Ionic Bonds. Chapter 8 Chemical Bonds (+VSEPR from Chapter 9) Li Be B C N O F Ne delocalized electron sea. 3. Introduction. Types of Chemical Bonds

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1 hapter 8: hemical Bonds (+ VSEPR) hapter bjectives: hapter 8 hemical Bonds (+VSEPR from hapter 9) Understand the principal types of chemical bonds. Understand the properties of ionic and molecular compounds. Draw Lewis dot structures for molecular compounds, including resonance structures. (hapter ) Predict the shape of a molecule using VSEPR theory, and determine whether the molecule is polar. Mr. Kevin A. Boudreaux Angelo State University EM 1411 General hemistry hemistry: The Science in ontext (Gilbert, 4 th ed, 2015) Introduction Most of the substances that we encounter in daily life are not elemental substances, but compounds (and frequently, complex mixtures of compounds). Why do elements form compounds in the first place? Bonding lowers the potential energy between the charged particles that compose atoms. There are three major ways of modeling bonds between compounds, with varying degrees of complexity: Lewis theory (Lewis dot structures) Valence Bond theory Molecular rbital theory 2 Types of hemical Bonds There are three types of bonds between elements: Ionic bonds result from a transfer of electrons from one species (usually a metal) to another (usually a nonmetal or polyatomic ion). ovalent bonds result from a sharing of electrons by two or more atoms (usually nonmetals). Metallic bonds result from a pooling of valence electrons by two or more metals into a Lewis Dot Structures for Elements In a Lewis dot structure [G. N. Lewis] for an element, the valence electrons are written as dots surrounding the symbol for the element. Place one dot on each side first; the remaining dots are paired with one of the first set of dots. A maximum of two dots are placed on each side. The unpaired dots indicate where covalent bonds can form, or where electrons can be gained or lost. Atoms tend to form bonds to satisfy the octet rule, having eight electrons in their valence shells. Li Be B N Ne delocalized electron sea. 3 4 Ionic Bonds Ionic bonds form when one atoms transfers one or more electrons to another atom, producing ions. Ionic Bonds 5 Ionic compounds are compounds that are held together by ionic bonds between positively-charged cations and negatively-charged anions. The ionic bond is the strong attraction between the cations and the anions. The cation and anion are not physically joined (i.e., they do not form a molecule). Ionic compounds generally result when a metal combines with a nonmetal: Metal + Nonmetal ionic compound Metal + Polyatomic ion ionic compound 6

2 hapter 8: hemical Bonds (+ VSEPR) ormation of Ionic Solids When an element which gives up electrons easily (i.e., that has a small ionization energy) comes in contact with an element that accepts an electron easily (i.e., that has a large negative electron affinity), an electron may be transferred, yielding a cation and an anion. ormation of Sodium hloride Sodium metal, Na hlorine gas, l 2 Sodium hloride, Nal Na + l Na + l Na [e] 2s 2 2p 6 3s 1 l [Ne] 3s 2 3p 5 Na + [e] 2s 2 2p 6 = [Ne] l - [Ne] 3s 2 3p 6 = [Ar] In both cases, we have formed an ion where the new valence shell is full, having eight electrons. 7 8 The ctet Rule The Na + and l - ion s electron configurations are the same as that of the nearest noble gas (the ions are said to be isoelectronic with the nearest noble gas). Atoms prefer to have a filled outermost shell because this is more electronically stable. This can be generalized into the octet rule: Maingroup elements tend to undergo reactions that leave them with eight outer-shell electrons. That is, main-group elements react so that they attain a noble gas e - configuration with filled s and p sublevels in their valence electron shell. There are many exceptions to the octet rule, but it is useful for making predictions about some chemical bonds. Lewis Structures of Ionic ompounds 1. Use Lewis structures to predict the formula of the ionic compound formed by the following elements. a. a and l b. a and S c. Na and S 9 10 Energy hanges in the ormation of Nal The energy changes in the formation of Nal from Na(s) and l 2 (g) are shown on the next slide: Step 1: sublimation of Na(s) to Na(g): Na(s) Na(g) D = +108 kj/mol Step 2: ionization of Na(g) into Na + (g) ions: Na(g) Na + (g) + e - D = +496 kj/mol Step 2: dissociation of l 2 (g) into l(g) atoms: ½l 2 (g) l(g) D = +122 kj/mol Step 4: addition of e - to l(g) to form l - (g) ions: l(g) + e - l - (g) D = -349 kj/mol The total so far: Na(s) + ½l 2 (g) Na + l - (g); D = +377 kj/mol Born-aber ycle for the ormation of Nal 11 12

3 hapter 8: hemical Bonds (+ VSEPR) Energy hanges in the ormation of Nal But we don t find Nal in the gas phase (unless it s really hot outside!): Step 5: formation of solid Nal crystals from isolated Na + and l - ions in the gas phase: Na + l - (g) Na + l - (s) D = -788 kj/mol Revised total: Na(s) + ½l 2 (g) Na + l - (s); D = -411 kj/mol Since this is a release of energy, this is a much more favorable energetic process. The energy change in step 5 drives the overall process of the formation of Nal. The energy changes are summarized in a diagram called a Born-aber cycle (previous slide). 13 Lattice Energy The energy change in step 5 is called the lattice energy (D lattice ), the energy change associated with the formation of a crystal lattice from isolated ions in the gas phase. It is the result of the electrostatic interactions between ions, and thus it is a measure of the strength of the ionic bonds in an ionic crystal. Ionic solids exist only because the lattice energy drives the energetically unfavorable electron transfer. The energy required for elements to gain or lose electrons is supplied by the electrostatic attraction between the ions they form. 14 Trends in Lattice Energy The force that results from the interaction of electric charges is described by oulomb s law: z1z2 k 2 d where k is a constant, z 1 and z 2 are the charges on the ions, and d is the distance between their centers. The lattice energy is the force times distance: z1z2 D lattice d k d As the size of the ions increases, d becomes larger, and the lattice energy decreases. As the magnitude of the cation and anion charge increases, the lattice energy also increases (Na = -923 kj/mol, Mg = kj/mol). Trends in Lattice Energy D lattice (kj/mol) Lil -834 Nal -787 Kl -701 sl -657 D lattice (kj/mol) Na -923 a Examples: Lattice Energies 2. Which compound in each of the following pairs of ionic substances has the most negative lattice energy? a. Nal or Kl b. Li or Lil Ionic Solids and rystals The ionic bond is the strong attraction between the cations and the anions. Unlike molecules, the cation and anion are not physically joined together. Thus, there is no molecule of Nal; ionic compounds instead form ionic solids, which contain equal amounts of positive and negative charge surrounding each other in a regular array called a crystal. c. Nal or Na 2 d. Mg or BaS 17 18

4 hapter 8: hemical Bonds (+ VSEPR) Properties of Ionic ompounds The oppositely charged ions are attracted to each other by electrostatic forces forming an ionic bond. The substance that forms is an ionic solid. The solid consists of a three-dimensional array called a crystal, which consists of cations surrounded in some fashion by anions Properties of Ionic ompounds A typical ionic compound, such as rock salt (Nal) is hard (doesn t dent), rigid (doesn t bend), and brittle (cracks without deforming). The attractive forces in ionic compounds hold the ions in specific positions; moving the ions out of position requires overcoming these forces, so the sample resists denting and bending Properties of Ionic ompounds Properties of Ionic ompounds Ionic compounds have very high melting points (Nal melts at 801º, Mg melts at 2852º), and extremely high boiling points, because it takes a lot of heat energy to overcome the electrostatic interactions between cations and anions. Solid ionic compounds do not conduct electricity, because the ions are not free to move. MP (º) BP (º) sbr NaI Mgl KBr al >1600 Nal Li K Mg When melted, or dissolved in water, the ions are free to move, allowing electricity to be conducted through the solution. 22 ovalent Bonds and Molecules ovalent bonds form when two or more nonmetals share their electrons. The electrons are at their lowest potential energy when they are between the two nuclei that are being joined. ovalent Bonds Share and Share Alike 23 Each atom in the bond holds on to the shared electrons, and the atoms are thus physically tied together. 24

5 hapter 8: hemical Bonds (+ VSEPR) The ormation of Diatomic ydrogen As two isolated atoms move closer together, the two positively-charged nuclei repel each other, and the two negatively-charged electrons repel each other, but each nucleus attracts both electrons. At some point, the attractions between the nuclei and the electrons are balanced against the repulsions between the nuclei and between the electrons. The shared electrons bind the two nuclei into an 2 molecule. The shared electrons act like they belong to both atoms in the bond. ovalent Bonding and Potential Energy When two isolated atoms approach each other, the potential energy is lowered energy is released when an 2 molecule forms. Pushing the nuclei any closer causes the potential energy to rise. 25 igure ovalent Bonding and Potential Energy The optimum distance between nuclei where the attractive forces are maximized and the repulsive forces are minimized is called the bond length. (or 2, the bond length is 74 pm.) In 2, the highest probability of finding electrons is in the space between the nuclei. The increased attractive forces in this area help to lower the potential energy of 2 relative to isolated atoms. Lewis Structures and the ctet Rule The sharing of electrons by nonmetals to form molecules is modeled by the use of Lewis structures, in which sticks ( ) represent pairs of shared electrons, and dots (:) represent unshared electrons (lone pairs). Atoms share electrons in such a way as to satisfy the octet rule, which gives each atom a total of eight valence electrons. ydrogen is an exception to the octet rule, since its 1s orbital can only hold two electrons. nce the Lewis structure has been drawn, the 3D shape of the molecule and the polarity of the molecule can be predicted using the VSEPR model Single ovalent Bonds + or + or + or + + or The shared pairs of electrons are bonding pairs. The unshared pairs of electrons are lone pairs or nonbonding pairs. Na + l - ionic bond covalent bond 29 Double and Triple ovalent Bonds Atoms can also share two pairs of electrons to form a double bond. Each oxygen atom in the last structure below has an octet of electrons: + Double bonds are shorter and stronger than single bonds. The sharing of three electron pairs forms a triple bond: N + N N N Triple bonds are even shorter and stronger than double bonds. 30

6 hapter 8: hemical Bonds (+ VSEPR) Lewis Structures Drawing the Line 31 ow To Write Lewis Structures 1. Determine the total number of valence electrons in the molecule or ion. Add one electron for each unit of negative charge. Subtract one electron for each positive charge. 2. Write the correct skeletal structure. or molecules of the formula AB n, place the atom with the greatest bonding capacity (largest # of unpaired electrons in the Lewis structure for that atom) in the center, and the other atoms on the outside. Draw single bonds between the outer atoms and the central atom. is NEVER a central atom, since it can only form one bond. 32 ow To Write Lewis Structures 3. Distribute the remaining valence electrons as lone pairs on the outer atoms first, making sure to satisfy the octet rule. nce all of the outer atoms have 8 electrons (or 2 for ), place any remaining electrons on the central atom. 4. ompare the number of valence electrons in the Lewis structure to the number in Step 1 to make sure you haven t miscounted electrons. 5. omplete the octet on the central atom. This is done by sharing lone pairs from the outer atoms with the central atom. The formal charge can be used as a guideline for placing double bonds. 33 ormal harges ormal charge = valence e - (½ bonding e - ) (lone pair e - ) ormal charge is the difference between the number of valence electrons on the free atom and the number of valence electrons assigned to the atom in the molecule. The sum of the formal charges must equal the charge on the species. Smaller formal charges are better (more stable) than larger ones. Like charges on adjacent atoms are not desirable. When a formal charge cannot be avoided, negative formal charges should reside on more electronegative atoms. 34 Examples: Lewis Structures 1. Write Lewis structures for the following molecules. a. 4 Examples: Lewis Structures 1. Write Lewis structures for the following molecules. f. l 2 2 b. N 3 g. 2 S c. N 4 + h. 2 6 d. l 4 i. 2 4 e. 2 2 j. N 2 (atoms connected in the order N N ) 35 36

7 hapter 8: hemical Bonds (+ VSEPR) Examples: Lewis Structures 1. Write Lewis structures for the following molecules. k. 2 Examples: Lewis Structures 1. Write Lewis structures for the following molecules. o. 3 l. al 2 m l 2 n. acetamide, 2 5 N N 37 Neither of the two 3 structures is correct by itself. 38 Resonance Structures When ne Lewis Structure Ain t Enough When there is more than one valid Lewis structure for a molecule, the actual electronic structure is an average of the different possibilities, called a resonance hybrid. Resonance forms differ only in the placement of the valence electrons, not the positions of the atoms! zone does not have a real double bond and a real single bond; the actual bond lengths in ozone are identical, and are between those of an bond and an = bond a one-and-a-half bond. Resonance Structures This molecule is shown more correctly with two Lewis structures, called resonance structures, with a two-headed arrow (T) between them: D NT USE TE STRAIGT, DUBLE-EADED ARRW (T) R ANYTING ELSE! Resonance structures are not real bonding depictions: 3 does not change back and forth between the two structures; the actual molecule is a hybrid of the two forms depicted. Bond order = 1 1 / Delocalized Electrons and harges In these resonance structures, one of the electron pairs (and hence the negative charge) is spread out or delocalized over the whole molecule. In contrast, the lone pairs on the oxygen in water are localized i.e., they re stuck in one place. Resonance delocalization stabilizes a molecule by spreading out charges, and often occurs when lone pairs (or positive charges) are located next to double bonds. (Resonance plays a big role in our understanding of structure and reactivity in organic chemistry.) A more accurate depiction of electron distribution is found in molecular orbital (M) theory. Resonance Structures in Benzene Another molecule in which resonance is important is benzene, 6 6, which has two resonance structures with alternating single and double bonds. The actual bond order of the carbon-carbon bonds in benzene is

8 hapter 8: hemical Bonds (+ VSEPR) A Resonance Analogy Examples: Lewis Structures With Resonance 2. Write Lewis structures for the following molecules, including resonance structures. a. N - b A mule is not sometimes a horse and sometimes a donkey; it s always one thing (a mule), just like purple is not sometimes red and sometimes blue. A real person can be described as having characteristics of two or more fictional characters. The fictional characters don t exist, but the real person does. c. N 3 - d Exceptions to the ctet Rule, Part 1 Electron deficient species, such as beryllium (Be) and boron (B), can have fewer than eight electrons around them, but have zero formal charge. ree radicals (or just radicals, or odd-electron molecules) contain an odd number of valence electrons. Radicals always have an unpaired electron, and are paramagnetic. These species are usually unstable, and are extremely reactive. Expanded octets are found on atoms that have more than eight electrons around them. Nonmetals from period 3 or higher, such as sulfur and phosphorus, can get around the octet rule by shoving extra electrons into empty d orbitals. [more later] Examples: Exceptions to the ctet Rule 3. Write Lewis structures for the following molecules, including resonance structures, if necessary. (BEWARE! The formula alone doesn t tell you about lone pairs on the central atom!) a. N b. N 2 c. Bel 2 d. B 3 45 e. S 6 46 Polar ovalent Bonds Electronegativity In reality, fully ionic and covalent bonds represent the extremes of a spectrum of bonding types. Electronegativity is the ability of an atom in a molecule to attract shared electrons to itself. [Linus Pauling, 1939; Nobel Prize 1954, 1963] Electronegativity is a periodic property, and increases from bottom to top within a group and from left to right across a period (inversely related to atomic radius)

9 hapter 8: hemical Bonds (+ VSEPR) Electronegativity Values (The Pauling Scale) The atom in a covalent bond with the larger electronegativity value has a partial negative charge (d - ). The atom in a covalent bond with the smaller electronegativity value has a partial positive charge (d + ). Nonpolar ovalent Bonds In an ionic bond, the electronegativity difference between the two atoms is so large that an electron is transferred from one atom to the other: A covalent bond in which there is no electronegativity different between the bonded atoms is a nonpolar covalent bond: igure Polar and Nonpolar ovalent Bonds Electronegativity and Bond Polarity A covalent bond in which there is an electronegativity difference between the bonded atoms (between 0.4 and 2.0) is a polar covalent bond: The polarity of a bond depends on the degree of the electronegativity difference: DEN = (small) Nonpolar covalent bond DEN = (intermediate) Polar covalent bond DEN > 2.0 (large) Ionic bond the electron cloud between the atoms is polarized; that is, it leans towards one side of the bond. The electrons are still shared, but unequally. 51 A polar covalent bond can be thought of as having partial covalent character and partial ionic character: the greater the DEN of the atoms, the larger the ionic character of the bond. 52 Electronegativity and Bond Polarity Electronegativity and Bond Polarity igure 8.4 igure

10 hapter 8: hemical Bonds (+ VSEPR) Electronegativity and Bond Polarity In the period 3 chlorides below, as the difference in electronegativity (DEN) decreases, the bond becomes more covalent; we move from tightly-bound ionic solids (Nal) to more weakly bound covalent liquids (Sil 4 ) to even more weakly bound gases (l 2 ). 55 Dipole Moment and Percent Ionic haracter Bond polarity is expressed numerically as a dipole moment,, which occurs when there is a separation between a positive and negative charge. The unit of dipole moment is the debye, D (1 D = m) By comparing the measured dipole moment of a bond to what the dipole moment would be if the electron were completely transferred from one atom to the other, the percent ionic character of the bond can be determined. A bond in which an electron is completely transferred would have 100% ionic character. In general, bonds with greater than 50% ionic character are considered to be ionic bonds. 56 Dipole Moment and Percent Ionic haracter Examples: Nonpolar / Polar / Ionic Bonds 2. Determine whether the bond formed between each of the following pairs of atoms is a nonpolar covalent bond, a polar covalent bond, or an ionic bond. a. Na and l A thin stream of a polar solvent, such as water, is deflected by a static electric charge, while a nonpolar molecule, such as hexane, is not. igure b. and l c. N and l d. N and e. Sr and f. l and l 58 Vibrating Bonds and the Greenhouse Effect Because 2 contains polar bonds, it can absorb and emit photons of infrared (IR) radiation. In some vibrational modes which change the electric fields associated with bond polarity, IR photons may be absorbed. This is why 2 is a greenhouse gas. igure 8.10 Because N 2 and 2 are nonpolar molecules, there is no vibrational mode of stretching or bending which changes the electric fields associated with bond polarity; they are IR inactive

11 hapter 8: hemical Bonds (+ VSEPR) You ve Just rossed ver into... The Lewis structures we ve seen so far only tell us how the atoms are connected, not what their arrangement in space is. VSEPR Theory Getting in Shape The approximate molecular structure, the threedimensional arrangement of the atoms in a molecule, can be predicted from the Lewis dot structure using the Valence-Shell Electron-Pair Repulsion (VSEPR) model, The 3D shape of a molecule is important in determining its overall polarity, chemical behavior, and is particular for biologically important molecules, which often have complex and specific 3D shapes The VSEPR Model In the VSEPR model, an atom s bonding and nonbonding electron pairs are positioned as far apart as possible to minimize electron-pair repulsions. Each bond or lone pair, or unpaired electron counts as one electron group (the book uses the term steric number). Multiple bonds count as one electron group. There are five electron-group shapes (or electron-pair geometries) in the VSEPR model, based on the positions of all of the electron groups around an atom (bond and lone pairs). The molecular shape (or molecular geometry) is the three-dimensional shape of the molecule, based on the positions of the atoms only. VSEPR Electron-Group Shapes A Balloon Analogy for the VSEPR Model Two Electron Groups EGS = Linear Two electron groups: electron-group shape = linear 2 bonds, 0 lone pairs: molecular shape = linear, l Be l Three Electron Groups EGS=Trigonal Planar Three electron groups: electron-group shape = trigonal planar 3 bonds, 0 lone pairs: molecular shape = trigonal planar, bonds, 1 lone pair: molecular shape = bent, < 120 linear trigonal planar <120 bent Electron clouds in lone pairs (b) take up more room than those in covalent bonds (a) ; this causes the other atoms to be 66 more squashed together.)

12 hapter 8: hemical Bonds (+ VSEPR) Three Electron Groups EGS=Trigonal Planar B our Electron Groups EGS = Tetrahedral our electron groups: electron-group shape = tetrahedral 4 bonds, 0 lone pairs: molecular shape = tetrahedral, bonds, 1 lone pair: molecular shape = trigonal pyramidal, < bonds, 2 lone pairs: molecular shape = bent, <109.5 S tetrahedral <109.5 trigonal pyramidal <109.5 bent 68 our Electron Groups EGS = Tetrahedral our Electron Groups EGS = Tetrahedral N Molecular Shape and Polarity or diatomic molecule, the polarity of the molecule depends on the polarity of the only covalent bond: if the bond is polar the molecule must be polar

13 hapter 8: hemical Bonds (+ VSEPR) Molecular Shape and Polarity In molecules with more than one bond, the polarity of the bonds and the overall shape determine whether the molecule is polar. 2 is nonpolar, because even though it has polar bonds, they point 180º from each other. Water is polar because it has polar bonds which point about 107º away from each other. Determining Molecular Polarity To determine whether a molecule is polar: Draw a Lewis structure for the molecule and predict its shape using VSEPR theory. Determine whether the molecule contains polar bonds. If the molecule contains polar bonds, superimpose a vector ( ), pointing towards the more electronegative atom of each bond. Determine whether the polar bonds add together to form a net dipole moment by adding the vectors corresponding to the polar bonds together. If the vectors sum to zero, the molecule is nonpolar, but if there is a net vector, the molecule is polar Determining Molecular Polarity l 4 l l 3 l l l l l l Physical Properties of Polar Molecules The polarity of molecules has a large effect on their chemical and physical properties: Polar molecules attract one another more strongly than nonpolar molecules do, and generally have higher boiling points. Polar molecules interact with each other more than they interact with nonpolar molecules, which is why a mixture of oil (nonpolar) and water (polar) separates into two layers. nonpolar molecule polar molecule Examples: Shape and Polarity 1. Write Lewis structures for the following molecules and determine their shape. State whether or not the molecules will be polar. a. B 3 b. N 3 c. S 2 Examples: Shape and Polarity 1. Write Lewis structures for the following molecules and determine their shape. State whether or not the molecules will be polar. f. P 3 g. N h. 2 d. 4 e

14 hapter 8: hemical Bonds (+ VSEPR) Expanded ctets: Lewis Structures, VSEPR, and Polarity Exceptions to the ctet Rule, Part 2 Expanded octets are found on atoms that have more than eight electrons around them. Nonmetals from period 3 or higher, such as sulfur and phosphorus, can get around the octet rule by shoving extra electrons into empty d orbitals. These atoms do not always violate the octet rule (e.g., the S in 2 S follows the octet rule, but the S in S 4 and S 6 has an expanded octet). Period 2 elements cannot have expanded octets. In general, nonmetals from period 3 or higher have expanded octets when they are bonded to strongly electronegative elements (,, l), or when an expanded valence shell reduces the formal charges on the atoms ive Electron Groups Trigonal Bipyramidal ive electron groups: electron-group shape = trigonal bipyramidal 5 bonds, 0 lone pairs: molecular shape = trigonal bipyramidal, 120 (equatorial), 90 (axial) 4 bonds, 1 lone pair: molecular shape = seesaw, <120 (equatorial), <90 (axial) 3 bonds, 2 lone pairs: molecular shape = T-shaped, <90 2 bonds, 3 lone pairs: molecular shape = linear, 180 ive Electron Groups Trigonal Bipyramidal eq eq = equatorial ax = axial ax ax <90 eq eq 90 trigonal bipyramidal <90 <120 seesaw 81 T-shaped linear 82 ive Electron Groups Trigonal Bipyramidal ive Electron Groups Trigonal Bipyramidal l l l P l l S Br Xe 83 84

15 hapter 8: hemical Bonds (+ VSEPR) Six Electron Groups ctahedral Six Electron Groups ctahedral Six electron groups: electron-group shape = octahedral 6 bonds, 0 lone pairs: molecular shape = octahedral, 90 5 bonds, 1 lone pair: molecular shape = square pyramidal, <90 4 bonds, 2 lone pairs: molecular shape = square planar, 90 S Br Xe 90 <90 90 octahedral square pyramidal square planar Examples: Exceptions to the ctet Rule Examples: Exceptions to the ctet Rule 1. Write Lewis structures for the following molecules, including resonance structures, if necessary. Predict the shape of the molecules, and whether or not the molecules are polar (BEWARE! The formula doesn t tell you how many lone pairs on the central atom!) 1. Write Lewis structures for the following molecules, including resonance structures, if necessary. Predict the shape of the molecules, and whether or not the molecules are polar (BEWARE! The formula doesn t tell you how many lone pairs on the central atom!) a. S 3 e. Pl 3 b. S 4 f. I 3 - c. S 6 g. Xe 2 d. S 4 2- h. Xe

16 hapter 8: hemical Bonds (+ VSEPR) Shapes of Larger Molecules The shapes of larger molecules are a composite of the shapes of the atoms within the molecule, each of which can be predicted using VSEPR theory Bond Lengths and Bond Energies Bond Lengths Bond order is the number of bonds between two atoms: a single bond has a bond order of 1 a double bond has a bond order of 2 a triple bond has a bond order of 3 As the bond order increases, the bond length decreases. Molecules in which where is resonance may have fractional bond orders if the double bond is spread out over more than one position: b.o. = 2 b.o. = 1.5 b.o. = 1 93 igure Bond Lengths Multiple bonds are shorter and stronger than their single-bond counterparts. When larger atoms are joined together, the bond becomes longer. Longer bonds are weaker than shorter bonds. Average ovalent Bond Lengths and Bond Energies igure

17 hapter 8: hemical Bonds (+ VSEPR) Bond Length & Bond Energy for Multiple Bonds Bond Bond Length (pm) Bond Energy (kj/mol) = t = t N N N=N NtN N =N tn Bond Energies The bond energy (or bond dissociation energy) is the amount of energy required to break 1 mole of the covalent bond in the gas phase. or the l l bond in l 2, the bond energy is 243 kj/mol: l 2 (g) 2l(g); D = +243 kj or the l bond in l, the bond energy is 431 kj/mol l(g) (g) + l(g); D = +431 kj The l bond is stronger than the l l bond, because it takes more energy to break the bond Bond Energies Bond energies are positive (endothermic) because it takes energy to break a bond. When a bond forms, the bond energy is the amount of energy that is released (exothermic). Table 8.3 contains a list of average bond energies, with values obtained by averaging the energies of that type of bond in a number of different compounds. D D Bond Energies and Enthalpy hanges Average bond energies can be used to estimate enthalpy changes for chemical reactions: rxn rxn D 's bonds broken positive values or D 's bonds broken D 's bonds formed negative values D 's bonds formed A reaction is exothermic when weak bonds break and strong bonds form. A reaction is endothermic when strong bonds break and weak bonds form Bond Energies and Enthalpy hanges Example: Use the table of average bond energies to estimate the change in enthalpy for the combustion of methane ( 4 ) Bonds Broken 4 4(+413 kj) 2 = 2(+495 kj) Bonds ormed 2 = 2(-799 kj) 4 4(-463 kj) D = ( ) + ( ) kj = (2642) + (-3450) kj = -808 kj Examples: Bond Length and Bond Strength 1. Using the periodic table, rank the bonds in each set in order of increasing bond length and increasing bond strength. a. S, S Br, S l b. =,, t

18 hapter 8: hemical Bonds (+ VSEPR) Examples: Estimating Enthalpy hanges 2. ydrogen gas can be made by the reaction of methane gas and steam. 4 (g) (g) 4 2 (g) + 2 (g) Use the bond energies in Table 8.3 to estimate D for this reaction. Properties of Molecular ompounds Ionic bonds are nondirectional, and hold together an entire array of ions in a crystal lattice. ovalent bonds are directional, and hold specific atoms together in a molecule. Molecular compounds are generally gases, liquids, or low-melting solids. The covalent bonds within molecules are very strong, but the attractive forces between the separate molecules are fairly weak. Answer: +162 kj Metallic Bonding: The Electron Sea Model In a solid metal, each metal atom donates one or more of its valence electrons to form an electron sea that surrounds the metal atoms. Metals conduct electricity because the electrons in the electron sea are free to move around. Metals conduct heat because the electrons help to disperse thermal energy throughout the metal Metals are malleable and ductile because there are no localized bonds between the metal atoms, allowing the metal to be deformed easily