1. Chem 400 Chapter 2

2. Going Further: The Structure of Atoms Dalton thought that atoms were the smallest particle of matter, but through a series of experiments starting in the late 1800 ’ s, this was proved to be incorrect. In a series of experiments by various scientists, the existence of electrons, protons, and neutrons were deduced and verified.

3. Going Further: The Structure of Atoms Electrons were discovered in 1897 by JJ Thomson, and they were found to have a negative electric charge. Protons were hypothesized in 1911 by Ernest Rutherford, and were verified in 1919 by Rutherford. They have a positive electric charge. Neutrons were discovered in 1932 by James Chadwick. They have no electric charge.

4. Going Further: The Structure of Atoms There are 2 important experiments that you should be aware of: –Robert Millikan ’ s Oil-Drop Experiment of 1909, which enabled him to measure the magnitude of the electric charge on electrons, and calculate their mass in grams. –Ernest Rutherford ’ s 1911 Gold-Foil Experiment, also called the Alpha-Scattering Experiment. This experiment enabled him to hypothesize the existence of protons in the nucleus of atoms.

5. Oil-Drop Experiment

7. The Structure of Atoms From this experiment, Millikan obtained the actual charge on an electron, -1.60x10 -19 C. And from this charge and Thomson ’ s charge/mass ratio, the exact mass of an electron was calculated to be 9.10x10 -28 g. So from these experiments, scientists deduced that atoms were made up of even smaller subatomic particles, one of which was the electron. Since electrons have a negative charge, while atoms are neutral, scientists also realized that there had to be at least 1 more subatomic particle with a positive charge.

8. The Structure of Atoms Where was this atomic particle & what did it look like?

9. Gold-Foil Experiment

10. The Structure of Atoms How could something in the atom cause such huge deflections in a massive positively charged particle like an alpha particle? So Rutherford proposed that atoms were composed of mostly empty space (where the electrons moved in circular orbits) with a very small, very massive, very dense center called the nucleus. The nucleus has a positive charge. This was Rutherford ’ s Atomic Model.

11. The Structure of Atoms

12. Rutherford then proved the existence of protons in 1919, and neutrons were discovered by James Chadwick in 1932. So what ’ s the overall picture of an atom, and what are the sizes, masses, charges, and densities of the particles and regions?

13. ParticleMass (g x10 -24 ) Mass (amu) Relative Mass Charge (C x10 -19 ) Relative Charge Location in Atom proton, p 1.6731.007311.6021nucleus electron, e or e - 0.0009110.0005490-1.602electron cloud neutron, n 1.6751.0087100nucleus The Structure of Atoms

14. The diameter of a typical atom is around 1x10 -10 m or 1 Å. The diameter of a typical nucleus is only 0.0001 Å. You can see that most of the mass of the atom is contained in a very small volume, so the nucleus is incredibly dense. The density of a typical nucleus is 1x10 13 to 1x10 14 g/cm 3, beyond our comprehension! If a matchbox had this density, it would weigh 2.5 billion tons!

15. The Structure of Atoms

16. Atomic Number, Mass Number and Isotopes Dalton thought that atoms of different elements differed mainly by mass, but we now know that atoms of different elements differ by the number of protons which they contain. The number of protons which an element contains is called the Atomic Number, Z. The Atomic Number is found on the Periodic Table above the elemental symbol.

17. Atomic Number, Mass Number and Isotopes

18. It is true that every atom of the same element contains the same number of protons. –So every H atom has 1 proton, and every C atom has 6 protons. So the number of protons defines the element. But it is NOT true that all atoms of the same element are identical. What ’ s different? Well, what else is there?

19. Atomic Number, Mass Number and Isotopes Although all atoms of the same element have the same number of protons, they DO NOT have the same number of neutrons! –and if it is a neutral atom, they all have the same number of electrons Atoms of the same element which have different numbers of neutrons are called isotopes.

20. Atomic Number, Mass Number and Isotopes So isotopes of the same element differ by the number of neutrons. And since neutrons are the same relative mass as protons, isotopes also differ by mass. Although this is not shown on the Periodic Table, every element has at least 2 isotopes (except some of the newly synthesized elements like Mt).

21. Atomic Number, Mass Number and Isotopes To show different isotopes, we have several different isotopic notations or isotopic symbols. They all use the Mass Number, A, which is the sum of the protons and neutrons in the nucleus. For example, H has 3 common isotopes, H-1, H-2, and H-3. Carbon also has 3 common isotopes, C-12, C-13, and C- 14. The number after the symbol or the superscript left number is the Mass Number.

22. Atomic Number, Mass Number and Isotopes Practice with Isotopic Notation: How many electrons, protons, and neutrons do the following isotopes have? Mass Number Atomic Number

23. Ions and Ionic Isotopic Notation A neutral atom has equal numbers of protons and electrons. Why? During chemical reactions, atoms may gain or lose (or share) electrons, e-. If an atom gains or loses 1 or more e-, there is an imbalance between protons and e-, and the result is a charged particle called an ion. So an ion is formed when an atom gains or loses e-.

24. Ions and Ionic Isotopic Notation If an atom loses 1 or more e-, then it has more protons than e-, so the ion has a + charge. It is called a cation. If an atom gains 1 or more e-, then it has less protons than e-, so the ion has a - charge. It is called an anion. Note that it is difficult to gain more than 3 e- or lose more than 4 e-. How do we show ions?

25. Ions and Ionic Isotopic Notation If we have an isotope which is an ion, we can show a complete isotopic notation for the ion. How many protons, e-, and neutrons do the following ionic isotopes have?

26. Elements and the Periodic Table Elements are fundamental substances. They can ’ t be broken down into smaller substances by chemical reactions. The Periodic Table arranges the known elements (114 of them). 90 of these are naturally occurring, while the rest have been synthesized in nuclear reactions.

27. Elements and the Periodic Table

28. Notice that the elements names have been given shorthand notations (called symbols) of 1 or 2 letters. Unnamed elements actually have a 3 letter designation until they are named. The first letter is ALWAYS capitalized, while the second letter is ALWAYS lowercase. What elements do you have to memorize (names and symbols)? 1-40; 42; 46-57; 76-90; 92; and 94.

29. Elements and the Periodic Table Although most of the symbols are obviously related to the name, like N for nitrogen, others seem to make no sense, like Pb for lead! This is because some of the symbols come from old Latin names or other languages. Plumbum was an old Latin name for lead. W for tungsten comes from the German name wolfram.

30. Elements and the Periodic Table Chemistry in some fashion has been around for centuries. Some elements were known thousands of years ago. But most elements were discovered and identified in the last 250 years.

31. Elements and the Periodic Table In the early to mid 1800 ’ s, chemists were trying to organize the 60-some known elements into some sort of pattern. Mendeleev designed a Periodic Table in 1869 which was based on the masses of the known elements (atomic weights) and the compounds they formed with hydrogen (hydrides) or oxygen (oxides). Today ’ s Table is similar, but the elements are arranged by atomic numbers (number of protons) instead of by atomic weights.

32. Elements and the Periodic Table If you look at a Periodic Table, there are 18 columns called Groups or Families. They are called families as they share common chemical properties or characteristics. The 7 rows are called Periods. The groups are numbered 2 ways on US Tables. The old US system uses numbers with A or B sections, while the internationally approved system simply numbers the groups from 1 to 18 going across from left to right.

33. Elements and the Periodic Table There are several basic regions on the Table: –Metals –Nonmetals –Semimetals (metalloids or semiconductors) –Main Block or Representative Elements –Transition Metals –Inner Transition Metals –Lanthanides –Actinides

34. Elements and the Periodic Table Several important Groups also have names: –Group 1, except hydrogen, are the Alkali Metals –Group 2 is the Alkaline Earth Metals –Group 17 are the Halogens –Group 18 are the Noble Gases

35. Elements and the Periodic Table Metals: lustrous, silvery, malleable, ductile, generally hard, solids except Hg, conductors, lose electrons to become cations, react with nonmetals to form ionic salts. Nonmetals: nonconductors, react with metals, gain electrons to form anions, brittle, most gases (1 l, 5 s) Metalloids: B, Si, As, Te, At, Ge, Sb: in between metals & nonmetals, semiconductors, solids

36. Elements and the Periodic Table Here are some shared characteristics in the regions and groups: Alkali Metals: Very reactive metals, soft, not found in nature as pure element Alkaline Earth Metals: same as Alkali metals but less reactive Halogens: most reactive nonmetals, corrosive, not found in nature as pure element

37. Elements and the Periodic Table Noble Gases: also called Inert gases as very nonreactive, don ’ t form compounds except Xe Lanthanides: f-fillers, rare earth metals, inner transition metals, reactive, silvery-grey Actinides: f-fillers, rare earth metals, inner transition metals, reactive, silvery-grey, radioactive, synthetic above 92 Why is hydrogen placed in Group 1 if it is NOT an Alkali Metal and is actually a nonmetal?

38. Elements and the Periodic Table

40. Atomic Mass and Weighted Averages of Elements As atoms have a very tiny mass in grams, scientists use another scale to state the masses of atoms, the atomic mass unit, amu. You see this in the table with the masses of p, e, and n given earlier. The conversion factor between mass in g and mass in amu is: 1 amu = 1.66054x10 -24 g OR 1 g = 6.02214x10 23 amu

41. Atomic Mass and Weighted Averages of Elements The average atomic mass of the elements is shown beneath the elemental symbol on the Periodic Table.

42. Atomic Mass and Weighted Averages of Elements But every element has different isotopes with different masses! That ’ s why the atomic masses on the Table are average masses: it is really the mass in amu of a single “ average ” atom of an element. But what does an “ average ” atom of an element look like? What does the “ average ” student look like? Does “ it ” exist?

43. Atomic Mass and Weighted Averages of Elements For H, 99.985% of all H atoms are H-1, while 0.015% are H-2 (there are basically 0% H-3). This is called the natural abundance or %-abundance of an isotope. So shouldn ’ t the “ average ” H atom look a lot like H-1, and shouldn ’ t the average atomic mass of H be very close to the mass of the H-1 isotope? Because the different isotopes do not have equal natural abundances, we calculate atomic masses of elements using a weighted average of all the isotopes:

44. Atomic Mass and Weighted Averages of Elements So if lead has 4 common isotopes with the following masses and %-abundance, what is the atomic mass of lead? Pb-204203.973020 amu1.40% Pb-206205.974440 amu24.1% Pb-207206.975872 amu22.1% Pb-208207.976627 amu52.4%

45. Molar Mass of Atoms & Avogadro ’ s Number The atomic mass is the mass in amu of a single “ average ” atom. Is this useful in the lab? Can we pick out and weigh an individual atom? Chemists weigh in g, which is a HUGE number of atoms. So we need a unit to express large numbers of atoms or molecules without using scientific notation. Chemists defined a counting unit to do this.

46. Molar Mass of Atoms & Avogadro ’ s Number They chose a unit called mole so that the atomic mass on the Periodic Table is also used for measuring grams. 1 mole = 1 mol = 6.02214x10 23 things So if you have 1 mol of pennies, how many pennies do you have? How many dollars is this? The number 6.02214x10 23 is Avogadro ’ s number.

47. Molar Mass of Atoms & Avogadro ’ s Number Avogadro ’ s number is very important as it is a conversion factor between number of things and moles of things. EX: if you have 2.5 mol of aluminum, how many atoms of aluminum do you have? Avogadro ’ s number is also very special as if we have 1 mol of an element, we have the atomic mass in g of the element.

48. Molar Mass of Atoms & Avogadro ’ s Number If we have 1 mol of Na atoms, prove that this is 22.99 g of Na:

49. Molar Mass of Atoms & Avogadro ’ s Number The molar mass is the mass in g of exactly 1 mol of an element. The atomic mass and the molar mass are the same number, they differ only by units!

50. Molar Mass of Atoms & Avogadro ’ s Number The molar mass is also a conversion factor: it converts between mol of an element and g of an element. –Ex: If you have 25.7840 g of gold, how many mol of gold is this? There are 6 types of simple calculations that we can do using molar masses and Avogadro ’ s number.