Unit 7 Notes Chemistry I CP Page 261 Your Key Chemistry Annotation Guide If you are NOT using the following annotation, put in your key to the left of each item. Mr. T s Key Items to be annotated Circle headings and subheadings. Box Write DEF Underline key content vocabulary next to definitions (sometimes #2 and #3 will be the same and in that case I expect to see a box AND DEF) important ideas (include captions for visuals). Mr. T s or Write Eq Underline Write Write in the margin next to an important equation. key steps to solving an example problem, or write those steps in the margin. any questions in the margin. a learning objective next to each subheading. Using this handout effectively The studying and annotations and other homework assignments are designed to require a maximum of 45 minutes to an hour of your time. This time limit, however, does not include things like test makup points or extra credit activities. Students are expected to plan their time. If students wait until the last minute to complete assignments that are assigned multiple days in advance, then the time required to complete the assignment may require more than the 1 hour limit. This document is a work in progress. I have made an attempt to put in as much about the knowledge you need for this unit of study as I can. This document, however, DOES NOT contain everything that you need to know to make 100% on the Unit 4 test. You must also rely on the knowledge that you should have acquired in previous units of study, classroom explanations, your notes, the textbook sections that were assigned for study, and in some cases creative thinking and problem-solving skills. Structural Particle Interactions & Gas Laws Unit 6: Chemistry I Honors The Standards in This Unit According to the South Carolina Science Standard C-5, students will demonstrate an understanding of the structure and behavior of the different phases of matter. Most of the material in this unit will come from standard. Under standard South Carolina Science Standards C-5 is a list of indicators and under each of those indicators are supporting documents. All the material in this section stems from these indicators and supporting documents. Indicator C-5.1 Indicator C-5.1 in the South Carolina Science Standards students should be able to explain the effects of the intermolecular forces on the different phases of matter. Key questions for this next section: 1) What does the term intermolecular forces mean? 2) What kind structural particles in an ionic compound? What kind of structural particles undergo dipoldipole attractions? What kind of structural particles undergo hydrogen bonding attractions? What kind of structural particles undergo London dispersion force attractions? 3) What does the term dipoldipole attraction mean? 4) What kind of structures undergo these dipoldipole attractions? Write your learning objective for this section here:
Unit 7 Notes Chemistry I CP Page 262 5) What are the 3 major factors that affect the state of matter (solid, liquid, or gas) of a substance? 6) What is hydrogen bonding? 7) What are the only elements that can bond with hydrogen and create a structure than can hydrogen bond? 8) For a structure that is capable of hydrogen bonding, what is the hydrogen attracted to when hydrogen bonding occurs? 9) What kind of illustration is used to represent hydrogen bonding between molecules? 10) What are London dispersion forces? 11) What kind of structures undergo the kind of attractions called London dispersion forces? 12) What must happen inside an atom or molecule for London dispersion forces to occur? 13) Describe the arrangement of structural particles in a solid? Describe the arrangement of structural particles in a liquid? Describe the arrangement of structural particles in a gas? 14) What is the term that describes the kind of attraction that occurs between oppositely charged ionic particles? 15) What is the name of the theory that describes the structural behavior of solid metals? Describe this theory. 16) Describe the structure of covalent network crystals, molecular solids, atomic solids, and amorphous solids. 17) What are the physical generalized physical properties of non-polar molecular compounds, polar molecular compounds, ionic compounds, metals, and covalent network compounds? 18) What does table 1 tell you about the effect of interparticle forces? 19) What does table 2 tell you about the effect of interparticle forces? 20) What does table 2 tell you about the effect of interparticle forces? 21) What do figures 6 and 7 tell you about the effect of interparticle forces? In this unit of chemistry students should: Diagram, describe, and give examples of the following intermolecular forces Dipoldipole attraction Hydrogen bonding London dispersion forces Compare intermolecular forces (dipoldipole interaction, and London dispersion forces) and ionic bonds, covalent bonds, and metallic bonds in terms of Nature of the attraction type of substance structural unit examples typical properties OK, so let s take these one at a time.
Unit 7 Notes Chemistry I CP Page 263 Students must be able to diagram, describe, and give examples of different kinds o fintermolecular forces Recall that the prefix inter- mean between, so intermolecular forces means the forces between molecules. As the remainder of the unit makes clear the real meaning of the material in this indicator refers to inter-particle forces. Intermolecular only refers to molecules, whether non-polar or polar. The structural particles in ionic compounds, metals, and covalent network compounds are not molecules at all. The structural particles in an ionic compound are typically ions, both monoatomic (a single atom that has lost or gained one or more electrons) and polyatomic (a covalently bonded set of atoms whose sum of electrons is more or less than the sum of all the protons of all the atoms). Students should know and be able to apply their knowledge of Dipoldipole attractions Dipoldipole attractions are the attractions between polar molecules. In a polar molecule there is an unequal sharing of electrons creating what are called partial positive and negative ends of the molecules. These partial positive ends of one molecule have an attraction to the partial negative ends of the next molecule. The 2 nd molecule is attracted to another and so forth and so on. This creates a kind of structure in which the molecules all kind of stick together with greater force than non-polar molecules. This molecular stickiness applies to molecules in a solid, liquid and even a gaseous form. This increased stickiness makes it harder to separate the molecules from each other. Molecules that stick together with enough force tend to the solids, molecules with a little less sticky force tend to be liquids, and molecules with still less stickiness tend to be gases. The force exerted by the stickiness of polarity cause the molecules to act somewhat like little magnets. The greater the overall electronegativity difference from one side of the molecules to the other the greater the sticky force. Whether a polar molecular substance is a solid, liquid, or gas depends a lot on how the polar stickiness is overcome by other factors and forces that affect the structural particles within the sample of substance. Two things that can overcome the polar stickiness are velocity and mass of the particles (which together make the momentum). If the molecules are moving fast enough (as when the substance is heated) then the motion of the molecules can shake them apart creating a liquid and then (if enough heat is added) creating a gas. If the polar molecules are massive enough (heavy enough) then it s harder to get them to move due to their momentum and polarity takes over and the molecules stick together in a liquid or a solid. So the state of matter is all a matter of balance between all the forces and factors that affect state of matter.
Unit 7 Notes Chemistry I CP Page 264 Strength of attraction between particles The size and shape of the particles. Factors that affect the state of matter of a substance. How fast the particles are moving. Figure 1. This Venn diagram illustrates some of the overlapping factors that affect whether a substance will be a solid, liquid, or gas. The greater the magnitude of the factors, the more likely the substance is to be a solid. The smaller the magnitude of the factors, the more likely the substance is to be a gas. Students should know and be able to apply their knowledge of Hydrogen Bonding Hydrogen bonding is an especially strong form of polarity that results from hydrogen being bonded to a small and very electronegative atom with lone pair electrons. Only 3 atoms are small enough and electronegative enough (and have lone pair electrons) to have hydrogen bonding potential when bonded with hydrogen: nitrogen (N), oxygen (O), and fluorine (F). The hydrogen on one molecule is attracted to a lone pair on the next molecule. Hydrogen bonding between one molecule and another is both quite strong for a polar attractions but it s also highly directional. By highly directional, I mean that a hydrogen bond has a very specific direction. From the molecule containing the hydrogen, the hydrogen bond falls on the same axis as the hydrogen s covalent bond with one of those 3 atoms. From the molecule with the lone pair, the hydrogen bond falls on the same axis as the direction of the lone pair from the nucleus of a nitrogen, oxygen, or fluorine. O H N H H Figure 2. An illustration of a water molecule (on the left) that is hydrogen bonded to an ammonia molecule (on the right). Hydrogen bonds have a specific length and direction. H H See these videos to help you understand this concept http://www.youtube.com/watch?v=lgwybeuvjhu http://www.youtube.com/watch?v=lkl5cbfqfrm
Unit 7 Notes Chemistry I CP Page 265 Students should know and be able to apply their knowledge of London dispersion forces London dispersion forces are weak forces between one molecule and another that come from a temporary polarity. London dispersion forces are also known as dispersion forces, London forces, or induced dipole dipole forces. London forces occur in nonpolar molecules because electrons move about a molecule s electron cloud randomly. At any given instant in time there is a high chance that the electron density will not be evenly distributed throughout a nonpolar molecule. When electrons are unevenly distributed, a temporary dipole exists. This dipole will interact with other nearby dipoles or will induce similar temporary polarity in nearby molecules. London forces are also present in polar molecules, but they are only a small part of the total interaction force. Electron density in a molecule may be shifted just by being close to another dipole. Electrons will gather on the side of a molecule that faces a positive charge and will retreat from a negative charge. Hence, a temporary dipole can be produced by a nearby polar molecule, or even by a temporary dipole in another nonpolar molecule. London forces are weaker than other intermolecular forces such as ionic interactions, hydrogen bonding, or permanent dipoldipole interactions. This phenomenon is the only attractive intermolecular force between neutral atoms like atoms of noble gases, and is the major attractive force between non-polar molecules such as nitrogen or methane. Without London forces, there would be no attractive force between noble gas atoms, and they couldn't exist in liquid form. London forces become stronger as the atom or molecule in question becomes larger and when the molecules shape is elongated. The larger the atom or molecule becomes the more electrons there are to shift around and elongated molecules have more surface area that can come in contact. This trend is exemplified by the halogens (from smallest to largest: F 2, Cl 2, Br 2, I 2 ). Fluorine and chlorine are gases at room temperature, bromine is a liquid, and iodine is a solid. The London forces also become stronger with larger amounts of surface contact. Greater surface area means closer interaction between different molecules. So long molecules, which can lay sidby-side, are more likely to become liquid or solid than are molecules of the same mass but a different shape. 10+ 10+ Figure 3. An illustration of 2 neon atoms with electrons randomly distributed.
Unit 7 Notes Chemistry I CP Page 266 London dispersion δ 10+ δ+ δ 10+ δ+ Figure 4. An illustration of 2 neon atoms with an uneven electron density in 2 adjacent molecules and a temporary dipole resulting in a temporary attraction between atoms. Figure 5. A diagram that illustrates the relationship between the different categories of interparticle attraction. Structural Particle Attractions (attractions between atoms, molecules or ions) Molecular Attractions (opposite poles attract) Temporary or permanent dipoles Ionic and Molecular Attractions (opposite charges attract) cations are attracted to negative poles of polar molecules and anions are attracted to positive poles of polar molecules Ionic Attractions (opposite charges attract) cations attract anions London forces weak attractions between nonpolar molecules or non-polar areas of molecules Dipoldipole forces stronger attractions between polar molecules or polar areas of molecules Dipolionic forces Even stronger attractions between polar molecules or polar areas of molecules and ions What you need to know about figure 5: You need to know how different kinds of structural particles are attracted. Students should know and be able to compare intermolecular forces (dipoldipole interaction, and London dispersion forces) and ionic bonds, covalent bonds, and metallic bonds in terms of Ionic, Metallic, Network atomic, Molecular Atomic, and Amorphous solids. To understand these forces, it is helpful to understand the differences between the 3 common states of matter: solid, liquid, and gas (we will not be considering plasma here). What s a solid? Solid refers to the state of matter in which the particles are locked into place without much freedom of movement. They can be locked into crystal lattices or just kind of stuck together with intermolecular forces so tightly that they can t really move around. As we ll see, it depends on the type of solid we re talking about. Solids differ from liquids in that the particles in liquids, while still stuck together, do have some freedom of motion. Solids differ from gases in that gas molecules really don t interact with each other much, flying all over the place. What are the different types of solids?
Unit 7 Notes Chemistry I CP Page 267 Ionic solids: Ionic solids are solids in which anions and cations (negatively and positively charged atoms or groups of atoms, respectively) stick together via electrostatic attraction. By electrostatic attraction, I basically mean that the opposite charges just like to stick to each other. When they do this, they form great big crystals in which each ion is surrounded by ions with the opposite charge. Such crystal lattices (as they re called) are really stable, requiring lots of energy (called the lattice energy ) to pull apart. Examples of ionic solids include sodium chloride (table salt) and Epsom salts (magnesium sulfate heptahydrate). Metallic solids: Metallic solids are solids in which the positively charged nuclei are held together by a bunch of valence electrons that kind of bind the whole mess together. These electrons are referred to as being delocalized because they don t stay between two atoms as they do in covalent bonds instead, they travel throughout the solid. This allows the atoms in solids to move around, as the bonding electrons can move around with them. A phrase commonly used to describe metallic bonding is electron sea theory, which describes the positive metal nuclei as floating around in an ocean of negative electrons that hold them together. Just about anything you know of that s a metal does this kind of bonding. Network atomic solids (also called covalent network crystals): Network atomic solids are great big crystals in which all of the atoms are stuck together using covalent bonds. Because the atoms are all locked into place, these solids usually have properties very similar to that of ionic compounds (high melting and boiling point, hard, brittle, and so forth) with the exception that they don t conduct electricity if you melt them. Typically, gemstones (such as amethyst, diamond, and ruby) are network atomic solids. Molecular solids: Molecular solids occur when covalent molecules are held together by intermolecular forces. In this type of bonding, which occurs in ice, the intermolecular forces between the molecules are strong enough to keep the molecules locked into place. Typically, these types of solids have much lower melting and boiling points than metallic, network atomic, or ionic solids, because the intermolecular forces holding the molecules together are much weaker than those of the bonds in the other compounds. Atomic solids: Atomic solids occur when noble gases are cooled to really low temperatures and lock themselves in place using very weak London dispersion forces. You won t run into these solids in the real world, because you need temperatures that are ridiculously low to see them. Amorphous solids: Amorphous solids, unlike the rest of these solids, have no particular crystal structure. In an amorphous solid, the particles are just kind of stuck all over the place, with no regular bonding pattern. Some amorphous solids are soft and rubbery (such as plastic and rubber) because they consist of long molecules which are just kind of tangled together and bound with intermolecular forces. Other molecular solids (called glassy solids) are a lot more like network atomic solids because they consist of atoms stuck together in an irregular fashion using covalent bonds. It won t be much of a surprise to find that glass is an example of such a solid! Students should know and understand the nature of the attraction between structural particles Intermolecular London dispersion forces - these are described in some detail above. Intermolecular dipoldipole forces - these interactions are the result of permanent polarity of molecules. Like little magnets, these polar molecules are attracted to each other by their opposite poles. These poles are not magnetic; they are electrostatic.
Unit 7 Notes Chemistry I CP Page 268 Dipoldipole and London forces apply to molecules. Ionic bonds - these substances are made from oppositely charged particles held together by those charges. Unlike magnets and polar molecules, ionic particles (ions) have no poles but their attractions are electrostatic. You should know that many ionic compounds will dissolve in molecular (covalent) compounds because positively charged ions (called cations) are attracted to the negative pole of a polar molecule or to the negative pole of a molecule that has a polar section AND because negatively charged ions (called anions) are attracted to the positive pole of a polar molecule or to the positive pole of a molecule that has a polar section. Metallic bonds these substances are made from metals atoms that hold their outer electrons loosely and share those electrons. The shared electrons are localized into the area between 2 atoms like covalently bonds, but are shared among many atoms. The structural particles (atoms) of a metal can be thought of as both an atom and an ion. These shared electrons flow freely through the atoms/ions in the metallic structure and form a kind of glue that holds everything together. Students should know and be able to apply the type of substance in which different particle attractions occur Intermolecular London dispersion forces - molecules and gaseous atoms (noble gases) that tend to be attracted to each other very loosely and therefore small structural particles (atoms or molecules) tend to be gases with a very low boiling and melting temperature (see table 1 below). Larger molecules with only London forces attractions might be liquids and even solids at room temperature and pressure if they are very large. These molecular substances have very low boiling and melting temperature (see table 1 below) compared to structural particles of roughly the same size. Intermolecular dipoldipole forces Molecules with dipoldipole forces (polar molecules) tend to be attracted to each other somewhat more strongly that those without polar attractions. Therefore even small structural particles (atoms or molecules) tend to be liquids at room temperature and pressure more often than those with no such attractions (London forces). These molecular substances have low to moderate boiling and melting temperature (see table 1 below) compared to structural particles of roughly the same size. Ionic bonds Ionic compounds are held together by oppositely charged particles (ions) and tend to be crystalline solids. These substances are almost always solids at room temperature and pressure and have a high boiling and melting point compared to molecular substances. Metals Pure metals and alloys of metals are held together by a shared sea of electrons in which the electrons flow freely around and between the atoms/ions of the metal (therefore, this model is called the electron sea model). With some notable exceptions, these substances are solids at room temperature and pressure and have a high boiling and melting point compared to molecular substances. Covalent Networks Very strong interlocking covalent lattices. Very high boiling and melting points.
Unit 7 Notes Chemistry I CP Page 269 Students should know and be able to apply typical properties of various types of substances Note: Recall that when we use the word typical that means that there will be exceptions. Non-polar molecular substances (those with intermolecular London dispersion forces) The general physical properties of non-polar molecular compounds are: (1) Can be solid, liquid or gas at room temperature but the smaller molecules or atoms (in the case of noble gases) are gases. (2) The solids are amorphous (that is to say that they are not crystalline). (3) They do not conduct electricity very well (with the notable exception of certain forms of carbon: diamond does not, but graphite and carbon nano-tubes do). (4) They have low boiling and melting points compared to similar sized structural units of other substances. (5) They are usually not very soluble in water. (6) They do not conduct heat very well. Polar molecular substances (those with intermolecular dipoldipole forces) The general physical properties of polar molecular compounds are: (1) Can be solid, liquid or gas at room temperature but even the smaller molecules can be liquid at room temperature if the polarity is high enough. (2) The solids are sometimes amorphous but the polarity tends to make them crystalline. (3) They do not conduct electricity very well. (4) They have low boiling and melting points compared to similar sized structural units of other ionic and metallic substances but their boiling and melting points are higher than similar sized non-polar molecules. (5) They are often very soluble in water. (6) They do not conduct heat very well. Ionic substances The general physical properties of ionic compounds are: (1) They tend to solid and even rigid at room temperature and pressure. (2) They tend to be crystalline. (3) Solids do not conduct electricity but the molten (melted) state of matter does. (4) They have high boiling and melting points. (5) They are very often quite soluble in water, although many have low solubility (refer to the solubility rules). (6) They do not conduct heat very well.
Unit 7 Notes Chemistry I CP Page 270 Metals The general physical properties of metals are: (1) They tend to be are strong and hard. (2) They are solids at room temperature (with the notable exception of mercury, which is liquid at room temperature) (3) They have a shiny luster when polished. (4) They make good heat conductors and electrical conductors. (5) They are dense. (6) They often ring when struck with a hard object. (7) They have high melting points. (8) They are malleable (they can be bent and made into wires without breaking). Covalent Networks The general physical properties of covalent networks are: (1) They tend to be are strong and hard. (2) They are solids at room temperature (3) They are crystalline. (4) They are generally poor electrical conductors. (5) They are dense. (6) They have high melting points and boiling points. Table 1. A comparison of examples of non-polar covalent, polar covalent, ionic, metallic, and covalent network substances. Type of Substance Structural Unit Inter-particle Force Substance Melting Point (1atm, C) Boiling Point (1atm, C) H 2-259 -253 Non-polar Covalent (molecular) molecule London Dispersion Forces O 2-218 -183 CH 4-182 -164 CCl 4-23 77 C 6 H 6 6 80 H 2 O 0 100 Polar Covalent (molecular) molecule Dipoledipole interaction H 2 S -85.5-61 HCl -114-85 NH 3-78 -33 Ionic ion Ionic bonds NaCl 801 1413 MgF 2 1266 2239 Cu 1083 2567 Metallic atom Metallic bonding Fe 1535 2750 Hg -39 357 W 3410 5660 Covalent Network atom Covalent bonds (SiO 6 ) x (quartz) 1610 2230 C x (diamond) 3500 3930 Students should be able to use a chart, such as the one above to compare the intermolecular forces present in substances with high, low, and moderate melting and boiling points.
Unit 7 Notes Chemistry I CP Page 271 Interpreting Table 1 Let s look 1 st at the non-polar covalent (molecule) category of structural units. You will need to recall that room temperature is between 20 and 25 degrees Celsius (20 and 25 C) and that room pressure is close to 1.00 atmospheres. Note that smaller molecules, such as H 2, O 2, and CH 4, at have boiling and melting temperatures that are way below 0 C, so these would be gases at room temperature. Larger non-polar covalent molecules, such as carbon tetrachloride (CCl 4 ) and benzene (C 6 H 6 ), are liquids at room temperature and pressure. This can lead us to some reasonable assumptions. Non-polar covalent molecules with a small molecular weight will be a gas at room temperature and pressure and while large ones will be liquid. The question is, how large must they be to be a liquid? Well, if we use methane and carbon tetrachloride as guides, methane has a molecular weight of ~ 16 u and it s a gas and carbon tetrachloride has a molecular weight of ~152 u and it s a liquid. So we could assume that the change from gas to liquid (at room temperature and pressure) should occur between these 2 weights. To refine it a bit further, we could say that non-polar molecules with a molecular weight of up to ~65 amu will be a gas at room temperature and pressure and those above will be a liquid. Now, this is just a guide. There are other factors beyond molecular mass, such as shape, that affect whether a substance will be a gas or a liquid or even a solid and you will need to take those into account when necessary. Table 2. A comparison of the molecular mass, boiling temperature, and physical state of some organic, non-polar, covalent substances. Notice that the physical state of these straight chains change from gas to liquid when the carbon chain increases from 4 to 5 carbons. Condensed Structural Formula Molecular Mass (in u) Boiling Temperature (in C) Physical State at Room Temperature (~22 C) CH 4 16.04-164 gas CH 3 CH 3 30.07-88.6 gas CH 3 CH 2 CH 3 44.1-42.09 gas CH 3 CH 2 CH 2 CH 3 58.12-0.5 gas CH 3 CH 2 CH 2 CH 2 CH 3 72.15 36.1 liquid CH 3 CH 2 CH 2 CH 2 CH 2 CH 3 86.18 69 liquid Now, let s look 2 nd category: the polar covalent (molecule) category of structural units. The polarity of a structural unit has a huge impact on the state of matter. Notice that water has about the same molecular mass as methane, but water s melting temperature is 0 C and while methane s is -182 C. The difference is that water is very polar and methane is non-polar. In fact, water molecules hydrogen bond with each other and hydrogen bonding can be thought of as a very strong form of polar attraction (it s not a bond in the sense that electrons are actually shared or that one water molecule gives up electrons to another). Notice also that while hydrogen sulfide (H 2 S), hydrogen chloride (HCl), and ammonia (NH 3 ) are polar, they are gases at room temperature and pressure. These last 3 substances are not nearly as polar as water. So, the more polar a substance is the more likely it is to be a liquid at room temperature and pressure. Look at this example: propane (condensed structural formula: CH 3 CH 2 CH 3 ) is a gas at room temperature and pressure. It s non-polar and its mass is is ~46 u. If we exchange a
Unit 7 Notes Chemistry I CP Page 272 methyl unit on one end with a hydroxyl unit, we will get ethanol (condensed structural formula: CH 3 CH 2 OH). The mass of this molecule is ~46 u but this substance is a liquid at room temperature and pressure. Both have a molecular mass that is less than 65 u, but the ethanol has an alcohol group (-OH) which allows the ethanol molecule to hydrogen bond to other ethanol molecules. Look at another example (one that is NOT found in table 1): Dimethyl ether (condensed structural formula: CH 3 OCH 3 ) is a gas at room temperature and pressure. It s polar and its mass is 46.0684 u. Ethanol (condensed structural formula: CH 3 CH 2 OH) has exactly the same mass because both substances have the same molecular formula (C 2 H 6 O). Again, the ethanol has an alcohol group (-OH) which allows the ethanol molecule to hydrogen bond to other ethanol molecules. On the other hand the dimethyl ether molecule is polar but is not capable of hydrogen bonding. At room temperature and pressure, dimethyl ether is a gas, but ethanol is a liquid. In fact, ethanol boils at 78.5 C, more than 100 C higher than dimethyl ether whose boiling temperature (or boiling point) is 23.6 C. You should take time to examine table 1 and make sure that you understand how to interpret the melting and boiling temperature behavior of the 5 categories of substances and the reason why those behaviors occur. Also use figures 5 and 6 to help understand this behavior. You may also want to go back and review the section on hydrogen bonding in an earlier part of this unit. Table 3. A comparison of the molecular mass, boiling temperature, and physical state of some organic, polar, covalent substances. Notice that the alcohol groups (ROH or OH) are much more polar and create much stronger hydrogen bonds than amine groups (RNH 2 or NH 2 ). As a result, methanol and ethanol are liquids at room temperature and pressure while methanamine and ethanamine are gases at room temperature and pressure Condensed Structural Formula Molecular Mass (in u) Boiling Temperature (in C) Physical State at Room Temperature (~22 C) CH 3 OH 32.04 64.7 liquid CH 3 CH 2 OH 46.07 78.4 liquid CH 3 CH 2 CH 2 OH 74.122 117.2 liquid CH 3 NH 2 31.06-6 gas CH 3 CH 2 NH 2 45.08 16.6 gas CH 3 CH 2 CH 2 NH 2 59.11 48 liquid CH 3 CH 2 CH 2 CH 2 NH 2 73.14 77 liquid
Unit 7 Notes Chemistry I CP Page 273 Figure 6. This illustration shows the range of melting/freezing temperatures of the representative structural particles in table 1. What you should know about figure 6. You are not expected to memorize the changes of temperatures. You ARE expected to know the relative comparisons of the categories of substances. For example, you are expected to know that covalent network compounds have a higher average melting point (the melting and freezing temperature) than metallic substances, that the melting temperatures are very high, but the ranges of the melting points in the 2 categories overlap significantly. You should apply this example to all the other categories.
Unit 7 Notes Chemistry I CP Page 274 Figure 7. This illustration shows the range of boiling/condensation temperatures of the representative structural particles in table 1. What you should know about figure 7. You are not expected to memorize the ranges of temperatures. You ARE expected to know the relative comparisons of the categories of substances. For example, you are expected to know that covalent network compounds have a higher average boiling point (the boiling and condensation temperature) than metallic substances, that the melting temperatures are very high, but the ranges of the boiling points in the 2 categories overlap significantly. You should apply this example to all the other categories. Practice Problems: 1. If you compare the structures of 2 molecular chemical compounds that have approximately the same molecular mass, and it would appear that chemical A will dipoledipole forces, and chemical B will have only London dispersion forces, which chemical is more likely to be a liquid and which one is more likely to be a gas? Explain why. 2. Explain how the halogens exemplify London dispersion force attractions.
Unit 7 Notes Chemistry I CP Page 275 3. Rank from strongest to weakest the following interparticle attractions: ionic, London dispersion forces, hydrogen bonding, and dipoldipole. Key questions for this next section: 22) What are the 5 major assumptions of the kinetic molecular theory as it applies to ideal gases? 23) Describe the model that explains the relationship between gas pressure and gas volume. 24) What is the name of the absolute temperature scale? Describe that scale. What happens to structural particle motion at zero on this absolute temperature scale? 25) What is the relationship between temperature and average kinetic energy? 26) Explain the difference between temperature and heat. 27) Define pressure. 28) Define volume. 29) How many liters of an ideal gas at 0.00 C and 1.00 atmosphere pressure are there in one mole? 30) What is the definition of temperature? 31) What are real gases? What are ideal gases? Indicator C-5.2 Write your learning objective for this section here: Indicator C-5.2 in the South Carolina Science Standards students should be able to explain the behaviors of gas; the relationship among pressure, volume, and temperature; and the significance of the Kelvin (absolute temperature) scale, using the kinetic-molecular theory as a model. Students should be able to use the Kinetic Molecular Theory as a model to explain the relationship between pressure, and volume in a gas sample. The Kinetic Molecular Theory (KMT) can be thought of as a very good model for most realworld applications. In effect, it allows us to perform calculations and make predictions that are a very close approximation of what really happens. It s only when gases are under very high pressures or very low temperatures that the assumptions of the KMT no longer work very well for us. When we treat gases as if they behave according to the assumptions of the KMT, then we call those gases ideal. When we talk about behavior of gases that significantly deviate from this ideal behavior then we call those real gases. There are 5 major assumptions of the kinetic molecular theory as it applies to ideal gases: The gas consists of very small particles, each of which is so far apart from the next that we can treat the particle as if it only has a mass or weight but no volume. The collisions of gas particles with the walls of the container holding them and with each other are perfectly elastic. These molecules are in constant, random motion. The rapidly moving particles constantly collide with each other and with the walls of the container. These particles, therefore, have kinetic energy (the energy of motion). There are no forces of attraction between the gas particles. The average kinetic energy of the gas particles depends only on the temperature of the system.
Unit 7 Notes Chemistry I CP Page 276 So, here s the behavior that you can observe or measure: so long as all other factors remain unchanged, any change in the pressure of a gas in a container that has variable volume will change the volume in inverse (opposite) proportion. The name of the model that can be used to explain the cause for this behavior is called the kinetic molecular theory. In that theory, gas particles (either atoms or molecules) are in constant random motion and that particle motion is creating pressure against the sides of any container. If the container can expand and contract then any change in the volume of the container will result in a inverse (or opposite) change in the pressure of the gas. Any change in the pressure will also result in an inverse change in the volume. Here s the model: The reason the pressure changes inversely with volume is that the number of gas particles doesn t change. If you decrease the size of the container (the volume) without changing the number of or velocity of the gas particles (which is the same as changing the temperature), then the particles will strike the sides of the container more often, and that creates more pressure. Use your imagination to figure out what happens to those particles when pressure or volume is changed by increasing the pressure (it decreases the volume; but why?), by decreasing the pressure, or by increasing the volume. This is the basis of Boyle s Law : PV = k. Before we get into Boyle s Law, let s look at a few of the assumptions listed above and see what happens when this model no longer works: The collisions of gas particles with the walls of the container holding them and with each other are perfectly elastic and there are no forces of attraction between the gas particles. (1) When gas molecules start sticking to each other that affects the proportions of the inverse relationship between pressure and volume. If the molecules stick to each other under high pressure, then the volume will drop far more than the pressure goes up and our inverse behavior no longer follows a pattern that can so easily be predicted. These molecules are in constant, random motion. The rapidly moving particles constantly collide with each other and with the walls of the container. These particles, therefore, have kinetic energy (the energy of motion). (1) It is the constant banging of molecules against the sides of the container that causes the pressure. Pressure is defined as force applied to a surface. So, the constant banging of those gas molecules apply a force to the inner surface of the container holding that gas. If the molecular motion stops being random say, for example, that they start traveling in one direction, then the pressure on the sides of the container will no longer be constant and inversely proportional to the volume. When a gas is traveling through a hose the molecular motion is mostly in one direction. In this case, the faster the gas travels down the hose, the lower the pressure on the insides of the hose will be. There are no forces of attraction between the gas particles. (1) The space between gas particles is so great (on average) that on average these particles are rarely close enough for the attractive forces to have much effect. So, unless the gas is under really high pressure (which greatly reduces the space between gas particles) or the gas is under really low temperatures (which greatly reduces the velocity and momentum of gas particles) the attractive forces between gas particles is so slight that we can ignore them. Students should be able to explain the significance of the absolute temperature scale in terms of the Kinetic Molecular Theory.
Unit 7 Notes Chemistry I CP Page 277 Since kinetic energy is dependent on temperature, if the temperature were to go below zero that would mean the kinetic energy would have to be less than zero. Either there is energy or there isn t; there cannot be negative energy. Therefore, there cannot be negative temperature. This is why the absolute temperature scale (the Kelvin scale) is necessary to gas law calculations and predictions. Explain the relationship between temperature and average kinetic energy. Temperature is not the same thing as heat. Heat is energy. Temperature is a way of determining how much random momentum (mass acceleration) a substance has. Temperature, then, is a relative comparison of the average random momentum of the particles in a substance. The higher the temperature is, the higher the average random momentum of the particles in the substance will be. An increase in temperature is the result of a transfer of energy into the substances where the temperature is being measured. Therefore, energy can be calculated by measuring the temperature change of a substance along with some other factors. In the next few indicators, we will be dealing exclusively with gases and their behaviors. Before going there, I would like for you to understand some terminology. Pressure: Pressure is force per unit area. The pressure of a gas is the force of collisions as the particles strike the walls of the container that contains the gas. If these collisions occur frequently, the gas pressure is high. If the collisions don t occur very often, the pressure is low. Any change in the conditions that results in more frequent collisions will increase the pressure. Volume: The volume of a gas is the empty space the particles travel through. Gas particles (the atoms or molecules) are so small compared to the distance between them that most of the time we can ignore the volume or size of the gas particle. So, the larger the volume, the greater the distance between particles. Any change in the conditions that results in a longer distance between particles is due to an increase in volume. Moles: When we talk about moles of a gas, we are talking about the gas particles. Again, those particles are the atoms or molecules that make up the gas. Noble gas particles are atoms because noble gases don t generally bond to other atoms. Oxygen and nitrogen particles however, are molecules because oxygen bonds to itself as does nitrogen. Those gases form diatomic molecules. The symbol that we use for number of moles in gas laws is the letter n, and that is the number of particles in groups of 6.02 10 23. In an ideal gas at 0.00 C and 1.00 atm, the volume of gas that equals one mole is 22.4 liters. Temperature: As was stated in the last section, the temperature of a gas is a relative measure of the average random momentum of the particles. The hotter the gas, the faster the particles are moving. The speeds of the individual gas particles vary, but we treat the speed (or more correctly, the momentum) as being the same throughout the sample of gas in the container. A note about real gases : Ideal gases don t really exist. As I stated earlier, however, for most of the work that we do with gases in the real world, ideal gas behavior is a close enough approximation that we can assume gases to be ideal and make the math easier. Recall that ideal gas behavior is described by the Kinetic Molecular Theory. This theory doesn t work in all cases and sometimes we have to take into account these gas behaviors that are not ideal. One of the assumptions that we make about ideal gases is that they don t stick to each other. Under indicator 5.1, we looked at behavior that is not ideal. We looked at all the ways that the structural particles of matter (atoms, ions, molecules, and even metals that
Unit 7 Notes Chemistry I CP Page 278 have a behavior that is somewhere between that of atoms and ions) are attracted to each other. In this unit we stated that for ideal gases collisions of gas particles are perfectly elastic. In other words, they don t stick to each other and they aren t repelled from each other by any forces like London forces or dipoldipole forces. In this unit, we also stated that (t)here are no forces of attraction between the gas particles. Both these statements are in opposition to what was stated in indicator 5.1. Indicator 5.1 was all about how particles are attracted. In fact, 2 entire units dealt in some details with the attractive forces between particles. Another of the assumptions that we make about ideal gases is that the structural particles in a gas (atoms or molecules) have no volume. Since gas particles are on average so far apart from each other, we can ignore the volume these particles take up. In most of the work that we do with gases, if we did include the volume of the particles in our calculations the effect would se so slight that when we rounded off for significant digits we d get the same answers as if we would have assumed the gas particles to have no volume. So, we just go ahead and assume that the particles have no volume right from the start, making our calculations and our lives easier. Still another of the assumptions that we make about ideal gases molecules is that they are in constant, random motion. This motion results in these particles hitting the sides of their container, creating pressure. If we start moving gas particles very rapidly in a single direction they stop hitting each other or the sides of their container as often. Therefore gases in motion in a single direction have a lower pressure than gases that aren t moving in any particular direction. This is the model that explains Bernoulli's principle. Bernoulli's principle states that an increase in the speed of a fluid occurs simultaneously with a decrease in pressure. Gases and liquids are both fluids. If slow the gas particles down a lot (lower the temperature) so that when they do come close to each other their momentum doesn t overcome their attractions or repulsions then the gas starts to exhibit non-deal or real gas behavior. Models: Recall that models describe how the universe or some small part of the universe works. Models describe underlying behaviors to explain why the things that we observe occurred as they did. If you take the time (and I always recommend that you do) and run a movie in your head, looking at gases atoms or molecules flying around and interacting, manipulating those particles to imagine those different conditions described above, it will help you to grasp these 2 models: the ideal and the real gas models. Practice Problems: 4. Explain how the kinetic molecular theory helps explain a model for the behavior of a gas in a container that has a flexible volume. 5. What happens to pressure when the volume is increased? Explain the model for the cause of that behavior. 6. What happens to volume when the pressure is increased? Explain the model for the cause of that behavior. 7. What happens to pressure when the volume is decreased? When pressure is decreased? Explain the model for the cause of those behaviors. Key questions for this next section:
Unit 7 Notes Chemistry I CP Page 279 32) What is the most often used form of Charles Law? 33) In what units must temperature be measured for Charles Law to work? 34) What are the symbols for temperature in Celsius and temperature in Kelvin? How are these temperature scales related? 35) What is the mathematical relationship between temperature in C and temperature in K? 36) What is the most often used form of Boyle s Law? 37) What gas law do you get when you combine Charles and Boyle s Laws? Write your learning objective for this section here: 38) What are the 5 steps that you MUST use when solving gas law problems in this class? =====================================================================
Unit 7 Notes Chemistry I CP Page 280 We are about to start learning to solve gas law problems. The format for these problem solutions will be as follows: 1) Identify known values and unknown values 2) Determine equation to be used 3) Substitute known values into equation 3) Isolate the unknown; cancel units and values 4) Solve. ===================================================================== Indicator C-5.3 Indicator C-5.3 in the South Carolina Science Standards students should be able to apply the gas laws to problems concerning changes in pressure, volume, or temperature (including Charles s law, Boyle s law, and the combined gas law). In this unit of chemistry students should: Students should be able to explain Charles law and Boyle s law in terms of the kinetic molecular theory Charles law: Recall the assumptions of the kinetic molecular theory of gases. As long as we can assume that a gas Gas molecules only have mass but no volume. Gas molecule collisions are perfectly elastic. Gas molecules are in constant, random motion (and therefore gases always have temperature and pressure). There are no forces of attraction between the gas particles. Temperature is the only thing that affects the average kinetic energy of gas. then, as long as the pressure and number of moles of gas don t change, the volume of a fixed mass of gas is directly proportional to the Kelvin temperature of the gas. This is Charles law. Charles law states: the volume of a fixed mass of gas is directly proportional to its Kelvin temperature if the pressure and moles (or mass) of the gas is kept constant. That mathematical relationship is: V T In practice, it is most often used in this form: = constant, or V = k. T V T V =. T 1 2 1 2 This direct relationship only works if the Kinetic Molecular Theory can be used so it doesn t work at very high pressures or very low temperatures. It also only works in a closed system and so long as the pressure doesn t change and there is no chemical reaction taking place that would change the number of moles of gas. For example, in a closed system increasing the pressure would either decrease the volume or increase the temperature. So, to use Charles law, you have to keep pressure constant and you cannot use this if there is a chemical change taking placed that would change the number of moles of gas in the system.
Unit 7 Notes Chemistry I CP Page 281 The absolute temperature scale is required for ALL gas law calculations, so any time temperatures are given in C (degrees Celsius or degrees centigrade) they must be converted to K (Kelvin): Boyle s law: Temperature in C + 273.15 = Temperature in K Again, recall the assumptions of the kinetic molecular theory of gases. So long as the pressure and temperature conditions are not too extreme, these assumptions are pretty reliable. In this situation, then, if the temperature and number of moles of gas don t change, the volume of a fixed mass of gas is inversely proportional to the volume of the gas. This is Boyle s law. Boyle s law states that for a given mass of gas at a constant temperature, the volume of the gas varies inversely with pressure. That mathematical relationship is: P V = constant, or PV = k. In practice, it is most often used in this form: P 1 V 1 = P 2 V 2. This inverse relationship only works if the Kinetic Molecular Theory can be used so it also doesn t work at very high pressures or very low temperatures. It also only works in a closed system and so long as the temperature doesn t change and there is no chemical or physical reaction taking place that would change the number of moles of gas. For example, since: N 2 (g) + 3H 2 (g) 2NH 3 (g) There are 4 moles of gas on the left and 2 on the right. If nitrogen (N 2 ) is put into a closed container with hydrogen (H 2 ) and pressure is applied, ammonia (NH 3 ) is produced. This reduces the number of moles of gas and the P 1 V 1 = P 2 V 2 relationship doesn t work any more. Students should be able to solve gas law problems concerning changes in gas pressure, volume, or temperature. This means that you should know how to not only use Charles law and Boyle s law to solve gas law problems, but you should also know how to use and apply the combined gas law: PV T = PV T 1 1 2 2 1 2 Solving gas law problems The format for these problem solutions will be as follows: 1) Identify known values and unknown values 2) Determine equation to be used 3) Substitute known values into equation 4) Isolate the unknown; cancel units and values 5) Solve. Example #1: If the pressure of a gas is initially 2.25 atm and the volume is 12.5 liters, what will the pressure become if the gas is compressed to 6.5 liters? Step 1: Identify known values and unknown values. The known values are 2.25 atm, 12.5 liters, and 6.5 liters. The problem states that 2.25 atm, 12.5 liters are the initial pressure and volume. So, we write P 1 = 2.25 atm