Honors Chemistry I 84.135

1. Honors Chemistry I84.135 Dr. Nancy De Luca Course web site: http://faculty.uml.edu/ndeluca/84.135

2. Text & Homework Information Chemistry, Structure and Properties by Nivaldo Tro. Published by Pearson. The text is bundled with Modified Mastering Chemistry, the online homework system and a student’s solution manual. You can also purchase the e-text and Modified Mastering Chemistry on line using a credit card.

3. Modified Mastering Chemistry You can access your homework assignments and also review specific areas of study on line. The link is: pearsonmylabandmastering.com You get 14 days of courtesy access before you must purchase access to the system.

4. Atoms The Particulate View of Matter

5. Chapter 1 Review Topics The following topics should be fairly familiar to you, and will not be covered in detail. Chapter 1: • Classification of Matter • Physical and Chemical Changes & Properties • Units of Measurement

6. Matter Matter is anything that has mass and occupies space. It includes everything around us, including the air that we breath, our skin and bones, and the earth underneath us. Matter is composed of particles. The structure of those particles determines the property of matter.

7. Matter Important structural properties include the molecular shape (linear, bent, etc.), and the distribution and location of atoms within the structure. For example, diamonds and graphite are both forms of pure carbon, yet their properties and structures are very different.

8. Diamond and Graphite

9. Classification of Matter

10. Properties of Matter Matter can be described by its physicalor chemicalproperties. Physical properties are a description of the substance, and include mass, color, physical state (solid, liquid or gas) at a specific temperature, density, melting or boiling point, odor, solubility, etc.

11. Properties of Matter • During a physical change, the chemical identity of the substance or substances does not change. • Examples of physical changes include evaporation, filtration, and changes of state.

12. Physical Changes When water boils, its chemical composition remains the same. The molecules are now farther apart.

13. Physical Change The dissolving of sugar in water is a physical change. The chemical identity of the water and the sugar remain unchanged.

14. Filtration During filtration, liquids are separated from solids by physical means. The liquid and solid maintain their chemical identity.

15. Distillation During distillation, liquids may be separated from other liquids, or from solids. The chemical identity of each component remains unchanged.

16. Properties of Matter Chemical properties are descriptions of how a substance reacts chemically. Examples include the rusting of iron in the presence of air and water, the souring of milk, or the burning of paper to form carbon dioxide and water vapor.

17. Chemical Changes As iron rusts, the iron atoms combine with oxygen in the air to form a new substance, rust, or iron (III) oxide.

18. Chemical Changes During a chemical change, atoms rearrange the way they are attached to each other, forming new substances with properties that are often quite different from the starting materials.

19. The Study of Matter and the Atom Early scientific work focused on the nature of matter. The concept of small indivisible particles of matter dates back to ancient Greece. Although many cultures performed chemical reactions (ceramics, metallurgy, etc.), the formal study of chemistry, initially using measurements of mass or volumes, is relatively recent.

20. Early Atomic History • The ancient Greek philosophers theorized that matter is discrete, rather than continuous. • Some, notably Demokritos, suggested that there is some small unit of matter that still retains the properties of the larger sample. It was thought that these smaller pieces of matter were indivisible, and were given the name atomos from which we get our modern word atoms.

21. Early Atomic Theory • During the next 2000 years, a lot was learned about matter. Several elements were discovered, metals were refined, acids prepared, etc. • In the mid-1600s, the scientific (rather than the philosophical or applied) study of the nature matter began to take shape.

22. Early Atomic Theory • Since most laboratories contained rudimentary equipment- burners and scales, many experiments involved the measurement of changes in volumes (for gases) and masses during chemical reactions. • Based on measurements and observations, several scientific laws were developed. These laws form the basis for our understanding of the composition of matter.

23. The Law of Conservation of Mass Antoine Lavoisier studied the masses of reactants and products during combustion with oxygen, and in 1789 proposed that In a chemical reaction, matter is neither created nor destroyed.

24. The Law of Definite Proportion • Joseph Proust (1754-1826) determined the chemical composition of many compounds. He found that a given compound always contains the exact same proportion of elements by mass. This is known as the law of definite proportion. For example, all samples of water contain 88.8% oxygen by mass, and 11.2% hydrogen by mass.

25. The Law of Multiple Proportions • This chemical law applies when two (or more) elements can combine to form different compounds. Common examples are carbon monoxide and carbon dioxide, or water and hydrogen peroxide. • John Dalton (1766-1844) conducted experiments on these types of compounds, and determined that there is a simple relationship between the masses of one element relative to the others.

26. The Law of Multiple Proportions • When two elements form a series of compounds, the ratios of the masses of one element that combine with a fixed mass of the other element are always in a ratio of small whole numbers. The meaning of this law is difficult to understand unless it is illustrated using a specific series of compounds.

27. The Law of Multiple Proportions • Consider the compounds of water and hydrogen peroxide. At this point in history, chemists knew the compounds were different, and that they both contain (or can be broken down into) the elements hydrogen and oxygen. They did not yet know the formulas for either compound, nor was the concept of atoms fully developed.

28. The Law of Multiple Proportions • Analysis of 100 grams of the compounds produced the following data:

29. The Law of Multiple Proportions • The Law of Multiple Proportions is illustrated when the numbers in the last column are compared. 15.8/7.93 = 2/1 The small whole number ratio suggests that there is twice as much oxygen in hydrogen peroxide as there is in water.

30. The Law of Multiple Proportions • The key feature is that small whole numbers are generated. The results support the hypothesis that molecules consist of various combinations of atoms, and that atoms are the smallest unit of matter. The ratio doesn’t produce fractions, since there is no such thing as a fraction of an atom. • For the example cited, we would propose that hydrogen peroxide contains twice as many oxygen atoms/hydrogen atoms than does water. We cannot, however, determine the actual formula of either compound.

31. Dalton’s Atomic Theory (1808) 1. Each element consists of tiny particles called atoms. 2. The atoms of a given element are identical, and differ from the atoms of other elements. 3. Compounds are formed when atoms of different elements combine chemically. A specific compound always has the same relative number and types of atoms. 4. Chemical reactions involve the reorganization of atoms, or changes in the way they are bound together.

32. Sub-Atomic Particles • The period from approximately 1880-1915 involved the study of the nature of the atom, using two relatively new tools: electricity and radioactivity. • Scientists knew that atoms of different elements had different relative atomic masses and different properties, and they wanted to find out the reasons for the differences.

33. Sub-Atomic Particles • In the late 1880s, J.J. Thomson (1856-1940) studied the properties of cathode rays. The rays are produced in partially evacuated tubes containing electrodes at either end. • The rays are invisible, unless a phosphorescent screen is used.

34. Sub-Atomic Particles Cathode Rays (Cathode) (Anode)

35. Sub-Atomic Particles Thomson made the following observations: 1. The cathode rays had the same properties regardless of the metal used for the cathode. 2. The rays traveled from the cathode (- charged) to the anode (+ charged). 3. The rays were attracted to the positive plate of an external electrical field, and repelled by the negative plate.

36. Sub-Atomic Particles Thomson concluded: 1. The cathode rays are a stream of negatively charged particles called electrons. 2. All atoms contain electrons, and the electrons from all elements are identical. 3. The atom must also contain matter with a positive charge, as atoms are neutral in charge.

37. Sub-Atomic Particles • Thomson also carried out deflection measurements, in which he applied a magnetic field to deflect the beam along with an external electrical field to straighten out the bent beam. www.youtube.com/watch?v=1iw0Plrk51Y

38. Sub-Atomic Particles • From his measurements, he was able to calculate the charge/mass ratio of the electron: e/m = -1.76x108 coulombs/gram

39. Sub-Atomic Particles • Around the same time as Thomson (1886), Eugen Goldstein observed that if a cathode ray tube contained very small amounts of gas, a glowing substance travelled toward the cathode. • Goldstein called the glowing substance “canal rays”, and observed that they were positive in charge, as they are attracted toward the negative cathode.

40. Sub-Atomic Particles • The apparatus contained a perforated cathode that contained many small holes. When an electric current is applied, the reddish glow forms in the stream of electrons, and travels toward the negative cathode and through the small holes. • http://www.youtube.com/watch?v=3WIjCtZLMDg

41. Sub-Atomic Particles It was several years before Goldstein could explain his observations. The rays were quite different from cathode rays: • Unlike cathode rays, canal rays were barely deflected by a magnetic field or external electric field. • The properties of the rays varied with the gas contained.

42. Sub-Atomic Particles • Later studies determined that hydrogen gas produced the “ray” with the largest charge to mass ratio (ie., the smallest mass). This particle, produced when a hydrogen atom loses its electron, was identified as the proton.

43. Sub-Atomic Particles • Robert Millikan (1868-1963) published the results of his Oil Drop Experiment in 1909. He designed an apparatus that could be used to determine the charge on an electron. The device used a fine mist of oil drops that had been exposed to ionizing radiation. The radiation caused some of the oil drops to take on one or more electrons.

44. Sub-Atomic Particles

45. The Charge of the Electron

46. The Charge of the Electron www.youtube.com/watch?v=XMfYHag7Liw

47. Sub-Atomic Particles • Millikin determined that the charge on the electron is -1.60 x 10-19coulombs. • Using Thomson’s value for the charge to mass ratio of the electron, the mass of the electron could be calculated. mass of e- = (-1.60 x 10-19 coulombs) (-1.76 x 108 coulombs/gram) = 9.11 x 10-28 grams = 9.11 x 10-31 kilograms

48. Early Atomic Models • J. J. Thomson had shown that all atoms contain negatively charged particles called electrons. Combined with the work of Millikan, they discovered that the electron has very little mass. • Thomson proposed that the bulk of the atom is a positively charged gel or cloud, with most of the atomic mass and all of the positive charge uniformly distributed throughout the gel.

49. Early Atomic Models • The electrons were viewed as discrete, very small particles that were stuck into the positively charged gel or cloud “like raisins in a pudding.” This model is often called the plum or raisin pudding model of the atom. • The electrons could be knocked out of the gel if enough energy is applied, and this is the source of the cathode rays.

50. Early Atomic Models One of the key features of Thomson’s atomic model is that most of the atomic mass and all of the positive charge is uniformly distributed throughout the atom.