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Date Created: 10/05/15
Attractive forces between particles 1 Covalent bonds involve the sharing of electrons between two or more atoms Strength very strong ranging from 200 kJmol to 1000 kJmol Seen in all molecular compounds all network solids all polyatomic ions 2 Ion ion attraction ionic bonds the electrostatic attraction between oppositely charged ions Strength very strong ranging from 400 kJmol to 3000 kJmol for an individual pair of ions The overall lattice energies are larger than these numbers because each ion is surrounded by several ions of the opposite charge Seen in all ionic compounds 3 Metallic bonds involve the sharing of electrons throughout a set of metal atoms Strength moderate to very strong ranging from 70 ldmol to 800 ldmol Seen in all metals T he following three types of attraction are referred to collectively as intermolecular forces 4 London dispersion forces LDF s involve the random uctuation of electron density in a molecule creating dipoles which attract one another The larger the atoms in a molecule the stronger the LDF s will be In addition the more atoms a molecule has the stronger the LDF s will be if you keep the same elements Strength weak to very weak ranging from ltl kJmol to 20 kJmol Seen in all molecular substances and the inert gases 5 Dipole dipole attraction involve the attraction of two molecules that have permanent dipoles Strength weak to very weak comparable to LDF s Seen in all polar molecular substances 6 Hydrogen bonds involve the attraction between a molecule that has a NH OH or FH bond which will be highly polar and a second molecule that has a negatively polarized N or 0 Strength weak ranging from 5 kJmol to 40 kJmol Seen in all substances that contain at least one N H OH or FH bond Note hydrogenbonded substances typically show noticeably higher melting and boiling points than substances that have similar composition but no ability to hydrogenbond Types of solids 1 Molecular solid Properties low melting and boiling points typically melt below 300 C although higher temperatures are seen for very large compounds Many are liquids or gases at room temperature Solubilities vary widely some dissolve in water and many can dissolve in nonpolar liquids no other type of substance dissolves in nonpolar solvents Do not conduct electricity and conduct heat poorly Bondingintramolecular covalent bonds Bondingintermolecular London dispersion forces all molecular solids Polar substances also show dipoledipole attraction Substances with NH 0H or FH bonds show hydrogenbonding a special case of dipoledipole All of these forces are relatively weak How to recognize first look at the physical properties particularly the melting point and solubility The chemical formula provides a good clue too any substance that is made entirely from nonmetals should be assumed to be molecular unless proven otherwise This includes acids all acids are molecular substances Examples H20 C02 NH3 CH4 HCl H2SO4 2 Ionic solid Properties high melting points and very high boiling points typical ionic solids melt above 500 C and boil above 1000 C All are solids at room temperature Many dissolve well in water and aqueous solutions conduct electricity well Ionic compounds conduct heat fairly well They do not conduct electricity when they are solid but they conduct electricity very well when melted Ionic compounds are normally crystalline in appearance and the crystals are very brittle they shatter when struck Bonding the ions are attracted to one another by ionion attraction Many ionic compounds contain polyatomic ions such as C03239 or NH4 which are held together by covalent bonds How to recognize any compound that contains a metal and one or more nonmetals should be assumed to be ionic until proven otherwise A few seemingly ionic compounds are actually molecular these are easily recognized by their low melting points An example is AlBI 3 which looks ionic Al is a metal and Br is a nonmetal but melts at 97 C far too low for an ionic compound Examples NaCl KN03 NaHC03 NH42SO4 3 Metallic solid Properties metals normally have high melting and boiling points but there is a very wide range from mercury melting point 39 C to tungsten melting point 34100C Mercury is the only metal that is a liquid at room temperature all other metals are solids Metals are insoluble in all common solvents they can only be dissolved in strong acids which convert the metal into an ionic compound Metals conduct heat and electricity very well Most metals are silvery or gray the exceptions are gold and copper and all are re ective when polished Metals are malleable they deform when struck rather than shattering and ductile they can be stretched into wires Bonding metallic bonds These metallic bonds are actually a special case of covalent bonding in which the valence electrons of the metal atoms are delocalized over the entire sample of metal How to recognize see if the element lies in the metal region of the periodic table A substance that contains nothing but metallic elements is normally a metal any substance that contains a nonmetal must be something else A few elements have two forms one of which is metallic these elements are found on the border between the metals and nonmetals on the periodic table and are called metalloids To confirm that a particular form of an element is metallic look for high electrical and thermal conductivity malleability and ductility Examples Fe Au Na Al brass a mixture of Cu and Zn 4 Network covalent solid Properties network solids have very high melting and boiling points They normally melt above 500 C and many melt well over 10000C All are solids at room temperature They are often very hard but this varies Network solids are brittle so they shatter when struck They normally conduct heat reasonably well but are poor conductors of electricity graphite being the exception as it conducts electricity well Network solids are insoluble in all solvents Bonding covalent bonds A diamond or a piece of quartz SiOz is actually a single enormous molecule How to recognize in general if a substance is made entirely from nonmetals it will be either molecular or network covalent these are easily distinguished by their melting points Examples C diamond or graphite B SiOz A GUIDE TO BORN HABER CYCLES We can measure the strength of a covalent bond because we can usually break a covalently bonded molecule into fragments and measure the energy required to do so For example if we want to measure the bond energy in HCl we can carry out the reaction HClg gt H Cl However we cannot measure the strength of an ionic bond this way because ionic compounds do not break apart into gaseous ions NaCls gt Nag Cl g won t occur BomHaber cycles are used to estimate the strength of ionic bonds in compounds such as NaCl In a BomHaber cycle we carry out the following sequence of reactions Elements gt Gaseous gt Gaseous gt Compound in their atoms ions standard states The energy of each step except the last one can be measured experimentally In addition we can measure the energy of the singlestep reaction below this is the heat of formation of the compound Elements in their standard states gt Compound By Hess s Law the energies of the reactions in the first sequence must add up to the heat of formation the energy of the singlestep reaction The basic concept is not hard The difficulty is in keeping the details straight Ionic compounds contain two or more elements each of which must be converted to gaseous atoms and then to gaseous ions For example here are the steps required to convert elemental calcium a solid at room temperature to gaseous calcium ions Sublime the solid calcium convert it to a gas Cas gt Cag Remove one electron from each atom Cag gt Cag 6 Remove a second electron from each atom Cag gt Ca2g e Nonmetals often form covalent molecules If so you must break the covalent bond as part of this process Here are the steps required to convert elemental bromine a diatomic liquid at room temperature to gaseous bromide ions Vaporize the liquid bromine convert it to a gas Brzl gt Brzg Break the covalent bond in Brz Brzg gt 2 Brg Add an electron to each atom Brg e gt Br g Your job in sorting out a BomHaber cycle has two parts The first is to be able to figure out exactly what reactions must occur when you convert the original element to a monatomic gas The second is to know how to identify the energy of each reaction type The reactions you might see in a Born Haber cycle 1 Heat of sublimation AHsubl this is the energy required to convert a solid to a gas For metals and most solid nonmetals sublimation produces a monatomic gas For iodine which is diatomic sublimation produces I2g NaS gt Nag SS gt Sg 12S gt 12g Heats of sublimation are always positive numbers 2 Heat of vaporization AHvap this is the energy required to convert a liquid to a gas The only elements for which this will come into play are bromine and mercury which are liquids at room temperature For bromine vaporization produces Br2g Hg1 gt Hgg Brz1 gt Brzg Heats of vaporization are always positive numbers 3 Bond dissociation energy BDE or AHBDE this is the energy required to break a covalent bond Bond dissociation energies only come into play for the diatomic nonmetals H2 N2 02 F2 C12 Br2 I2 and At2 In these cases sublimation or vaporization gives us diatomic molecules not individual atoms H2g a 2 mg N2 2 M9 Bond dissociation energies are always positive numbers 4 Ionization energy IE or AHIE this is the energy required to remove one electron from a gaseous atom Since many ionic compounds contain metals that have lost two or more electrons we often need to consider two or more successive ionization energies For example if we need to make aluminum ions we must remove three electrons from Alg so we must consider the rst three ionization energies of aluminum Alg gt Alg 6 AH rst ionization energy IEl Alg gt A12g 6 AH second ionization energy IE2 Alzg gt Al3g 6 AH third ionization energy IE3 Ionization energies are always positive numbers and they increase as you remove more electrons so IEl lt IE2 lt IE3 for any given element 5 Electron affinity EA or AHEA this is the energy absorbed or released when you add one electron to a gaseous atom In general only the rst electron affinity can be measured directly because negative ions repel electrons However the second and third if necessary electron af nity can be estimated using a variation on the BornHaber cycle Here are the reactions that must be considered if you need to make oxide ions Og e gt O g AH rst electron affinity EAl or simply EA O g e gt 027g AH second electron af nity EA2 The rst electron affinity is usually a negative number a few elements have positive EA s The second EA and beyond is always positive 6 Crystal lattice energy CLE or AHCLE this is the energy change that takes place when gaseous ions are combined to form a solid ionic compound For example the lattice energy of aluminum uoride corresponds to the following reaction Al3g 3 F g gt AlF3s AH crystal lattice energy of AlF3 Lattice energies are always negative numbers and are normally very large In the BornHaber cycle these reactions are the last step since the nal product of a BomHaber cycle is the ionic compound The reverse of lattice energies for ionic r 39 is 39 to bond 139 39 quot energies for covalent bonds the energy required to separate atoms into the gas phase 7 Heat of formation AHf this is the energy absorbed or released when you make an ionic compound from its constituent elements as they normally appear at room temperature and 1 atm pressure For aluminum uoride the corresponding reaction is Als 32 Fzg gt A1F3s Heats of formation are almost always negative numbers for ionic compounds Putting the reactions together Here is an example of how these reactions can be tted together to make the complete Bom Haber cycle for KF We start with solid K and gaseous F2 the standard states of these elements at room temperature and 1 atm pressure AH IE1 Ks Squotb gt we gt mg CLE Km 1 XBDE 12F2g2 gtFg i Fg T AHf Energy bookkeeping AHsubl IE1 12 x BDE EA CLE AHf For Can the energy bookkeeping would look like this see if you can figure out why AHsub1 IE1 IE2 BDE 2 x EA CLE AHf