Chapter 6 Molecular Geometry Section 5 Molecular Geometry Chapter 6 Molecular Geometry The properties of molecules depend not only on the bonding of atoms but also on molecular geometry: the three-dimensional arrangement of a molecule’s atoms. The polarity of each bond, along with the geometry of the molecule, determines molecular polarity, or the uneven distribution of molecular shape. Molecular polarity strongly influences the forces that act between molecules in liquids and solids. A chemical formula, by itself, reveals little information about a molecule’s geometry.

Section 5 Molecular Geometry Chapter 6 VSEPR Theory As shown at right, diatomic molecules, like those of (a) hydrogen, H2, and (b) hydrogen chloride, HCl, can only be linear because they consist of only two atoms. To predict the geometries of more-complicated molecules, one must consider the locations of all electron pairs surrounding the bonding atoms. This is the basis of VSEPR theory.

Section 5 Molecular Geometry Chapter 6 VSEPR Theory The abbreviation VSEPR (say it “VES-pur”) stands for “valence-shell electron-pair repulsion.” VSEPR theory states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible. example: BeF2 The central beryllium atom is surrounded by only the two electron pairs it shares with the fluorine atoms. According to VSEPR, the shared pairs will be as far away from each other as possible, so the bonds to fluorine will be 180° apart from each other. The molecule will therefore be linear:

VSEPR Theory, continued Section 5 Molecular Geometry Chapter 6 VSEPR Theory, continued VSEPR theory can also account for the geometries of molecules with unshared electron pairs. examples: ammonia, NH3, and water, H2O. The Lewis structure of ammonia shows that the central nitrogen atom has an unshared electron pair: VSEPR theory postulates that the lone pair occupies space around the nitrogen atom just as the bonding pairs do.

VSEPR Theory, continued Section 5 Molecular Geometry Chapter 6 VSEPR Theory, continued Sample Problem E Solution, continued Boron trichloride is an AB3 type of molecule. Its geometry should therefore be trigonal-planar.

VSEPR Theory, continued Section 5 Molecular Geometry Chapter 6 VSEPR Theory, continued Unshared electron pairs repel other electron pairs more strongly than bonding pairs do. This is why the bond angles in ammonia and water are somewhat less than the 109.5° bond angles of a perfectly tetrahedral molecule.

VSEPR and Molecular Geometry Section 5 Molecular Geometry Chapter 6 VSEPR and Molecular Geometry

VSEPR and Molecular Geometry Section 5 Molecular Geometry Chapter 6 VSEPR and Molecular Geometry

Polar vs. Non-Polar Molecules Polarity in a molecules determines whether or not electrons in that molecule are shared equally. When determining the polarity of a molecule, it is all about symmetry. Asymmetric molecules tend to be polar. Symmetric molecules are always non-polar.

When determining the polarity of a molecule, follow these steps: Draw the Electron Dot structure of the molecule. Using the electronegativity chart determine the difference in electronegativity for each bond. 0—0.4 = Non-polar 0.5—1.7 = Polar

What is electronegativity? Why are the noble gases not included?

Nonpolar Molecules BF3 Dipole moments are symmetrical and cancel out.

The molecule is non-polar if : each bond in the molecule is non-polar and there are no unbonded electron pairs. each bond in the molecule has the same polarity and there are no unbonded electron pairs on the central atom. There is no net dipole moment (all moments cancel out)

Polar Molecules H2O Dipole moments are asymmetrical and don’t cancel . Molecule has a net dipole moment. H2O H O net dipole moment

The molecule is polar if: There is a net dipole moment each bond in the molecule is non-polar, but there are unbonded electron pairs on the central atom. bonds in the molecule have different polarities and/or there are unbonded electron pairs on the central atom.

Polar Bonds vs. Polar Molecules The effect of polar bonds on the polarity of the entire molecule depends on the molecule shape carbon dioxide has two polar bonds, and is linear = nonpolar molecule!

Polar molecules The effect of polar bonds on the polarity of the entire molecule depends on the molecule shape water has two polar bonds and a bent shape; the highly electronegative oxygen pulls the e- away from H = very polar!

Determining Molecular Polarity Therefore, polar molecules have... asymmetrical shape (lone pairs) or asymmetrical atoms CHCl3 H Cl net dipole moment

Intermolecular Forces The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together.

Intermolecular Forces They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.

Intermolecular Forces These intermolecular forces as a group are referred to as van der Waals forces.

Dipole-dipole interactions Hydrogen bonding London dispersion forces van der Waals Forces Dipole-dipole interactions Hydrogen bonding London dispersion forces

Ion-Dipole Interactions A fourth type of force, ion-dipole interactions are an important force in solutions of ions. The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.

Dipole-Dipole Interactions Molecules that have permanent dipoles are attracted to each other. The positive end of one is attracted to the negative end of the other and vice-versa. These forces are only important when the molecules are close to each other.

Dipole-Dipole Interactions The more polar the molecule, the higher is its boiling point.

London Dispersion Forces While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom.

London Dispersion Forces At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.

London Dispersion Forces Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1.

London Dispersion Forces These forces are present in all molecules, whether they are polar or nonpolar. The tendency of an electron cloud to distort in this way is called polarizability.

Electronegativity An atom or ion’s electronegativity is its ability to pull electrons towards itself in a covalent bond. The most electronegative elements are found towards the top right corner of the periodic table.

Electronegativity Which covalent bonds would be the most polar? Atom F 4.0 O 3.4 Cl 3.2 N 3.0 Br I 2.7 S 2.6 C H 2.2

Hydrogen Bonding The three types of bonds which give molecules significant hydrogen bonding are; (i) N – H (ii) O – H (iii) F – H These three bonds all have; A strong permanent dipole A hydrogen atom An atom with lone pair electrons

Hydrogen Bonding Hydrogen bonding in water results in some unusual properties; Higher than expected boiling point High specific heat capacity (absorbs a lot of heat energy with only a small change in temperature) Ice is less dense than water

This section of water is frozen This section of water is liquid

The ice structure has large empty spaces which gives it a lower density than water.

Hydrogen Bonding in Hydrogen Fluoride .. H F .. .. Fluorine atoms have three electron lone pairs for bonding to other HF molecules

Hydrogen Bonding in Hydrogen Fluoride .. H F .. H F .. H F ..

Ice Both lone pairs are involved in hydrogen bonds Both hydrogen atoms are involved in hydrogen bonds

Results of Hydrogen Bonding Wool and nylon fibres can hydrogen bond to water – these fabrics can absorb water Polythene has no hydrogen bonding – polythene clothes would get very sweaty and sticky Ice floats on water making life possible

Try it: CO3-2 C: 4 e- , O: 6 x 3 e- , (-2): e- = total 24 e- O O C O

Try it: CO3-2 C: 4 e- , O: 6 x 3 e- , (-2): e- = total 24 e- O O C O

Try it: CO3-2 C: 4 e- , O: 6 x 3 e- , (-2): e- = total 24 e- O O C O

Try it: CO3-2 All 24 e- have been used, buy C does NOT have a complete octet….. Need to make a double bond….Does it matter which O it makes the double bond with? O O C O

Try it: CO3-2 The bond can go here but it could also go between the C and either of the other Os O O C O

Finished Product  -2 C C C