MOLECULAR SHAPEChapter 8

THE SHAPE OF SMALL MOLECULES8-1

Molecular Geometry

• The behavior of atoms is determined chiefly by their electron configurations

• The behavior of molecules also depends on their structural characteristics

• In this section you will look at the shapes of molecules and what characteristics of their bonds produce those shapes

Molecular Geometry• Two ways of looking at the structure of molecules to account for their shapes:

• The first model takes into account the repulsive forces of electrons pairs around an atom

• The second model considers ways in which atomic orbitals can overlap to form orbitals around more than one nucleus.• The electrons in these combined orbitals then serve to bind the

atoms together

Molecular Geometry• In order to describe the shape of a molecule or polyatomic

ion, it is useful to draw a Lewis electron dot diagram

• For all atoms tat form covalent bonds, except hydrogen, eight electrons represent a full outer level.

Molecular Geometry• Take a water molecule for example:

It is the only arrangement of electrons in which all three atoms can achieve a full outer level.

Notice, that 2 electrons in the outer level of Oxygen are involved in bonding the Hydrogens.

These are called shared pairs.

The other 2 pairs of electrons are not involved in the bonding, they are called unshared pairs or lone

pairs.

Molecular Geometry• Notice that the shared electrons contribute to a full outer

level for both atoms sharing the electrons

Molecular Geometry• Electron Pair Repulsion:

• One way of looking at molecules is to consider electron repulsion.

• Each bond and each unshared pair in the outer level of atom form a charge cloud which repels all other charge clouds.

• In part, this repulsion is due to all electrons having the same charge.

Molecular Geometry• Another more important factor is the Pauli exclusion

principle.

• Although electrons of opposite spin may occupy same volume of space, electrons of the same spin may not do so.

Molecular Geometry• The repulsions resulting from the Pauli principle are

greater than the electrostatic ones at small distances.

• Because of these repulsions, atoms cannot be compressed.

Molecular Geometry• The repulsions between the charge clouds in the outer

level of atoms determine the arrangement of the orbitals.

• The orbital arrangement, in turn, determines the shape of molecules.

Molecular Geometry• Structural formula: ammonia

• Gives number of atoms and unpaired electrons.• Does not indicate the shape of an ammonia molecule.• Only identifies bonds, but not how bonds are arranged in

space.

Molecular Geography

• Molecular models• Spheres = nucleus and inner-level electrons• Sticks = bonds• Three dimensional• Symmetrical

Molecular Geometry• Why are atoms arranged

symmetrically?• Valence electrons are found in pairs• Valence electrons repel other electron pairs

because of similar electric charges.

• VSEPR theory – valence-shell electron pair repulsion theory.• In a small molecule, the pairs of valence

electrons are arranged as far apart from each other as possible.

Molecular Shapes

• As result, the following rule may be stated:

• Electron pairs spread as far apart as possible to minimize repulsive forces.

• If there only two electron pairs in the outer level, they will be on opposite sides of the nucleus.

• The arrangement is called linear.

Molecular Shapes

• Linear• Atoms are in a straight line• All molecules that contains only 2 atoms. • Bond angle – 180°

• Because atoms are arranged as far apart as possible.

• Example: CO2

Molecular Shape

• Trigonal Planar• Triangular, flat shape• Bond angle = 120• Usually have a central atom

that is bonded to three other atoms and the central atom has no unshared pairs of electrons.

• Example: Boron trichloride (BCl3)

Molecular Shape

• Tetrahedral• Tetra = 4• A shape with 4 surfaces

• Three dimensional• Bond angle = 109.5• Example: Methane (CH4)

Molecular Shape• Pyramidal

• Represents a shape with unshared pair of electrons.

• All pairs of valence electrons repel each other equally.

• Unshared pairs exert a greater repulsion force (take up more room).

• Bond angle = 107°• Usually have a central atom bonded

to 3 other atoms and an unshared pair of valence electrons.

• Example: Ammonia (NH3)

Molecular Shape

• Bent• Example: Water H2O

• Oxygen in the central atom with two bonds to hydrogen and two pair of unshared electrons.

• The two unshared pairs around the oxygen atom exert a greater repulsion force that the two electron pairs in the bonds.

• 105° bond angles

Molecular Shape• The bonds and unshared electron pairs determine the shape a molecule.

• An unshared pair is acted upon by only one nucleus.

• It's charge cloud is like a very blunt pear, Figure 13-2, with its stem end at the nucleus.

Molecular Shape

•A shared pair of electrons moves within field of two nuclei.

•The cloud is more slender.

Molecular Shape

•The electron pair repulsions in a molecule may not all be equal.

•The repulsion between two unshared pairs is greatest when they occupy the most space.

Molecular Shape

•The repulsion between shared pairs is least because they occupy the least space.

•The repulsion between an unshared pair and a shared pair is an intermediate case.

Molecular Shape• unshared-unshared repulsion >

unshared-shared repulsion >

shared-shared repulsion

• Electron pair repulsion strengths may not be equal.

Molecular Shape• Let us look at the molecular shapes of the compounds CH4, ,H2O, and NH3 to illustrate this repulsion.

• In each of these compounds, the central atom has four clouds around it.

• We expect the axes of all four charge clouds to point approximately to the corners of a tetrahedron.

Molecular Shape

• In methane molecules all clouds are shared pairs, so their sizes are equal and each bond angle is in fact 109.5o

•Methane is therefore a perfect tetrahedron

Molecular Shape

• In NH3 molecules, there are one unshared pair and three shared pairs

•The unshared pair occupies more space than any of the other three, so the bond clouds are pushed together and form an angle of 107o with each other

Molecular Shape

•Although the electron clouds form a tetrahedron one cloud is not involved in bonding.

•Therefore, the atoms composing the molecule form a trigonal pyramid

Molecular Shape

• In H2O molecules, two unshared pairs are present

•Both of these clouds are larger than the bond clouds

•This additional cloud size results in a still greater reduction in the bond angle which is, 104.5o

Molecular Shape

•Note that the electron clouds are tetrahedral but the molecule is “V” shaped, or bent

Molecular Shape

•Note that in the 3 molecules discussed, each has 4 electron clouds.

•The differences in molecular shape result from the unequal space occupied by the unshared pairs and the bonds

Molecular Shape• In most compounds, the outer level is considered full with four paris or 8 electrons

• The outer level in some atoms can contain more than eight electrons• If the outer level is the third or higher level

• Some nonmetals, but mainly halogens form compounds in which the outer level is expanded to 10, 12, or 14 electrons

Molecular Shape

• Other shapes:• T-shaped• Square Planar• Trigonal bipyramidal• Octahedral

Practice Problems

• What is the molecular shape of nitrogen trifluoride? (NF3)• What are the bond angles?

• What is the molecular shape of carbon tetrachloride? (CCl4)• What are the bond angles?

Hybrid Orbitals• Electrons are found in orbitals around the nucleus

• 1s, 2s, 2p….

• Orbitals do not explain the electrons in bonds of a molecule.

• When atoms bond, the electrons are found in hybrid orbitals.• Atomic orbitals of different atoms “mix” together. • They have a combination of the properties of the atomic orbitals

that formed them.

Hybrid Orbitals• In nature there are many different possibilities.

• Linear – mix of s and p orbital = sp orbital• Trigonal planar – mix of s and 2 p orbitals = sp2 orbital

• Hybrid orbitals are often used to categorize molecular shape.

Hybrid Orbitals

•Methane is the bonding of 4 H to 1 C

•The bonds involve the overlap of the s orbital of each H atom with one of the sp3 hybrid orbitals of a C atom

Hybrid Orbitals

•There is an angle of 109.5o between each carbon-hydrogen bond axis

Hybrid Orbitals

•When 2 carbon atoms bond their sp3 overlap. The 3 remaining sp3 orbitals may bond with the s orbital of 3 hydrogen atoms

Hybrid Orbitals

•A covalent bond is formed when an orbital of one atom overlaps an orbital of another atom and they share the electron pair the bond.

•For example, a bond may be formed by the overlap two s orbitals.

Hybrid Orbitals

• A bond formed by the direct overlap of two orbitals is called a sigma bond, and is designated σ.

Hybrid Orbitals

•A sigma bond is also formed by the overlap of an s orbital of one atom with a p orbital of another atom,

• the overlap of 2 p orbitals, • the overlap of 2 hybrid orbitals, • or the overlap of a hybrid orbital with an s orbital

Hybrid Orbitals

•Because p orbitals are not spherical, when 2 half-filled p orbitals overlap, one of two types of bonds can form

• 1. If 2 p orbitals overlap along an axis in and end-to-end fashion, a sigma bond forms

Hybrid Orbitals

•2. If the 2 p orbitals overlap sideways (parallel), they form a pi bond, designated π.

Hybrid Orbitals

•Pi bonds are always formed by the sideways overlap of unhybridized p orbitals

Hybrid Orbitals

• Double bonds are 2 pairs of electrons that are shared between the bonding atoms.

• A double bond always consists of one sigma bond and one pi bond

Hybrid Orbitals

• In a triple bond 3 pairs of electrons are shared between the bonded atoms

• 2 sp hybrid orbitals, one from each carbon, overlap to form 1 sigma bond

• The 2 p orbitals from atom overlap to form 2 pi bonds

Hybrid Orbitals

• Both double and triple bonds are less flexible than single bonds are, and they are also shorter

• Pi bonds are more easily broken than sigma bonds are because the electrons forming pi bonds are farther from the nuclei of the 2 atoms

Hybrid Orbitals

• So molecules containing multiple bonds are usually more reactive than are similar molecules containing only single bonds

• Compounds that contain double or triple bonds between carbon atoms are called unsaturated compounds

Hybrid Orbitals

• If atoms share more than one pair of electrons, all atoms in the molecule can have full outer levels

Hybrid Orbitals

•How does the electron-pair repulsion theory predict the shapes of molecules containing multiple bonds?

Hybrid Orbitals

• Remember that double bonds consist of 4 electrons occupying the space between the bonded atoms

• The resulting cloud will occupy more space than a single bond

• The triple bond occupies still more space than the double bond because it has 6 electrons being shared

Hybrid Orbitals

•How is molecular shape affected by the presence of multiple bonds?

Hybrid Orbitals

•The methanal molecule below has a double bond and 2 single bonds

•This would form a trigonal planar shape

Hybrid Orbitals• Bond types effect the bond angles:

• H – C – H 116• H – C = O 122• C = C = O 180• H – C = O 120• 180

Hybrid Orbitals

• Because the pi electrons are shared equally among all the carbon atoms and not confined to one atom or bond, they are delocalized.

• This delocalization of pi electrons among the carbon atoms in benzene results in greater stability of the compounde

Hybrid Orbitals

•Whenever multiple p orbital overlap can occur, the molecule is said to contain a conjugated system

• This can also occur in rings

Bond Length• Different pairs of atoms form bonds of different lengths. • Trends:

• Moving down a group in the periodic table – atoms form longer bonds.• Atoms become larger as you move down a group.

• Multiple bonds are shorter than single bonds. • The more electrons in a bond, the stronger that bond attracts the

positively charged nuclei of the bonding atom.

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POLARITY8-2

Polarity

• In a polar bond electrons are shared unequally between 2 atoms.

• Electrons are pulled closer to the more electronegative atom giving it a slight negative charge and the other atom a slight positive charge.

• In a nonpolar bond, electrons are shared equally.

Polarity

• Molecules can also be polar or nonpolar.

• A polar molecule has one end with a positive charge and another end with a negative charge. • Dipoles – polar molecules

• Polarity gives molecules different properties:• Align in electric fields• Attracted to or deflected by a magnetic

field

Determining Polarity• Any molecule that is composed of only one kind of atom is

a nonpolar molecule.• Only have nonpolar bonds.• H2, O2

• A molecule that contains polar bonds is not necessarily a polar molecule.• Example: CO2

Determining Polarity

• To determine whether a molecule is polar, you need to look at its shape.

• The shape of a molecule and the polarity of its bonds together determine whether the molecule is polar or nonpolar.

Determining Polarity

•Because no 2 elements have exactly the same electronegativities, in a covalent bond between different elements, one of the atoms attracts the shared pair more strongly than the does the other.

Determining Polarity

• The resulting bond is said to be polar covalent.

• In this bond, the atom with higher electronegativity attracts the electrons more strongly, and that end of the bond will have a partial negative charge.

• The bond at the other end of the bond will have a partial positive charge

Water

• Very important molecule• Liquid state at room temperature and part of almost every liquid on earth.

• Liquid because:• Positive hydrogen end of one water

molecule attracts to the negative oxygen end of another water molecule. • Loosely bonds molecules together.

• Only compounds found in nature as solid, liquid, or gas.

Determining Polarity• Partial charges within a molecule are indicated by δ

(delta)

• Water molecules are a good example of this.

Determining Polarity

• Polar bonds may produce polar molecules

• To be polar the charges must be unequal.

• To be a nonpolar molecule the charges must be pulling in equal strength and therefore, cancel each other out

Determining Polarity

• What about the water molecule, is it polar or nonpolar?

• What about the carbon dioxide molecule, is it polar or nonpolar

Determining Polarity

• In carbon dioxide molecules the carbon-oxygen bond is polar because oxygen has a greater electronegativity than carbon does.

• However, the polarities of the two bonds are in exactly opposite directions and so they cancel each other out.

• This does not occur with water.

Carbon Dioxide

• CO2

• Linear molecule• Two carbon-oxygen double bonds. • Carbon-oxygen bonds are polar. • But carbon dioxide is not a polar molecule:

• Positive charge is concentrated in the center. • Negative charge is divided equally on both sides.

• Being nonpolar gives carbon dioxide important properties:• Molecules have little attraction to each other making

carbon dioxide a gas at room temperature.

Determining Polarity

•Water has a bent/angular geometry so the bonds aren’t exactly opposite from each other.

•Therefore, they don’t cancel each other out.

•So water is polar

Formaldehyde• Used to preserve biological specimens.• CH2O

• Carbon forms bonds with 3 other atoms.

• Oxygen atom has highest electronegativity.• Electrons in the C-O bond are attracted

more towards the oxygen.• Oxygen becomes partially negative and

carbon partially positive.

• Carbon more electronegative than hydrogen.

• Difference in negative and positive partial charges makes molecule polar.

Determining Polarity

• Polar bonds are a necessary but not a sufficient condition for polar molecules.

• In a polar molecule, the polar bonds cannot be symmetrically arranged.

Determining Polarity

• Because it has both a positive and a negative pole, a polar molecule, such as water, is also said to be dipole, or to have a dipole moment.• Not to be confused with a shiny hair moment

• A dipole moment is a measure of the strength of the dipole and is a property that results from the asymmetrical charge distribution in a polar molecule

Dipoles

• The dipole moment depends upon the size of the partial charges and the distance between them.

• μ is the dipole momement, q is the size of the partial charge in coulombs and r is the distance in meters between the partial charges.

• The units are in coulomb x meters

Dipoles• The higher the dipole moment, the stronger the

intermolecular forces; and, consequently, the higher the melting point and boiling point for molecules of similar mass.

Dipoles

• Van der Waals forces are sometimes referred to as weak forces because they are much weaker than chemical bonds.

• Weak forces involve the attraction of the electrons of one atom for the protons of another.

Dipoles

• Intramolecular forces are forces within a molecule that hold atoms together, that is, covalent bonds

• Intermolecular forces are forces between molecules that hold molecules to each other, that is, van der Walls forces

Dipoles

• The first van der Waals force is the dipole-dipole force.

• With dipole-dipole forces, two molecules of the same or different substance that are both permanent dipoles, will be attracted to each other.

Dipoles

• A dipole can also attract a molecule that is ordinarily not a dipole.

• When a dipole approaches a nonpolar molecule, its partial charge either attracts or repels the electrons of the other particle.

Dipoles

• For instance, if the negative end of the dipole approaches a nonpolar molecule, the electrons of the nonpolar molecule are repelled by the negative charge.

• The electron cloud of the nonpolar molecule is distorted by bulging away from the approaching dipole as shown in

• Figure 14-3.

Dipoles

• As a result, the nonpolar molecule is itself transformed into a dipole.

• We say it has become an induced dipole.

• Since it is now a dipole, it can be attracted to the permanent dipole.

Dipoles

• Interactions such as these are called dipole-induced dipole forces.

• An example of this force occurs in a water solution of iodine.

• The I2 molecules are nonpolar while the water molecules are highly polar.

• The case of two nonpolar molecules being attracted must also be taken into account.

Dipoles

• For instance, there must be some force between hydrogen molecules; otherwise it would be impossible to form liquid hydrogen.

• Consider a hydrogen molecule with its molecular orbital including both nuclei.

Dipoles

• We know intuitively that the electrons occupying that orbital must have a specific location.

• If they are both away from one end of the molecule for an instant, then the nucleus is exposed for a short time.

Dipoles• That end of the molecule has a partial positive charge for an instant; a temporary dipole is set up.

• For that time, the temporary dipole can induce a dipole in the molecule next to it and an attractive force results as shown in Figure 14-4.

• The forces generated in this way are called dispersion forces.

Dipoles

• The various kinds of interactions making van der Waals forces affect each other, but we are only interested in the net result.

• The liquid and solid states of many compounds exist because of these intermolecular forces.

Dipoles

• These forces are effective only over very short distances.

• They vary roughly as the inverse of the sixth power of distance.

• In other words, if the distance is doubled, the attractive force is only 1/64 as large.

Dipoles

• Of the three contributing factors to van der Waals force, dispersion forces are the most important.

• They are the only attractive forces that exist between nonpolar molecules.

Dipoles

•Even for most polar molecules, dispersion forces account for 85% or more of the van der Waals forces.

•Only in some special cases, such as NH3 and H20, do dipole-dipole interactions become more important than dispersion forces.

Large Molecules• Small molecules – the shape

helps to determine polarity.• Large molecules – the polarity

often helps to determine its shape.

• Example: Protein• Essential to all living things. Build and

repair cells and components of many cell structures.

• Extremely large molecules. (thousands of atoms)

• Composed of individual subunits into a chain.

Large Molecules• Subunits have polar

sidechains• Molecule bent and

twisted because polar sections attracted to each other.

• Large molecules have a large variety of shapes. • Geometry around

individual atoms is identical to small molecules.

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