CHAPTER 2 Atom and the Periodic Table

1. CHAPTER 2 Atom and the Periodic Table

2. Elements There are millions of different pure chemical substances that are observed in the world. Early scientists (philosophers) suggested that all of these substances were composed of a much smaller number of substances, called elements. Element - A substance that cannot be separated into more simple substances by chemical means. Compound - A substance composed of two or more elements. Example - calcium oxide calcium carbonate  calcium oxide + carbon dioxide calcium oxide  calcium + oxygen carbon dioxide  carbon + oxygen

3. Early Theories of Matter The ancient Greeks discussed two possibilities for the essential property of matter. Matter is continuous (no “particles” of matter) - Plato, Aristotle, and a majority of Greek philosophers. Matter is discrete (composed of particles) - Democritus, Leucippus, and a small number of Greek philosophers. Does this process have an end? yes – particles of matter exist no – matter is continuous

4. Dalton’s Atomic Theory A comprehensive theory that accounted for the above obser-vations was proposed by John Dalton, an English chemist, in 1808. There were three parts to the theory (the first of which we will discuss further in this chapter). 1) Elements are composed of particles, called atoms. a) All atoms of the same element are identical in size, mass, and chemical properties. b) Atoms of different elements differ in their size, mass, and chemical properties.

5. Dalton’s Atomic Theory (continued) 2) Chemical substances (compounds) are composed of atoms of more than one element. a) In any particular pure chemical compound the same kinds of atoms are present in the same relative numbers. 3) Chemical reactions can rearrange atoms, but atoms cannot be created, destroyed, or converted from atoms of one element to atoms of a different element. We now know that some of the hypotheses in Dalton’s atomic theory are not completely correct; however, the theory represents a good starting point in understanding the composition of matter.

6. Atomic Structure In Dalton’s atomic theory the smallest particles (atoms) could not be further broken down. However, a series of experiments, beginning in the mid-19th century, demonstrated that atoms themselves were composed of smaller particles.

7. Radioactivity In 1895, Antoine Becquerel discovered that some substances (such as radium and uranium) spontaneously emit “radiation”, a process called radioactivity. Three types of radioactivity were found: alpha () radiation: positively charged particles, now known to be He2+ nuclei (2 protons + 2 neutrons) beta () radiation: negatively charged particles, now known to be electrons gamma () radiation: uncharged, now known to be high energy photons (particles) of light

8. Electrons J. J. Thompson (1897) found that when a high voltage was applied across two electrodes at low pressure a beam of particles moved from the negative to the positive electrode. The particles, named electrons, were negatively charged and the same regardless of the gas between the electrodes or the metal used in the electrodes.

9. Charge and Mass of the Electrons Thompson’s experiment made it possible to find the value for the charge to mass ratio for an electron. Charge to mass ratio = charge of an electron = 1.76 x 108 C/g mass of an electron The charge of an electron (and therefore the mass of the electron) was first determined experimentally by Millikan (1909). Charge of electron = - 1.6022 x 10-19 C (C = Coulomb) Mass of electron = 9.11 x 10-28 g This mass is small compared to the mass of an atom of hydrogen, the lightest element (m1H = 1.67 x 10-24 g)

10. Consequences of the Discovery of the Electron Based on the work of Thompson and Millikan, it was known that atoms contained electrons. This suggested the following: 1) Atoms must also contain a positively charged particle to balance the negative charge of the electron. 2) The mass of the positively charged particle is most likely much larger than the mass of the electron.

11. The “Plum Pudding” Model To account for the predictions about the positive particles in an atom, Thompson suggested that most of the space within an atom consisted of a positively charged substance, with electrons embedded within, the “plum pudding” model.

12. Rutherford and the Nuclear Atom To test Thompson’s plum pudding model, Ernest Rutherford (1909) carried out an experiment where a beam of positively charged particles (alpha particles) were directed at a thin sheet of gold metal.

13. The results of this experiment were inconsistent with the plum pudding model. Rutherford proposed a new model, called the nuclear model of the atom, that did account for the experimental results.

14. Subatomic Particles particlechargemass Coulombs elementary kg amu proton (p+) + 1.60 x 10-19 +1 1.673 x 10-27 1 neutron (n) 0 0 1.675 x 10-27 1 electron (e-) - 1.60 x 10-19 - 1 9.11 x 10-31  0 ________ 1 amu = 1.6605 x 10-27 kg mp/me = 1836.

15. Atomic Structure nucleus electron charge cloud 1) The protons and neutrons of the atom are found in a small region in the center of the atom, called the nucleus. This region contains most of the mass of the atom, and all of the positive charge. 2) Electrons in the atom form a diffuse cloud of negative charge centered on the nucleus and occupying most of the volume of the atom. 3) The size of the charge for the proton and electron is the same. The charge for the proton is positive, and the charge for the electron is negative. Neutrons have no charge.

16. 4) The type of element for an atom is determined by the number of protons in the atomic nucleus. Element (new definition) - An element is a pure chemical substance composed of atoms, each of which has the same number of protons in the nucleus. Hydrogen - one proton per atom Helium - two protons per atom Lithium - three protons per atom . . . . Similarly, we can now define a compound (new definition) as a pure chemical substance composed of two or more different kinds of atoms.

17. The Periodic Table

18. Atomic Number and Mass Number 1) The atomic number (Z) is equal to the number of protons in the atom. 2) Since atoms are electrically neutral, the number of electrons in an atom is also equal to Z, the atomic number. 3) The mass number (A) is equal to the number of protons + neutrons in the atom. a) Because protons and neutrons have a mass of approximately 1 (in amu) and electrons have a mass of approximately 0 (in amu) the mass number is equal to the approximate mass of the atom in amu. b) Based on the above, the number of neutrons in an atom is equal to A - Z. So for an atom: # protons = Z # electrons = Z # neutrons = A - Z

19. Notation For Atoms We use the following general notation to represent isotopes of atoms. mass number atomic number symbol for element Since we can use the symbol for the element and the periodic table to determine Z, the atomic number, we often omit Z in giving the symbol for the atom. Example: 3416S = 34S We can omit the subscript because all sulfur atoms contain 16 protons.

20. Isotopes The atomic number determines the number of protons and electrons in an atom. This does not place any restrictions on the number of neutrons in the atom. It is possible for atoms of the same element to have different numbers of neutrons. These different types of atoms are called isotopes. Isotopes of Hydrogen normal hydrogen deuterium tritium 1H 2H 3H Note that to a very good approximation isotopes of a particular element are chemically identical to one another.

21. Example of Notation for Isotopes As an example of using the above notation, consider the follow-ing naturally occurring isotopes of oxygen (Z = 8). protons neutrons electrons mass number symbol 8 8 8 16 16O 8 9 8 17 17O 8 10 8 18 18O Example: How many protons, neutrons, and electrons are there in one atom of 56Fe? What is the approximate mass of one atom of 56Fe in amu and in kg?

22. Example: How many protons, neutrons, and electrons are there in one atom of 56Fe? What is the approximate mass of one atom of 56Fe in amu and in kg? # protons = Z = 26 # neutrons = A – Z = 56 – 26 = 30 # electrons = Z = 26 approximate mass (amu) = A = 56 approximate mass (kg) = 56 amu 1.6605 x 10-27 kg 1 amu = 9.30 x 10-26 kg Note that the actual mass of one atom of 56Fe is 55.934939 amu.

23. Nuclear Stability There are many factors that determine whether a particular nucleus will radioactively decay (is unstable) or not. Based on observations, the following has been observed: 1) Nuclei with an even number of both protons and neutrons are generally more stable than those with an odd number of protons and/or neutrons. 2) Nuclei containing 2, 8, 20, 50, 82 or 126 protons or neutrons usually have a large number of stable isotopes. 3) All nuclei with more than 83 protons are radioactive.

24. Stable Nuclei

25. Atomic Mass Units (amu) The mass of a single atom of an element, expressed in SI units, is an extremely small number. For example, the mass of a single atom of 16O is 2.6560 x 10-26 kg. For convenience, we often express values for atomic mass in terms of atomic mass units (amu). Atomic mass units are defined as follows 12.00 amu = mass of one atom of 12C (exact) From this we get 1. amu = 1.6605 x 10-27 kg (approximate) The mass of any other atom (or particle) is found relative to the ratio of its mass to the mass of a 12C atom, which can be measured experimentally. Mass of particle (amu) = mass particle • (12.00 amu) mass 12C atom

26. Example: A mass spectrometer is a device for determining values for mass for atoms (and molecules). In a particular experiment, the ratio (mass X/mass 12C) is measured and found to be equal to 2.581. What is the mass of the atom X (in amu)?

27. Example: A mass spectrometer is a device for determining values for mass for atoms (and molecules). In a particular experiment, the ratio (mass X/mass 12C) is measured and found to be equal to 2.581. What is the mass of the atom X (in amu)? mass X = 2.581 mass 12C atom Mass X = 2.581 (mass 12C atom) = 2.581 (12.00 amu) = 30.97 amu Note that because of the way we define atomic mass units, the only isotope whose mass is exactly equal to its mass number is 12C. isotope mass (amu) 1H 1.007825 12C 12.000000... (exact) 238U 238.0508

28. Atomic Mass in the Periodic Table Because different isotopes of an element have different masses, the question arises as to which mass should be given in the periodic table. For short lived radioactive elements the mass number of the most stable isotope of the element is listed. Element Z A technetium (Tc) 43 98 radon (Rn) 86 222 plutonium (Pu) 94 244

29. Average Atomic Mass For naturally occurring elements, the value for mass given in the periodic table is the average atomic mass, based on the natural abundance of the isotopes that is observed. Isotope protons neutrons electrons mass (amu) 12C 6 6 6 12.0000 (exact) 13C 6 7 6 13.0034 14C 6 8 6 14.0032 Average mass in periodic table = 12.0107 amu

30. Average Atomic Mass If we know the percentage of each isotope of an element that is found in nature, and the mass of each isotope, than the average atomic mass of the element can be found as follows: Mave = f1 M1 + f2 M2 + f3 M3 + … = i=1n fi Mi where f1, f2,...are the fractions of each isotope observed in nature M1, M2,…are the corresponding masses for each isotope (in amu) Note the following: f1 + f2 + f3 + …= 1 fx = % X 100 %

31. Non-chemical Example A person has a box of sandwiches. Half of the sandwiches are 6.0 ounces, and half of the sandwiches are 10.0 ounces. What is the average weight of a sandwich? Average weight = (0.50)(6.0 oz) + (0.50)(10.0 oz) = 8.0 ounces We use the same procedure in finding the average mass of an atom. We multiply the fraction of each isotope by the mass of that isotope, and then add the results to find the average mass.

32. Chemical Example There are three naturally occurring isotopes of the element magnesium. Based on the information below, find the average atomic mass of a magnesium atom. Isotope percent f M(amu) 24Mg 78.99 % 23.98504 25Mg 10.00% 24.98584 26Mg 11.01% 25.98259

33. Chemical Example There are three naturally occurring isotopes of the element magnesium. Based on the information below, find the atomic mass of a magnesium atom. Isotope percent f M(amu) 24Mg 78.99 % 0.7870 23.98504 25Mg 10.00% 0.1003 24.98584 26Mg 11.01% 0.1117 25.98259 So Mave = (0.7899)(23.98504 amu) + (0.1000)(24.98584 amu) + (0.1101)(25.98259 amu) = 24.31 amu, the value given in the periodic table.

34. Periodic Table The periodic table is an arrangement of the chemical elements based on similarities in their physical and chemical properties The periodic table contains a large amount of useful information about the chemical elements. Organization There are several ways in which the elements in the periodic table may be classified. Rows = Periods Columns = Groups This is the more important classification. Elements in the same group usually have similar physical and chemical properties.

35. Simplified Periodic Table 1A 2A 3A 4A 5A 6A 7A 8A

36. 1A 2A 3A 4A 5A 6A 7A 8A You are responsible for knowing the names/symbols for elements 1-57, 72-86, and 92.

37. Major Groups in the Periodic Table 1A 2A 3A 4A 5A 6A 7A 8A

38. Metals, Nonmetals, and Metalloids Metals: Usually solid at room temperature (exceptions Cs, Fr, Hg) Shiny metallic luster Good conductors of electricity and heat Malleable (can be hammered into thin sheets) Ductile (can be drawn into thin wires) Nonmetals: Can be solid, liquid, or gas at room temperature Dull colored (as solids) Poor conductors of electricity and heat Not malleable, not ductile Metalloids (semimetals): Intermediate between metals and nonmetals

39. Metals, Nonmetals, Metalloids in the Periodic Table 1A 2A 3A 4A 5A 6A 7A 8A

40. States of Elements (solid, liquid, gas) 1A 2A 3A 4A 5A 6A 7A 8A

41. Examples of Elements (as found in nature) nickel germanium sulfur (metal) (metalloid) (nonmetal)

42. The Mole It is nearly impossible to work with individual atoms in the laboratory because of their small size and mass. For example, one 12C atom has a mass of 1.993 x 10-26 kg, far to small to measure directly with an analytical balance. It is convenient to have a unit representing a number of atoms that can easily be measured by the usual techniques in the laboratory. This unit is called the mole.

43. Avogardo’s Number The unit mole is simply a number. By definition, one mole of anything is equal to Avogadro’s number of things. Avogadro’s number, an experimentally determined quantity is NA = 6.022 x 1023. Note that there is no difference between the concept of moles and more common terms used for specific numbers of things. 1 dozen = 12 1 score = 20 1 thousand = 1000 1 mole = 6.022 x 1023

44. Example (Avogadro’s Number) If I have three dozen eggs, then how many eggs do I have? 1 dozen eggs = 12 eggs (a conversion factor) number of eggs = 3 dozen eggs 12 eggs = 36 eggs 1 dozen If I have three moles of carbon atoms, then how many carbon atoms do I have 1 mole carbon atoms = 6.022 x 1023 carbon atoms number of carbon atoms = 3 moles C 6.022 x 1023 C atoms 1 mole = 1.81 x 1024 carbon atoms Moles is the fundamental SI unit for quantity of substance. The symbol for moles is n, and the abbreviation for the unit moles is mol.

45. Significance of Avogadro’s Number There must be something special about the number 6.022 x 1023 (Avogadro’s number). The significance is as follows. Consider a collection of identical objects. The following relationship will apply. If one object has a mass of X amu… …then one mole of objects has a mass of X g. This is a subtle point. Consider the unit “thousand”. We could make the following statement If one object has a mass of X g… …then one thousand objects has a mass of X kg. Example: The mass of one penny is 2.50 g. The mass of one thousand pennies is mass one thousand pennies = 1000 (2.50 g) = 2500. g 1 kg = 2.50 kg 1000 g

46. The same principle applies to the concept of moles mass of one 19F atom = 19.00 amu, so mass of one mole of 19F atoms = 19.00 g. Check (Note that 1 amu = 1.6605 x 10-27 kg) mass one mole 19F = (6.022 x 1023) (19.00 amu) = 1.1442 x 1025 amu 1.6605 x 10-27 kg 1 amu = 0.01900 kg = 19.00 g

47. Use of the Mole Concept We may use the concept of moles to reinterpret atomic mass. Atomic mass of S = 32.06 amu = 32.06 g/mol We interpret the atomic mass of sulfur (or any other element in the periodic table) in two equivalent ways: The average mass of one sulfur atom is 32.06 amu. The mass of one mole of sulfur atoms is 32.06 g.

48. Use of the Mole Concept (Example) A chemist has a 14.38 g sample of copper (Cu). a) How many moles of copper does she have? b) How many atoms of copper does she have?

49. Use of the Mole Concept (Example) A chemist has a 14.38 g sample of copper (Cu). a) How many moles of copper does she have? moles Cu = 14.38 g Cu 1 mol = 0.2260 mol Cu 63.546g b) How many atoms of copper does she have? atoms Cu = 0.2260 mol Cu 6.022 x 1023 atom Cu 1 mol Cu = 1.361 x 1023 atom Cu

50. End of Chapter 2 “…the ultimate particles of all homogeneous bodies are perfectly alike in weight, figure, and so forth.” - John Dalton, A New System of Chemical Philosophy (1808) “Elements arranged according to the size of their atomic weights show clear periodic properties.” - D. I. Mendeleev (1869) “I don’t believe that atoms exist!” - Ernst Mach (1897)