Polar Molecules. Textbook pg Molecules in which the charge is not distributed symmetrically among the atoms making up the molecule

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1 Textbook pg Polar Molecules Molecules in which the charge is not distributed symmetrically among the atoms making up the molecule Electronegativity and Polar Molecules Pauling realized that electron pairs could be shared evenly or unevenly there are differences in attraction of nuclei for a shared pair of electrons in a covalent bond Created electronegativity to explain and predict polarity of molecules F has the highest electronegativity (EN) at 4.0 When you compare the electronegativities between two atoms, the greater the ΔEN (difference in electronegativity) the greater the polarity of the chemical bond The smaller the ΔEN the more non-polar the bond Polar bond results from a difference in electronegativity between the bonding atoms; one end of the bond is, at least partially, positive and the other end is equally negative o A very polar bond is an ionic bond a bond in which the bonding pair of electrons is mostly with one atom/ion o Polar bonds result when two different kinds of atoms (non-metals) form a bond o The electrons are spending more time closer to one of the atom than to the other o The end where the electrons stayed was slightly more negative (or partially negative and represented as δ - ), and the other end where the electrons did not stay was more positive (or partially positive and represented as δ + ) Non-polar bond results from zero difference in electronegativity between the bonded atoms; a covalent bond with equal sharing of bonding electrons o A non-polar bond is a covalent bond a bond in which electrons are shared somewhat equally

2 General Rule when the ΔEN is greater than 1.7, the percent ionic character exceeds 50% Examples: o H H, ΔEN = = 0, non-polar o P Cl, ΔEN = = 0.9, polar covalent bond o Na Br, ΔEN = = 1.9, ionic bond Polar Molecules Polar bonds in a molecule do not necessarily mean you have a polar molecule Ex. CO 2 is non-polar even though the C=O bonds are considered polar o According to Lewis structures and VSEPR theory, CO 2 is linear o Can use electronegativities to predict the polarity of each of the bonds Bond dipole proportional to the ΔEN, shown by an arrow pointing from the lower EN (δ + ) to the higher EN (δ - ) o A vector quantity where direction is important two dipoles of equal magnitude, but opposite direction will cancel each other out In the CO 2 molecule, dipoles cancel each other out so the molecule exhibits no polarity and is said to be non-polar Ex. H 2 O is polar o The Lewis structure and VSEPR theory predict that the molecule will be V- shaped o The dipoles do not cancel and instead the vertical components add together to produce a non-zero molecular dipole o The molecule has definite overall polarity and is said to be a polar molecule (partially negative at the O end and partially positive at the H end) It should now be evident that both the shape of the molecule and the polarity of the bond are necessary to determine if a molecule is polar or non-polar Ex. CH 4 is non-polar o The outer part of the molecule is positive on all sides and none of the ends are charged differently o This is because the CH 4 molecule is symmetrical Therefore, in all symmetrical molecules the sum of the bond dipoles is zero and the molecule is non-polar

3 Steps to Predicting Molecule Polarity 1. Draw the Lewis structure for the molecule 2. Use the number of electron pairs and VSEPR rules to determine the shape around each central atom 3. Use electronegativities to determine the polarity of each bond. 4. Add the bond dipole vectors to determine if the final result is zero (non-polar molecule) or nonzero (polar molecule). Sample Problem Predict the polarity of the ammonia, NH 3, molecule, including your reasoning. Draw the Lewis structure. Based on the Lewis structure, draw the shape of the molecule. Add the electronegativities of the atoms, from the periodic table, and assign δ + and δ - to the bonds. Draw in the bond dipoles. Reasoning The ammonia molecule is polar because it has polar bonds that do not cancel to zero. The electron pairs are in a tetrahedral arrangement, but one of these pairs is a lone pair and three are bonding pairs. Therefore, the bond dipoles do not cancel. Intermolecular Forces The forces of interaction that may exist between molecules Much weaker than covalent bonds (ie. if covalent bonds are assigned a strength of about 100, ten intermolecular forces are generally to 15) Types of intermolecular forces can be classified as dipole-dipole forces, London forces, and hydrogen bonding

4 Dipole-Dipole Forces Molecules can be classified as polar or non-polar The dipole-dipole force is an attractive intermolecular force resulting from the tendency of polar molecules to align themselves such that the positive end of one molecule is near the negative end of the other The strength of the dipole-dipole force is dependent on the polarity of the molecule London (Dispersion) Forces Fritz London (1930) accounted for the weak attraction between any two molecules (specifically non-polar molecules) by recognizing that electrons orbiting the nucleus of an atom may be on any one side at any point in time As a result, there may be a small, instantaneous dipole, with one side having a partial negative charge and the other having a partial positive charge If another atom is present nearby, the partial negative charge of the first atom will repel the electrons of the second atom, creating a partial positive charge on the second atom o The partial positive charge and the partial negative charge of each atom will result in an attractive force between two atoms This weak attraction occurs instantaneously because the electrons are in constant motion nonetheless, the motion of the electrons in one atom will influence the motion of electrons of other atoms London forces (or dispersion forces) are the weak attractive forces between molecules resulting from the small, instantaneous dipoles that occur because of the varying positions of the electrons during their motion about the nuclei London forces increase as the molecular weight increases due to: o The increased number of electrons in motion o The increased size of the molecule, which permits electrons to move further from the nucleus (more polarisable) Van der Waals Forces Van der Waals forces is a general term for those intermolecular forces that include dipole-dipole and London forces These forces will affect the ease or difficulty with which a molecule leaves a liquid Can observe the effect of intermolecular forces by looking at molecular weight o An increase in molecular weight generally results in increase in London forces Can also observe the effect of intermolecular forces by looking at boiling points of liquids Surface tension is also dependent on intermolecular forces o Surface tension is the energy needed to increase the surface area of the liquid o To increase the surface area, it is necessary to pull molecules apart against the intermolecular forces of attraction Viscosity of a liquid depends in part on the intermolecular forces o Increasing attractive forces between molecules increases the resistance to flow o Viscosity if also dependent on other factors, such as the possibility of molecules tangling together (long molecules)

5 Hydrogen Bonding Comparing fluoromethane (CH 3 F) and methanol (CH 3 OH) will show that their boiling points are significantly different o CH 3 F = -78ºC (gas under normal conditions) o CH 3 OH = 65ºC (liquid under normal conditions) o Both molecules have the same molecular weight and polarity (dipole) o Suggests another intermolecular forces is at work (not just van der Waals) o Chemical structure shows that fluoromethane has a C-F bond and methanol has a C-OH bond o The OH group creates additional attractive forces between molecules Hydrogen bonding is a weak to moderate attractive force that exists between a hydrogen atom covalently bonded to a very electronegative atom (N, O or F), and a lone pair of electrons on another small, electronegative atom (N, O or F) in adjacent molecules Compare boiling points of hydrides with Group 16 elements: H 2 O (100ºC), H 2 S (- 60ºC), H 2 Se (-40ºC), H 2 Te (0ºC) o If London forces were the intermolecular forces present one would expect the boiling points to increase from H 2 O to H 2 Te based on molecular weight o This does not occur, as water has a much higher boiling point o Consistent with the view that hydrogen bonding exists in H 2 O but is virtually nonexistent in the other molecules Can also see the same pattern when looking at other hydrides of: o Group 17 = HF, HCl, HBr, HI o Group 15 = NH 3, PH 3, AsH 3, SbH 3 However, the pattern is not evident with Group 14 hydrides = CH 4, SiH 4, GeH 4, SnH 4 The Structure and Properties of Solids All solids have a definite shape and volume, are incompressible, and do not flow readily The hardness, melting point, mechanical characteristics and conductivity can vary between solids due to the forces between particles in elements and compounds

6 Ionic Crystals Arrangement of ions forms a crystal lattice (a 3D arrangement of ions in a crystal structure) Wide variety of shapes means that there is a wide variety of internal structures Ionic compounds are hard and brittle solids at SATP, conduct electricity in liquid form but not in solid state, form conducting solutions in water and have high melting points These qualities are interpreted to mean that ionic bonds are strong and directional In general, ionic bonding is stronger than all intermolecular forces The properties of ionic crystals are explained by a 3D arrangement of positive and negative ions held together by strong, directional ionic bonds Metallic Crystals Metals are shiny, silvery, flexible solids with good thermal and electrical conductivity X-ray diffraction shows that all metals have a continuous and very compact crystalline structure All metals have a closely packed structure (with few exceptions) Current theory to explain this is the electron sea model where the fixed positive nuclei are bonded to loosely held, mobile electrons Ideas of this model are: o Low ionization energy of metal atoms to explain loosely held electrons o Empty valance orbitals to explain electron mobility o Electrostatic attractions of (+) centers and (-) electron sea to explain the empirical properties of metals Properties of metallic crystals are explained by a 3D arrangement of metal cations held together by strong, nondirectional bonds created by a sea of mobile electrons

7 Molecular Crystals Explain molecular solids which are not hard, have relatively low melting points, and are nonconductors (as liquids or solids) X-ray analysis shows a crystal-like structure with the molecules packed very tightly together and more complex than in ionic compounds The properties of molecular crystals are explained by a 3D arrangement of neutral molecules held together by relatively weak intermolecular forces Covalent Network Crystals Are brittle, very hard, high melting points, insoluble and non-conductors of electricity Are usually much harder and have much higher melting points than ionic and molecular crystals The network of covalent bonds leads to the common name for these crystals as covalent network Properties of hardness and high melting point provide evidence that the overall bonding is very strong The interlocking structure gives it strength, and is why individual atoms are not easily displaced and the sample is so hard To break a covalent network crystal, many bonds must be broken which would require lots of energy this is why their melting points are so high Electrons are not free to move so these substances do not conduct electricity Include diamond, quartz (used in making emery sandpaper) o The shape and X-ray diffraction analysis shows that diamonds have carbon atoms in a large tetrahedral network with each C bonded to 4 other C atoms o Each diamond is a crystal and can be described as a single macromolecule with a chemical formula of C (s) Properties of network covalent crystals are explained by a 3D arrangement of atoms held together by strong, directional covalent bonds

8 Other Covalent Networks of Carbon Carbon is extremely versatile and can bond to itself to form many pure carbon substances 3D structures (diamonds, Figure a), layers of sheets (graphite, Figure b), large spherical molecules (buckyballs, Figure c), and long thin tubes (carbon nanotubes, Figure d) Read Semiconductors on Pg.272 and make notes

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