CHEM. 1 HONS. UNIT 8 CH. 11 INTERMOLECULAR FORCES AND LIQUIDS AND SOLIDS LIQUIDS AND SOLIDS

CHEM. 1 HONS UNIT 8 CH. 11 IMF s and Liquids and Solids 1 CHEM. 1 HONS. UNIT 8 CH. 11 INTERMOLECULAR FORCES AND LIQUIDS AND SOLIDS READ P. 418-454 OMIT P. 428 Crystal Structure) to P. 439. ASSIGNMENTS: #1 P. 9 notes #1-10 Intermolecular Forces #2 P. 10 notes #1-13 Interparticle forces #3 P. 12 notes #1-8 Liquids #4 P. 16 notes #1-12 Forces in Solids #5 P. 17 notes #1-4 Let the Forces Be With You #6 P. 18 notes #1-28 Phase Changes #7 P. 24-25 notes #1-10 Vapor Pressure and Boiling Point #8 P. 27 notes #1-8 Specific Heat #9 P. 30-31 notes #1-11 Phase Changes CH. 11 LIQUIDS AND SOLIDS Gas Low density High compressibility Completely fills its container Indefinite shape, indefinite volume Solid High density Only slightly compressible Rigid Maintains its shape Definite shape, definite volume Liquid Properties lie between those of solids and gases H 2 O(s) H 2 O(l) H fusion = 6.02 kj/mol (The amount of heat to melt 1 mole (18.02 g) of ice.) H 2 O(l) H 2 O(g) H vaporizati on = 40.7 kj/mol (The amount of heat to vaporize 1 mole of water to steam.) vaporization The large value of H suggests a much greater change in structure in going from a liquid to a gas than from a solid to a liquid. This also suggests that there are extensive attractive forces among the molecules in liquid water, similar but not as strong as those in the solid state. Similar densities of the solid and liquid state indicate similarities in the structure of liquids and solids: Densities of the three states of water: H 2 O(g) D = 3.26 x 10 4 g/cm 3 (400 C) H 2 O(l) D = 0.9971 g/cm 3 (25 C) H 2 O(s) D = 0.9168 g/cm 3 (0 C) January 25 Test: February 10 07/12/2006

CHEM. 1 HONS UNIT 8 CH. 11 IMF s and Liquids and Solids 2 PHASES OF MATTER AND THE KINETIC MOLECULAR THEORY Solids and liquids are grouped together separate from gases because they have some common properties that can be examined. The densities, compressibilities and heats of phase change all indicate that the forces that hold solids and liquids together are similar. Gases have no such forces. The kinetic molecular theory of gases assumes that 1. Gas molecules or atoms are widely separated 2. No forces of attraction exist between them 3. The molecules are in continual, random and rapid motion 4. Their kinetic energy is determined only by the gas temperature. Any two gases at the same temperature will have the 3 same average kinetic energy. INFORMATION ONLY: (KE) average = RT 2 The first two points only are the ones that distinguish gases from liquids and solids. Because gases have no forces attracting molecules to one another, gases can expand infinitely to fill their containers uniformly and completely. In contrast, stronger forces exist in liquids and solids, and these attractive forces in a liquid or solid prevent a significant expansion of a liquid or solid. LIQUIDS AND SOLIDS In this unit, you will study liquids and solids. Liquids and solids are substances which have forces holding their INDIVIDUAL PARTICLES together. For example, a glass of water is made of zillions of water MOLECULES that are attracted to each other by some type of attractive force. One molecule of water is attracted to another water molecule and if there are enough of them, they form a quantity (like a glass full) of liquid water. These attractive forces exist between all particles that make up solids and liquids. Without these forces, everything would be in the gas state even you! In this unit you will study what the different particles are that make up solids and liquids and what forces hold those particles together. The type of force depends on the type of particles that make up the liquid or solid. The chart below is a summary of what you will learn: Particles being held together to form the liquid and/or solid Ions Atoms Metal atoms Carbon atoms, SiC, silicon with oxygen (SiO x ), Group 8A atoms Molecules Hydrogen bonding Dipole-dipole London Dispersion Forces (LDF) Force(s) holding those particles together Ionic bonds Metallic bonds Network covalent bonds London Dispersion Forces IMF s Intermolecular Forces Polar molecules (H is a bridge between F, O, and/or N atoms Polar molecules All molecules The properties of liquids and solids are determined by the forces that hold the components of the liquid or solid together. You will notice from the chart above that the forces may be summarized as: Covalent bonds Ionic bonds Metallic bonds Weaker intermolecular forces between molecules THE STUDY OF INTER PARTICLE FORCES Forces between particles of matter hold the individual particles together to form solids and liquids. Consider water for example. The bonds holding oxygen atoms to hydrogen atoms within a water molecule are covalent bonds and they are called INTRAMOLECULAR FORCES (intra means within ). We have already studied this in previous chapters so we are not going to do it again here. What we want to know now is what kind of force(s) hold individual water molecules to other water molecules to form liquid and solid water. These forces holding particles of matter together to form solids and liquids are called INTERPARTICLE FORCES (inter means between ). January 25 Test: February 10 07/12/2006

CHEM. 1 HONS UNIT 8 CH. 11 IMF s and Liquids and Solids 3 LOOK AT THE WATER MOLECULES IN THE DIAGRAM TO THE RIGHT. Intermolecular Forces (IMF s) Intramolecular forces are covalent bonds between atoms to form molecules. These are forces within a molecule. (Water the intramolecular forces are covalent bonds that hold the O and H atoms to each other within the water molecule. Intermolecular forces mean forces between molecules. These are the forces that hold the components of the liquid or solid together. (Forces holding one water molecule to another water molecule.) Intermolecular forces are weaker than covalent and ionic bonds. That is why it is easy to separate water molecules from one another but very difficult to separate a water molecule into H and O atoms. It is important to realize that when a substance such as water changes from solid to liquid to gas, the molecules remain intact. The changes in states are due to changes in the forces between the molecules rather than in those within the molecules. This is an example of the types of forces holding water molecules together. Water is made of molecules and intermolecular forces hold them together to other water molecules. There are many different particles that form liquids and solids other than just molecules. You will be studying the different types of particles that make up liquids and solids and the forces that hold them together as liquids and solids. ****IN THIS SECTION, WHEN CONSIDERING THE TYPES OF FORCES BETWEEN PARTICLES OF LIQUIDS AND SOLIDS (INTERPARTICLE FORCES), YOU MUST ALWAYS AND FOREMOST ASK WHAT KIND OF PARTICLES ARE BEING ATTRACTED TO EACH OTHER? THE POSSIBLE TYPES OF PARTICLES ARE: IONS ATOMS BY THEMSELVES: o GROUP 8A ATOMS (INERT GASES) o METAL ATOMS o SPECIFICALLY CARBON ATOMS AND SILICON ATOMS (WITH OXYGEN AND/OR CARBON) MOLECULES TYPES OF INTERPARTICLE FORCES In order to understand the forces between solid and liquid particles (interparticle forces), we will first determine the types of forces that can exist between particles and then move to a description of the liquid and solid phases. There are different forces of attraction involved in solids and liquids. From strongest to weakest they are: 1. Network Covalent Bonds bonds between single carbon atoms; bonds with silicon atoms (explained later). 2. Ionic Bonds bonds between ions 3. Metallic Bonds bonds between metal atoms 4. Intermolecular forces bonds between molecules (and Group 8A atoms). Three types: (a) Hydrogen bonds (b) Dipole-dipole IMF s (c) London dispersion forces (Induced dipole-induced dipole IMF s) STRENGTH OF INTERPARTICLE FORCES STRONG BONDS (strongest to weakest) 1. Network Covalent Bonds 2. Ionic Bonds 3. Metallic Bonds WEAK BONDS (strongest to weakest) Intermolecular Forces which include: 1. Hydrogen Bonding (a very strong dipole-dipole force) 2. Dipole-Dipole Forces (about 1% as strong as covalent or ionic bonds) 3. London Dispersion Forces (Induced Dipole-Induced Dipole Forces) ** London Dispersion Forces exist between all molecules in the liquid and solid phase. The strong forces are so strong compared to LDF s that LDF s are ignored as a force in particles attracted by the 3 strong forces above. All these forces will be explained first briefly. January 25 Test: February 10 07/12/2006

CHEM. 1 HONS UNIT 8 CH. 11 IMF s and Liquids and Solids 4 THE STRONG FORCES 1. NETWORK COVALENT BONDS Strongest strong bond. Occurs only in solids In a network solid, atoms are covalently bonded in a crystal (lattice) structure. Network solids are very hard and their strong bonds result in very high melting and boiling point temperatures. Network solids have localized electrons that are in fixed positions in the covalent bonds. The shared electrons stay localized between the 2 atoms that they bond together. This makes network solids poor conductors of electricity because the electrons cannot move. Examples of network covalent solids include o Diamond a form of carbon in which its atoms have an sp 3 hybridization and the atoms are arranged in a tetrahedral network. Diamond consists of millions of carbon atoms bonded to each other by covalent bonds. The tetrahedral structure and the network of covalent bonds makes diamond very strong. Each carbon is surrounded by a tetrahedral arrangement of other carbon atoms to form a huge molecule. The term network is used to indicate that the structure is stabilized by the covalent bonds in a tetrahedral pattern that gives strength to the structure. o Graphite o Quartz crystal (rock) SiO 2 and silicon carbide SiC. 2. IONIC BONDS (Between ions) Medium strong bond. In an ionic solid, adjacent ions have electrostatic attractions (attraction of + ions to ions) and are arranged in a crystal lattice structure. Ionic solids have strong bonds resulting in high melting and boiling point temperatures. Ionic solids have localized electrons which are in fixed positions around the atoms. This makes ionic solids poor conductors of electricity. Ionic liquids however (ex. molten NaCl) do conduct electricity because the electrons are free to move about. Ionic solids that dissolve in water are also good conductors of electricity. Strength of the ionic bond The strength of an ionic bond depends on 2 factors: (a) The product of the charges of the 2 ions and (b) The radius of the ions When considering the strength of an ionic bond, use the product of the charges first. If they are the same, then use radius secondly. Product of the charges: The charges on the ions in an ionic solid are directly proportional to the strength of the ionic bonds, that is the higher the charge on the ion, the stronger the ionic bond. The strength of the bond between 2 ions is equal to the product of their charges. Ex: The ionic bonds in MgO (charges +2 and 2) are stronger (product of their charges is 4) than those in NaF (charges +1 and 1). The product of their charges is 1. Ion radius: When the products of the charges is the same, then consider the sizes of the ions that make up the ionic solid. The sizes of the ions in an ionic solid are inversely proportional to the strength of the ionic bonds, that is the larger the ion, the weaker the ionic bond. Ex: The ionic bonds in LiF are stronger than those in KF (because Li is much smaller than K). Ex: Which ionic solid in each pair would have a higher melting point? (a) AlCl 3 and NaBr (d) (b) CaO and Na 2 S (e) (c) LiF and MgCl 2 LiCl and NaCl CaO and MgO January 25 Test: February 10 07/12/2006

CHEM. 1 HONS UNIT 8 CH. 11 IMF s and Liquids and Solids 5 3. METALLIC BONDS (bonds between metal atoms) Weakest strong bond Metallic solids are often described as a group of nuclei surrounded by a sea of mobile electrons. In other words, the electrons in metallic solids are delocalized and are free to move about; making metals good conductors of electricity. The metal atoms give off electrons and become ions. It is the attraction of the cations to the sea of electrons that keeps the metal atoms together in the solid state. Metallic bonds are still relatively strong forces of attraction and most metals have high melting and boiling point temperatures. Mercury is the only metal that is not solid at room temperature. Though most metals are very hard, the unrigid structure of the electrons makes them malleable (can be shaped) and ductile (can be made into wires). The electrons are delocalized they move amongst all the metal atoms and do not stay between just 2 of them. Like in ionic solids, the charge of the ions formed and the radius of the atom can be used to describe the relative strengths of metallic solids. For example, which metal is stronger and has a higher melting point, lead or sodium? When considering the interparticle force between metal atoms, the first determining factor is the size of the ionic charge. The larger the ionic charge of the cation, the more strongly held they are and the higher the melting point. Melting point Pb 4+ > melting point Na + even though Na + has a smaller ionic radius. Ex.: Which metal is stronger and has a higher melting point, potassium or sodium? Both have ionic charges of +1 so you must consider the radius of the atoms. The size of the metal atoms is inversely proportional to the strength of the metallic bonds. In other words, the smaller the atoms, the stronger the force. Smaller size allows for the positively charged nucleus of one metal atom to be closer to the negatively charged electrons of another, increasing the strength of the attraction between them. Forces of Na + > forces of K + because Na + has a smaller radius. Therefore, Na metal has a higher melting point than K metal. THE WEAK FORCE INTERMOLECULAR FORCES (the weak forces) are forces between molecules. 1. DIPOLE-DIPOLE INTERMOLECULAR FORCES (medium weak bond) A dipole-dipole force exists between neutral polar molecules. Polar molecules attract each other when the positive end of one molecule is attracted to the negative end of another molecule. Dipole-dipole forces are about 1% as strong as covalent or ionic bonds. They become weaker with distance. The greater the polarity (difference in electronegativity of the atoms in the molecule), the stronger the dipole-dipole attraction. Dipole-dipole attractions are very weak and substances held together by these forces have low melting and boiling point temperatures. Generally, substances held together by dipole-dipole attractions are gases at room temperature. In general, the melting point and boiling point for a polar compound are higher than those for a nonpolar compound of approximately equal molar mass. Substance Molar mass (g/mole) Boiling point (K) Propane (non polar) 44 231 Acetaldehyde (polar) 44 294 The presence of permanent dipoles influences how soluble a substance is. Polar molecules dissolve in polar solvents and nonpolar molecules dissolve in nonpolar solvents giving rise to the saying LIKE DISSOLVES LIKE. Water (polar) dissolves alcohol (polar). Water and oil (non-polar) do not mix but oil and gasoline mix because both are non-polar. January 25 Test: February 10 07/12/2006

CHEM. 1 HONS UNIT 8 CH. 11 IMF s and Liquids and Solids 6 2. LONDON DISPERSION FORCES (Induced dipole-induced dipole IMF s - weakest weak bond) London dispersion forces (LDF) exist in all molecules and are the only forces that exist between neutral, nonpolar molecules. LDF occur due to the random motion of electrons. Non-polar molecules must exert some kind of force or they would never solidify. They are the most common of all IMF s and are found in all molecular substances. They are electrostatic in nature and arise from attractions involving induced dipoles. The fact that non-polar gases can be liquefied tells us that there must be some kind of attractive interactions between the particles. Dispersion forces are attractive forces that arise as a result of temporary dipoles induced in atoms or molecules. O 2 is nonpolar and isn t expected to be very soluble in water which is highly polar. The very fact that it dissolves in water at all is due to the ability of polar molecules to induce (to make happen) a dipole in a nonpolar molecule. The approach of the polar molecule causes a shift in the electron density of the nonpolar molecule thus inducing a temporary or instantaneous dipole in the otherwise nonpolar substance. All the electrons will shift to one side of the molecule. The process of inducing a dipole is called polarization and the degree to which the electron cloud of an atom or molecule can be distorted is called its polarizability. The more electrons there are in atom or molecule, the more easily the electrons can be polarized, because they are further from the restraining forces of atomic nuclei. Polarizability therefore increases with increasing size, and increasing size is usually accompanied by increasing molar mass. Polarizability is thought of to be the measure of the squashiness of its electron cloud. DISPERSION FORCES EXIST BETWEEN ALL MOLECULES (whether polar or non-polar) AND GROUP 8A ATOMS IN THE LIQUID OR SOLID STATE. THE ONLY ATTRACTIVE IMF S THAT EXIST BETWEEN NONPOLAR MOLECULES AND GROUP 8A ATOMS ARE DISPERSION FORCES. DISPERSION FORCES INCREASE IN STRENGTH WITH INCREASING MOLECULAR WEIGHT. (See chart below). This table shows the increased boiling point as the molar mass (and number of electrons increases. MELTING POINTS OF SIMILAR NONPOLAR COMPOUNDS COMPOUND MELTING POINT ( C) CH 4 182.5 CF 4 150.0 CCl 4 23.0 CBr 4 90.0 CI 4 171.0 In general, dispersion forces are weaker forces than dipole-dipole IMF s. Contradiction: In many cases, dispersion forces are the same or even greater than dipole-dipole forces between polar molecules. This is because some nonpolar molecules contain many more electrons than the electrons in the molecule with a dipole. The electrons are held very loosely so a very strong induced dipole is created. In fact, dispersion forces between polar molecules can contribute more to overall attractive forces than do dipole-dipole forces. For example, the B.P. of HBr is 206.2K. The B.P of HCl is 189.5K. The stronger attractive forces for HBr cannot be attributed to greater dipole-dipole forces because HCl is more polar so it should have a higher B.P. The higher B.P of HBr (a heavier molecule) is due to the stronger LDF s due to the increased molecular weight and size. In this case the LDF attractive forces contribute more to the attraction between molecules than does the dipole-dipole interaction. Here is another example: CH 3 F polar (M.P. 141.8 C) IMF s: dipole-dipole, dispersion forces CCl 4 nonpolar (M.P. 23 C) IMF s: dispersion forces only. The M.P. of CH 3 F should be higher because dipole-dipole forces are stronger. BUT, CCl 4 has a higher M.P. because it contains more electrons (higher molar mass). As a result, the dispersion forces between CCl 4 molecules are stronger than the dipole-dipole forces between CH 3 F molecules. January 25 Test: February 10 07/12/2006

CHEM. 1 HONS UNIT 8 CH. 11 IMF s and Liquids and Solids 7 LDF s are very weak and short-lived. To form a solid when only LDF s exists requires very low temperatures. The molecules or atoms must be moving slowly enough for the LDF s to hold the molecules or atoms together in a solid unit. Notice that as the molar mass of the noble gas increases, the freezing point increases. This implies that the LDF between the atoms is stronger as the molar mass increases. Element Freezing Point ( C) helium 269.7 neon 248.6 argon 189.4 krypton 157.3 xenon 111.9 3. HYDROGEN BONDING (strongest weak bond after ion-dipole force) A hydrogen bond is an especially strong dipole-dipole IMF. It exists between the hydrogen atom in a polar bond and is a bridge between hydrogen and highly electronegative atoms namely F, O, and N. (FON) Because H is very small and the other atoms are highly electronegative, a very strong dipole is produced. Ex. H 2 O, HF, NH 3. Example 1: Which of the following can form hydrogen bonds with water? (Hint: Draw the structures first and determine the shape. (a) CH 3 OCH 3 (b) CH 4 (c) F (d) HCOOH (e) Na + Example 2: Which of the following species are capable of hydrogen bonding among themselves? (Hint: Draw the structures first.) (a) H 2 S (b) C 6 H 6 (c) CH 3 OH ONE MORE WEAK IMF: - ION-DIPOLE FORCES Useful for explaining soluble ionic compounds in water (NaCl dissolving in water). An ion-dipole interaction that is of particular interest to chemists involves the interaction of ions with water molecules, because these interactions are involved in the dissolving of ionic solids such as salt NaCl dissolved in water. A dipole consists of separated positive and negative charges, so an ion can be attracted to one end of the dipole and be repelled by the other. Polar molecules are dipoles; they have a positive end and a negative end. Ion-dipole forces are especially important in solutions of ionic substances in polar liquids. January 25 Test: February 10 07/12/2006

CHEM. 1 HONS UNIT 8 CH. 11 IMF s and Liquids and Solids 8 TRENDS IN THE PERIODIC TABLE Boiling points and melting points have a tendency to increase as molecular weight increases. As you go left to right across a period, the boiling points of similar substances containing elements in the same period should increase with increasing molecular mass. (See chart below for CH 4, NH 3, HF, H 2 0.) Contrary to the trend, the lightest compound has the highest boiling point contrary to our expectations based on molar mass (should be higher molar mass = higher B.P.). Expected B.P.: HF > H 2 O > NH 3 > CH 4 Actual B.P.: H 2 O > HF > NH 3 > CH 4 Why this exception? All are polar except CH 4. It has only dispersion forces, therefore CH 4 has the lowest B.P. as expected. H 2 O has extensive H bonding due to 2 unshared pairs of electrons on O. HF and NH 3 both have H bonding but HF is more polar due to the high electronegativity of F. Another exception: As you go down a family or group, the B.P. is expected to increase due to increased molecular mass. In group 4A the trend is followed. In groups 5A, 6A, and 7A the trend is not followed due to the effect of strong hydrogen bonding. (See figure below.) LINEAR AND BRANCHED MOLECULES - FORCES IN MOLECULES WITH THE SAME MOLECULAR WEIGHT Two molecules have the same molecular formula C 4 H 10 (M.W. 58.14 amu) but different structures: n-butane (B.P. 0.5 C) methylpropane (B.P. 11.7 C) Linear structures are held more strongly by IMF s because they can be stacked making more bonds. (See the diagrams below.) January 25 Test: February 10 07/12/2006

CHEM. 1 HONS UNIT 8 CH. 11 IMF s and Liquids and Solids 9 DO ASSIGNMENT #1 P. 9 NOTES #1-10 ASSIGNMENT #1: INTERMOLECULAR FORCES (Do these on a separate piece of paper) 1. Indicate the most important types of interparticle forces that exist in each of the following substances in the liquid or solid phase: (a) BaSO 4 (b) H 2 S (d) C 2 H 6 (e) Cs (g) P 4 (h) SO 2 (j) H 2 O (c) Xe (f) Hg (i) P 8 2. Which of the following substances should have the strongest intermolecular forces and why? N 2, Ar, F 2, Cl 2 3. HCl has a higher dipole moment than HCN. Which one should have the highest boiling point and why? 4. Put the following molecules in order of increasing polarizability: GeCl 4, CH 4, SiCl 4, GeBr 4. 5. Predict which would have the greater interparticle force: (a) CO 2 or OCS (b) PF 3 or PF 5 (c) SO 2 or SO 3 (d) SF 2 or SF 6 6. Indicate the types of interparticle forces that exist between molecules (or basic units) in each of the following: (a) PCl 3 (b) CO 2 (c) Cl 2 (d) ICl (e) KCl 7. Which of the following should have the strongest intermolecular forces? (a) CH 4 (b) Cl 2 (c) CO (d) CS 2 8. Which member of each pair should have the higher boiling point? (a) CH 3 F and CH 4 (c) CH 3 CH 2 O H and CH 3 O CH 3 (b) HF and HCl (d) CH 4 and GeH 4 9. Which member of each pair should have the higher boiling point? (a) O 2 or CO (c) H 2 O or H 2 S (b) Br 2 or ICl (d) PH 3 or AsH 3 10. What type of forces must be overcome in order to: (a) Vaporize water (b) Dissociate H 2 into H atoms (c) Boil liquid O 2 DO ASSIGNMENT #2 P. 10 NOTES #1-13 January 25 Test: February 10 07/12/2006

CHEM. 1 HONS UNIT 8 CH. 11 IMF s and Liquids and Solids 10 ASSIGNMENT #2: INTERPARTICLE FORCES 1. (a) Explain the term polarizability. What kind of molecules tend to have high polarizabilities? (b) What is the relationship between polarizability and intermolecular forces? 2. Give some evidence that all atoms and molecules exert attractive forces on one another. 3. Which elements can take part in hydrogen bonding? 4. Why is hydrogen unique in this kind of interaction? 5. The compounds Br 2 and ICl have the same number of electorons, yet Br 2 melts at 7.2 C and ICl melts at 27.2 C. Explain. 6. Which member of each of the following pairs of substances would you expect to have a higher boiling point? (a) O 2 and N 2 (b) SO 2 and CO 2 (c) HF and HI 7. Which substance in each of the following pairs would you expect to have the higher boiling point? Explain why. (a) Ne or Xe (b) CO 2 or CS 2 (c) CH 4 or Cl 2 (d) F 2 or LiF (e) NH 3 or PH 3 8. Explain in terms of intermolecular forces why: (a) NH 3 has a higher boiling point than CH 4 and (b) KCl has a higher melting point than I 2 9. What kind of attractive forces must be overcome in order to (a) Melt ice (b) Boil molecular bromine (c) Melt solid iodine (d) Dissociate F 2 into F atoms. 10. Ammonia is both a donor and an acceptor of hydrogen in hydrogen-bond formation. Draw a diagram showing the hydrogen bonding of an ammonia molecule with two other ammonia molecules. 11. Which of the following species are capable of hydrogen-bonding among themselves? (a) C 2 H 6 (c) KF (b) HI (d) BeH 2 (e) CH 3 COOH 12. Arrange the following in order of increasing boiling point: RbF CO 2 CH 3 OH CH 3 Br. Explain your reasoning. 13. Identify the most important types of interparticle forces present in the solids of each of the following substances: (a) NH 4 Cl (e) CHCl 3 (b) Steel (f) Ge (c) Teflon, CF 3 (CF 2 CF 2 ) 8 CF 3 (g) NO (d) Polyethylene, CH 3 (CH 2 CH 2 ) 8 CH 3 (h) BF 3 January 25 Test: February 10 07/12/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 11 PROPERTIES OF LIQUIDS A typical liquid might best be viewed as containing a large number of regions where the arrangements of the components are similar to those found in the solid, but with more disorder, and a smaller number of regions where holes are present. SURFACE TENSION IMF s pull the molecules into the liquid and cause the surface to tighten like an elastic film. To increase a liquid s surface area, molecules must move from the interior of the liquid to the surface. (A water drop changing from a round spherical drop to being flat and spread out) This requires energy, since some IMF s must be overcome. The resistance of a liquid to an increase in its surface area is called the SURFACE TENSION of a liquid. LIQUIDS WITH LARGE IMF S (LIKE POLAR MOLECULES) HAVE HIGH SURFACE TENSIONS. Because of hydrogen bonding, water has a considerably greater surface tension than most other liquids. CAPILLARY ACTION The spontaneous rising of a liquid in a narrow tube. This is an example of surface tension. A thin film of water sticks (adheres) to the walls of the glass tube. The surface tension of water causes the film to contract and as it does it pulls the water up the tube. Two types of forces are responsible for this property: 1. COHESIVE FORCES The IMF S among molecules of the liquid (attraction between like molecules like water). This is called COHESION. 2. ADHESIVE FORCES Forces between the liquid molecules and their container (attraction between unlike molecules like water and glass). This is called ADHESION. If adhesion is stronger than cohesion, the contents of the tube will be pulled upward (water in a thin tube). In mercury, cohesion is greater than the adhesion between mercury and glass so that when a capillary tube is dipped in mercury, the result is a depression or lowering at the mercury level. The height of the liquid in the capillary tube is below the surface of the mercury. Water Mercury H 2 O in glass Hg in glass Question: When water beads up on a freshly waxed car, which is stronger, cohesive or adhesive forces? VISCOSITY Is a measure of the resistance to flow of a liquid. Syrup flows more slowly than water therefore syrup has a higher viscosity. Liquids whose molecules have stronger IMF s have greater viscosities than liquids with weak IMF S. (Water strong H bonds; oil has dispersion forces only but oil is formed from very large molecules and that makes them more viscous.) January 25 Test: February 10 07/11/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 12 PROPERTIES OF WATER. Unusual properties of water: Water has a considerably greater surface tension than most other liquids. This is because each atom of oxygen in water can form 2 hydrogen bonds. The boiling point of water is about 200 C higher than might be reasonably expected. Boiling point tends to be related to the molar mass. Here are some that are similar in molar mass (see table below) Liquid Molar mass Boiling Pt. (g/mole) ( C) H 2 O 18 100 CH 4 16 159 NH 3 17 33 If water did not have the ability to form hydrogen bonds, it would be a gas at room temperature Density versus temperature. As liquids cool, they become more and more dense (the same mass in a smaller volume). As water is cooled from 100 C down to 4 C it does get denser like other liquids. From 4 C down to 0 C, the opposite happens: it gets less dense. Water has its greatest density at 4 C. The density of water @ 4 C = 1.00 g/ml; density at 0 C = 0.92 g/ml. Density of liquid versus density of solid. Of over 7 million chemical compounds almost all have their densities in the solid state greater than the density in the liquid state, as expected. Only about 12 chemicals are exceptions; density of the liquid is greater than the density of the solid. Because of this property, ice floats on water. DO ASSIGNMENT #3 P. 12 NOTES #1-8 ASSIGNMENT #3: LIQUIDS 1. Phosphorus trichloride PCl 3 is more volatile than arsenic trichloride, AsCl 3 at 25 C. (a) Which substance has the greater intermolecular forces? Explain. (b) Which substance has the higher vapor pressure at 25 C? (c) Which substance will have the higher boiling point? 2. Two pans of water are on different burners of a stove. One pan of water is boiling vigorously, while the other is boiling gently. What can be said about the temperature of the water in the two pans. 3. Water is a special substance with special properties. (a) Why does water expand when it freezes? (b) Why do snowflakes have a hexagonal shape? (c) Give an example of the biological significance of the fact that ice is less dense than water. 4. How do the viscosity and surface tension of liquids change as the intermolecular forces become stronger? 5. How do the viscosity and surface tension of liquids change as temperature increases? 6. Explain the following observations: (a) The viscosity of the covalently bonded molecule ethanol CH 3 CH 2 O H is greater than that of the covalently bonded molecule ether, CH 3 CH 2 O CH 2 CH 3. (b) In contact with a narrow capillary tube made of polyethylene, water forms a concave-downward meniscus like that of mercury in a glass tube. 7. Explain the following observations: (a) The surface tension of CHBr 3 is greater than that of CHCl 3. (b) As temperature increases, oil flows faster through a narrow tube. (c) Rain drops that collect on a waxed automobile hood take on a nearly spherical shape. 8. Consider the property of viscosity. (a) What is viscosity? (b) What is the relationship between intermolecular forces and viscosity? January 25 Test: February 10 07/11/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 13 THE SOLID STATE We will now study how the properties of solids relate to their structures and bonding. Solids can be either: o crystalline or o amorphous (noncrystalline). AMORPHOUS SOLIDS These substances appear to be solids but do not have a well-defined crystalline arrangement. They can be thought of as very slow flowing liquids. Glass (SiO 2 silicon dioxide) is an example. Over the years glass will be thicker on the bottom than on the top. Even though glass is made of SiO 2 like in quartz crystal, the glass cools too quickly and does not have a chance to form the well-ordered network covalent bonds. Another example of an amorphous solid is candle wax. Crystalline Solids: A crystalline solid is a solid whose atoms, ions, or molecules are ordered in well-defined arrangements called crystals. These solids usually have flat surfaces or faces that make definite angles with one another. The orderly stacks of particles that produce these faces also cause the solids to have higher regular shapes. Quartz and diamond are examples of crystalline solids. The components (atoms, ions, or molecules) have a regular arrangement and the positions of the components are represented by a lattice. A lattice is a three-dimensional system of points designating the positions of the components (atoms, ions, or molecules) that make up the substance. The unit cell (right) is the smallest repeating unit of the lattice; the individual particles are located at the lattice points (the corners of the unit cell) Crystalline solids can be one of the following: Covalent network (atomic) crystals made up of non-metal atoms. They have atoms at the lattice points. Ionic crystals (ionic compounds). They have ions at the lattice points. Metallic crystals (metal atoms). They have metal atoms at the lattice points. Molecular crystals (made up of molecules and atoms of Group 8A). They have molecules at the lattice points; Group 8A atoms are included because they have the same forces as molecular compounds (LDF s.)). ICE A Common Crystalline Solid Ice is a crystalline solid. Crystalline metals and compounds possess a rigid order; its atoms, molecules, or ions occupy specific positions in an orderly crystal lattice. The arrangement of such particles in a crystalline solid is such that the net attractive intermolecular forces are at their maximum. The forces responsible for the stability of a crystal can be ionic forces, covalent bonds, metallic bonds, hydrogen bonds, or a combination of these forces. Amorphous solids such as glass lack a well-defined arrangement or molecular order. Crystal lattices are built of unit cells that are the smallest repeating unit that has all of the symmetry characteristics of the overall structure. Therefore, the external appearance of a crystalline solid is a reflection of its unit cell (example a salt crystal NaCl). METALLIC SOLIDS (see figure right) Consist entirely of metal atoms. Bonding in metals is too strong to be due to London dispersion forces and yet there are not enough valence electrons for covalent bonding. The bonding is due to delocalized valence electrons that are throughout the entire solid. Visualize the metal as an array of positive ions immersed in a sea of delocalized electrons. Metals vary in strength of bonding as shown by their wide range of physical properties such as hardness and melting point. The strength of the bonding increases as the number of electrons available for bonding increases. Sodium: 1 valence electron M.P. 98 C. Chromium: 6 valence electrons M.P. 1890 C. The melting point of a metal depends on its charge. If the ions have the same charge, the size of the ion will determine the melting point. Sodium metal melts at 98 C whereas potassium metal melts at 63 C. This is due to the smaller ionic radius of the sodium ion and the higher melting point. The movement of the electrons makes metals good conductors of heat and electricity. January 25 Test: February 10 07/11/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 14 NETWORK COVALENT SOLIDS These are macromolecules (very large molecules) in which atoms are held together in large networks or chains by covalent bonds (diamond is one big molecule with millions of carbon atoms joined held together by covalent bonds. The lattice points (points of the crystal) are occupied by atoms held together by a network of covalent bonds. Formed by elements of group 4A (C, Si,). Best examples are diamond, graphite, quartz crystal (SiO 2 ), silicon carbide SiC. In SiO 2, (like in diamond) Si atoms are bonded such that there are oxygen atoms between each pair of Si atoms. In SiC, carbon and silicon atoms are bonded covalently in a tetrahedral shape. Have high melting points and are extremely hard because of the large number of covalent bonds that have to be broken to melt or break up the crystal. The entire crystal is thought of as one giant molecule. Extremely high M.P. and B.P. (M.P. diamond 3350 C.) Diamond a network covalent solid. Diamond is the hardest substance known to man. Each carbon is bonded to 4 other carbon toms using sp 3 hybrid orbitals. This forms a tetrahedral structure which makes the solid very strong. Graphite Consists of 2-dimensional sheets of carbon atoms. It is a good lubricant because the layers can slide over each other. This is because the attractive forces between layers are relatively weak intermolecular attractions. Each carbon is covalently bonded to 3 other carbon atoms in a trigonal planar sp 2 geometry that gives graphite its structure of flat sheets. The extra p electron that is not used in the sp 2 bonding holds these sheets together. The weak bonding of the p electrons allows the flat sheets to slide over each other easily and is responsible for the slippery feel of graphite to conduct electricity. The 2p orbitals form π molecular orbitals above the plane of the rings. The electrons are delocalized in these π molecular orbitals. These delocalized electrons allow for electrical conductivity. Graphite is slippery, black, and a conductor. SiO 2 is the empirical formula for sand and quartz. Silicon dioxide forms a covalent crystal with each silicon forming bonds to 4 oxygen atoms and each oxygen bonding to 2 silicon atoms with a tetrahedral geometry (just like diamond). Silicon carbide is another network crystal similar to diamond with alternating tetrahedral silicon and carbon atoms. It is very hard and is used as an industrial substitute for diamonds. Network crystals, like ionic substances, are represented by their empirical formulas. Two-dimensional representation of (a) crystalline quartz and (b) noncrystalline quartz glass. The small spheres represent silicon, in reality, the structure of quartz is threedimensional. Each Si atom is tetrahedrally bonded to four O atoms. January 25 Test: February 10 07/11/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 15 Glass an amorphous solid Formed when silica is heated and cooled rapidly. More closely resembles a viscous solution than a crystalline solid Adding different substances to the melted silica results in different properties for the glass Add B 2 O 3 to produce glass for lab ware (PYREX). It undergoes very little expansion or contraction with large temperature changes. Add K 2 O to produce a very hard glass that can be ground for eyeglasses or contacts. IONIC SOLIDS Are 3-dimensional crystalline arrangements of positive and negative ions at the lattice points. The opposite charged ions are held together by electrostatic attractions (ionic bonds) These bonds are relatively strong so ionic solids have high M.P. s. Each ion in an ionic crystal is surrounded by ions of opposite charge. The charge and size of the ions affects the M.P. of the solid. NaCl M.P. 801 C, MgO M.P. 2800 C. The ions of Mg and O with charges of +2 and 2 (product of charges is 4) have a much stronger attraction for one another than the ions of Na and Cl which are +1 and 1 (product of charges is 1). When solid, are non-conductors of electricity but will be a conductor when melted or dissolved in water. Examples: NaCl, CsCl, MgO, Al 2 O 3. MOLECULAR SOLIDS Are made of molecules held together by weak IMF s (dipoledipole, hydrogen, dispersion). Group 8A (Noble Gases) solids are included here because they have the same forces as molecular solids (LDF s). Group 8A solids are atomic solids. You may find the Noble Gases included as molecular solids. Examples are dry ice (solid CO 2, iodine I 2, phosphorous P 4, sulfur S 8, H 2, O 2 ). You would expect low M.P. due to the weak forces but the size of some of these molecules makes the M.P. higher than expected. (CO 2 gas at room temperature but P 8 is a solid at room temperature as is iodine I 2 ). Other examples are ice and water. Are poor conductors of heat and electricity. (Grp. 8A atoms) (crystal, not glass) graphite, SiC DO ASSIGNMENT #4: P. 16 Notes #1-12 FORCES IN SOLIDS January 25 Test: February 10 07/11/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 16 ASSIGNMENT #4 FORCES IN SOLIDS 1. Give and describe examples of the following types of crystals: (a) Ionic crystals (b) Molecular crystals (c) Metallic crystals (d) Network covalent crystals. 2. A solid is hard, brittle, and electrically nonconducting. Its melt (the liquid form of the substance) and an aqueous solution containing the substance conduct electricity. Classify the solid. 3. A solid is soft and has a low melting point (below 100 C). The solid, its melt, and an aqueous solution containing the substance are all nonconductors of electricity. Classify the solid. 4. A solid is very hard and has a high melting point. Neither the solid nor its melt conducts electricity. Classify the solid. 5. Why are metals good conductors of heat and electricity? Why does the ability of a metal to conduct electricity decrease with increasing temperature? 6. Which of the following are molecular solids and which are covalent solids? (a) Se 8 (d) CO 2 (b) HBr (e) C (c) Si (f) P 4 O 6 (g) SiH 4 7. Classify the solid state of the following substances as ionic crystals, covalent crystals, molecular crystals, or metallic crystals: (a) CO 2 (b) B1 2 (c) S 8 (d) KBr (e) Mg (f) SiO 2 (g) LiCl (h) Cr 8. Explain why diamond is harder than graphite. Why is graphite an electrical conductor but diamond is not? 9. What is an amorphous solid? How does it differ from a crystalline solid? 10. Which of the following substances has the highest polarizability? CH 4, H 2, CCl 4, SF 6, H 2 S. 11. Which of the following statements are false? (a) Dipole-dipole interactions between molecules are greatest if the molecules possess only temporary dipole moments. (b) All compounds containing hydrogen atoms can participate in hydrogen-bond formation. (c) Dispersion forces exist between all atoms, molecules, and ions. (d) The extent of ion-induced dipole interaction depends only on the charge on the ion. 12. Carbon and silicon belong to Group 4A of the periodic table and have the same valence electron configuration (ns 2 np 2 ). Why does silicon dioxide SiO 2 have a much higher melting point than carbon dioxide CO 2? DO ASSIGNMENT #5 P. 17 Notes #1-4 LET THE FORCES BE WITH YOU January 25 Test: February 10 07/11/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 17 NAME DATE PERIOD ASSIGNMENT #5 LET THE FORCES BE WITH YOU!! 1. What type of crystal will each of the following substances form in its solid state? Choices to consider are metallic, ionic, covalent, or molecular crystals. (a) C 2 H 6 (b) CO 2 (c) Al (d) Na 2 O (e) C (diamond) (f) SiO 2 (quartz) (g) NaNO 3 (h) SO 2 (i) BaSO 4 (j) Xe (k) P 4 (l) Si (m) Mo (n) Li 2 O (o) NaHSO 4 (p) S 8 2. Circle all the compounds in the following list which would be expected to form intermolecular hydrogen bonds in the liquid state: (a) CH 3 OCH 3 (dimethyl ether) (b) CH 4 (c) HF (d) CH 3 CO 2 H (e) Br 2 (f) H 2 S 3. Specify the forces involved for each substance in the space immediately following the substance, then in the last column, indicate which member of the pair you would expect to have the higher boiling point. Substance #1 Force(s) Substance #2 Force(s) Substance with Higher Boiling Point (a) HCl(g) I 2 (b) CH 3 F (c) H 2 O CH 3 OH H 2 S (d) SiO 2 SO 2 (e) Fe (f) CH 3 OH Kr CuO (g) NH 3 CH 4 (h) HCl(g) NaCl (i) SiC Cu 4. Distinguish between the following pairs of terms: (a) Polarizability and polarity (b) London dispersion forces and dipole-dipole forces (c) Intermolecular forces and intramolecular forces January 25 Test: February 10 07/11/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 18 DO ASSIGNMENT #6 P. 18 NOTES #1-28 ASSIGNMENT #6: PHASE CHANGES READ TEXT: P. 442 453 OMIT Clausius-Clapryon Equation P. 446, 447 Define the following terms: 1. Phase change 2. Vaporization 3. Boiling 4. Boiling point 5. Normal boiling point 6. Evaporation 7. Melting 8. Fusion 9. Freezing 10. Condensation 11. Sublimation 12. Deposition 13. Melting point 14. Freezing point 15. Dynamic equilibrium (short form: equilibrium) 16. Equilibrium vapor pressure (short form: vapor pressure) 17. Molar heat of vaporization ( H vap ) 18. Molar heat of fusion ( H fus ) 19. Critical temperature (T c ) 20. Critical pressure (P c ) 21. Specific heat of a substance 22. What is the relationship between boiling, evaporation, and vaporization? 23. What is the similarity between boiling and evaporation? 24. What is the difference between boiling and evaporation 25. At what temperature does water freeze? 26. At what temperature does ice melt? 27. What is the relationship between freezing point and melting point? 28. Label the diagrams HEATING AND COOLING CURVES on page 19. January 25 Test: February 10 07/17/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 19 HEATING AND COOLING CURVES A heating curve is a convenient way to summarize the solid-liquid-gas change for a compound. A typical heating curve from solid phase to gas is shown below. This is a typical heating cooling curve for any substance. The flat lines indicate there is no temperature change. Two phases co-exist in dynamic equilibrium (water and ice; water and water vapor). The heat going into the substance is being used to break the IMF s. The molecules are not increasing in kinetic energy. The flat lines are H fus and H vap. The sloped lines indicate a temperature change and this is where a single phase is being heated or cooled (heating water). January 25 Test: February 10 07/17/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 20 PHASE CHANGES Phase change: change of state. In any phase change, energy (in the form of heat) is added or removed. Molecular order: solids most order, gas, least order or most randomness. When a change of state involves going from a more ordered state to a less ordered state, energy must be supplied to overcome IMF s ex.: melting ice needs heat. Vaporization change of phase from liquid to gas. Evaporation vaporization that takes place at the surface of the liquid. Boiling vaporization that takes place everywhere in the liquid. As temperature, kinetic energy of molecules ; more molecules leave the liquid. Gases condense to liquids when the molecules no longer have sufficient kinetic energy to overcome the IMF s. Melting is also called fusion. The enthalpy (heat) change for melting a solid is called the enthalpy of fusion or heat of fusion H fus. Molar heat of fusion is the amount of heat required to melt 1 mole of a solid. Ex. H fus of ice = 6.01 kj/mole. This means it takes 6.01 kj of heat to melt 1 mole (18.02 grams) of ice. Molar heat of vaporization H vap is the heat required to vaporize 1 mole of a liquid. H vap water = 40.67 kj/mole. H vap as IMF s. Melting, vaporization and sublimation are endothermic processes (heat is taken in). Freezing, condensation and deposition are exothermic. Steam causes sever burns. When it comes in contact with skin, it condenses, releasing considerable heat. A steam burn is worse than a burn from boiling water. WHY? January 25 Test: February 10 07/17/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 21 Critical temperature and pressure Gases can change into liquids by either cooling them or applying pressure. As temperature rises, gases become more difficult to liquefy because of increasing kinetic energies of their molecules. For every substance there exists a temperature above which the gas cannot be liquefied regardless of the pressure. This is the critical temperature T c. Critical temperature T c : The temperature of a substance above which its gas phase cannot be made into a liquid no matter how great the pressure. (It is the highest temperature at which a substance can exist as a liquid.) Critical pressure T p : is the pressure required to bring about liquefaction (change a gas to a liquid) at the critical temperature. Significance of T c and T p : Above T c, molecules contain so much kinetic energy that they cannot be made into a liquid. Engineers and people that work with gases need to know T c and T p. It is useless to try to liquefy a gas by applying pressure if the gas is above its T c. O 2 has a T c = 154.4K ( 118.8 C). It must be cooled below this temperature before it can be liquefied by pressure. Above T c, molecules have so much kinetic energy that they are able to break away from any attraction or IMF s and they cannot be liquefied. At the temperature in the room today, is it possible to convert oxygen gas into liquid oxygen? Why or why not? Vapor Pressure (pressure of vapor): Vapor pressure of a substance is the measure of the tendency of its molecules to leave the liquid phase and enter the vapor phase at a given temperature. It is the pressure exerted by the gas phase of a liquid at a given temperature. The more volatile a substance is, the higher the vapor pressure at a given temperature. More volatile liquids evaporate more quickly because of weaker IMF s (alcohol is more volatile than water and has a greater vapor pressure at a fixed temperature.). Liquids placed in closed containers attain a state of dynamic equilibrium with their vapors. Dynamic equilibrium is a state in which the rate of a forward process is exactly balanced by the rate of the reverse process. They occur at the same rate. Here are some examples of dynamic equilibrium: At 0 C for water: rate of freezing = rate of melting In a saturated solution of NaCl with solid NaCl at the bottom of the container, rate of dissolving = rate of crystallization In this room: if 5 students leave and 5 students come in at the same time, rate of leaving = rate of entering Equilibrium vapor pressure is the vapor pressure measured when the rate of vaporization is equal to the rate of condensation. Equilibrium vapor pressure: RATE vaporization = RATE condensation Explanation: When a liquid evaporates, the gaseous molecules exert a vapor pressure. In a closed container, the process of evaporation does not continue indefinitely. At first, molecules are moving from the liquid to the empty space. As the concentration of molecules in the vapor phase increases, some molecules condense (return to the liquid state). At a fixed temperature, a state of dynamic equilibrium is eventually reached where the rate of evaporation is equal to the rate of condensation. As temperature, vapor pressure. For a particular liquid, vapor pressure increases only with an increase in temperature. If IMF s are strong, H vap is high; therefore, a liquid with high H vap has a low vapor pressure (molecules cannot escape very easily). January 25 Test: February 10 07/17/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 22 BOILING POINT AND HEAT OF VAPORIZATION The value of H vap will tell you the strength of the IMF s of a liquid. If IMF s are strong, it takes a lot of energy to free the molecules from the liquid phase. A liquid with low vapor pressure has a high H vap and strong IMF s. Boiling point is the temperature at which the vapor pressure of a liquid is equal to the external pressure (atmospheric pressure). The normal boiling point of a liquid is the temperature at which it boils when the external pressure is 1 atm. (See the diagram below) The normal boiling point of water is 100 C, mercury 357 C. The higher the boiling point, the greater the H vap. The boiling point decreases with increasing elevation and decreasing atmospheric pressure so foods take longer to cook at elevations higher than sea level. Foods take less time to cook in pressure cookers because the pressure inside the cooker is greater than atmospheric pressure and water boils at a higher temperature (water is hotter so less cooking time). Liquid solid equilibrium Freezing Point (or melting point) is the temperature at which the solid and liquid phases of a substance coexist at equilibrium. Normal melting (or freezing) point of a substance is the temperature at which a substance melts (or freezes) at 1 atm pressure. If a glass of water with ice cubes in it is held at a temperature of 0 C, the ice and water will be in dynamic equilibrium, that is: RATE melting = RATE freezing January 25 Test: February 10 07/17/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 23 INTERMOLECULAR FORCES Strong intermolecular forces lead to: Weak intermolecular forces lead to: Nonvolatile substances Volatile substances High boiling points Low boiling points Low evaporation rates High evaporation rates Low vapor pressures at room temperature High vapor pressures at room temperature DO ASSIGNMENT #7 P. 24-25 #1-10 Notes Vapor Pressure and Boiling Point January 25 Test: February 10 07/17/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 24 NAME DATE PERIOD ASSIGNMENT #7: VAPOR PRESSURE AND BOILING POINT 1. What is the definition of boiling point? 2. What is the definition of normal boiling point? 3. What effect does increasing temperature have on vapor pressure? 4. What is the normal boiling point of substance A? ; substance B? 5. Which of the two substances will boil at 70 C if the atmospheric pressure is 20 kpa? 6. Give the relative strength of the intermolecular forces in these two liquids. 7. Name all the possible phase changes that can occur between different states of matter. Which of these are exothermic and which are endothermic? January 25 Test: February 10 07/17/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 25 8. Explain the following observations: (a) During the cold winter months, snow often gradually disappears without melting. (b) The heat of fusion for any substance is generally lower than its heat of vaporization. (c) Ethyl chloride C 2 H 5 Cl boils at 12 C. When liquid C 2 H 5 Cl under pressure is sprayed on a surface at atmospheric pressure, the surface is cooled considerably. (d) When heated above 279 C, carbon disulfide CS 2 cannot be liquefied regardless of how great the pressure exerted on the gas. 9. Explain why the boiling point of a liquid varies substantially with pressure, whereas the melting point of a solid depends little on pressure. 10. Explain how each of the following affects the vapor pressure of a liquid (increase, decrease, no change): (a) Surface area (b) Temperature (c) Volume of the liquid. January 25 Test: February 10 07/17/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 26 SPECIFIC HEAT (or Specific Heat Capacity) (designated s ) The specific heat of a substance is the quantity of heat required to raise one gram of a substance one degree Celsius. The SI units of specific heat are J/g C. (See table below) The specific heat of water is 4.184 J/g C and of iron 0.444 J/g C. This means it takes 4.184 J of heat to raise the temperature of 1 gram of water 1 C. Iron heats up faster because it only takes 0.444 J of heat to raise the temperature of 1 gram of iron 1 C. This means with the same amount of heat applied, iron heats up faster than water. o The heating (or cooling) of a substance involves two different processes depending on whether a phase change is involved. change in temperature: t = t final - t initial Two formulas will be used: Heating (or cooling) a substance with no phase change: (change in temperature) q = ms t where q = heat m = mass (in g) s = specific heat (in J/g C)(from table below) t = change in temperature C t = t final t initial. Use the absolute value of t because a change has no sign. Ex: My weight changed by 10 lb. You do not know if I gained or lost 10 lb. If a sign is assigned, a gain in weight is +10 lb and if weight was lost it is 10 lb. Heating (or cooling) a substance to change phase. (no temperature change) q = n H fus where n = moles of the substance H fus = heat of fusion (from table below) q = n H vap H vap = heat of vaporization (from table below) Ice 2.03 Steam 1.99 C 2 H 5 OH (ethanol) 2.46 January 25 Test: February 10 07/17/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 27 **In the following problems, always determine if there are phase changes involved. **Use the absolute value of t. **Heat lost during cooling is the same as heat gained in heating. Ex. 1: A 466-gram sample of water is heated from 8.50 C to 74.60 C. Calculate the amount of heat absorbed by the water in kj. (129 kj) Ex. 2: A piece of iron initially at a temperature of 25 C absorbs 10.0 kj of heat. If its mass is 500. grams, calculate the final temperature of the piece of iron. The specific heat of iron is 0.444 J/g C. (70. C) Ex. 3: Calculate the amount of energy (in kj) needed to heat 346 g of liquid water from 0 C to 182 C. Assume that the specific heat of water is 4.184 J/g C and the specific heat of steam is 1.99 J/g C). (984 kj) Solution: The total heat involved must be added from the following processes: 1. Heating water from 0 C to 100 C. (use q = ms t) 2. Change of state water at 100 C to steam at 100 C. (use q = n H vap ) 3. Heating steam from 100 C to 182 C. (use q = ms t) Watch your units when using the tables. ****NO ROOM HERE. DO THIS PROBLEM ON A SEPARATE PIECE OF PAPER. DO ASSIGNMENT #8 P. 27 Notes #1-8 SPECIFIC HEAT ASSIGNMENT: #8: SPECIFIC HEAT 1. When a 50.0 g piece of nickel absorbs 350.0 J of heat, the temperature of the nickel changes from 20.0 C to 36.0 C. What is the specific heat of nickel? (0.438) 2. Calculate the heat (in kj) required to change 40.0 g of ice at 30.0 C to steam at 150.0 C? (127.0) 3. The specific heat of ethanol is 2.46 J/g C. How many joules of heat are required to heat 193 g of ethanol from 19.0 C to 35.0 C? (7600) 4. It takes 585 J of energy to raise the temperature of 125.6 g of mercury form 20.0 C to 53.5 C. Calculate the specific heat of mercury. (0.139) 5. The specific heat capacity of aluminum is 0.900 J/ C g. (a) Calculate the energy needed to raise the temperature of (a ) 8.50 x 10 2 gram block of aluminum from 22.8 C to 94.6 C. (b) Calculate the molar heat capacity of aluminum. ((a) 54.9 (b) 24.3) 6. How much heat in kj is lost when 100.0 g of steam at 200. C is cooled down until it forms ice at 30.0 C? (328) 7. How much heat is evolved when 10.0 grams of steam condenses at 100.0 C? (22.7) (You may see the answer given in #7 as 22.7 kj. This just means that this was an exothermic reaction and heat was given off.) 8. Calculate the amount of hear liberated (in kj) from 366 g of mercury when it cools from 77.0 C to 12.0 C.(3.31 kj) January 25 Test: February 10 07/17/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 28 PHASE DIAGRAMS The phase in which a substance exists depends on its temperature and pressure. In addition, two phases may exist at certain temperatures and pressures. A phase diagram as shown below summarizes information about the stable phases for a specific compound. The phase diagram for water. T m represents the normal melting point; T 3 and P 3 denote the triple point; T b represents the normal boiling point; T c represents the critical temperature; P c represents the critical pressure. The negative slope of the solid/liquid line reflects the fact that the density of ice is less than that of liquid water. NOTE: Phase diagrams are not drawn to scale. January 25 Test: February 10 07/17/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 29 A phase diagram is a graph of pressure versus temperature that shows the conditions under which the phases of a substance exist. A phase diagram also reveals how the states of a system change with changing temperature or pressure. The diagram is divided into three regions; one for the solid phase, one for the liquid phase, and one for the gas phase. The line between the solid and liquid regions is made up of points of P and T at which the solid and liquid phases are in equilibrium. The line between the solid and gas regions is made up of points of P and T at which solid and gas are in equilibrium. The line between the liquid and gas regions gives temperatures and pressures at which liquid and gas are in equilibrium. Since all three lines intersect at the triple point, this one point describes the conditions under which gas, liquid, and solid are all in equilibrium. (For water 0.01 C and 0.006 atm) The phase diagram shows at a glance several properties of a substance: normal melting point, normal boiling point, and triple point. If a point (P and T) describing a system falls in the solid region, the substance exists as a solid. If the point falls on a line such as that between liquid and gas regions, the substance exists as liquid and vapor in equilibrium. The vapor-pressure curve ends at the critical point which is at the critical temperature and critical pressure of the substance. The critical temperature T c is the temperature above which the substance cannot exist as a liquid. No matter how much pressure you subject the liquid to, it will not become a liquid above the critical temperature. The critical temperature of water is 373.99 C. Above this temperature, water cannot be liquefied no matter how much pressure is applied. The critical pressure P c is the lowest pressure at which the substance can exist as a liquid at the critical temperature. The critical pressure of water is 217.75 atm. Beyond the critical point, the liquid and gas phases become indistinguishable. The liquid gradually changes into a vapor but goes through an intermediate fluid region, which is neither true liquid nor vapor. This substance is called a supercritical gas. A supercritical gas is one at pressure so high that its density resembles a liquid s while its viscosity (ability to flow) remains close to that of a gas. **You will recognize the phase diagram for water because the solid-liquid line has a negative slope. Any substance in which the liquid phase is more dense than the solid phase (like water) will have a solid-liquid equilibrium line with a negative slope. Most other substances have a positive slope because the density of the liquid phase is less than that of the solid phase. The phase diagram of water. Each solid line between two phases specifies the conditions of pressure and temperature under which the two phases can exist in equilibrium. The point at which all three phases can exist in equilibrium (0.006 atm and 0.07 C) is called the triple point. (b) This phase diagram tells us that increasing the pressure on ice lowers its melting point and that increasing the pressure of liquid water raises its boiling point. January 25 Test: February 10 07/17/2006

CHEM. 1 HONS. UNIT 8 CH. 11 IMF s and Liquids and Solids 30 The phase diagram for carbon dioxide Similar to water with one exception: the slope of the curve between solid and liquid in CO 2 is positive. Because water expands when it freezes, the slope of the curve between solid and liquid in water is negative. this is how to distinguish the phase diagram between water and other substances. The triple point of CO 2 is 5. 2 atm and 57 C. Notice that from the phase diagram, carbon dioxide does not melt under normal conditions. it is not stable below 5.2 atm so that only the solid and vapor phases can exist under atmospheric conditions. DO ASSIGNMENT #9 P. 30-31 Notes #1-11 PHASE DIAGRAMS ASSIGNMENT #9 PHASE DIAGRAMS 1. What is a phase diagram? What useful information can be obtained from the study of a phase diagram? 2. Explain how water's phase diagram differs from those of most substances. What property of water causes the difference? 3. The blades of ice skates are quite thin, so that the pressure exerted on ice by a skater can be substantial. Explain how this fact enables a person to skate on ice. 4. A length of wire is placed on top of a block of ice. The ends of the wire extend over the edges of the ice, and a heavy weight is attached to each end. It is found that the ice under the wire gradually melts, so that the wire slowly moves through the ice block. At the same time, the water above the wire refreezes. Explain the phase changes that accompany this phenomenon. 5. The boiling point and freezing point of sulfur dioxide are 10 C and 72.7 C (at 1 atm), respectively. The triple point is 75.5 C and 1.65 x 10 3 atm, and its critical point is at 157 C and 78 atm. On the basis of this information, draw a rough sketch of the phase diagram of SO 2. (Do not draw to scale.) 6. A phase diagram of water is shown on the right. Label the regions. Predict what would happen as a result of the following changes: (a) Starting at A, we raise the temperature at constant pressure. (b) Starting at C, we lower the temperature at constant pressure. (c) Starting at B, we lower the pressure at constant temperature. 7. Why does the line on a phase diagram that separates the gas and liquid phases end rather than going to infinite pressure and temperature? 8. Refer to the phase diagram of water (a) on the previous page and describe all the phase changes that would occur in each of the following cases: (a) Water vapor originally at 1.0 x 10 3 atm and 0.10 C is slowly compressed at constant temperature until the final pressure is 10.0 atm. (b) Water originally at 10 C and 0.30 atm is heated at constant pressure until the temperature is 80.0 C. January 25 Test: February 10 07/17/2006