4.3.1 Basic Material from Physics and Chemistry

1. 4.3.1 Basic Material from Physics and Chemistry 4.3.1.1 Atoms and Molecules

2. Basic Material from Physics and Chemistry • In this section we will look at basic background material from physics and chemistry needed to understand the nature of the chemical hazards we will be modeling.

3. Atoms • Matter is made up of atoms and molecules. • Atoms are the smallest components of the basic chemical elements, which include hydrogen, carbon, iron, uranium, etc. • As of November 2011, 118 elements have been identified – the first 98 occur naturally on earth, 80 are stable, and the others are radioactive and decay into other elements. • Periodic Table of the Elements: http://www.atomic-elements.info/

4. Atoms • Atoms are made up of a nucleus, surrounded by orbiting electrons. • The nucleus is composed of protons and neutrons – usually an atom has the same number of protons as neutrons. • The number of protons in an atom determines its atomic number, so for example, carbon has 6 protons, hence an atomic number of 6. • An atom’s atomic weight is determined by its total number of protons and neutrons, which is approximately twice its atomic number.

5. Atomic Weight • Looking at Table 4.2 on page 111 of our textbook, we see that an element’s atomic weight is not an integer – for example, carbon has an atomic weight of 12.011. • The reason for this is that atoms of the same element may have different numbers of neutrons – for example carbon atoms usually have 6 neutrons, but may have 7 or 8 neutrons, leading to an atomic number of 13 or 14. • Carbon atoms with more than 6 neutrons are radioactive – the different types of carbon are called isotopes of carbon.

6. Atomic Weight • Radioactive carbon 14 is used for carbon 14 dating to determine the age of fossils or old artifacts. • To get an element’s atomic weight, scientists have determined how much of each isotope of an element occurs in the universe and have computed a weighted average.

7. Molecules • Most substances are made up of a combination of atoms. • A molecule is a collection of atoms bound together in particular combinations and structures. • For example, water molecules are made up of two oxygen and one hydrogen atom – we denote water by H2O. • Another example of a molecule is CH4, natural gas (methane), which is used for cooking – it is made up of 4 hydrogen atoms and one carbon atom. • Ozone molecules (O3) are comprised of three oxygen atoms!

8. Molecular Weight • The molecular weight of a molecule is the sum of the atomic weights of the atoms making up the molecule. • For example, the molecular weights of the molecules on the last slide are: • Water: 2H + 1O = 2(1.008) + 1(15.9994) = 18.0154 • Methane: 1C + 4H = 1(12.011) + 4(1.008) = 16.043 • Ozone: 3O = 3(15.9994) = 47.99982

9. Molecular Weight • What would be the molecular weights of these hydrocarbons: • Acetylene (C2H2) • Trichloroethylene (C2HCI3) • Propane (C3H8) • Butane (C4Hl0) • Ethanol (C2H5OH). • Note that many hazardous materials such as these which are used as fuels, solvents, etc., are made up of hydrocarbons!

10. 4.3.1 Basic Material from Physics and Chemistry 4.3.1.2 Physical Properties of Matter

11. States of matter • There are three common states of matter – solid, liquid, and gas. • We will mostly be concerned with liquids or gases, especially the transition from liquids to gases as spilled hazardous liquids evaporate or react to form gases. • These gases can be toxic or flammable and may move from the scene of an accident to a location with an unprepared or unsuspecting population.

12. Density • Definition: The density of a substance is its mass per unit volume. • Typical units are: lb/ft^3, gm/cm^3, etc. • For example, • Water has a density of 62.4 lb/ft^3 or one g/cm^3. • Solid rock has a density of about 200 lb/ft^3. • Air has a density of 0.004 lb/ft^2 at sea level.

13. Specific Gravity • A quantity related to density is specific gravity. • Definition: The specific gravity of a substance is the ratio of its density to the density of water. • For example, the specific gravity of solid rock with a density of 200 lb/ft^3 would be: • Specific gravity of rock = (density of rock)/(density of water) = (200 lb/ft^3)/(62.4 lb/ft^3) = 3.21 • In metric units, the same rock would have a specific gravity of 3.21 g/cm^3, since the density of water is 1 g/cm^3. • For this reason, one may encounter instances of the terms “density” and “specific gravity” being used interchangeably.

14. Density and Temperature • Most substances will expand when heated. • Thus, the same amount of material which is heated will require more volume, which means it will have a lower density! • Can you think of a substance that doesn’t behave this way? • http://www.ehow.com/info_8272150_types-materials-shrink-heated.html • http://phys.org/news/2011-11-incredible-material-reveal-scandium-trifluoride.html • http://www.ncnr.nist.gov/AnnualReport/FY1998/rh1.pdf

15. Evaporation • Consider an open container of some chemical liquid, such as water, antifreeze, alcohol, etc. • What would you expect to happen over time? • Which would you expect to evaporate more quickly? • Intuitively, evaporation is the process of a liquid turning into a gas!

16. Evaporation • We will need to understand the process of evaporation, since it will play a major role in spills of hazardous materials! • Here are two principles to keep in mind when considering evaporation: • 1. The rate of evaporation is proportional to the surface area (all other factors being equal). • 2. The rate of evaporation increases as the temperature of the liquid increases.

17. Evaporation and Surface Area • Physically, evaporation is a process in which molecules close to the surface of a liquid that have sufficient kinetic energy to break through the surface do so and escape individually into the space above the liquid. • The molecules that escape become part of the gaseous or vapor component of the chemical. • It follows that if the liquid has a larger surface area, then more molecules will be able to escape – for example, doubling the surface area will double the rate of evaporation!

18. Evaporation and Surface Area Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

19. Evaporation and Temperature • The temperature of a substance such as a liquid is a measure of the average kinetic energy of the substance’s molecules. • Thus, when heat is applied to a liquid, the liquid will gain more energy, causing the liquid’s molecules to increase their movement, in turn increasing the liquid’s average energy. • It follows that a proportion of the molecules near the liquid’s surface will have higher energy, leading to an increase in the evaporation rate!

20. Evaporation and Temperature Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

21. Vapor Pressure • Let’s consider a “simple experiment”! • Place a beaker of chemical under a larger closed glass cover. • Initially, all of the chemical is inside of the beaker. Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

22. Vapor Pressure • The rest of the space inside the glass cover is filled with some other material such as air that doesn’t react with the chemical (i.e. it is “inert” with respect to the chemical. Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

23. Vapor Pressure • Now, suppose the material in the beaker begins to evaporate. • Some of the molecules from the beaker join those in the vapor space inside the cover, outside of the beaker. Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

24. Vapor Pressure • Since none of the original gas molecules have any place to go, there is an increase in the total number of molecules in the same space. • This results in an increase in the pressure under the cover. Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

25. Vapor Pressure • This process will eventually slow down and reach an equilibrium point once there are so many chemical molecules in the vapor space that the chemical is no longer able to evaporate. Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

26. Vapor Pressure • A similar effect happens on hot days when there is high relative humidity – there is so much water vapor in the air that sweat produced on a body is unable to evaporate! Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

27. Vapor Pressure • Technically, what is really happening is that at all times there are molecules with sufficient energy leaving the liquid and entering the vapor space as well as molecules in the gas space colliding with and rejoining the liquid. • Initially, more molecules leave the liquid than enter, but as time increases, the rates even out until evaporation no longer is able to take place – at this point the system is at equilibrium. Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

28. Vapor Pressure • Once equilibrium is reached, since there are more total gas molecules in the vapor space than there were initially, the total pressure will be higher. • The amount of pressure increase, denoted P, is called the vapor pressure of the chemical at the system’s current temperature. Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

29. Vapor Pressure • Since increasing the temperature would increase evaporation, it follows that vapor pressure would also increase. • Thus, vapor pressure is an increasing function of temperature! Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

30. Vapor Pressure • Note that if the gas in the vapor space is truly inert with respect to the chemical in the beaker and undergoes no interactions with the chemical (in either liquid or gaseous form), then the vapor pressure P is independent of initial pressure! Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

31. Vapor Pressure • While in practice, idealized conditions like this never occur, it turns out that the only factor that significantly impacts vapor pressure is the system’s temperature. • For this reason, when working with vapor pressure, one must know the temperature of the system in question. Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

32. Boiling Process • Suppose a beaker that contains a liquid chemical is gradually heated to higher and higher temperatures, in a fashion similar to heating a pot of water on a stove. • Keep in mind that ordinary atmospheric pressure of the air is always pushing down on the liquid – in the figure this is represented by an imaginary piston. Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

33. Boiling Process • If the vapor pressure of the liquid is increased enough so that it is greater than the atmospheric pressure, then vapor bubbles of the chemical can expand rapidly, causing the effect known as boiling. • At this point, the chemical can enter the vapor form throughout the liquid, not just at the surface, since vapor bubbles create their own vapor space where they develop! Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

34. Boiling Process • Under boiling conditions, much larger quantities of the chemical can move into the vapor state rapidly – the primary limiting factor is the heat provided. • The reason for this is that it takes a certain amount of heat energy to change a fixed amount of a given chemical from liquid to gaseous form. • This is true for both evaporation and boiling. Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

35. Boiling Process • As an example, 540 calories of heat energy are needed to convert one gram of water from liquid to gas under normal conditions. • This amount of energy is called water’s heat of vaporization. Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

36. Boiling Process • Once a liquid reaches its boiling point, all of the heat energy being applied to it is used up to convert more of the liquid to gaseous form. • At this point, the liquid’s temperature essentially stays constant right at the boiling point, rather than continuing to rise higher. Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

37. Boiling Process • To illustrate these ideas, consider the process of boiling an egg. • In New York, NY (at sea level), if it takes 4 minutes to boil the egg at 100 degrees Celsius, will it take the same amount of time at the same temperature in Denver, CO? Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

38. Boiling Process • Since Denver’s elevation is one mile, the atmospheric pressure is lower, so to heat the water to boiling requires less heat. • Thus, the egg will take longer to cook than 4 minutes. Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

39. Boiling Process • Pressure cookers are specially designed pots with a screw-on lid designed to build up the pressure inside the pot to a pressure higher than atmospheric pressure. • Then, to reach the boiling point, more heat needs to be applied to get the pot’s ingredients to boil, which means the food cooks more quickly! Image Courtesy Charles Hadlock: Mathematical Modeling in the Environment

40. Boiling Point • We tend to equate “boiling point” with “high temperature,” most likely because when cooking, this is usually the case. • For many dangerous chemicals, it turns out that the boiling point is already achieved at room temperature because the chemicals have very high vapor pressures. • For example, propane gas is stored in small metal cylinders or bottles under high pressure. • At room temperature, under normal atmospheric conditions, propane will boil.

41. Boiling Point • Similar to a pressure cooker, by holding the propane liquid at a pressure at or slightly above vapor pressure, the liquid doesn’t boil. • But what if a tanker truck full of propane crashes and breaks open? • The propane liquid would pour on the ground, boil just like water on a hot griddle, and form a highly flammable expanding gas cloud that can float off into a surrounding neighborhood. • The goal of this chapter is to analyze situations like this!

42. Mixtures vs. Pure Substances • All of the scientific concepts discussed so far have been for “pure substances,” i.e. materials consisting of a single chemical. • The chemical may be made up of molecules that consist of more than one element, but each molecule is identical. • Water, methane, or ozone would be a pure substance. • A “mixture” of water and other chemicals such as acetone or gasoline, or a mixture of gasoline and oil, etc. is not a pure substance. • Many hazardous materials are in fact a mixture of chemicals.

43. Mixtures vs. Pure Substances • A natural question to ask is: How does a mixture’s chemical properties relate to the chemical properties of the mixture’s constituents? • For example, how are boiling points, vapor pressures, etc. affected? • It turns out that mixtures are much more complicated than pure substances – instead of a boiling point, a mixture may have a range of temperatures through which they boil. • Also, chemicals within a mixture may interact in ways that alter the chemical properties of the individual chemicals in the mixture.

44. Resources • http://www.atomic-elements.info/ • http://www.ehow.com/info_8272150_types-materials-shrink-heated.html • http://phys.org/news/2011-11-incredible-material-reveal-scandium-trifluoride.html • http://www.ncnr.nist.gov/AnnualReport/FY1998/rh1.pdf • Charles Hadlock – Mathematical Modeling in the Environment, Section 4.3