Chapter 8 Concepts of Chemical Bonding

Chapter 8 Concepts of Chemical Bonding

Chapters 8,9 Concepts of Chemical Bonding and Bonding Theories Chemical Bonding Chemical Bonds Three basic types of bonds: Ionic Electrostatic attraction between ions Covalent Sharing of electrons Metallic Metal atoms bonded to

several other atoms Chemical Bonding Ionic Bonding Ionic compounds are formed when electrons are transferred from one atom to another to form ions with complete outer shells of electrons. In an ionic compound the positive and negative ions are attracted to each other by strong electrostatic forces, and build up into a strong lattice. Ionic compounds have high melting points as considerable energy is required to overcome these forces of attraction. Defined by formula units. Chemical Bonding

[Ne] 3s1 11p 11e [Ne] 11p 10e [Ne] 3s2 3p5 Na Na+ + + 17p 17e

Cl Cl - [Ne] 3s2 3p5 17p 18e Chemical Bonding Chemical Bonding Chemical Bonding

Energetics of Ionic Bonding As we saw in the last chapter, it takes 495 kJ/mol to remove electrons from sodium. Chemical Bonding Energetics of Ionic Bonding We get 349 kJ/mol back by giving electrons to chlorine. Chemical Bonding

Energetics of Ionic Bonding But these numbers dont explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic! 495 - 349= +146 kJ Chemical Bonding Energetics of Ionic Bonding There must be a third piece to the puzzle.

What is as yet unaccounted for is the electrostatic attraction between the newly formed sodium cation and chloride anion. Chemical Bonding Lattice Energy This third piece of the puzzle is the lattice energy: The energy required to completely separate a mole of a solid ionic compound into its gaseous ions. The energy associated with electrostatic

interactions is governed by Coulombs law: Q1Q2 Eel = r2 Chemical Bonding Lattice Energy Lattice energy, then, increases with the charge on the ions. It also increases with decreasing size of ions. Chemical Bonding

Formation of Ionic Compounds Small ions with high ionic charges have large Coulombic forces of attraction. Large ions with small ionic charges have small Coulombic forces of attraction. Al2O3 > CaO > KCl Use this information, plus the periodicity rules to arrange these compounds in order of increasing attractions among ions KCl, Al2O3, CaO Chemical Bonding Properties of Ionic compounds Strong bonds due to very strong electrostatic forces.

High melting points. The higher the charge and the smaller the ions, the higher the melting point: Coulombs law Very hard Low volatility Cleave along planes Brittle 3D structure Ions line up in a repetitive pattern that maximizes attractive forces and minimizes repulsive forces. Not malleable or ductile impact + - repulsion of like charges

+ - + - - + - + -

+ + - + - + - + -

- + + - + - - + +

- - + - + - + Chemical Bonding Properties of Ionic compounds Solubility and conductivity Most are soluble in polar solvents They conduct electricity only when molten or dissolved in a polar solvent, as the

charged particles are free to move. The higher the concentration of ions in a solution, the higher the electrical conductivity. Chemical Bonding Recall Dissociation When an ionic substance dissolves in water, the solvent pulls the individual ions from the crystal and solvates them. This process is called dissociation. Water is a polar molecule. Its positive ends (hydrogen sites) are attracted by the negative ions in the ionic compound, while the negative ends (oxygen sites) are attracted to

the positive ions. Ionic compounds dissolve in water because the attraction of the ions to water is stronger than the attraction between the ions themselves. Chemical Bonding Problem 1: Which solution of each pair exhibits the stronger electrical conductivity? Explain a) 1.0 M Na2CO3 or 1.0 M NaCl b) 1.0 M K2SO4 or 1.5 M KI c) 2.0 M C2H5OH or 2.0 M LiF Chemical Bonding Melting point

When making predictions about melting point, Coulombic forces must be taken into account. Problem 2: Which compound from each set has the greatest melting point? Why? a) LiF LiI b) MgCl2 MgO c) NaF MgI2 Chemical Bonding Problem 3: Which of the following metal fluorides would have the highest melting point? Why? a) b Metal fluoride b) will have the highest melting point as the ionic radius of the metal is shorter than the

ionic radius of the metal in metal fluoride a). According to Coulombs law, the force of attraction between opposite charged ions increases as the ionic radius of the ions decreases. Therefore a higher temperature will be needed to separate the ions in metal fluoride b) Chemical Bonding Covalent Bonding In these bonds atoms share electrons. There are several electrostatic interactions in these bonds: Attractions between electrons and nuclei Repulsions between electrons

Repulsions between nuclei Chemical Bonding This is how colleges explain covalent bond !!! Chemical Bonding Electronegativity: The ability of atoms in a molecule to attract electrons to itself. On the periodic chart, electronegativity increases as you go

from left to right across a row. from the bottom to the top of a column. Chemical Bonding Polar Covalent Bonds Although atoms often form compounds by sharing electrons, the electrons are not always shared equally. Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does. Therefore, the fluorine end of the molecule has more electron density than the

hydrogen end because electrons spend more time around fluorine. Chemical Bonding Polar Covalent Bonds When two atoms share electrons unequally, a bond dipole results. The dipole moment, , produced by two equal but opposite charges separated by a distance, r, is calculated: = Qr It is measured in debyes (D). Chemical Bonding

Bond Polarity In diatomic molecules containing the same element (e.g. H2 or Cl2) the electron pair will be shared equally, as both atoms exert an identical attraction because they have the same electronegativity. When atoms are different, the more electronegative atom exerts a greater attraction for the electron pair. One end of the molecule will thus be more electron rich than the other end, resulting in a polar bond. This relatively small difference in electronegativity is represented by + and -. The bigger the difference in electronegativities the more polar the bond. Chemical Bonding Polar Covalent Bonds

The greater the difference in electronegativity, the more polar is the bond. Chemical Bonding Covalent or Molecular Compounds Defined by molecules two or more non metals bonded together to form a compound. Polyatomic ions Bonds between atoms are non polar covalent or polar covalent Chemical

Bonding Lewis Structures Lewis structures are representations of molecules showing all electrons, bonding and nonbonding. Chemical Bonding Writing Lewis Structures PCl3 5 + 3(7) = 26 1. Find the sum of valence electrons of all atoms in the

polyatomic ion or molecule. If it is an anion, add one electron for each negative charge. If it is a cation, Chemical Bonding subtract one electron for each Writing Lewis Structures 2. The central atom is the least electronegative element that isnt hydrogen. Connect

the outer atoms to it by single bonds. Chemical Bonding Writing Lewis Structures 3. Fill the octets of the outer atoms. Chemical Bonding Writing Lewis Structures 4. Add up the electrons you have used and subtract them from the total number of

valence electrons (step 1). Attach leftover electrons to the central atom as lone pairs. Chemical Bonding Writing Lewis Structures 5. If you run out of electrons before the central atom has an octet form multiple bonds until it does. Chemical Bonding

Formal Charges In some cases, the atoms in a molecules can be assembled in different ways when drawing a Lewis diagram. Formal charges are assigned to identify the most stable or likely structure. For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms. Subtract that from the number of valence electrons for that atom: The difference is its formal charge. Chemical Bonding Writing Lewis Structures The best Lewis structure is the one with the fewest charges.

puts a negative charge on the most electronegative atom. Chemical Bonding Problem 4: Which is the most likely structure? Cl N Cl | Cl or N Cl Cl | Cl Chemical

Bonding Problem 5: Which is the most likely structure? [C = S = N]or [S = C = N] - Chemical Bonding Problem 6: Write Lewis structures for: CCl4 H2O SCl2 CO32-

NH2- SO2 Chemical Bonding Resonance This is the Lewis structure we would draw for ozone, O3. + -

Chemical Bonding Resonance But this is at odds with the true, observed structure of ozone, in which both OO bonds are the same length. both outer oxygen atoms have a charge of 1/2. the effective number of bonds between each pair of oxygen atoms is 1.5 Chemical Bonding

Resonance One Lewis structure cannot accurately explain a molecule such as ozone. We use multiple structures, resonance structures, to describe the molecule. Chemical Bonding Resonance Just as green is a synthesis of blue and yellow ozone is a synthesis of

these two resonance structures. Chemical Bonding Resonance In truth, the electrons that form the second CO bond in the double bonds below do not always sit between that C and that O, but rather can move among the two oxygens and the carbon. They are not localized, but rather are delocalized. Chemical Bonding Resonance Structures, Bond Length and Bond Energy

Due to the delocalization of the electrons, each C O is in between a double and a single bond. The bond order is 1.5. Chemical Bonding Resonance The organic compound benzene, C6H6, has two resonance structures. It is commonly described as a hexagon with a circle inside to signify the delocalized

electrons in the ring. Chemical Bonding Problem 7: Write all possible Lewis structures for the nitrate ion. Calculate the formal charge to determine the most likely structure. Then write all resonance structures of the ion and state the effective number of bonds between the nitrogen and each oxygen atom. Chemical Bonding Exceptions to the Octet Rule There are three types of ions or molecules that do not follow the octet rule: Ions or molecules with an odd number of electrons.

Ions or molecules with less than an octet. Ions or molecules with more than eight valence electrons (an expanded octet). Chemical Bonding Odd Number of Electrons Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons. Chemical Bonding Fewer Than Eight Electrons Just for You

Consider BF3: Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine. This would not be an accurate picture of the distribution of electrons in BF3. Chemical Bonding Fewer Than Eight Electrons Chemical Bonding More Than Eight Electrons The only way PCl5 can

exist is if phosphorus has 10 electrons around it. It is allowed to expand the octet of atoms on the 3rd row or below. Presumably d orbitals in these atoms participate in bonding. Chemical Bonding More Than Eight Electrons Just for You Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, the better structure puts a

double bond between the phosphorus and one of the oxygen atoms. Chemical Bonding Molecular models Space-Filling Model Shows differences in the atomic radii of bonded atoms. Does not show the 3D positions of atoms very well. Does not show double and triple bonds. Ball-and Stick Model Shows the 3D position of atoms well. Shows single, double and triple bonds. Balls are not proportional to the size of the atoms. Sticks are not proportional to and greatly exaggerate the bond length. H

C C H Chemical Bonding Covalent Bond Strength Energy must be absorbed to break a chemical bond, therefore the strength of a bond is measured by determining how much energy is required to break the bond. The amount of energy required to break a bond is equal to the amount of energy released when that same bond forms. This is the bond enthalpy.

Chemical Bonding The bond enthalpy for a ClCl bond, Average Bond Enthalpies This table lists the average bond enthalpies for many different types of bonds. Average bond enthalpies are positive, because bond breaking is an endothermic process. Chemical Bonding Average Bond Enthalpies NOTE: These are average bond enthalpies, not absolute bond enthalpies; the CH bonds in methane, CH4, will be a bit different than the

CH bond in chloroform, CHCl3. Chemical Bonding Enthalpies of Reaction Energy is required to break bonds. Endothermic process. Energy is released when new bonds are formed. Exothermic process. If the energy absorbed in breaking bonds is greater than the energy released when new bonds are formed, then the enthalpy of the reaction is endothermic. (positive) If the energy absorbed in breaking bonds is less than the energy released when new bonds are formed, then the enthalpy of the reaction is exothermic. (negative) Chemical Bonding

Enthalpies of Reaction In other words, Hrxn = (bond enthalpies of bonds broken) (bond enthalpies of bonds formed) Chemical Bonding Enthalpies of Reaction CH4(g) + Cl2(g) CH3Cl(g) + HCl(g) In this example, one CH bond and one ClCl bond are broken; one CCl and one HCl bond are formed. Chemical Bonding

Enthalpies of Reaction So, bonds broken - bonds formed Hrxn = [BE(CH) + BE(ClCl) [BE(CCl) + BE(HCl) = [(413 kJ) + (242 kJ)] [(328 kJ) + (431 kJ)] = (655 kJ) (759 kJ) = 104 kJ Chemical Bonding Bond Length Most atoms have lower potential energy when they are bonded to other atoms than they have as they are independent particles. The figure below shows potential energy changes during the formation of a hydrogen-hydrogen bond.

The distance between two bonded atoms at their minimum potential energy (the average distance between two bonded atoms) is the bond length. Chemical Bonding Bond Enthalpy and Bond Length As the number of bonds between two atoms increases, the bond length decreases, the bond energy (strength) increases and the PE decreases. This is because as the electron density between the positive nuclei increases, the attractive forces between the Chemical protons and the bonding electrons increase. Bonding

Problem 8: Consider the space-filling models below. a) Which molecule has the least amount of PE associated with its bond? b) Which bond has the lowest bond energy? I II Lowest PE, as it has the shortest bond length Lowest bond energy, as it has the longest bond length Chemical

Bonding Bond Order Bond order is the number of bonds between two atoms. When the bond order increases: the bond length decreases the PE associated with the bond decreases the bond energy increases Bond type Bond order Single 1 Double

2 triple 3 Chemical Bonding Problem 9: Consider the ball-and-stick models below. a) The carbon-carbon bond in which molecule has the lowest PE? b) The carbon-carbon bond in which molecule has the lowest bond energy? c) What are the weaknesses of this model? I H C

C II H H H C C H Lowest PE, as it has the shortest bond length

H Lowest bond energy, as it has the longest bond length Weaknesses: Balls are not proportional to the size of the atoms. Sticks are not proportional to the bond length and greatly exaggerate the bond length Chemical Bonding Metallic Bond A metallic solid can be represented as positive kernels (or cores) consisting of

the nucleus and inner electrons of each atom surrounded by a sea of mobile valence electrons. Chemical Bonding Metallic Bond Properties of metals are linked to their bonding. Good conductors of electricity: electrons are delocalized and relatively free to move. Malleable and ductile: deforming the solid does not change the environment surrounding each metal core. Chemical Bonding

Metallic Bond. Alloys Metals can be mixed into alloys. The properties of alloys are linked to the size of the components atoms. Interstitial alloy: Formed between metal atoms of different radius. The smaller atom fills the interstitial spaces between the larger atoms. The interstitial atoms make the lattice more rigid, decreasing the malleability and ductility. Example: steel, where C occupies the interstices in Fe. Substitutional alloy: Formed between atoms of comparable radius. One of the atom substitutes the other in the lattice. The density lies between those of the component elements, but the malleability and ductility remains. Example: brass, where Cu is Chemical substituted with Zn. Bonding

Covalent Network solids Atoms are covalently bonded together into a two or a three-dimensional network. (diamond, graphite, silicon dioxide, silicon carbide). The highest melting point of all bonding types since all atoms are covalently bonded. Three-dimensional covalent networks tend to be rigid and hard. The covalent angles are fixed. Example: diamond. Chemical Bonding Covalent Network solids Graphite is an allotrope of C that forms sheets of two-dimensional network. High melting point because the atoms are

covalently bonded. Soft because adjacent layers can slide past each other relatively easy. London dispersion forces between the layers. Chemical Bonding Covalent Network solids Silicon is a covalent network solid and a semiconductor. It forms a three-dimensional network similar to diamond. Conductivity increases as the temperature increases. N-type semiconductor: doping silicon with an element with extra valence electron. P-type semiconductor: doping silicon with

an element with one less valence electron. Chemical Bonding Molecular Structure & Covalent Bonding Theories Chemical Bonding 72 Stereochemistry Stereochemistry is the study of the three dimensional shapes of molecules. Some questions to examine in this chapter

are: 1. 2. 3. 4. Why are we interested in shapes? What role does molecular shape play in life? How do we determine molecular shapes? How do we predict molecular shapes? Chemical Bonding 73 Stereochemistry These questions can be answered by looking at how drugs of any kind work. If the molecule of the drug does not have the right shape to plug or

attach into the specific neuron or protein they are designed for, the drug will not work. This is the basis for how most pharmaceutical drugs work. Drugs are organic compounds that bind to proteins because they have the right shape like a lock and key. Lewis structures do not tell us shape, but it gives some information we need to figure the shape. Chemical Bonding 74 Two Simple Theories of Covalent Bonding Valence Shell Electron Pair Repulsion Theory Commonly designated as VSEPR Principal originator R. J. Gillespie in the 1950s

Valence Bond Theory Involves the use of hybridized atomic orbitals Principal originator L. Pauling in the 1930s & 40s Chemical Bonding 75 VSEPR Theory Valence Shell Electron Pair Repulsion Theory Charged clouds (bonding or lone pairs of electrons) repel each other due to Coulombic repulsions. Terminal atoms move as far away from one another as possible to minimize that repulsion. This results in distinctive geometric shapes. Chemical

Bonding Steps to identify the shape The same basic approach will be used in every example of molecular structure prediction: 1. Draw the correct Lewis dot structure. Identify the central atom. 2. Count the number of charge clouds (regions of high electron density) around the central atom. a) Each item below counts as a single cloud:

One single bond (consist of 2 electrons) One double bond (consist of 4 electrons) One triple bond (consist of 6 electrons) One lone pair (consist of 2 electrons) One single unpaired electrons (consist of 1 lone electron) Chemical Bonding 77 Steps to identify the shape b) Each item below is considered ONE BOND: A single bond (consist of 2 electrons) A double bond (consist of 4 electrons) A triple bond (consist of 6 electrons)

c) Each item below is considered a LONE PAIR: One lone pair (consist of 2 electrons) One single unpaired electrons (consist of 1 lone electron) Chemical Bonding 78 Problem 10: Identify the number of charge clouds in the following compounds: a) NH3 b) SF4 c)

NO2 d) BF3 Chemical Bonding Overview of Chapter 3. Predict the shape - There are 15 shapes You must know them all by name You must know the bond angles for each. Chemical Bonding 80

VSEPR Theory Charge clouds Bonds Lone Pairs Shape Examples 2 2 0

Linear CO2, BeF2 Chemical Bonding 81 VSEPR Theory Charge clouds Bonds Lone Pairs Shape

Examples 3 3 0 Trigonal planar CO32-, BeF3 Chemical Bonding 82

VSEPR Theory Charge clouds Bonds Lone Pairs Shape Examples 3 2 1

Bent NO2- Bent less than 120 The lone pair exerts a greater repelling force on the bonding electrons as it is sitting closer to the nitrogen. . .. . Chemical Bonding

VSEPR Theory Charge clouds Bonds Lone Pairs Shape Examples 4 4 0

Tetrahedral CCl4 Chemical Bonding 84 VSEPR Theory Charge clouds Bonds Lone Pairs Shape

Examples 4 3 1 Trigonal Pyramidal NH3 Trigonal pyramidal Around 107 The lone pair exerts a greater repelling

force on the bonding electrons as it is sitting closer to the nitrogen. .. .. Chemical Bonding 85 VSEPR Theory Charge clouds Bonds

Lone Pairs Shape Examples 4 2 2 Bent H2O Bent

104.5 for water. The two lone pairs exert a greater repelling force on the bonding electrons as it is sitting closer to the oxygen. .. .. .. .. Chemical Bonding 86

VSEPR Theory Charge clouds Bonds Lone Pairs Shape Examples 5 5 0

Trigonal Bipyramidal PCl5 Chemical Bonding 87 VSEPR Theory Charge clouds Bonds Lone Pairs

Shape Examples 5 4 1 See Saw SF4 See saw Ideally 90 and120. The lone pair occupies the equatorial position

and it exerts a greater repelling force on the bonding electrons as it is sitting closer to the central atom. .. .. Chemical Bonding 88 VSEPR Theory Charge clouds Bonds

Lone Pairs Shape Examples 5 3 2 T-shaped ICl3 ..

.. .. .. T-shaped Ideally 90 and120. The two lone pairs occupy the equatorial position and they exert a greater repelling force on the bonding electrons as they are sitting closer to the central atom. Chemical Bonding 89

VSEPR Theory Charge clouds Bonds Lone Pairs Shape Examples 5 2 3 Linear

XeF2 .. .. .. .. .. Linear Exactly 180. The three lone pairs occupy the equatorial position and they cancel each other dipoles.

.. Chemical Bonding 90 Charge clouds 6 VSEPR Theory Bonds Lone Pairs Shape 6 0

Octahedral Examples SF6 Chemical Bonding 91 Charge clouds 6 VSEPR Theory Bonds Lone Pairs Shape

5 Square pyramidal Less than 90. The lone pair exerts a greater repelling force on the bonding electrons as it is sitting closer to the central atom. 1 .. Square Pyramidal Examples

IF5 .. Chemical Bonding 92 Lone Pairs Shape Examples 4 2 Square

Planar XeF4 Square planar Exactly 90 and 180. The lone pairs cancel each other out. .. 6 Bonds .. ..

Charge clouds VSEPR Theory .. Chemical Bonding 93 Example of Molecules with More Than One Central Atom Look at each central atom on its own and

everything bonded to it is considered a terminal atom. Count the charge clouds and predict the shape around it. Chemical Bonding 94 Problem 11: Predict the shape around each carbon atom in this compound. H H C=C=C H H

Chemical Bonding Valence Bond Theory It combines Lewis theory of filling octets by sharing pairs of electrons with the electron configuration of atomic orbitals. It states that bonding occurs when atomic orbitals overlap. Example: Building BF3 with Valence Bond Theory. B 1s 2s

2p 1s 2s 2p 1s 2s 2p 1s 2s

2p F Chemical Bonding Valence Bond Theory Example: Building BF3 with Valence Bond Theory. B 1s 2s 1s 2s

2p 1s 2s 2p 1s 2s 2p F 2p

But this does not explain all three identical B-F bonds The bond angles are wrong since 1 bond is between an s orbital from B and a p orbital from F. Hybrid orbitals will explain the identical angles. Chemical Bonding Hybridization Is the morphologic mixture of 2 or more atomic orbitals 1 orbital s + 3 orbitals p = 4 hybrid sp3 orbitals 1 orbital s + 2 orbitals p = 3 hybrid sp2 orbitals 1 orbital s + 1 orbital p = 2 hybrid sp orbitals hybridization sp3 hybridization Charge clouds 4

sp2 hybridization 3 sp hybridization 2 see animation Chemical Bonding Hybridization hybrid orbitals explain bonding electrons and lone pair of electrons. all single bonds are sigma bonds.

double bonds consist of a sigma bond and a pi bond. triple bonds consist of a sigma bond and 2 pi bonds. bond results from a direct head-to-head overlap of orbitals. bond results from the side-to-side attractive forces between unhybridized p orbitals. bonds are stronger as they have shorter bond lengths and greater bond energies. Chemical Bonding Problem 12: Use the Valence Bond and Hybridization theories to explain the trigonal planar geometry at each carbon atom H H C=C

H H Chemical Bonding Chemical Bonding 101 Chemical Bonding 102 Problem 13: Use the Valence Bond and Hybridization theories to explain the linear geometry at each carbon atom HCCH

Chemical Bonding A bond results from the head-on overlap of two sp hybrid orbitals. Chemical Bonding 104 The unhybridized p orbitals form two p bonds. Note that a triple bond consists of one and two p bonds. Chemical Bonding 105

Extended Bonds Example: Benzene CH in every corner Each p-orbital can overlap with two different p-orbitals This leads to delocalization of electrons Can be used to explain resonance in Lewis Chemical Bonding structures Polar Molecules: The Influence of Molecular Geometry Molecular geometry affects molecular polarity. Due to the effect of the bond dipoles and how they either cancel or reinforce each

other. A B A linear molecule nonpolar A B A angular molecule polar Chemical Bonding 107 Polar Molecules: The Influence of

Molecular Geometry Polar Molecules must meet two requirements: 1. One polar bond or one lone pair of electrons on central atom. 2. Neither bonds nor lone pairs can be symmetrically arranged that their polarities cancel. Chemical Bonding 108 Molecular Shapes and Polarity Charge clouds

Bonding / lone pairs Shapes Polarity 2 2/0 linear Non polar * 3 3/0

trigonal planar Non polar * 3 2/1 bent Polar 4 4/0 tetrahedral

Non polar * 4 3/1 trigonal pyramidal Polar 4 2/2 bent Polar

5 5/0 trigonal bipyramidal Non polar * 5 4/1 see saw Polar 5

3/2 T shaped Polar 5 2/3 linear Non polar * 6 6/0

octahedral Non polar * 6 5/1 square pyramidal Polar 6 4/2 square planar

Non polar Chemical * Bonding 109 Problem 14: For each of the molecules below identify shape, angle, hybridization, and polarity a) H2O b) SiO2 c) SF4 Chemical Bonding

Molecular Orbital (MO) Theory In the Hybrid Orbital theory each atom in the compound retains its associated orbitals and electrons. The MO theory views a molecule as a whole instead of a collection of individual atoms. MOs are similar to atomic orbitals. Similarities with hybridization: They both have specific energy levels. They both have specific sizes an shapes. They can both hold a maximum of two electrons that spin in opposite directions. Chemical Bonding Molecular Orbital (MO) Theory Atomic orbitals combine to form MOs. When two atomic orbitals combine, two Mos are formed.

Orbitals are always conserved. Example: Formation of H2. + + + + *1s anti-bonding + 1s +

+ bonding Chemical Bonding Molecular Orbital (MO) Theory Bonding Orbital Is a MO that is lower than any atomic orbital from which it was derived. Electrons that occupy these orbitals cause stability. Anti-Bonding Orbital Is a MO that is higher than any atomic orbital from which it was derived. Electrons that occupy these orbitals cause instability. Non-Bonding Orbital Is a MO that is at the sane energy level as the one atomic orbital that it was derived from.

Electrons that occupy these orbitals do not cause stability or instability. These are orbitals that contain lone pair of electrons. Chemical Bonding Molecular Orbital (MO) Diagram H2 * 1s antibonding MO E n e r g y 1s atomic orbital on

1st H atom 1s atomic orbital on 2nd H atom 1s bonding MO Chemical Bonding Molecular Orbital (MO) Diagram O2 * 2p E n

e r g y * 2p 2p 2p 2p 2p * 2s 2s 2s 2s

Chemical Bonding Problem 15: Develop the Molecular orbital diagram for a molecules of F 2. * 2p * 2p E n e r g y 2p 2p

2p 2p * 2s 2s 2s 2s Chemical Bonding

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