomic ions may also be represented in this way For the ammonium ion, only 8 valence e-s are shown even though N has 5 valence e-s and each H has 1 (a total of 5 +4 = 9) The charge of +1 implies the species has one less e- than the original atoms • Lewis formula is an e- bookkeeping method that is useful as a 1st guide to bonding schemes • They only show: • # of valence e-s • The # and kinds of bonds • Order in which atoms are connected • They are not intended to show 3D shape of the molecule p. 262 Guide to Writing Lewis Formulas See Handout Resonance A molecule or polyatomic ion for which two or more Lewis formulas with the same arrangements of atoms can be drawn is said to exhibit resonance. Resonance does not mean: The molecules changes from one structure to the next Experiments show that the C-O bond is neither a double nor a single bond but has intermediate bond length and strength. p. 274 Resonance Another way to represent this situation is by delocalization of bonding electrons: The dashed lines indicate that some of the e-s shared b/w C and O are delocalized (spread across) all four atoms p. 275 Resonance Example: Draw two resonance structures for SO2. Solution: S O N = 1(8) + 2(8) = 24eA = 1(6) + 2(6) = 18eS=N–A = 6e- shared The resonance structures are: Or showing the delocalization of the electrons as follows: p. 275 In Summary… Chemical bonding can be described as a continuum that may be represented as: ∆E for the bonding atoms Bonding type zero non-polar covalent intermediate large polar covalent ionic p. 279