Chapter 1: Carbon and Its Compounds PDF

Chapter 1 Carbon and Its Compounds Assigned Reading: Sections 1.1-1.4 Additional Sections: 5.5-5.11(Resonance)  Chapter Objectives 1. Understand and explain why hybrid orbitals are often used to describe bonding in molecules rather than pure atomic orbitals....

Chapter 1 Carbon and Its Compounds Assigned Reading: Sections 1.1-1.4 Additional Sections: 5.5-5.11(Resonance)  Chapter Objectives 1. Understand and explain why hybrid orbitals are often used to describe bonding in molecules rather than pure atomic orbitals. 2. Write hybridization schemes for the formation of sp, sp2, and sp3 ; Determine the hybridization of an atom in a molecule, radical or ion; approximate bond angles in the molecule and bond lengths. From this information, determine the shape of a molecule, and the 3D position of its component atoms. 3. Determine the presence of lone pairs or charges on atoms 4. Determine the presence, and direction of a dipole in a molecule 5. Identify the different types of bonds in compounds and recognize ionic and covalent bonds in different organic and inorganic molecules. Draw proper Lewis structure diagrams of compounds. 6. Understand the concept of delocalized electrons. 7. Know & understand the idea of resonance contributors and resonance energy; Be able to draw resonance contributors and a resonance hybrid of a structure.  INTRODUCTION Structure and bonding H H H C C H C C H C H H H H H Nature of the atom: hydrogen and hydrogen-like orbitals Bonding in molecules: covalent bonds, polar covalent bonds, ionic bonds Lewis Structures/VSEPR: bonding, shapes of molecules Valence bond theory, Molecular orbital theory Hybridization: sp3, sp2, sp  Structure of an atom  Positively charged nucleus (very dense, protons and neutrons) and small (10-15 m diameter)  Negatively charged electrons are in a cloud (distance: 10-10 m) around nucleus  Diameter of atom is about 2  10-10 m (200 picometers (pm)) [the unit angstrom (Å) is 10-10 m = 100 pm]  Atomic Number and Atomic Mass  The atomic number (Z) is the number of protons in the atom's nucleus  The mass number (A) is the number of protons plus neutrons  All the atoms of a given element have the same atomic number  Number of electrons = number of protons  Isotopes are atoms of the same element that have different numbers of neutrons and therefore different mass numbers  The atomic mass (atomic weight) of an element is the weighted average mass in atomic mass units (amu) of an element’s naturally occurring isotopes  The molecular weight of a compound is the sum of the atomic weights of all the atoms in the molecule Atomic Orbitals - where the electrons live (defined by Ψ) An electron in orbital is described by 4 quantum numbers. No two electrons can have the same 4 quantum numbers. 1) principal quantum number (n) n = 1, 2, 3,... size, energy 2) angular momentum quantum number (  )  = 0, 1, … , n-1 shape of orbitals = 0 s = 1 p  =2 d 3) magnetic quantum number (m) m = -, -+1, … , -1,  orientation of orbital 4) electron spin quantum number (ms) ms = +1/2, -1/2 orbital can hold two electrons with opposing spin  energy of the orbital shape of the orbital 1) principal quantum number (n) 2) angular momentum quantum number ( ) n = 1, 2, 3,...  = 0, 1, … , n-1 direction of the orbital 3) magnetic quantum number (m or m) m = -, -l+1, … , -1,  Subshell Number of orbitals n  Notation m in subshell 1 0 1s 0 1 2 0 2s 0 1 1 2p -1,0,1 3 3 0 3s 0 1 1 3p -1,0,1 3 2 3d -2,-1,0,1,2 5 4 0 4s 0 1 1 4p -1,0,1 3 2 4d -2,-1,0,1,2 5 3 4f -3,-2,-1,0,1,2,3 7  s orbitals All s-orbitals are spherical. As n increases, the s-orbitals get larger. As n increases, the number of nodes increase. A node is a region in space where the probability of finding an electron is zero. 1s 0 nodes For an s-orbital, the number of nodes is n - 1. 2s 1 node 3s 2 nodes  p Orbitals There are three p-orbitals px, py, and pz. The three p-orbitals lie along the x-, y- and z- axes of a Cartesian system. The letters correspond to allowed values of ml of -1, 0, and +1. The orbitals are dumbbell shaped. As n increases, the p-orbitals get larger. All p-orbitals have a node at the nucleus.  d Orbitals There are five d orbitals. Three of the d-orbitals lie in a plane bisecting the x-, y- and z- axes. Two of the d-orbitals lie in a plane aligned along the x-, y- and z-axes. Four of the d-orbitals have four lobes each. One d- orbital has two lobes and a collar. There are five 3d orbitals. Each one has two sets of nodes.  Electron Configurations  Ground-state electron configuration of an atom lists orbitals occupied by its electrons from lowest to highest energy = lowest energy arrangement. Rules:  1. Lowest-energy orbitals fill first: 1s 2s 2p 3s 3p 4s 3d [Aufbau (build-up) principle]  2. Electron spin can have only two orientations, up  and down . Only two electrons can occupy the same orbital, and they must have opposing spins (Pauli exclusion principle) to have unique wave equations  3. If two or more empty degenerate orbitals are available, electrons occupy each singly with spins parallel, until all orbitals have one electron (Hund's rule). The Periodic Table-order of orbitals Group # = # of electrons in outermost electron shell (valence shell) Principal quantum number “l” Principal quantum number “n” # of electrons in Group # subshell C: 1s2 2s2 2p2 C valence shell Ar Electron Octet = eight electrons in the valence shell-highly stable example: noble gases (Ar, Ne, Kr…)  Practice Question Which of the following atom(s) have at least one electron in the 4p subshell? 18 20 32 35 Ar Ca Ge Br 39.948 40.078 72.61 79.904 a. Ar, Ca and Ge only d. Ar only b. Ar and Br only e. Ge and Br c. Ar and Ca only 4s Ca ([Ar]4s2) [Ar] 4s 3d 4p Ge ([Ar]4s23d104p2) [Ar]  BONDING Why: Bonds forms because the molecule formed has a lower energy than the atoms apart-molecule more stable than atoms apart. How: Generally, electrons are shared or transferred so that each atoms has a filled valence shell-outermost shell of electrons.  Ionic Bonding: attraction between cations and anions Electrostatic attractive forces between oppositely charged ions, formed from electron transfer. (usually metal with non-metal) unpaired electron, unfilled valence shell filled valence shell _ _ Li F Li+ F Li+ F low I.E. high E.A. ionic bond electron transfer Covalent Bonding: Electrons shared between atoms. Usually between two non-metals  Multiple Covalent Polar Covalent Bonds Bonds Polar Covalent Bonds – electron distribution between atoms is not symmetrical H H Non-polar covalent bond (equal sharing) H F Polar covalent bond (non-equal electron sharing) F “hogs” electrons. δ+ δ- H F H F Direction of bond polarity partial partial indicated by arrow: positive negative Head = negative end charge charge Tail = positive end Electronegativity: ability of an atom in a molecule to attract electrons to itself Increasing EN O CH3 CH2 F CH3 CH3 In organic molecules, electronegative atoms are said to be electron withdrawing.  ΔEN < 2 ΔEN  2 ΔEN = 0 Electronegativity Inductive Effect: shifting of electrons in a bond by an atom in response to EN of nearby atoms Dipole Moment – measure of the separation of charge in a polar covalent bond Molecules with a dipole moment align in an electric field. δ+ δ- δ+ H δ+ H O H + δ-O δ+ H - O H H δ+ δ- O H H H O H δ+ H δ-O H δ+ Polar bonds have a “dipole” Size of dipole indicated by dipole moment Dipole Moment – measure of the separation of charge in a polar covalent bond  = Q (charge on atom)  r (distance between two charges) 1 D = 3.34  10-30 coulomb  meter O H H Net dipole (or measure dipole) is the vector sum of all of the local dipole moments. What do dipoles say in passing each other? Have you got a moment?!  Molecular Dipole Moment The vector sum of the magnitude and the direction of the individual bond dipole determines the overall dipole moment of a molecule  Practice Question Determine the dipole moments for each bond and the net dipole for each molecule shown below: F H C C F Cl F Cl F Cl Cl H Cl H C C C C Cl H H Cl  Lewis Theory: An Overview  Valence e- play a fundamental role in chemical bonding.  e- transfer leads to ionic bonds.  Sharing of e- leads to covalent bonds.  e- are transferred or shared to give each atom a noble gas configuration  the octet (ns2 np6).  Lewis Symbols  A chemical symbol represents the nucleus and the core e-.  Dots around the symbol represent valence e-. Si N P As Sb Bi Al Se I Ar HINT s-block elements: group number = # of valence electrons p-block elements: group number-10 = # of valence electrons  Lewis Structures for Ionic Compounds Writing Lewis Structures of Ionic Compounds. Write Lewis structures for the following compound: Li2O + 2- Li2O 2Li O Binary ionic compound Note the use of the “fishhook” arrow to denote a single electron movement. A “double headed” arrow means that two electrons move. Single electron Electron pair Covalent Bonding: An Introduction  Multiple Covalent Bonds O C O O C O O C O O C O N N N N N N N N  Writing Lewis Structures  All the valence e- of atoms or ions must appear-# of valence e- equal to group number for main group elements.  Negative ion: add one e- for each negative charge to count of valence electrons  Positive ion: subtract one e- for each positive charge from total count of valence electrons  Usually, the e- are paired.  Usually, each atom requires an octet.  H only requires 2 e-.  Multiple bonds may be needed.  Readily formed by C, N, O, S, and P. Strategy for Writing Lewis Structures  Rules to determine terminal and central atoms in skeletal structure  Hydrogen atoms are always terminal atoms.  Central atoms are generally those with the lowest electronegativity.  Carbon atoms are always central atoms.  Generally structures are compact and symmetrical.  Formal Charges Formal charge = number of valence electrons – (number of lone pair electrons +1/2 number of bonding electrons) Nitrogen has five valence electrons Carbon has four valence electrons Hydrogen has one valence electron and halogen has seven  Important Bond Numbers Neutral Cationic Anionic  DRAWING LEWIS DIAGRAMS: example acetic acid 1. Count and add the number of valence electrons contributed by each atom of the molecule or ion. (2 C + 2 O + 4 H) = 24 electrons for acetic acid (CH3COOH) 2. Draw skeletal structure (identify terminal and central atoms). Join atoms by SINGLE covalent bonds-subtract 2 e- for each bond from total number of e-. H H C C O O H H  DRAWING LEWIS DIAGRAMS: example acetic acid (cont’d) 3. Complete the octets of terminal atoms (H needs 2). Subtract the electrons used from the total. Then, complete the octets of the central atoms. H H H C C O O H H C C O H H H O 24 electrons – 14 = 10 electrons 24 electrons – 14 = 10 electrons 4. If any central atoms do not have full valence shell, use lone-pairs on one or more terminal atoms to form multiple covalent bonds to central atoms. H H C C O H H O  LEWIS RULES 1. OCTET RULE: In a completed Lewis Diagram, Period 2 or Period 3 atoms will have a completed octet : ( 8 electrons). Hydrogen atoms will have a “duet” : ( 2 electrons). 2. BONDS. BONDS Bonds are made by sharing a pair of electrons between two atoms. Single Bonds ( 1 shared pair ) Double Bonds ( 2 shared pairs ) Triple Bonds ( 3 shared pairs ) are all allowed in constructing a Lewis Diagram. Hydrogen, is always singly bonded. LEWIS RULES...... continued 3. ELECTRON PAIRS. PAIRS Electrons not involved in forming bonds (non-bonded or unshared electrons) are arranged in pairs. 4. CORRECTNESS. CORRECTNESS The final structure must have the correct number (total) of valence electrons.  To summarize: C is tetravalent O is divalent C C C H and halogens are monovalent  Important: Do not draw any structure with more than 4 bonds on a carbon 4 bonds C O H Br (tetravalent) (divalent) (monovalent) C Never draw 5 (or more) bonds on carbon  Exceptions to the Octet Rule a) Odd-Electron Species  Odd electron species are paramagnetic  Most common odd-electron species: free radicals: H H—C—H O—H  Exceptions to the Octet Rule b) Incomplete Octet Structures SOME ATOMS DO NOT FOLLOW THE OCTET RULE GROUP THREE ELEMENTS OFTEN FORM INCOMPLETE OCTET STRUCTURES. Boron often makes structures with.B. an incomplete octet. It can only form three bonds! BF3 = F B F F  EXPANDED OCTET STRUCTURES GROUP 5A OR 6A ELEMENTS OFTEN FORM EXPANDED-OCTET STRUCTURES Phosphorous can form up to 5 bonds P Cl Cl Cl PCl5 = P Cl Cl GROUP 5A Sulfur can form up to 6 bonds F S F F SF6 = S F F 3d orbitals GROUP 6A F are available  VALENCE SHELL OF SULFUR Sulfur can use 3d orbitals TWO BONDS 3s 3p 3d promotion (x2) SIX BONDS 3s 3p 3d both ways of bonding are common VALENCE SHELL OF SULFUR Electrostatic Potential Maps  The Shapes of Molecules H O H Bond length – distance between nuclei. Bond angle – angle between adjacent bonds  Terminology  VSEPR Theory  Electron pairs repel each other whether they are in chemical bonds (bond pairs) or unshared (lone pairs). Electron pairs assume orientations about an atom to minimize repulsions.  Electron group (domain) geometry – geometrical distribution of e- pairs- geometry of molecule  Molecular group geometry – geometrical arrangement of atomic nuclei- shape of molecule  VSEPR: Valence Shell Electron Pair Repulsion The VSEPR model provides an understanding of the shapes of molecules in terms of electrostatic repulsion between electron pairs(domains) in the valence shell. 2 e- pairs 3 e- pairs 4 e- pairs linear 180o planar 120o tetrahedral 109.5o The shape of the molecule is defined by the atoms: e.g. tetrahedral arrangement of electrons A: central atom X: covalently bonded atom E: non-bonding e- pair AX4 AX3E AX2E2 H  H   C N H O  H  H     H H  H   H H H trigonal tetrahedral angular pyramid O (bent) C N H H H H H H H Atoms and lone pairs around central atom: steric # or electron group geometry Balloon Analogy  Applying VSEPR Theory  Draw a plausible Lewis structure.  Determine the number of e- groups and identify them as bond or lone pairs.  Establish the e- group geometry.  Determine the molecular geometry.  Multiple bonds count as one group of electrons.  More than one central atom can be handled individually.  - localized model of bonding VSEPR, Lewis and hybrid orbital theory collectively called Valence Bond Theory  bond = sigma bond H H  Valence bond (VB) theory Two s-orbitals can overlap HEAD ON to form a  (sigma) bond: + H H 1s 1s  bond = sigma bond Two p-orbitals can overlap SIDEWAYS to form a  (pi) bond: H H + C C H H 2p 2p  bond = The double bonds of ethylene and other pi bond alkenes are  bonds.  In-phase overlap forms a bonding MO; out-of-phase overlap forms an antibonding MO: Note: omit MO rules, equations, symmetries, LCAO theory in Section 1.2.2F Sigma bond () is formed by end-on overlap of two p orbitals: A  bond is stronger than a  bond-WHY?? Pi bond () is formed by sideways overlap of two parallel p orbitals:  From the valence atomic orbitals of carbon… 2s 2px 2py 2pz Why not just 2 bonds for C? 1s 2s 2p CH2 is known, but very unstable with very fleeting existence:  Bonding in Methane how can we explain this? Electron promotion requires 96 kcal/mol of energy; making 4 covalent bonds releases 420 kcal/mol of energy; overall there is 324 kcal/mol increased stability. 1s 2s 2p The concept of hybridization is needed to explain the geometry of molecules.  Valence bond (VB) theory H H C H H methane has four identical bonds unhybridized 25% s 75% p tetrahedral sp3 hybridization: s+3p  sp3 hybridization: s+3p H tetrahedral C H H sp3 hybridized H Structure of Ethane C2H6 hybridized orbitals in ethene C 2s __ 2p __ __ __ C2H4 C 2s __ 2p __ __ __ C sp2 __ __ __ p __  Valence bond (VB) theory H H C C H H unhybridized 33% s 67% p trigonal planar sp2 hybridization: s+2p  sp2 hybridization: s+2p trigonal planar H H C C H H sp2 hybridized hybridized orbitals in acetylene C2H2 C 2s __ 2p __ __ __ C 2s __ 2p __ __ __ C sp __ __ p __ __  Valence bond (VB) theory unhybridized 50% s 50% p linear sp hybridization: s+p H C C H O C O H C C H O C O The two C of acetylene and the C of CO2 are sp hybridized.  Orbitals of Acetylene  Two sp hybrid orbitals from each C form sp–sp  bond  p orbitals from each C form a p –p  bond by z z z sideways overlap and py orbitals overlap similarly  ETHANE ETHYLENE ACETYLENE bond angle 109.5o 120o 180o H H H H H H C C C C H C C H H H H H Sigma bond sp 3 sp 2 sp bond length 154 pm 133 pm 120 pm bond energy 376 kJ/mol 611 kJ/mol 835 kJ/mol The double bond in ethylene is not quite twice as strong as the single bond. The C-H bond increases in strength: sp3 < sp2 < sp. C-H bond energy 420 kJ/mol 444 kJ/mol 552 kJ/mol Number of electron groups on the central atom of molecule dictates number of atomic orbitals that must be hybridized.  Hybridized orbitals in methyl cation, anion and radical CH3+: CH3.: CH3-:  Hybridization of Nitrogen and Oxygen  Elements other than C can have hybridized orbitals  H–N–H bond angle in ammonia (NH3) 107.3°; 109.5° in ammonium ion  N’s orbitals (spxpypz) hybridize to form orbitals  One sp3 orbital is occupied by two nonbonding electrons, and three sp3 orbitals have one electron each, forming bonds to H (NH3)  Hybridization of Oxygen in Water  The oxygen atom is sp3-hybridized  Oxygen has six valence-shell electrons and forms only two covalent bonds, leaving two lone pairs  The H–O–H bond angle is 104.5° (lone pairs take up more space and repel each other) Practice: Draw how the hybrid orbitals for O in water form. Bonding in Hydrogen Halides  Summary The shorter the bond, the stronger it is The greater the electron density in the region of orbital overlap, the stronger is the bond The more s character, the shorter and stronger is the bond The more s character, the larger is the bond angle  What is Resonance?? (READ Chapter 5: SECTIONS 5.5-5.11) Resonance structures are Lewis drawings that differ ONLY in the position of bonding and non-bonding electrons (not in connectivity of atoms)  RESONANCE When it is possible to draw more than one valid Lewis Diagram for a molecule or ion, that species is said to have resonance. The molecule or ion is said to be a resonance hybrid of the individual resonance forms with characteristics of all resonance forms. For species with resonance, no single Lewis Diagram will suffice to describe them correctly. In drawing resonance structures, the arrangement of the electrons changes but the nuclei do not move. RESONANCE OCCURS ONLY IN THE PI SYSTEM. Resonance occurs only between pi bonded electrons, unshared pairs, and atoms with an incomplete octet. That is, in a pi system (). Sigma bonded electrons are not involved in resonance. Sigma bonds form the backbone of the molecule. Breaking sigma bonds, breaks the molecule apart.  RESONANCE IN CARBONATE ION CO32- “contributing structures” YOU CAN DRAW THREE EQUIVALENT STRUCTURES “RESONANCE HYBRID”.. _ 2/3 RESONANCE HYBRID: composite of all resonance contributors _ O: 2/3.. real :O C molecule O :_.. 2/3 resonance hybrid of..... _.... _ : O: _ O: : O:.... _.. :O C :O.. C :O.. C :O.. :_ :O.. :_ O :.. … imaginary structures RESONANCE LOWERS THE ENERGY OF A MOLECULE OR ION RESONANCE STRUCTURES -O O -O C O C O- C O- E -O -O O N E R G Y the real molecule has lower energy REAL than any contributing MOLECULE structure would suggest Many Resonance Structures Consist of Two-Center and/or Three-Center Systems (Patterns) Two-Center Systems: Only  or lone pair electrons move! Three-Center Systems: No  bonds broken!  Three-Center Bonding: Amide Resonance Move lone-pair electrons toward the sp2 oxygen: Amide resonance responsible for the strength of amide bonds found in hair, skin, muscle, Kevlar vests, etc.  Multi-Center Allylic Resonance Composed of allylic centers: Allylic Allylic Benzylic cation:  Drawing Resonance Structures Move electrons away Charge moves - ANION from the negative charge. no new charges... :O C C :O C C.... C C C C C C Move electrons toward Charge moves - CATION the positive charge. no new charges.  GROUPS THAT ACCEPT ELECTRONS TO STABILIZE AN ANION Carbonyl Groups:.. Aldehydes, Ketones, Esters, Acids C X=Y -.. C C O:.. - C C O:.... Q Q ( Q = H, R, OR, OH ) Nitro.. - -....: O..: O C N.. C N.. + O: + O:..-.. - Cyano  GROUPS THAT DONATE ELECTRONS TO STABILIZE A CATION.. methoxy C-Y +.. + C.. O CH3 C.. O CH3 amino +.. + C N CH3 C N CH3 In drawing resonance structures, a (+) charge is Charged OK on oxygen or nitrogen as long as they retain oxygen is an octet. Conversely, this is not OK : left without + + C C O: C C O: an octet -.. NO !.. this is unacceptable.  SUMMARY ANIONS Move electrons away from the charge. CATIONS Move electrons toward the charge. NEUTRAL MOLECULES Start with an unshared pair or a double bond and look for a place to push the electrons. New charges will usually result... + - N C C N C C  Rules of Resonance 1) Resonance structures are not real. (The actual structure is a composite of all of the resonance structures.) 2) Resonance structures differ ONLY in the placement of the π-electrons and lone-pair electrons. H H Do not break sp3() bonds H C C C H H C C C H (Movement of electrons shown C C C C by curved arrows) H C H H C H H H 1) Individual resonance structures are not always equivalent. 1) More covalent bonds = more stable structure 2) Structures with atoms that have complete octet are more stable 3) Charge separation decreases stability; - charge should be on most EN atom; + charge on least EN atom  Rules of Resonance 4) The resonance hybrid is more stable than any individual resonance form-resonance leads to STABILITY. 5) All resonance structures must be valid Lewis structures and obey rules of valency. So an sp3 carbon cannot accept electrons; it already has an octet. O O O ? C - CH3 C CH3 C - CH3 H3 C C H3C C H3C C H H H 6) Any 3 atom grouping with a multiple bond has two resonance forms.  Rules of Resonance 7) Stability is decreased in contributors with: 1) Atoms that are an incomplete octet 2) Negative charge not on the most electronegative atom or positive charge not on the most electropositive atom 3) Charge separation vs a resonance form with no charge present 8) All resonance contributors must have the same net charge. 9) The greater the number of stable resonance contributors, the greater the delocalization (resonance) energy. 10) The more equivalent the different resonance contributors are, the greater the delocalization energy. Drawing all possible resonance structures Drawing all possible resonance structures  Localized vs. Delocalized Electrons allylic lone pairs of electrons on an atom adjacent to a double bond are delocalized vinylic lone pairs of electrons on an atom are localized. How do you determine localized e- vs delocalized e- ? Pyridine:  Localized vs. Delocalized Electrons How do you determine localized e- vs delocalized e- ? Imidazole:  Localized vs. Delocalized Electrons Determine the localized e- vs delocalized e- in the following molecule. Draw any relevant resonance structures.  Think, Pair, Share Question In 1996, the FDA approved the use of Ritonavir, developed by Abbott and Merck pharmaceutical companies, for the treatment of HIV and AIDS. Ritonavir is an HIV protease inhibitor, and binds the active site of the HIV protease with 10,000 times higher affinity than the natural ligand. The structure of Ritonavir is shown at the bottom: 1) Determine the hybridization, EGG and MGG of the atoms labeled in the molecule. 2) Determine the bond angles around the ritonavir atoms labeled.  Condensed Structures Lewis structures can be simplified by showing only the atoms in a molecule as a “list” and eliminating the lines showing covalent bonds These are called “condensed” structures: CH3 CH2 CH2 CH2 CH3 = pentane C-C and C-H bonds are not shown C with 3 hydrogens = CH3 C with 2 hydrogens = CH2 HONC Rule:  Skeletal or Line Structures = Each vertex and end of each line is a carbon. Each unspecified valence is a hydrogen. Rule 1 Rule 2 Rule 3 Rule 1: Carbon atoms are not shown Each line segment represents a C at each end. Rule 2: H’s are omitted, except when connected to an explicitly drawn atom (i.e. O, N, S, etc). H’s connected to C atoms are implied. Rule 3: Atoms other than C or H (example: O, N, S, etc) are shown  Key Concepts  Atomic number = number of protons; mass number = sum of protons and neutrons; isotopes = same atomic number, different mass number  Atomic orbital: region where electrons are found. Electrons assigned to orbitals according to Aufbau principle, Pauli exclusion principle and Hund’s rule.  Electronic configuration: describes orbitals occupied by atom’s electrons.  Lewis structures: used to show bonding between atoms; show lone pairs and valence electrons.  Ionic bond: transfer of electrons; covalent bond: sharing of electrons. Polar covalent bond: bond between atoms with different electronegativities; has dipole measured by dipole moment.  Valence bond orbital theory: covalent bonds are considered to form by the overlap of atomic orbitals of the bonded atoms. End-to-end overlap of orbitals leads to the formation of cylindrically symmetrical sigma ( ) bonds. Side-to-side overlap of two p orbitals produces a pi ( ) bond.  Key Concepts  Hybrid orbitals are needed to describe geometry of more complex molecules, in which simple overlap of orbitals does not describe the proper geometry. The hybridization scheme is one that produces an orientation of hybrid orbitals that matches the electron group geometry predicted by the VSEPR theory; so the number of electron groups determines how many atomic orbitals hybridize to form the hybrid orbitals:  2 electrons groups = sp hybrid  3 electron groups = sp2 hybrid  4 electron groups = sp3 hybrid  5 electron groups = sp3d hybrid  6 electron groups = sp3d2 hybrid  Single covalent bonds are  bonds. A double bond consists of one  bond and one  bond. A triple bond is one  bond and two  bonds.  In the molecular orbital theory, atomic orbitals combine to give the same number of molecular orbitals. The electrons in the original atomic orbitals are redistributed into the new molecular orbitals according to the Aufbau principle (lowest to highest energy).  Each two atomic orbitals that combine from two atoms give one bonding molecular orbital and one antibonding molecular orbital. Electron density is high in bonding molecular orbitals, which is lower in energy than the original atomic orbitals. Antibonding molecular orbitals have low electron density and are higher in energy than the atomic orbitals.