CHEM 200/202

Professor Jing GuOffice: EIS-210

All emails are to be sent to:chem200@sdsu.edu

My office hours will be held on zoom on Tuesday from 9:00 to 11:00 am or by appointment (https://SDSU.zoom.us/s/

99415148959)

https://sdsu.zoom.us/s/99415148959https://sdsu.zoom.us/s/99415148959https://sdsu.zoom.us/s/99415148959https://sdsu.zoom.us/s/99415148959https://sdsu.zoom.us/s/99415148959https://sdsu.zoom.us/s/99415148959

UPCOMING IMPORTANT DATES

• Lab Report Due: Atomic Emission, Nov.8th-11:59 pm

• OWL Pre-Assignment: Al-Zn alloy, Nov. 8th-11:59 pm

• Pre-Lab: Atomic Emission, Nov. 8th-11:59 pm

• SIM: Flame Test, Nov.8th-11:59 pm

VSPER

IDEAL BOND ANGLES

BOND ANGLE DISTORTIONS

• The bond angles expected from the models (ideal) are not always meet in reality. Difference in electron group types will alter the bond angles.

• Relative size of electron groups:

• Lone pairs > multiple bonds > single bonds

• VSPER theory can predict if an actual bond will be smaller or larger than the ideal bond angle.

QUESTIONWhat is the molecular shape of NH3 and the

H-N-H bond angle?

Tetrahedral (109.5°)Trigonal pyramidal (109.5°)

Trigonal planar (120°)Trigonal pyramidal (>109.5°)Trigonal pyramidal (

MOLECULAR POLARITY• Individual, covalent bonds, are polar when there is a difference in the electronegativity of the bonded atoms (e.g. Si-O, H-F...)

• Molecules can also have polarity, however, this is a product of the aggregate polarity of all the bonds.

• Molecular polarity is influenced by both the bond polarities and the molecular shape.

• The polarity of a molecule is termed the dipole moment (µ), and is measured in units debyes.

• Non-polar molecules do not have dipole moments (µ=0)

POLAR MOLECULES• Polar molecules are influenced by electric fields; non-polar molecules are not affected by electric fields.

• Polar molecules will self-orient to the applied electric field.

Electric field ON Electric field OFF

MOLECULAR POLARITY• Molecular polarity is the result of

synergistic bond polarities.

• If the polarity of the bonds around the central atom result in an unequal distribution of electronegativity the molecule will be polar.

• The VSEPR structure of the molecule must be considered in determining molecular polarity.

Net polarity:

Net polarity:

PRACTICE QUESTION

CHAPTER 8Advanced Theories of Covalent Bonding

BONDING THEORIES

• The Lewis structures and VSPER models do a good job of depicting the arrangement of the atoms accurately around central atoms.

• But they do not explain how bonds are formed or why carbon forms four separate, but equal, bonds from its valence electrons (2s22p2).

TWO QUANTUM MECHANICAL DESCRIPTIONS OF CHEMICAL BONDING

• Valence Bonding (VB) Theory

• Describes the bonding in molecules using atomic orbitals (AOs).

• Explains bonding through localized overlap of AOs.

• Molecular Orbital (MO) Theory

• Describes the bonding in molecules using molecular orbitals.

• Explains bonding using electron orbitals delocalized over the the entire molecule.

VALENCE BOND THEORY• Basic Principles

• A covalent bond forms when the partially filled (one electron) atomic orbitals of two atoms overlap.

• The overlapping region becomes occupied by the two electrons (one from each atom).

• The two electrons (which have opposing spins) can be described by two wave functions that are in phase with each other, thus increasing the amplitude of the wave function where the orbitals overlap.

https://www.youtube.com/watch?v=vHXViZTxLXo

https://www.youtube.com/watch?v=vHXViZTxLXohttps://www.youtube.com/watch?v=vHXViZTxLXo

VALENCE BOND THEORY• Themes

• The set of overlapping orbitals has a maximum of two electrons which must have opposing spins.

• The greater the orbital overlap, the strong the bond - the bond is more stable.

ORBITAL SHAPES & OVERLAPs orbital pz orbital•Bonds will form between

all orbital types.•H-F is formed by an overlap between the 1s and 2p orbitals of H & F

ORBITAL SHAPES & OVERLAP• Though the direct overlap of atomic orbitals can

explain how the bonds are formed, it does not help explain the shape of the resulting molecules.

• The three p orbitals are at 90° from each other, so how does water have a H-O-H bond angle of 104.5°?

• How does carbon, which only has two 2 electrons in s orbital and 2 electron in p orbital, form four bonds with hydrogen to give methane, which has bond angles of 109.5°?

all p orbitals

tetrahedral

HYBRID ORBITALS• The bonds for water, methane and many other compounds

arise from the mixing of different atomic orbitals on an atom.

• The averaging of the wave functions of the overlapped orbitals give rise to new shapes, which match the molecular shapes.

• The number of overlapped orbitals is equal to the number of hybrid orbitals formed (e.g. 1 s orbital + 2 p orbitals = 2 sp orbitals).

LECTURE OBJECTIVES

• Chapter 8.2

• Determine orbital hybridization from the VSEPR molecular structures.

• Chapter 8.3

• Describe multiple bonds (i.e. double and triple bonds) in terms of atomic orbital overlap.

ORIENTATION OF HYBRID ORBITALS

Period 2 elements have small distortions from ideal angles.Period 3-6 elements have larger distortions.

180° 120° 109.5°

SP HYBRID ORBITALS• Example: BeCl2 (linear)

• Be needs to hybridize to have unpaired electrons.

• Forms two sp orbitals from one s and one p orbital.

• Chlorine does not change its orbitals.

SP2 HYBRID ORBITALS• Example: BF3 (trigonal planar)

• Boron forms three sp2 orbitals from one s orbital and two p orbitals - one p orbital remains unhybridized.

• Fluorine does not alter its orbitals.

SP3 HYBRID ORBITAL • Example: CH4 (tetrahedral)

• Carbon forms four sp3 hybrid orbitals from one s orbital and three p orbitals.

SP3D HYBRID ORBITAL • Example: PCl5 (trigonal bipyramidal)

• Phosphorus forms five sp3d hybrid orbitals from one s orbital, three p, and one d orbitals.

SP3D2 HYBRID ORBITAL • Example: SF6 (octahedral)

• Phosphorus forms six sp3d2 hybrid orbitals from one s orbital, three p, and two d orbitals.

LONE PAIRS & HYBRID ORBITALS• Not all hybrid orbitals need to be bonding pairs of electrons

- some are filled with a lone pair of electrons (non-bonding).

ELECTRON GROUP ARRANGEMENT & HYBRIDIZATION

• The hybridization of the orbitals is most easily identified by the VSPER model of the molecule.

• Remember to account for lone pairs along with the bonding electrons.

Figure 8.21 The shapes of hybridized orbital sets are consistent with the electron-pair geometries. For example, anatom surrounded by three regions of electron density is sp2 hybridized, and the three sp2 orbitals are arranged in atrigonal planar fashion.

It is important to remember that hybridization was devised to rationalize experimentally observed moleculargeometries. The model works well for molecules containing small central atoms, in which the valence electron pairsare close together in space. However, for larger central atoms, the valence-shell electron pairs are farther from thenucleus, and there are fewer repulsions. Their compounds exhibit structures that are often not consistent with VSEPRtheory, and hybridized orbitals are not necessary to explain the observed data. For example, we have discussed theH–O–H bond angle in H2O, 104.5°, which is more consistent with sp3 hybrid orbitals (109.5°) on the central atomthan with 2p orbitals (90°). Sulfur is in the same group as oxygen, and H2S has a similar Lewis structure. However,it has a much smaller bond angle (92.1°), which indicates much less hybridization on sulfur than oxygen. Continuingdown the group, tellurium is even larger than sulfur, and for H2Te, the observed bond angle (90°) is consistent withoverlap of the 5p orbitals, without invoking hybridization. We invoke hybridization where it is necessary to explainthe observed structures.

416 Chapter 8 | Advanced Theories of Covalent Bonding

This OpenStax book is available for free at http://cnx.org/content/col11760/1.9

PROBLEMWhat is the hybridization of carbon in

each of the following compounds?

Compound Electron Group Hybridization

CH3Cl Tetrahedral sp3

CO32- Trigonal Planar sp2

CO2 Linear sp

QUESTIONIn the molecule below, which central atom has an

orbital hybridization of sp3?

OH •••• N•• C O••••

Answer:A - OxygenB - NitrogenC - Carbon

LECTURE OBJECTIVES

• Chapter 8.2

• Determine orbital hybridization from the VSEPR molecular structures.

• Chapter 8.3

• Describe multiple bonds (i.e. double and triple bonds) in terms of atomic orbital overlap.

Answer: C

Answer:B

Answer:C

Answer:D

SIGMA BONDS• Sigma (σ) bonds are the result

of direct orbital overlap.

• Any two orbitals that overlap can form σ bonds (e.g. s-p, p-p, p-d, s-spx,…)

• σ bonds have very high electron density along the axis of the bond.

DOUBLE BONDS

• If bonds arise from the overlap of atomic and/or hybrid orbitals, how are double and triple bonds formed?

• What orbitals are interacting and how do they interact for form multiple bonds?

C C

H

HH

H

MULTIPLE BONDS• The heart of every multiple bond is a σ bond.

• The additional “bonds” in a multiple bond arise from pi (π) bonding.

• π bonds arise from two single (unpaired) electrons in p orbitals on atoms bound by a σ bond.

•π-bonds are formed by the side-to-side overlap of partially filled (one electron) p atomic orbitals on adjacent atoms.

• They are typically not as strong as corresponding σ-bonds - π-bonds overlap above and below the bond axis.

•π-bonds react differently than σ-bonds.

π-Bonds

ETHYLENE (C2H4) BONDING• Carbon hybridization: sp2

• Bonds: 4 single bonds (4×σ) and 1 double bond (1×σ + 1×π)

Basic structure

σ-bonding

Electron density

Basic structure

σ-bonding

π-bonding

σ Bonds π Bonds

Electron density

MULTIPLE PI BONDS• Acetylene forms a triple bond, with a single σ-bond and two π-bonds.

Single Bond σ-bond

Bond Bond Types Hybridization Shape

Double Bond

Triple Bond

σ-bond & π-bond

σ-bond & 2 π-bonds

sp3

sp2

sp

QUESTIONWhich hybrid orbital, and what number, and

type of bonds are used by all the central atoms in C2F4?

sp2; 4 σ & 1 π bondssp2; 4 σ & 2 π bondssp2; 5 σ & 1 π bondssp3; 4 σ & 1 π bondssp3; 4 σ & 2 π bonds

Answer:ABCDE

LIMITATIONS OF VALENCE BOND THEORY

• Does not adequately explain the magnetic and spectral properties of molecules.

• Understates the importance of delocalized electrons (resonance structures).

• Does not provide a satisfactory explanation of the bonding in hypervalent molecules.

The paramagnetic properties of O2

Lycopene

The color of tomatoes.

LIMITATIONS OF VALENCE BOND THEORY

https://www.youtube.com/watch?v=KcGEev8qulA