Water Vs. Hydrocarbons

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Water vs. HydrocarbonsBy: Jasmine Gilbert, Matt Huber, Michael Wild, and Dr. Faith YarberryIn this module the student will:Become familiar with the structure of water and methane.Understand how the structure influences the polarity of water and methane.Be able to identify the intermolecular force associated with each substance.Understand how the intermolecular force relates to the properties associated with eachsubstance.Understand how the intermolecular force influences the solubility of a solute in a givensolvent.Water vs HCPage 1

Lesson 1: Structure of water and methane.Water, H2O, and methane, CH4, are both covalent molecules. A covalent molecule is achemical compound that contains covalent bonds. A covalent bond is a bond that arisesthrough the sharing of electrons between two atomic nuclei. In reality the bond is anelectrostatic attraction between the protons of one atom and electron(s) of an adjacent atom.A covalent bond will form when the attractions just barely outweigh the repulsions. Covalentbonds are found between two non-metals or a non-metal and a metalloid. These elementsparticipate in covalent bonds because too much energy would be required for either element tolose electrons to become a cation, which is required for the formation of ionic compounds. Inthe structure of water you will find that the oxygen atom shares a pair of electrons with each oftwo hydrogen atoms. Additionally, two lone pairs of electrons are found on the oxygen atom.In the structure of methane a carbon atom shares a pair of electrons with each of four hydrogenatoms.Water vs HCPage 2

The electrons found in the lone pairs and within the covalent bonds are referred to as valenceelectrons. It is the valence electrons that are responsible for the shape of a covalent molecule.Electrons are negatively charged subatomic particles. What is observed when the negative endsof two magnets are brought in close contact with one another? The repel each other. The same isobserved for a molecule’s valence electrons. When they are brought in close proximity to eachother they will repel one another. VSEPR Theory, Valence Shell Electron Pair Repulsion,states that the valence electrons, being negatively charged, will repel each other so that theelectrons are as far from one another as possible. For four electron groups, minimumrepulsion is obtained when the pairs of valence electrons point to the corners of a tetrahedron.Demonstration: Use four balloons to show that 4 sets of electrons obtain minimal repulsionwhen the groups of electrons point to the corners of a tetrahedron. Point out that in a tetrahedralstructure the groups are approximately 109.5o from each other.After demonstration, pass around a model of methane and water made from Styrofoam balls andpopsicle sticks. This visual will allow the student to observe the spacing of the atoms and lonepairs in the molecules to prepare for completion of Activity #1.Activity #1 – Perfect for homework.Although all electrons (bonding groups and lone pairs) participate in repulsion, the shape of themolecule is described by the atoms present. Methane has four bonding groups, its structure isthat of the tetrahedron. Water has two bonding groups and two sets of lone pairs, its structure isreferred to as bent.Water vs HCPage 3

Lesson 2 – Non-polar Covalent Bonds and Polar Covalent BondsEven though a covalent bond, which arises through the sharing of electrons, is present incovalent molecules, it does not mean that the electrons in that bond are shared evenly. In realitysome elements prefer that the electrons spend a greater amount of time in their vicinity than doother elements. This distribution of the electrons between two atoms can be determined byevaluating each element’s electronegativity. Electronegativity is the ability of an atom in amolecule to attract the shared electrons in a covalent bond. Below is a table ofelectronegativities.A bond is considered to be a non-polar covalent bond when the difference in theelectronegativities of the atoms that make up the bond is between 0 and 0.4. A polarcovalent bond arises when this difference is calculated to be between 0.5 and 1.9. Below isan electron density map of a non-polar covalent bond and a polar covalent bond.Water vs HCPage 4

Note: in the non-polar covalent bond, the density map is evenly distributed between the twoatoms of chlorine, but, in the polar covalent bond the electron density is much greater around thechlorine atom due to the significant differences in atomic electronegativity. Again, this diagramillustrates where the electrons are most likely found within the bond. Since electrons arenegative and since they spend a greater amount of time around the chlorine atom, the result isthat the chlorine atom takes on a partially negative charge. In the same regard, the electrons arepulled away from the hydrogen atom therefore giving the hydrogen atom a partially positivecharge. The difference in the electronegativities of the elements create a dipole. A dipole is theformation of a negative and positive pole. An arrow notation is commonly used to illustrate thedirection of the electron density within the molecule. The head of the arrow will always bepointed in the direction of the more electronegative atom. The opposite end of the arrow iscrossed to represent the partial positive charge possessed by the least electronegative atom.Water vs HCPage 5

Lesson 3 – Polar Molecules, Non-Polar Molecules, and Intermolecular ForcesThe geometry of a molecule as well as the difference in the electronegativities of the atoms thatmake up the molecule, determine whether a molecule will have a dipole moment or not. Adipole moment arises when polar bonds reinforce each other causing the molecule to bepolarized. If no dipole moment exists, the molecule is considered to be non-polar. Below arediagrams of water and methane.The electronegativity difference between oxygen and hydrogen is 1.4, meaning that the electronswithin each H-O bond spend a greater amount of time around the oxygen atom. Given the bentshape of a water molecule, the polarities of these two bonds reinforce one another. Water,therefore, has an overall dipole moment in the direction of the oxygen atom. In the methanemolecule, each C-H bond has an electronegativity difference of 0.4. The electrons within theC-H bond spend a little more time around the carbon than the hydrogen, causing the bond to beonly slightly polarized in the direction of carbon. However, instead of the bonds reinforcing oneanother they cancel each other out causing methane to be a non-polar molecule.The polarity of a molecule as well as the connectivity of the atoms within the moleculedetermines the intermolecular force present. The intermolecular force(s) present in a covalentmolecule can be one of three types: hydrogen bonding, dipole-dipole, or London dispersion.A hydrogen bonding intermolecular force is an attractive force found between polarmolecules that contain the criteria of hydrogen directly bonded to a nitrogen, oxygen, orfluorine atom.A dipole-dipole intermolecular force is an attractive force found between polar molecules.A London dispersion intermolecular force is an attractive force found between non-polarmolecules that results from an instantaneous, temporary dipole caused by electron motion.The strongest intermolecular force, of the forces listed, is that of hydrogen bonding. Below is atable that compares the strength of the three forces.Water vs HCPage 6

ForceStrength ofIntermolecular Force(energy required to break onemole of molecules apart thatcontain that force)Hydrogen Bond10-40 kJ/molDipole-Dipole3-4 kJ/molLondon Dispersion1-10 kJ/molCharacteristicsOccurs between polarmolecules with O-H, N-H, andF-H bondsOccurs between polarmoleculesOccurs between all molecules;strength depends on size andpolarizabilityWater contains the intermolecular force - hydrogen bonding given that the molecule is polar andit contains O-H bonds. The hydrogen bond occurs between the partially negative oxygen of onewater molecule and the partially positive hydrogen on an adjacent water molecule.Methane contains London dispersion forces since it is a non-polar molecule. The diagram belowillustrates London forces within chlorine, Cl2. Illustration (a) represents the chlorine molecules intheir standard state. Because the electrons are in constant motion, a time will come when thechlorine molecule on the left has a temporary dipole. This temporary dipole has a domino effecton its nearest neighbors forcing them to have a temporary dipole as illustrated in diagram (b).The polarity of the molecules will be short lived and soon the molecules will return to theirstandard state (a).Water vs HCPage 7

Activity #2 – builds on the molecules from Activity #1.Water vs HCPage 8

Lesson 4 – The effect of intermolecular forces on molecular properties.Most solids are substances whose constituent particles have an ordered arrangementextending over a long range. Molecular solids are held together by the intermolecular forcesdescribed in the previous lesson. The melting point of a solid describes the amount of energyneeded to overcome some of these attractions and allow the particles to move more freely inthe form of a liquid. Liquids are substances whose particles are still relatively closetogether but are allowed to move around more freely due to weakened intermolecularforces. Even though these intermolecular forces are weakened, they still exist. The boilingpoint describes the amount of energy needed to break the remaining attractions and allowthe particles to move very freely as gas particles. Because gas particles are very smallcompared to the amount of space between them, gaseous substances do not containintermolecular forces.Many of the properties associated with water can be explained by its intermolecular forces.Why does water, with a molar mass of 18 g/mol, melt at 0oC while methane, with a fairly similarmolar mass of 16 g/mol, melt at -182.5oC? Why does water boil at 100oC while methane boils at-161oC? The answer is found in the intermolecular forces. Water contains hydrogen bondingwhich is a much stronger intermolecular force than methane’s London forces. Since watercontains the stronger intermolecular force it means that a greater amount of energy will need tobe added to break two water molecules apart.Activity #3 – Hands on discovery of additional propertiesIntermolecular force also explains cohesion, adhesion, surface tension, and capillary action.Cohesion, the ability of like molecules to stick together, will increase with strongerintermolecular forces. Why can a glass be over-filled without the water overflowing? The answeris cohesion. The individual water molecules are attracted to four other water molecules throughhydrogen bonding. Because hydrogen bonding is quite strong, it is more difficult to separate themolecules from one another, hence it forms a convex surface. Adhesion, the ability ofdissimilar molecules to stick together due to attractive forces, increases when the attractionbetween the dissimilar substances is stronger than the attraction between like molecules. Whywill water form a concave surface when in a half-filled glass? The answer is adhesion. The watermolecules are attracted to the silicon dioxide of the glass through hydrogen bonding. ThisWater vs HCPage 9

attraction between the water and the glass is stronger than the attraction between the watermolecules themselves.Whether the surface of the liquid is concave or convex it is referred to as a meniscus, a curve atthe surface of a molecular substance in response to the surface of the container. Finally,capillary action describes the behavior of liquids in thin tubes. Capillary action is also relatedto intermolecular forces. Why does water travel up a small capillary tube? The answer is foundwithin the attraction of the water to the silicon dioxide. This attraction is stronger than thehydrogen bonding between water molecules. Because the attraction is stronger the water creepsup the sides of the capillary tube.Since methane contains London forces, these observations will be very different.Water vs HCPage 10

OverheadsWater vs HCPage 11

Definitions – Lesson #1Covalent molecule - a chemical compound that containscovalent bondsCovalent bond - a bond that arises through the sharing ofelectrons between two atomic nucleiVSEPR Theory - Valence Shell Electron Pair RepulsionTheory - states that the valence electrons, being negativelycharged, will repel each other so that the electrons are as farfrom one another as possibleWater vs HCPage 12

Definitions – Lesson #2Electronegativity - the ability of an atom in a molecule toattract the shared electrons in a covalent bondNon-polar covalent bond – a bond in which the difference inthe electronegativities, of the atoms that make up the bond,is between 0 and 0.4Polar covalent bond – a bond in which the difference inelectronegativities, of the atoms that make up the bond, iscalculated to be between 0.5 and 1.9Dipole – formation of a negative and positive poleWater vs HCPage 13

Definitions – Lesson #3Dipole moment - arises when polar bonds reinforce eachother causing the molecule to be polarizedHydrogen bonding - an intermolecular force found betweenpolar molecules that contain the criteria of hydrogen directlybonded to a nitrogen, oxygen, or fluorine atom.Dipole-dipole – an intermolecular force found between polarmoleculesLondon dispersion – an intermolecular force is foundbetween non-polar molecules that results from aninstantaneous, temporary dipole caused by electron motionWater vs HCPage 14

Definitions – Lesson #4Solids – substances whose constituent particles have anordered arrangement extending over a long rangeMelting point – describes the amount of energy needed toovercome some of the intermolecular forces found in solidsso that the particles are allowed to move more freely in theform of a liquidLiquids – substances whose particles are relatively closetogether, contain intermolecular forces, but, whose forcesare weaker than those found in solidsBoiling point – describes the amount of energy needed tobreak the intermolecular forces found in liquids so that theparticles are allowed to move very freely as gas particlesWater vs HCPage 15

Cohesion – the ability of LIKE molecules to stick togetherdue to attractive forcesAdhesion – the ability of DISSIMILAR molecules to sticktogether due to attractive forcesCapillary action – describes the behavior of liquids in thintubes associated with adhesionConvex – curving out or bulging outwardConcave – curving in or hallowed inwardSurface Tension – attraction of one surface molecule in aliquid to additional surface molecules of the liquid associatedwith cohesionWater vs HCPage 16

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ForceStrength ofIntermolecular Force(energy required to break onemole of molecules apart thatcontain that force)Hydrogen Bond10-40 kJ/molDipole-Dipole3-4 kJ/molLondon Dispersion1-10 kJ/molWater vs HCCharacteristicsOccurs between polarmolecules with O-H, N-H, andF-H bondsOccurs between polarmoleculesOccurs between all molecules;strength depends on size andpolarizabilityPage 19

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ActivitiesWater vs HCPage 23

Activity #1Introduction:The purpose of this activity is to develop a visual understanding of the structure of water and ofhydrocarbons.Materials:Marshmallows (Both Large and Small)ToothpicksRedhotsProcedure:1. Prepare 3 water molecules. Use the toothpicks to represent bonds, the largemarshmallows to represent the oxygen atom and the small marshmallows to represent thehydrogen atoms of elements, and the redhots to represent the lone pairs (redhots will needto be licked to stick to the marshmallow). The angles between each external groupshould be approximately 109o from one another.2. Prepare 2 methane molecules. Use the toothpicks to represent bonds, the largemarshmallows to represent the carbon atom and the small marshmallows to represent thehydrogen atoms. The angles between each external group should be approximately 109ofrom one another.Water vs HCPage 24

3. Prepare 2 propane, C3H8, molecules. Use the toothpicks to represent bonds, the largemarshmallows to represent the carbon atom and the small marshmallows to represent thehydrogen atoms. The angles between each external group should be approximately 109ofrom one another.Water vs HCPage 25

Activity #2Introduction:The students will develop a visual understanding of the difference between hydrogen bondingand London forces. They will be able to describe the effect that the intermolecular force has onthe properties of melting points and boiling points.Materials:Molecules developed previouslyPipe cleanersRibbonProcedure:1. Attach the oxygen in the one water molecule to the hydrogen in a second water moleculeusing a pipe cleaner.2. Attach the carbon in one methane molecule to the carbon in a second methane moleculeusing ribbon.3. Attach the three carbons in a propane molecule to the three carbons in a second propanemolecule using ribbon.4. Try to pull the molecules apart.Observations:Which is the flimsiest intermolecular force?List the compounds (water, methane, propane) in order of increasing difficulty to separatecompletely.What was the difference between methane and propane, each of which contain London forces?How do you think that the intermolecular force relates to melting points and boiling points?Melting point and boiling point describes the energy required to separate molecules.Water vs HCPage 26

Procedure:1. Attach the oxygen in the one water molecule to the hydrogen in a second water moleculeusing a pipe cleaner.2. Continue this process until all water molecules are connected together in a circle.Sketch your result.Ice floats on water because it is less dense than water. Give an explanation for this statement,taking into account the presence of air in the voids of your drawing above.Water vs HCPage 27

Activity #3Introduction:The students will understand the properties of cohesion, adhesion, and surface tension and relatethese properties to intermolecular forces.Materials:Stop Watch (students should spend 10 minutes at each station)Procedure for each stationStation 1:Fluted wine glass with string tied to the stemPiece of light plastic larger than the opening of the glassWaterSinkStation 2:2 regular drinking glassesWaterSeveral pieces of small corkStation 3:3 to 4 capillary tubes of different diametersSmall beakerFood coloringWaterStation 4:Paper card cut according to diagramLiquid detergentDropperMetal Pie pan or shallow trayWaterStation 5:Small bottles of different size openingsCheese cloth or screenRubberbandsWaterWater vs HCPage 28

Procedures:Station 1:1. Fill the glass mostly full of water.2. Place the plastic across the opening of the glass3. While holding the plastic snuggly to the opening of the glass gently turn the glass upsidedown.4. Slowly remove your hand from the plastic.Station 2:1. Fill one glass half way with water2. Float a cork on the water surface. Observe.3. Fill the second glass full of water (all but overflow the water)4. Float a cork on the water surface. Observe.Station 3:1. Fill the beaker with water and place a few drops of food coloring in it.2. Hold three or four capillaries, of different diameters, close to each other and dip them inthe water.3. Observe the water level in each capillary.Station 4:1. Fill a shallow tray or pie pan with water.2. Gently place a boat on top of the water.3. Add a drop of liquid detergent in the center of the opening of the boat.Station 5:1. Use a rubber band to fasten the screen over the open end of the bottle.2. Pour water through the screen.3. Invert the bottle and observe.Water vs HCPage 29

Data and ConclusionsStation 1:Did the water form a convex or concave meniscus?If convex, t

Water contains the intermolecular force - hydrogen bonding given that the molecule is polar and it contains O-H bonds. The hydrogen bond occurs between the partially negative oxygen of one water molecule and the partially positive hydrogen on an adjacent water molecule. Methane contains L

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