Chapter 6 Section 5 Molecular Geometry Modern Chemistry Chemical Bonds Chapter 6 Section 5 Molecular Geometry Modern Chemistry

6.5 Chemical Bonding & Molecular Geometry Molecular geometry (shape) is integral to molecular function. Think back to what you learned about enzymes and enzyme-substrate interactions in your biology class If a molecule didn’t have the proper shape it couldn’t do the proper job Trying to understand what influences chemical shape and, therefore, chemical behavior is an important goal in chemistry

6.5 Chemical Bonding & Molecular Geometry What factors, including shape, help determine the chemical and physical properties of molecular compounds? Type of bonding between atoms Three dimensional arrangement of a molecule’s atoms in space (= shape) Molecular polarity

6.5 Chemical Bonding & Molecular Geometry What theories or models are used to explain molecular shape? Valence Shell Electron Pair Repulsion Theory helps explain the bond angles between atoms Hybrid Orbital Theory describes the orbitals that contain the valence electrons of a molecule’s atoms

6.5 Chemical Bonding & Molecular Geometry Objectives Explain VSEPR theory Predict the shapes of molecules or polyatomic ions using VSEPR theory Explain how the shapes of molecules are accounted for by hybridization theory Cont’d next slide …

6.5 Chemical Bonding & Molecular Geometry Objectives cont’d Describe dipole-dipole forces, hydrogen bonding, induced dipoles, and London dispersion forces and their effects on properties such as boiling and melting points Explain what determines molecular polarity

10.1 VSEPR Model The valence shell electron pair repulsion (VSEPR) model predicts the shapes of molecules from their Lewis structures The main premise of the model is that electron pairs about an atom repel each other A molecule has a shape that minimizes electrostatic repulsions between valence shell electron pairs Minimum repulsion results when the electron pairs are as far apart as possible

6.5 Chemical Bonding & Molecular Geometry The following slides show molecules that have a central atom surrounded by either two, three, four, five, or six other atoms and their associated geometries/molecular shapes & bond angles Please note that central atoms with five or six bonded atoms are exceptions to the octet rule The only central atoms that you might see with five or six bonded atoms are phosphorous (P) or sulfur (S) Note that in the notation used in the following slides A = the central atom and B = the # of atoms bonded to the central atom

6.5 Chemical Bonding & Molecular Geometry Central atom with two bonded atoms = AB2 Shape Linear BeCl2 Central atom with three bonded atoms = AB3 Trigonal Planar BF3 Central atom with four bonded atoms = AB4 Tetrahedral CH4

6.5 Chemical Bonding & Molecular Geometry Central atom with five bonded atoms = AB5 Trigonal Bipyramidal Examples PF5 Central atom with six bonded atoms = AB6 Octahedral SF6

6.5 Chemical Bonding & Molecular Geometry The Lewis dot structures of some molecular compounds, indicate the presence of lone pairs of electrons in addition to bonded atoms Ex: H2O NH3 H O: H N H H H Shorthand Notation: AB2E2 AB3E a lone pair of electrons : :

6.5 Chemical Bonding & Molecular Geometry Note that in the shorthand notation, lone pairs are denoted by the letter E and the subscript tells how many lone pairs are present Please note - only the lone pairs around the CENTRAL atom are counted! When lone pairs of electrons are present, they influence the shape of the molecule although they are not part of the shape

6.5 Chemical Bonding & Molecular Geometry Therefore, water (H2O) is not linear as would be expected from an AB2 molecule. Rather it is bent in shape since it is an AB2E2 molecule Bent 2 bonds, 2 lone pairs

6.5 Chemical Bonding & Molecular Geometry Similarly, ammonia (NH3) is not trigonal planar as would be expected from an AB3 molecule. Rather it is trigonal pyramidal in shape since it is an AB3E molecule Trigonal pyramidal 3 bonds, 1 lone pair

6.5 Chemical Bonding & Molecular Geometry Lewis Dot Structure Shorthand Notation Molecular Geometry (Shape) Approximate Bond Angles Example AB2 linear 180° BeCl2 AB3 trigonal planar 120° BH3 AB2E bent or V-shaped 115° O3 AB4 tetrahedral 109° CH4 AB3E trigonal pyramidal 107° NH3 AB2E2 105° H2O AB5 trigonal bipyramidal 180°/120°/90° PCl5 AB6 octahedral 180°/90° SF6

6.5 Chemical Bonding & Molecular Geometry Molecular Polarity Molecular geometry (shape) can influence whether a molecule is classified as polar Polar molecules contain an unequal distribution of charge Generally: polar molecules interact with other polar molecules, and nonpolar molecules react with other nonpolar molecules Ex: sugar dissolves in water/oil does not

6.5 Chemical Bonding & Molecular Geometry Molecular Polarity Whether a molecule is polar or nonpolar depends on the difference in electronegativities between bonded atoms, and the overall shape of the molecule The electronegativity differences or dipoles for each bonded pair can be represented by arrows pointing toward the more electronegative atom. Longer arrows represent more polar bonds H F H Br

Electronegativity: The ability of an atom in a molecule to attract shared electrons to itself.

6.5 Chemical Bonding & Molecular Geometry Molecular Polarity The dipoles in CO2 cancel because the linear shape orients the equal magnitude bond dipoles in exactly opposite directions. The CO2 molecule is, therefore, nonpolar

6.5 Chemical Bonding & Molecular Geometry Molecular Polarity The bond dipoles do not cancel in COSe; they are oriented in the same direction and are of unequal length. They do not cancel in OF2 because the V-shape of the molecule does not orient them in opposite directions.

6.5 Chemical Bonding & Molecular Geometry Molecular Polarity The bond dipoles in BCl3 and CCl4 cancel because of the regular shape and equal magnitude.

6.5 Chemical Bonding & Molecular Geometry Molecular Polarity The bond dipoles in BCl2F and CHCl3 do not cancel because they are not of the same magnitude.

6.5 Chemical Bonding & Molecular Geometry Molecular Polarity Based on the given examples: Note that a molecule with polar bonds can be nonpolar if the geometry causes the bond polarities to sum to zero. Note that molecules are nonpolar when there are no lone pairs on the central atom and all of the atoms bonded to the central atom are identical Note that molecules with lone pairs of electrons on a central atom are generally polar, but there are exceptions

6.5 Chemical Bonding & Molecular Geometry Molecular Polarity Some problem solving….. Predict which of the following molecules are polar and which are non-polar after drawing their Lewis dot structures. HCN PF3 SiBr4 1) polar 2) polar 3) nonpolar

6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces In addition to the chemical bonds that exist within a molecule, there are also interactions between molecules The interactions or “forces of attraction” between molecules are called intermolecular forces (IMFs) and include Hydrogen Bonding Dipole-Dipole Forces London Dispersion Forces

6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces The number and strength of the intermolecular forces affect the properties of the substance. The stronger the IMF’s are, the higher the melting and boiling points of a given substance The stronger the IMF’s, the more energy is required to melt, evaporate or boil. IMF’s are broken to go from solid  liquid and from liquid  gas. Breaking IMF’s requires energy.

6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces To highlight the difference between intramolecular bonds (within molecules) and intermolecular bonds (between molecules) let’s consider what happens when water evaporates? Which bonds remain intact? Which bonds are broken?

6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces The function of many biologically important molecules is determined by their intermolecular forces as well (think H-bonding in DNA) Although there are several different types of intermolecular attractions, all intermolecular attractions depend on electrostatic interactions, the attraction between charges of opposite sign

6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces/London Dispersion Forces Electrons move around the nuclei. They could momentarily all “gang up” on one side This lop-sidedness of electrons creates a partial negative charge in one area and a partial positive charge in another. All molecules have electrons. + Positively charged nucleus - Negatively charged electron + - + - Electrons are fairly evenly dispersed. + - As electrons move, they “gang up” on one side.

6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces/London Dispersion Forces The instantaneous “lopsidedness” of electron density can induce a near neighbor to behave similarly.

6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces/London Dispersion Forces All molecules have electrons…all molecules can have London Dispersion Forces The more electrons that gang-up, the larger the partial negative charge. The larger the molecule, the stronger the London Dispersion Forces Larger molecules have more electrons Larger molecules have stronger London Dispersion Forces than smaller molecules.

6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces/London Dispersion Forces Electrons can gang-up and cause a non-polar molecule to be temporarily polar The electrons will move again, returning the molecule back to non-polar The polarity was temporary, therefore the molecule cannot always form LDF. London Dispersion Forces are the weakest of the intermolecular forces because molecules can’t form them all the time.

6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces / Dipole-Dipole Forces Dipole-dipole attractions exist due to the partial positive and negative charges on atoms produced by unequal sharing of electrons due to electronegativity differences in polar molecules The larger the dipole moment of a molecule, the stronger the dipole-dipole attraction. - +

6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces / Dipole-Dipole Forces Polar molecules always have a partial separation of charge. Polar molecules always have the ability to form attractions with opposite charges Dipole forces are stronger than London Dispersion Forces

6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces / Hydrogen Bonding Hydrogen bonding is an extreme example of a dipole-dipole interaction Hydrogen bonding occurs between a hydrogen atom that is bonded to a small, highly electronegative atoms such as nitrogen, oxygen or fluorine and a second nitrogen, oxygen or fluorine on a neighboring atom

6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces / Hydrogen Bonding Many unusual properties of water can be explained due to extensive hydrogen bonding Ice is less dense than liquid water and therefore floats (most solids are denser than their corresponding liquids) Water has a high specific heat that modulates Earth’s daytime and nighttime temperatures Water has a high heat of vaporization that allows for evaporative cooling through sweating Water is a versatile solvent Water can travel by capillary action against gravity

6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces The types of bonding forces vary in their strength as measured by average bond energy. Ionic or (>200 kcal/mol) Covalent bonds Strongest to Weakest Hydrogen bonding (12-16 kcal/mol ) Dipole-dipole interactions (2-0.5 kcal/mol) London forces (less than 1 kcal/mol)

6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces Although intermolecular forces such as hydrogen-bonding, dipole-dipole interactions, and London dispersion forces are much weaker than ionic or covalent bonds, they are still important forces that explain much chemical behavior

6.5 Chemical Bonding & Molecular Geometry Hybrid Orbital Theory Valence shell electron pair repulsion theory explains much about molecular geometry but not everything… Let’s look at bonding in a molecule of CH4 to explain why hybrid orbital theory was developed

What Proof Exists for Hybridization? We have studied electron configuration notation and the sharing of electrons in the formation of covalent bonds. Lets look at a molecule of methane, CH4. Methane is a simple natural gas. Its molecule has a carbon atom at the center with four hydrogen atoms covalently bonded around it.

Carbon ground state configuration What is the expected orbital notation of carbon in its ground state? Can you see a problem with this? (Hint: How many unpaired electrons does this carbon atom have available for bonding?)

Carbon’s Bonding Problem You should conclude that carbon only has TWO electrons available for bonding. That is not enough! How does carbon overcome this problem so that it may form four bonds?

Carbon’s Empty Orbital The first thought that chemists had was that carbon promotes one of its 2s electrons… …to the empty 2p orbital.

However, they quickly recognized a problem with such an arrangement… Three of the carbon-hydrogen bonds would involve an electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a hydrogen atom.

This would mean that three of the bonds in a methane molecule would be identical, because they would involve electron pairs of equal energy. But what about the fourth bond…?

The fourth bond is between a 2s electron from the carbon and the lone 1s hydrogen electron. Such a bond would have slightly less energy than the other bonds in a methane molecule.

This bond would be slightly different in character than the other three bonds in methane. This difference would be measurable to a chemist by determining the bond length and bond energy. But is this what they observe?

The simple answer is, “No”. Measurements show that all four bonds in methane are equal. Thus, we need a new explanation for the bonding in methane. Chemists have proposed an explanation – they call it Hybridization. Hybridization is the combining of two or more orbitals of nearly equal energy within the same atom into orbitals of equal energy.

Hybridization - The Blending of Orbitals + = Poodle + Cocker Spaniel = Cockapoo + = s orbital + p orbital = sp orbital

In the case of methane, they call the hybridization sp3, meaning that an s orbital is combined with three p orbitals to create four equal hybrid orbitals. These new orbitals have slightly MORE energy than the 2s orbital… … and slightly LESS energy than the 2p orbitals.

sp3 Hybrid Orbitals Here are some other ways to look at the sp3 hybridization and energy profile…

sp Hybrid Orbitals While sp3 is the hybridization observed in methane, there are other types of hybridization that atoms undergo. These include sp hybridization, in which one s orbital combines with a single p orbital. This produces two hybrid orbitals, while leaving two normal p orbitals

sp2 Hybrid Orbitals Another hybrid is the sp2, which combines two orbitals from a p sublevel with one orbital from an s sublevel. One p orbital remains unchanged.

Summary of Hybridization and Molecular Geometry Forms Overall Structure Hybridization of “A” AX2 Linear sp AX3, AX2E Trigonal Planar sp2 AX4, AX3E, AX2E2 Tetrahedral sp3 A = central atom X = atoms bonded to A E = nonbonding electron pairs on A