Bonding Theory Overview Lewis structure Formal Charge Bond polarity; Electrostatic Potential Map Valence bond theory vs. Molecular orbital theory VSEPR theory Polarity and Solubility

Atomic Structure A review from General Chemistry Protons (+1) and neutrons (neutral) reside in the nucleus Electrons (-1) reside outside the nucleus. High energy electrons are far away from the nucleus and others are far away Electrons at the highest shell (far away from nucleus) are the valence electrons Valence electrons are important as they participate bonding

Electrons in Atom: Atomic Orbitals Quantum physics description of electron in Atom: Wavefunctions ψ Wavefunction may have sign of (–), (+), or ZERO, depending on the location. Ψ2 describes the probability of electrons (electron density) around the nucleus. Nodal plane: Ψ2 = 0. Zero electron density. The sign of the wavefunction will be important for orbital overlapping in bonds.

Shapes of Atomic Orbitals s Orbital p Orbitals: px , py , pz d Orbitals

Covalent Bonding A covalent bond is a PAIR of electrons shared between two atoms. For example…

Polar Covalent Bonds Covalent bonds are either polar or nonpolar Nonpolar Covalent –bonded atoms share electrons evenly Polar Covalent – One of the atoms attracts electrons more than the other Electronegativity - how strongly an atom attracts shared electrons

Polar Covalent Bonds Electrons tend to shift away from lower electronegativity atoms to higher electronegativity atoms. The greater the difference in electronegativity, the more polar the bond.

Isomerism: different substances with the same formula

Common atoms in Organic molecules Organic compounds are based on hydrocarbon compounds (CxHy). Atoms that are most commonly bonded to carbon include N, O, H, S and halides (F, Cl, Br, I). With some exceptions, each element generally forms a specific number of bonds with other atoms (HONC)

Lewis Structures For simple Lewis Structures… Draw the individual atoms using dots to represent the valence electrons. Put the atoms together so they share PAIRS of electrons to make complete octets. HONC rule (#bonds: 1234): preference of atom as center

Formal Charge FORMAL charge: Atoms with an unbalanced electrons and protons Molecules with more formal charges tend to have higher energy, less stable, more reactive Electrons in the molecule: Bonding electrons (BE) vs. Lone Pair electrons (LPE) FORMAL charge = #VE – (½ #BE + #LPE)

Example: To Find Formal Charge Calculate the formal charge on each atom. or Carbon (4 single bonds): FC = 4 – ½(4x2) = 0 Hydrogen (1 single bond): FC = 1 – ½(1x2) = 0 Oxygen (1 single bond): FC = 6 – (½ x 1 x 2 + 6) = -1

Practice: Lewis Structures and Formal Charges Draw Lewis structure for N2O: two double bonds Find formal charges for all atoms in this structure

Practice: Lewis Structures and Formal Charges Draw Lewis structure for N2O: B. Single + triple bonds Find formal charges for all atoms in this structure

Valence Bond Theory Chemical bonds form when the orbitals (wave functions of electrons in the atoms) on those atoms interact (overlap) The kind of overlap depends on whether the orbitals align along the axis between the nuclei (sigma bond), or outside the axis (pi bond) Tro: Chemistry: A Molecular Approach, 2/e

H2 according to VB Theory The bond for a H2 molecule results from constructive interference Where do the bonded electrons spend most of their time?

VB theory: Hybridized Atomic Orbitals The bond angle in methane can not be explained by simple bonding between the 2s or 2p orbital of carbon with 1s orbital of hydrogen.

Hybridized Atomic Orbitals The carbon must undergo hybridization to form 4 equal atomic orbitals The atomic orbitals must be equal in energy to form four equal-energy symmetrical C-H bonds

Hybridized Atomic Orbitals Should the shape of an sp3 orbital look more like an s or more like p orbital?

Hybridized Atomic Orbitals To make CH4, the 1s atomic orbitals of four H atoms will overlap with the four sp3 hybrid atomic orbitals of C

Tro: Chemistry: A Molecular Approach, 2/e

Practice: Hybridized Atomic Orbitals Draw a picture that shows the necessary atomic orbitals and their overlap to form ethane (C2H6). Draw a picture that shows the necessary atomic orbitals and their overlap to form water.

sp2 hybridization Consider ethene (ethylene). Each carbon in ethene must bond to three other atoms, so only three hybridized atomic orbitals are needed

sp2 Hybridized Atomic Orbitals An sp2 hybridized carbon will have three equal-energy sp2 orbitals and one unhybridized p orbital

σ Bond The sp2 atomic orbitals overlap to form sigma (σ) bonds Sigma bonds provide maximum HEAD-ON overlap Free rotation along C-C single bond

π bond The unhybridized p orbitals in ethene form pi (π) bonds, SIDE-BY-SIDE overlap Rotation along C=C bond requires high energy

sp Hybridized Atomic Orbitals Consider ethyne (acetylene). Each carbon in ethyne must bond to two other atoms, so only two hybridized atomic orbitals are needed

sp Hybridized Atomic Orbitals The sp atomic orbitals overlap HEAD-ON to form sigma (σ) bonds while the unhybridized p orbitals overlap SIDE-BY-SIDE to form pi bonds Practice with SkillBuilder 1.7

Comparison: C-C, C=C, and CC bonds Note the different strengths and lengths below.

Combination of Wavefunctions A bond occurs when constructive atomic orbitals overlap. Overlapping orbitals is like wave interference. Only constructive interference results in a bond

Molecular Orbital (MO) Theory Atomic orbital wavefunctions overlap to form MOs that extend over the entire molecule. MOs can be both constructive (Bonding MO, lower energy) and destructive (Antibonding MO, higher energy) interference. The number of MOs created must be equal to the number of AOs that were used. H2 MOs

MO Theory: Examples Consider TWO of the many MOs that exist for CH3Br There are many areas of atomic orbital overlap Notice how the MOs extend over the entire molecule Each picture below represents ONE orbital. bonding MO antibonding MO

About MO bonds For each MO, up to TWO electrons can fit into the orbital. In the ground state, electrons occupy some MOs and not others, depending on the relative energy level Depending on the circumstances, we will use both MO and valence bond theory to explain phenomena

π bond in MO theory The unhybridized p orbitals in ethene form pi (π) bonds, SIDE-BY-SIDE overlap of p-orbitals giving both bonding and anti-bonding MO orbitals MO theory shows the orbitals that result. Remember, red and blue regions are all part of the same orbital

Molecular Geometry Valence shell electron pair repulsion (VSEPR theory) Valence electrons (bonded and lone pairs) repel each other To determine molecular geometry… Determine the Steric number

Molecular Geometry Valence shell electron pair repulsion (VSEPR theory) Valence electrons (bonded and lone pairs) repel each other To determine molecular geometry… Predict the hybridization of the central atom If the Steric number is 4, then it is sp3 If the Steric number is 3, then it is sp2 If the Steric number is 2, then it is sp

sp3 Geometry For any sp3 hybridized atom, the 4 valence electron pairs will form a tetrahedral electron group geometry Methane has 4 equal bonds, so the bond angles are equal lone pair electrons repel stronger, smaller bond angles More lone pair electrons causes even smaller bond angle

sp3 Geometry The molecular geometry is different from the electron group geometry. HOW?

sp2 Geometry Calculate the Steric number for BF3 Electron pairs that are located in sp2 hybridized orbitals will form a trigonal planar electron group geometry molecular geometry: _____________

sp2 Geometry Analyze the steric number, hybridization, electron group geometry and molecular geometry for this imine?

sp Geometry Analyze the Steric number, the hybridization, the electron group geometry, and the molecular geometry for the following molecules BeH2 CO2

From Hybridization to Geometry

Molecular Polarity Electronegativity Differences cause induction Induction (shifting of electrons WITHIN their orbitals) results in a dipole moment. Dipole moment = (the amount of partial charge) x (the distance the δ+ and δ- are separated)

Molecular Polarity Polarity of some other common bonds Percent ionic character of covalent bond depends on difference in electronegativity.

Molecular Polarity For molecules with multiple polar bonds, the dipole moment is the vector sum of all of the individual bond dipoles

Molecular Polarity Electrostatic potential maps are often used to give a visual depiction of polarity

The dipole moment for pentane = 0 D?

Polarity of Pentane : single bond dipole moment : locally combined dipole moment

Intermolecular Forces (IMF) Many properties such as solubility, boiling point, density, state of matter, melting point, etc. are affected by the attractions BETWEEN molecules Neutral molecules (polar and nonpolar) are attracted to one another through… Dipole-dipole interactions Hydrogen bonding Dispersion forces (a.k.a. London forces or fleeting dipole-dipole forces)

Dipole-Dipole Dipole-dipole forces result when polar molecules line up their opposite charges. Note acetone’s permanent dipole results from the difference in electronegativity between C and O The dipole-dipole attractions BETWEEN acetone molecules affects acetone’s boiling and melting points. HOW?

Dipole-Dipole Why do isobutylene and acetone have such different MP and BPs?

Hydrogen Bonding Hydrogen bonds are an especially strong type of dipole-dipole attraction Hydrogen bonds are strong because the partial + and – charges are relatively large

Hydrogen Bonding Only when a hydrogen shares electrons with a highly electronegative atom (O, N, F) will it carry a large partial positive charge The large δ+ on the H atom can attract large δ– charges on other molecules Compounds with H atoms that are capable of forming H-bonds are called protic

Practice: Hydrogen Bonding in Solvent Which of the following solvents are protic (capable of H-bonding), and which are not (aprotic)? Acetic acid Diethyl ether Methylene chloride (CH2Cl2) Dimethyl sulfoxide

Practice: Hydrogen Bonding and Boiling Point among the isomers in boiling points Dispersion force Dipole dipole force Hydrogen bonding

Hydrogen Bonding H-bonds are among the forces that cause DNA to form a double helix and some proteins to fold into an alpha-helix

London Dispersion Forces The constant random motion of the electrons in the molecule will sometimes produce an electron distribution that is NOT evenly balanced with the positive charge of the nuclei Such uneven distribution produces a temporary dipole, which can induce a temporary dipole in a neighboring molecule

London Dispersion Forces The result is a fleeting attraction between the two molecules Such fleeting attractions are generally weak. But like any weak attraction, if there are enough of them, they can add up to a lot

London Dispersion Forces The greater the surface area of a molecule, the more temporary dipole attractions are possible Consider the feet of Gecko. They have many flexible hairs on their feet that maximize surface contact The resulting London dispersion forces are strong enough to support the weight of the Gecko

Molar Mass affects Dispersion Force Molecules with more mass generally have higher boiling points Boiling point: I2 > Br2 > Cl2 > F2 CH3OH < C2H5OH

Dispersion Forces: Shape of Molecule Explain why more highly branched molecules generally have lower boiling points

Solubility The “like-dissolves-like” rule Polar compounds generally mix well with other polar compounds If the compounds mixing are all capable of H-bonding and/or strong dipole-dipole, then there is no reason why they shouldn’t mix Nonpolar compounds generally mix well with other nonpolar compounds If none of the compounds are capable of forming strong attractions, then no strong attractions would have to be broken to allow them to mix

Solubility We know it is difficult to get a polar compound (like water) to mix with a nonpolar compound (like oil) We can’t use just water to wash oil off our dirty cloths To remove nonpolar oils, grease, and dirt, we need soap

How does Detergent help cleaning? Soap molecules organize into micelles in water, which form a nonpolar interior to carry away dirt.

Additional Example Problems Give all formal charges in the structures below.

Additional Example Problems Analyze the geometry, polarity and types of intermolecular attractions for the following molecules.