1. Atoms and Elements Mr. Hollister Holliday Legacy High School Regular & Honors Chemistry

2. Experiencing Atoms Atoms are incredibly small, yet they compose everything. Atoms are the pieces of elements. Properties of the atoms determine the properties of the elements. 2

3. 3 Early Theories of Matter The ancient Greeks believed that all matter was composed of only four elements: earth, air, water, and fire.

4. 4 First Atomic Theory Democritus (c.a. 460 BC) was the first thinker ever to reason that matter was not infinitely divisible. He called the smallest piece of matter ‘atomon’, which in Greek means ‘not cutable’. Democritus

5. Aristotle The Ancient Greeks Aristotle (c.a. 380 BC) reasoned, however, that matter was infinitely divisible. Aristotle’s ideas agreed more closely with those of the Christian church, therefore, they were accepted as fact for over two millenia.

6. 6 Enlightenment Thinking Robert Boyle (c.a. 1661) wrote a book entitled The Sceptical Chymist. In it, Boyle argues that the four classical elements of the Greeks were actually not elements at all. He also argued that matter was composed of indivisible particles or atoms. Robert Boyle

7. 7 The First Modern Atomic Theory After two hundred years of experimental results, the first modern atomic theory was developed by John Dalton (c.a. 1805). Dalton reasoned that indivisible, spherical particles made up all of matter. He also argued that these particles were rearranged during a chemical reaction. John Dalton

8. 8 Dalton’s Original Theory of the Atom and of Reactions

9. Dalton’s Atomic Theory 1. Each Element is composed of tiny, indestructible particles called atoms. Tiny, hard, indivisible, spheres. 2. All atoms of an element are identical. They have the same mass, volume, and other physical and chemical properties. So, atoms of different elements are different. Every carbon atom is identical to every other carbon atom.  They have the same chemical and physical properties. However, carbon atoms are different from sulfur atoms.  They have different chemical and physical properties. 9

10. 10 Dalton’s Atomic Theory 3. Atoms combine in simple, whole-number ratios to form molecules of compounds. Because atoms are unbreakable, they must combine as whole atoms. The nature of the atom determines the ratios in which it combines. Each molecule of a compound contains the exact same types and numbers of atoms.  Chemical formulas

11. 11 Discovery of the Subatomic Particles J. J. Thompson (c.a. 1897) performed experiments involving cathode rays that proved that atoms could be broken down into smaller particles. Thompson called the particle that he discovered an ‘electron’, which means unit of electricity in Greek. J. J. Thomson

12. 12 Cathode Ray Tube (Discovery of the Electron)

13. 13 Thomson’s Results the cathode rays are made of tiny particles these particles have a negative charge because the beam always deflected toward the + plate the amount of deflection was related to two factors, the charge and mass of the particles every material tested contained these same particles Thompson called these particles ‘electrons’, because they were the components of all electricity

14. 14 Discovery of the Nucleus Ernest Rutherford (c.a. 1908) utilized a radioactive ( α particles) piece of gold foil to prove that the plum- pudding model of the atom was wrong. In so doing, he discovered the atomic nucleus and the proton. Ernest Rutherford

15. 15 Rutherford’s Conclusions Atoms were mostly empty space because almost all the particles went straight through Atoms contain a dense particle that were small in volume compared to the atom but large in mass because of the few particles that bounced back This dense particle were positively charged because of the large deflections of some of the particles

16. 16 Rutherford’s Interpretation – the Nuclear Model 1) The atom contains a tiny dense center called the nucleus the amount of space taken by the nucleus is only about 1/10 trillionth the volume of the atom 2) The nucleus has essentially the entire mass of the atom the electrons weigh so little they give practically no mass to the atom 3) The nucleus is positively charged the amount of positive charge balances the negative charge of the electrons 4) The electrons are dispersed in the empty space of the atom surrounding the nucleus (electron cloud)

17. 17 Structure of the Atom Rutherford proposed that the nucleus had a particle that had the same amount of charge as an electron but opposite sign based on measurements of the nuclear charge of the elements these particles are called protons charge = +1.60 x 10 19 C mass = 1.67262 x 10 -24 g since protons and electrons have the same amount of charge, for the atom to be neutral there must be equal numbers of protons and electrons

18. 18 The Nuclear Atomic Model

19. 19 atomic radius ~ 100 pm = 1 x 10 -10 m nuclear radius ~ 5 x 10 -3 pm = 5 x 10 -15 m Rutherford’s Model of the Atom 2.2 “If the atom is the Cowboys Stadium, then the nucleus is a marble on the 50-yard line.”

20. 20 Atomic Theory from 1800-1911

21. 21 Some Problems How could beryllium have 4 protons stuck together in the nucleus? shouldn’t they repel each other? If a beryllium atom has 4 protons, then it should weigh 4 amu; but it actually weighs 9.01 amu! Where is the extra mass coming from? each proton weighs 1 amu remember, the electron’s mass is only about 0.00055 amu and Be has only 4 electrons – it can’t account for the extra 5 amu of mass

22. 22 There Must Be Something Else There! to answer these questions, Rutherford proposed that there was another particle in the nucleus – it is called a neutron neutrons have no charge and a mass of 1 amu mass = 1.67493 x 10 -24 g  slightly heavier than a proton no charge

23. 23 The Sub-Atomic Particles

24. 24 Elements each element has a unique number of protons in its nucleus the number of protons define the element the number of protons in the nucleus of an atom is called the atomic number the elements are arranged on the Periodic Table in order of their atomic numbers Each element also has a corresponding mass number. Protons + neutrons.

25. 25

26. 26 Elements Each element has a unique name and symbol. The symbol is either one or two letters  One capital letter or one capital letter + one lower case letter.  H = Hydrogen = “water-former”  Br = Bromine = ‘stench’ Liquid Bromine

27. 27 The Periodic Table of Elements Atomic number Atomic mass Element symbol

28. Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons X A Z U 238 92 Mass Number Atomic Number Element Symbol 2.3 Elemental & Isotope Notation Isotope Notation = Uranium-238

29. Example: How many protons, electrons, and neutrons are in an atom of ? for most stable isotopes, n 0 > p + Check: Z = 24 = # p + # e - = # p + = 24 Solution: in neutral atom, # p + = # e - mass number = # p + + # n 0 Concept Plan: Relationships: therefore A = 52, Z = 24 # p +, # e -, # n 0 Given: Find: symbol atomic number # p + # e - symbol atomic & mass numbers # n 0 A = Z + # n 0 52 = 24 + # n 0 28 = # n 0 29

30. Isotopes: Isotopes = atoms of the same element with different masses. Isotopes have different numbers of neutrons. 11 6 C 12 6 C 13 6 C 14 6 C 30 Carbon-11 Carbon-12 Carbon-13 Carbon-14

31. Neon 9.25%221210Ne-22 or 0.27%211110Ne-21 or 90.48%2010 Ne-20 or Percent natural abundance A, mass number Number of neutrons Number of protonsSymbol 31

32. Mass Spectrometer Tro: Chemistry: A Molecular Approach, 2/e 32

33. Mass Spectrum A mass spectrum is a graph that gives the relative mass and relative abundance of each particle Relative mass of the particle is plotted in the x-axis Relative abundance of the particle is plotted in the y-axis 33

34. By definition: 1 atom 12 C “weighs” 12 amu 1 amu = 1.6606 x 10 -24 g On this scale 1 H = 1.008 amu 16 O = 16.00 amu Atomic mass is the mass of an atom in atomic mass units (amu) 3.1 Atomic Mass

35. Average Atomic Mass Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations. Average atomic mass is calculated from the isotopes of an element weighted by their relative abundances. 35

36. Natural lithium is: 6 Li (6.015 amu) = 7.42% 7 Li (7.016 amu) = 92.58% (7.42 x 6.015) + (92.58 x 7.016) / 100 = 6.941 amu Average atomic mass of lithium: 36

37. 37 Charged Atoms The number of protons determines the element. All sodium atoms have 11 protons in the nucleus. In a chemical change (aka: a reaction), the number of protons in the nucleus of the atom doesn’t change, however the number of electrons may. Atoms in a compound that gain or lose electrons and have a charge, these are called ions.

38. Types of Ions Cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. Anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. Na 11 protons 11 electrons Na + 11 protons 10 electrons Cl 17 protons 17 electrons Cl - 17 protons 18 electrons 2.5

39. 39 Ions of Hydrogen

40. Ions 40 A monatomic ion contains only one atom Na +, Cl -, Ca 2+, O 2-, Al 3+, N 3- A polyatomic ion contains more than one atom OH -, CN -, NH 4 +, NO 3 - Nitrate Polyatomic Ion

41. 41 Practice—Fill in the Table.

42. 42 Practice—Fill in the Table, Continued.

43. Example —Find the Number of Protons and Electrons in Ca 2+.

44. 44 Review What is the atomic number of boron, B? What is the atomic mass of silicon, Si? How many protons does a chlorine atom have? How many electrons does a neutral neon atom have? Will an atom with 6 protons, 6 neutrons and 6 electrons be electrically neutral? Will an atom with 27 protons, 32 neutrons, and 27 electrons be electrically neutral? Will an Na atom with 10 electrons be electrically neutral?