The Structure & Stability of Atoms

1. The Structure & Stability of Atoms

2. Early Atomic History • There have been many different theories, reflecting different times and cultures, to explain the composition of matter. • In addition, chemical reactions, refinements of ores, purification of salt, etc. have been carried out for thousands of years.

3. 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.

4. 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.

5. 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.

6. The Law of Conservation of Mass • Antoine Lavoisier (1743-1794) measured the masses of reactants and products for a variety of chemical reactions. He determined that matter is neither created nor destroyed during a chemical reaction. This is known as the law of conservation of mass.

7. 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.

8. 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.

9. 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.

10. 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.

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

12. 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.

13. 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.

14. The Law of Multiple Proportions

15. 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.

16. Sub-Atomic Particles • The period from approximately 1900-1915 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.

17. Sub-Atomic Particles • 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.

18. Sub-Atomic Particles

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

20. 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.

21. 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.

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

23. 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.

24. Sub-Atomic Particles

25. The Charge of the Electron

26. 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

27. 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.

28. 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.

29. 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.

30. Early Atomic Models • Thomson had a graduate student, Ernest Rutherford, working for him. In 1911, Rutherford, Geiger and Marsden performed an experiment to confirm Thomson’s atomic model. • They bombarded a thin gold foil with alpha (α) particles. The α particles have twice the charge of an electron and are positive in charge, with a mass that is 7300 times greater than the mass of an electron.

31. Early Atomic Models • The α particles can best be thought of as a positively charged, fast traveling atomic sized bullet. They created a thin beam of α particles and directed the beam at a very thin gold foil.

32. Early Atomic Models If Thomson’s model is correct, most of the α particles should pass right through the gold atoms. Some slight deflection might occur if the positively charged α particle travels near an electron.

33. Gold Foil Experiment

34. Early Atomic Models • The film that lined the apparatus showed that most α particles went through the foil with little or no deflection. However, some of the particles were deflected at great angles.

35. Early Atomic Models • The deflection of the α particles was consistent with a large concentration of positive charge and atomic mass. This very small extremely dense positively charged area is called the nucleus.

36. Early Atomic Models • The atom is mostly empty space, with the electrons found outside of the nucleus. If the nucleus was the size of a pea, it would have a mass of 250 million tons, and the electrons would occupy a volume approximately the size of a stadium.

37. Atomic Nucleus

38. Sub-Atomic Particles • We now know that the positive charge of an atom, contained in the nucleus, is due to particles called protons. • Protons have a charge equal in magnitude to that of an electron, but positive in charge. • The mass of a proton is roughly 1800 times greater than the mass of an electron.

39. Sub-Atomic Particles • The nuclei of atoms also can contain neutrons. Neutrons are neutral in charge, with a mass similar to that of a proton. • Neutrons are found in the nucleus of atoms, along with protons.

40. Sub-Atomic Particles

41. Sub-Atomic Particles • During chemical reactions, atoms may lose or gain electrons to form charged particles called ions. • Atoms of a given element may have differing numbers of neutrons. These forms of the same element are called isotopes. • It is the number of protons or the atomic number that defines the identity of the atom.

42. Atomic Symbols • The periodic table lists the elements in order of increasing atomic number (the number of protons). • The atomic number, represented by the letter Z, is linked with the atomic symbol. For example, oxygen is atomic number 8, and any atom containing 8 protons, regardless of the number of neutrons or electrons, is represented by the symbol O.

43. Atomic Symbols • To indicate a specific isotope, the atomic symbol must also contain the mass number. • The mass number is the number of neutrons plus protons for a particular isotope. The mass number is never found on the periodic table. • Since the mass number is the number of particles (neutrons + protons) in the nucleus, it is always an integer.

44. Isotopes of Sodium Mass number Atomic number

45. Atomic Symbols • For example, there are three isotopes of carbon: 12C, 13C and 14C The mass number, if specified, appears in the upper left corner of an atomic symbol. Since all carbon atoms have 6 protons (carbon is atomic number 6 on the periodic table), atoms of carbon may have 6, 7 or 8 neutrons in the nucleus. The isotopes are called carbon-12, carbon-13 and carbon-14.

46. Atomic Symbols • If the atom has lost or gained electrons, the charge is written in the upper right corner of the atomic symbol. • The atomic number, though optional, may be written in the lower left corner of the symbol. 37Cl1- This ion of chlorine contains 17 protons, 20 neutrons, and 18 electrons.

47. Relative Atomic Masses • Once the formulas of simple gases and compounds could be determined, scientists could also determine the relative masses of the elements. • For example, since equal volumes of gases contain equal numbers of particles (at the same T and P), the masses of gases could be compared to hydrogen, the lightest gas.

48. Stoichiometry • Stoichiometry is a Greek word that means using chemical reactions to calculate the amount of reactants needed and the amount of products formed. • Amounts are typically calculated in grams (or kg), but there are other ways to specify the quantities of matter involved in a reaction.

49. Relative Atomic Masses • As the early chemists explored the nature of matter, they discovered that atoms of the elements had different masses. Avogadro’s Hypothesis which states that under the constant temperature and pressure equal volumes of gases contain an equal number of particles could be used to determine relative atomic masses for gaseous elements.

50. Relative Atomic Masses • Equal volumes of gases contain an equal number of particles. Although the number of particles (atoms or molecules) in a liter of gas (at a specific T and P) wasn’t known, Avogadro’s Hypothesis said that a liter of any other gas under the same conditions would contain the same number of particles.