Module 1: Electronic Structure and Bonding - PDF

Module 1: Electronic Structure and Bonding Let’s review some basics about the structure of an atom: of neutrons determine Atoms Made up of isotope protons neutrons salactrons Protons are...

Module 1: Electronic Structure and Bonding Let’s review some basics about the structure of an atom: of neutrons determine Atoms Made up of isotope protons neutrons salactrons Protons are positively charged. rocaf Neutrons have no charge. Electrons are negatively charged. 39rad nucleus Atomic number = # of protons Atomic number of carbon = 6 6protons 6neutrons opposites weight attract of c determines of protons gives identity charge  Review of Atomic Orbitals While protons and neutrons are found in an atom’s nucleus, electrons are found outside the nucleus in atomic orbitals. Atomic orbitals exist in concentric shells that sit further and further from the nucleus. The farther the orbital is from the nucleus, the higher in energy it is. Each row of the periodic table is a new orbital shell. New energyaval The position on the periodic table controls the number and type of orbitals an atom has. First row elements (H, He): Second row elements (Li, Be, B, C, N, O, F, Ne): § One 1s orbital ᵗEa § One 1s orbital 2P degenerat § One 2s orbital 29 and § Three 2p orbitals is energa 009Eur NOT nucleus  Review of Electron Configuration and Valence Shell An atom’s electron configuration describes which orbitals have electrons. Neutral atoms have the same number of electrons as protons. Carbon 19225228 When filling orbitals, follow three key rules: 2p 1 1 Aufbau principle: fill the lowest energy orbital first 2s 1L Pauli exclusion principle: only 2 electrons per orbital w we oppositespin 1s V Hund’s rule: put one electron into each degenerate orbital before pairing in outer Electron configurations of the first 10 elements valancee c Mostshall le Lc The most important 30 electrons are those 45 in the valence shell. be be  Review of Electron Configuration and Valence Shell We can also use the periodic table to help us find the valence electrons of an atom. Count from left to find of valence e lve Nove Cl Te Li B 3vc Obve C 4 ve N 5 v0  Bonding and the Octet Rule: Ionic bonding 1 strive to achieve a full valence shell. This is done through transferring or sharing LeeAtoms electrons to make chemical bonds. For atoms of the second period, a full valence shell has 8 electrons (an octet). There are two types of chemical bonding – Ionic bonding and Covalent bonding Ionic Bonding electrostatic attraction between charged ions requires full transfer of an electron no sharing ftp.t s occurs between elements with large differences in electronegativity forms a salt Far right far left elements NOT a molecule Atoms on the left-side of the periodic table easily lose an electron. becomes a cation t charged Atoms on the right side of the periodic table easily gain an electron becomes an anion charge Needs 1 mode directional  dicco a Bonding and the Octet Rule: Covalent bonding bonding For atoms that are closer in electronegativity, electrons are shared to fill their valence shells, rather than a full transfer of an electron. Covalent Bonding the sharing of electrons between two nuclei to complete octet creates discrete bonds rather than a crystal lattice of both form molecules with a distinct shape atoms Examples: Cl2 N2 Ieds 3more CH4 C 4ve H ve Ve needs 1 Hot more N N c c Polar Covalent Bonding the unequal sharing of electrons between two nuclei the bond is created between two atoms with a large SFiser electronegativity difference 919m a polar covalent bond creates a bond dipole Examples: HF H2O g 87 dipolearrow pointstowards Fst  al Covalent Bonds and Organic Structures Each element will generally form a predictable number of covalent bonds when incorporated into a molecule How many covalent bonds does an atom typically make in a neutral molecule? Atoms with 1–47valence electrons: volonapairs have lone pairs Atoms with 5 or more valence electrons: bonds of Va of bonds 8 OF Vc Ex C 40 4bonds Ex N 5 ve 8 5 3bonds B 30 3 bonds Important elements to remember: Neutral H N O C abonds bonds 4bonds 3 bonds 2 lonepairs 1 lonepair F Cl Br I Tue Halogens IMPORTANT RULE: I bond 3 lone pairs and period elements C's row NEVER exceeds 80 in valence shell Nevermorethan 4 shells  Covalent Bonds and Formal Charge LCC 3 When an atom deviates from its predicted number of bonds, a formal charge is created Formal Charge: The charge we assign to an atom s are assuming e shared caually Formal charge = valence electrons – (# of bonds + # of unshared electrons) Should does Examples: eeitipecte H ttiifi.ir 6 3 2 1 ononas 6 5 elf expected se s of bonds Noform charge of iii isE  Drawing Lewis Structures Lee 4 Given a molecular formula, we need to know where the electrons are and how the atoms are connected – this can be represented in Lewis Structures How to draw Lewis structures: CH2O 12Va total 1. Determine valence electrons, keeping overall charge in mind (add an electron for a negative charge, subtract an electron for a positive charge). Place atom that makes most bond in middle 2. Draw a skeleton that places atoms that form more than one bond in the center. Add hydrogen and halogens around the outside. Use electrons to form bonds between the atoms. Avoid O–O bonds. 0 0 bonds as reactive unstable Avoid and period claments Follow octet rule He  Drawing Lewis Structures 3. Add lone pairs to complete octets, starting with the most electronegative atoms, until you run out of valence electrons. 4. If there are atoms that do not have full octets, fill the remaining octets by making multiple bonds (using lone pairs of electrons). 5. Add any formal charges, if required #1 rule of Lewis structures: 2nd period elements cannot have more than eight electrons  Practice with Lewis Structures Let’s find the Lewis structure for the molecules below: CH4O If.EE If ao CH3N EEEIE H it full octets ÑH mnt Ytif a C one Eore  Beyond Lewis Structures: Kekule and Condensed Structures Lec 5 There are multiple ways to draw chemical structures beyond Lewis structures. We will discuss two now and a third in module 3. Kekule Structures: Kekule structures are very similar to Lewis structures with the only difference being lone pairs are not drawn in. of lone Formal 6 2 27 pairs implied by H H charge Lewis: H C O H Kekule: H C O H H H 2 x lonepairs Condensed Structures: Condensed structures are quite different in that the bonds are not drawn in, only the atoms are shown. You will need to remember the bonding patterns of atoms to interpret the structures correctly. Atoms connected to each carbon are written beside in a CXn pattern. Double and triple bonds are implied to fulfill octets. impied S abotnatt.ee atom Ebona CH3CHBrCH3 CH3(CH2)2CHO (CH3)2CHCH2OH CH3CO2H it Iiiii a ff ftp.iie.tt aft  How Atoms Form Covalent Bonds How do atoms form covalent bonds? Let’s use H2 as an example. Each atom has 1 electron in a 1s orbital. A bond forms when the two orbitals physically overlap in space. When this happens, each electron has access to both nuclei. This creates a sigma (s) bond H H H H H–H bring atomic overlapping orbitals orbitals together share electrons  Bonding and Shape of Molecules: Orbital Hybridization LCC 6 The covalent bonding of organic molecules will determine the shape a molecule takes; therefore, our theory of bonding must explain all experimental observations. Let’s consider the simplest organic molecule: methane, CH4 Methane has four identical bonds to hydrogen but does not have four identical orbitals ofPg M inplane fg H C Going H H inty from H coming of Tetraggade by How do we create four identical bonds from s and p orbitals? Orbital Hybridization tetrahedral creates 109 go bond fp stop 4 sp 80 go penor ANGIA app 4 sp3 60 08 orbitalhightailed The of orbitalsmixedMustequalthe of  Bonding and Shape of Molecules: Orbital Hybridization Let’s consider a molecule with a C–C bond: ethane, CH3CH3 Ethane carbonis 9tl Tnf5fized bond sp H H H 109.91 C C H 1s H 983 H sp obond andon strong overlap AllowForrotation aroundC Cbond bond make 1 f m s's'll nona to C Video 6 knowledge Check 1 Findallthesp3 C atoms in themolecule Hsc en  Bonding and Shape of Molecules: Orbital Hybridization LCC 7 Now let’s consider the bonding in a molecule with a C–C double bond: ethylene, C2H4 s aretrigonal Planar H H All 3 atoms 22900 aroundcentralatom C C areinthesame 120 spa spa off Plane H H overlap 98 E How does this differ from our previous examples? Carbon only has 3atomssurroundingit What is the hybridization of the carbons above? unhybridized P orbital p Flap 3 5 3 sp2 58 Tibondcannot A C–C double bond consists of: rotate Tipondis sideonoverlap not Fulloverlap of Porbital likeIndia weakerbond  Bonding and Shape of Molecules: Orbital Hybridization Finally, let’s consider a molecule with a C–C triple bond: acetylene, C2H2 PP bonded only a Needs 8 15 spsp H C C H 290 1A Geometry Linear 180 bondangle p psideon What is the hybridization of the carbons above? overlap 180 oppndto bondto c 5 5hsit on oppositeside of C atom A C–C triple bond consists of: 8ntyh.gr 1 o bond 2 ITbond video 7 knowledge check Findall the spa ssp carbon atoms inthemoleculebelow cÉn He c É d  obond Bonding and Shape of Molecules: Atoms with Lone Pairs Lec 8 Lonepairs stillneed orbitals Let’s consider some atoms with lone pairs: water and ammonia spacearound them H All 3 atomshav 109.97493 C 0PH Tetrahedral N O 3 4 things H H around central H H 4109.50 Atom H Tag ammonia H water H methane Ammonia (NH3): Water (H2O): 20 bonds 2 lonepairs Needs 4hybrid need 4 orbitals 30 bonds S1 lonepair orbitals nonfirs 11 mn.ae P mar 11111 4 sp3 4 503 1L It 5 count o bonds atoms lonepairs  Orbital Hybridization: A Summary Hybridization rules: P2 SP 5058 p3 cannot create or destroy 1. # of atomic orbitals in = # of hybridized orbitals out orbitals 2. All hybridized orbitals are identical 3. Not all atomic orbitals have to hybridize still there if NOT hybridized becomes P orbitals for Ti bonds # of “things” Hybridization Bond Angle Geometry around an atom 4 sp3 4 sp3 109.5 tetrahedral 3 sp2 3 sp2 P 120 trigonal planar 2 sp 2 Sp 2 P 180 linear  1 14 Summary Hybridization bond lengths strengths angles All H bonds are E.IE All doublebonds cec All triplebondsare single o bonds are one o bond one Ti bond one s oneTI bond o bond To determine hybridization of C N or 0 it's easiest to count the OF TT bonds IF no Ti bonds are formed it's sep3 hybridization IF 1 Tibond is formed it's Sp2 hybridization IF 2 TY bonds are Formed it's sp hybridization Carbocations carbonradicals are the exception which are spahybridization w o Ti bonds because it has partiallyFilled or empty p orbital CH3 CH2 a COOCH CAHIN in gifted Hiiiii  Bonding and Shape of Molecules: Orbital Hybridization Determine the hybridization and geometry of each of the following: w HCN CH2O H 2 things Lonepairin 3 things SP H IN orbital linearSP t 2 things a + CH3 –CH 3 E.IE QEiingsn 1 spa Trigonal things Planar it Sp3 Tetrahedral In module 1, we will review some concepts from general chemistry involving electron structures of atoms and bonding to make molecules. Speci cally, we will discuss the following topics with a perspective towards organic molecules: atoms and electron con gurations ionic, covalent and polar covalent bonding Lewis structures hybridization and hybrid orbitals Module 1 Learning Objectives After learning the material in this chapter you should be able to: Determine the electron con guration for an element Determine the number of valence electrons for an atom Identify an ionic or covalent bond Identify bond dipoles and predict the direction of the dipole with a dipole arrow or partial charges Draw a Lewis structure for an organic molecule, including formal charges and lone pairs Determine the formal charge or the number of lone pairs on an atom Interpret condensed structures of molecules State the hybridization of atoms in organic molecules Predict molecular shapes and bond angles based on hybridization Describe the bonding in a molecule in terms of the overlapping orbitals and the types of orbitals holding lone pairs