Ionic vs. Covalent Compounds UNIT 3: BONDING. Covalent or Ionic? Ionic Compounds. Ionic Bonding NaCl example. Ionic Nomenclature 9/30/2011
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1 Ionic vs. Covalent Compounds Ionic compounds contain a metal formula units Covalent compounds only non-metals molecules UNIT 3: BONDING Covalent & Ionic Covalent or Ionic? H 2 O covalent NaCl ionic HgSO 4 ionic PF 4 covalent Ionic Compounds EN difference is high (> 2.1) electrons are transferred NOT shared cation atom (or group of atoms) that loses electrons anion atom (or group of atoms) that gains electrons #electrons lost = # electrons gained charges for entire compound must add up to zero Ionic Nomenclature monoatomic cations name of element Na + sodium ion monoatomic anions root of element name + ide ending Cl - chloride ion polyatomic ions see chart check memory work! compound name = cation name + anion name NO PREFIXES Ionic Bonding NaCl example 495 kj is required to remove one mole of electrons from one mole of Na (ionization energy) creates Na + endothermic process 349 kj is released when adding one mole of electrons to one mole of Cl (electron affinity) creates Cl - exothermic process This would be an overall endothermic process, but 1
2 Ionic Bonding NaCl example Naming Acids Lattice energy the amount of energy required to separate ionic compounds energy is required to separate the ions energy is released when ions get together Net energy = Ionization Energy Electron Affinity Lattice Energy Net energy change is negative, more stable Watch ending of anion -ide becomes hydro-root-ic acid HCl hydrochloric acid -ate becomes root-ic acid HNO 3 nitric acid -ite becomes root-ous acid HNO 2 nitrous acid Naming Hydrates Valence Electrons salt name prefix-hydrate CuClO 3 5 H 2 O copper chlorate pentahydrate outer shell electrons involved in bonding Remember A group numbers 1 mono 2 di 3 tri 4 tetra 5 penta 6 hexa 7 hepta 8 octa 9 nona 10 - deca Covalent Nomenclature Ionic Bonds transfer electrons one or more electrons leave on atom and join another a cation (positive) and an anion (negative) form cation sticks to anion like a magnet Covalent Bonds share electrons neither atom loses or gains electrons MEMORIZE THIS! 1 mono 2 di 3 tri 4 tetra 5 penta 6 hexa 7 hepta 8 octa 9 nona 10 - deca 2
3 Naming First element: prefix (except mono) + name of element Second element: prefix + root of element name + ide ending Name the following SeI 4 selenium tetraiodide N 2 O 3 dinitrogen trioxide CSe 2 carbon diselenide P 2 Cl 5 diphosphorus pentachloride Write formulas for the following phophorus pentabromide PBr 5 disulfur trioxide S 2 O 3 aresenic trifluoride AsF 3 dichlorine monoxide Cl 2 O Diatomic Elements elements not found alone in nature memorize these! (think 7) H 2 N 2 O 2 F 2 Cl 2 Br 2 I 2 Lewis Dot Structures Drawings that show how electrons are shared between atoms in covalent bonds Use only valence electrons Lewis Structures for Atoms Na C B S I Ne Special H He 3
4 Octet Rule Covalent Bonds All atoms want 8 valence electrons Atoms will bond to satisfy this rule Exceptions! H and He only want two electrons! B is satisfied with 6 electrons Can have expanded octets Single bond one pair of shared electrons Double bond two pair of shared electrons Triple bond three pair of shared electrons Drawing Lewis Structure for Molecules 1. Find the total number of valence electrons 2. Determine the central atom 3. Place single bonds between atoms 4. Place lone pairs 5. Check octet rule & total number of electrons 6. If needed add double or triple bonds one at a time 7. The best arrangement has the lowest formal charge or the negative formal charge on the most EN element Determining Formal Charge C f = E v (E u + ½ E b ) C f = formal charge E v = valence electrons E u = unshared electrons E b = bonding electrons Example Practice Drawing Lewis Structures CH 3 OH vs CH 2 OH 2 CH 4 NH 3 H 2 O CSe 2 CO HCN 4
5 Polyatomic Ions Add or subtract the appropriate number of electrons Resonance When more than one Lewis Structure is possible NH 4 + SO 4-2 O 3 NO 3 - CO 3-2 Free Radical Molecular Geometries Sometimes there is an odd number of electrons NO describe the 3-D shape of a molecule VSEPR valence shell electron pair repulsion Draw Lewis Structure Count bonds (doesn t matter what type) Count lone pairs (unshared pairs of electrons) Molecular Geometries Molecular Geometries for Expanded Octets Geometry Bonds on central atom Lone Pairs on central atom Angle Hybridization Linear sp Trigonal sp2 Planar Bent 2 1 Tetrahedral sp 3 Trigonal 3 1 Pyramidal Bent 2 2 Geometry Bonds on central atom Lone Pairs on central atom Angle Hybridization Trigonal , 90 sp 3 d Bipyramidal Seesaw 4 1 T-shaped 3 2 Linear 2 3 Octahedral sp 3 d 2 Square Pyramidal Square Planar
6 Drawing 3-D Lewis Structures Bond Angles (solid line) in the plane of the page (dotted line) away from the reader (wedge) toward reader Depend on the amount of repulsion Lone pairs > triple bond > double > single CO 2 SO 2 CH 2 O CH 4 NH 3 H 2 O Polarity of Molecules Electronegativity Molecules are polar if they are asymmetrical Look for Lone pairs on the central atom Differing atoms attached to the central atom the ability of an atom to attract electrons involved in bonding with another atom values predict the type of bonding that will occur Polarity of Bonds What type of bond is formed? EN Difference < > 2.1 Type of Bond non-polar covalent (even sharing) polar covalent (uneven sharing) ionic (electrons transfer) KCl = 2.4 ionic bond NH = 0.8 polar covalent bond PS = 0.4 non-polar covalent 6
7 Dipole Draw Lewis Structure & Indicate Net Dipole A polar bond Strength is called the dipole moment Ex. H 2 O BrF 5 ICl 4 - ClF 3 AsF 5 NH 3 A molecule can have polar bonds but be symmetrical (non-polar) overall (CCl 4 ) Valence Bond Theory Sigma ( ) Bond Combines Lewis structures and atomic orbital theory Orbitals from two different atoms overlap Pi ( ) Bond Overlapping two p orbitals single bond = bond double bond = + bond Triple bond = 2 bond + bond 7
8 Hybrid Orbitals sp hybridization Created from existing orbitals forming a new sublevel E.g. HCl H = 1s 1 Cl = [Ne]3s 2 3p 5 Unfilled s and unfilled p orbitals overlap forming 2 sp orbitals ( bond) sp 2 hybridization sp 3 hybridization (no lone pairs) sp 3 hybridization (one lone pair) sp 3 d hybridization 8
9 sp 3 d 2 hybridization Hybrid Orbitals determine Geometry Hybrid Orbital sp sp2 sp 3 sp 3 d sp 3 d 2 Geometry linear trigonal planar tetrahedral trigonal bipyramidal octahedral Intermolecular Forces attractions between molecules that hold them together forming a liquid or a solid ion-dipole: occurs between an ionic compound and polar molecules H-bonding (not real bonding) occurs between the H ( +) on one molecule and the lone pairs on the F, O, or N of another molecule dipole-dipole: occurs between two polar molecules London forces: occur because of a disturbance creating temporary dipoles (occur in ALL molecules) Strength of IM forces increase as polarizability increases left to right on previous chart increases with increasing molecular weight larger atoms have larger electron clouds and are easier to polarize Intermolecular Forces affect Properties Melting and boiling points as IM forces increase, MP & BP increase viscosity resistance to flow as IM forces increase, viscosity increases as temperature increases, viscosity decreases surface tension a measure of the inward forces that must be overcome to expand the surface area as IM forces increase 9
10 List the following in order of increasing boiling points. barium chloride neon hydrogen carbon monoxide hydrogen fluoride H 2 <Ne<CO<HF<BaCl 2 BaCl 2 ionic highest Ne, CO, HF all have similar molecular weights HF hydrogen bonding CO dipole-dipole Ne nonpolar H 2 non-polar & lowest molecular weight lowest Solid Liquid fixed volume and shape incompressible does not flow diffusion within a solid is very, very slow molecules move but do not leave their location shape of container, fixed volume non-compressible fluid other substance diffuse slowly molecules move around each other Gas State depends on energy & IM forces shape & volume of container compressible fluid other substance diffuse quickly within a gas molecules move in roughly a straight line until they run into something 10
11 Phase Changes Heat = Enthalpy = Energy H fus = heat (enthalpy) of fusion the amount of energy required to change a solid into a liquid at its melting point H vap = heat (enthalpy) of vaporization the amount of energy required to change a liquid into a gas at its boiling point H vap > H fus H sub = H vap + H fus Heating Curve Notice: energy can be used to raise temperature or change state, but not both at the same time Calculating H for Temperature/Phase Changes Calculate the enthalpy change upon converting 1.00 mole of ice at -25 C to water vapor (steam) at 125 C under a constant pressure of 1 atm. specific heats ice = 2.09 J/g-K water = 4.18 J/g-K ice = 1.84 J/g-K For water H fus = 6.01 kj/mol H vap = kj/mol answer: H = 56.0 kj Vapor Pressure Critical temperature the highest temperature at which a liquid can exist Critical pressure the pressure required to bring about liquefication at the critical temperature molecules on the top layer of a liquid gain enough energy from the environment to escape these gas molecules sitting on top of the liquid exert a pressure on the surface of the liquid as the amount of vapor increases, the likelihood that a gas molecule will hit the surface of the liquid increases dynamic equilibrium will be established liquid molecules escaping to become gas at the same rate gas molecules condense back to liquid 11
12 Vapor Pressure affects BP Phase Diagram boiling occurs when the vapor pressure of a liquid is equal to the external (atmospheric) pressure as external pressure decreases, liquid surface experiences less pressure allowing more molecules to turn into gas form Triple Point Solids the temperature and pressure at which a substance is in equilibrium among all three phases (solid, liquid, gas) Amorphous no order Crystalline highly ordered Ionic Solids Metals Pack themselves to maximize ionic attractions & minimize repulsions Metal atoms bond because electrons are easily delocalized forms an electron sea 12
13 Covalent Network Solids Covalent Network Solids Diamonds (left) are examples of solids where molecules are covalently bonded Graphite is an example of a solid where molecules are held together by van der Waals forces Types of Bonding in Crystalline Solids 13
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