1. Chapter 4: The Structure of the Atom

2. Atoms are the fundamental blocks of matter… Chapter Big Idea

3. Section 1: Early Ideas About Matter

4. What are the similarities and differences of the atomic models of Democritus, Aristotle, and Dalton? How was Dalton’s theory used to explain the conservation of mass? Section 1: Essential Questions & Vocabulary Dalton’s atomic theory Theory Vocabulary

5. Section 1: Big Idea The ancient Greeks tried to explain matter, but the scientific study of the atom began with John Dalton in the early 1800’s.

6. How was science thousands of years ago different from science now? Lacked controlled experimentation & tools for scientific investigations Intellectual thought as truth Roots of Atomic Theory

7. Philosophers – scholarly thinkers Speculated about the nature of matter & formulated their own explanations based on their own life experiences Common Conclusions: Matter composed of things such as earth, water, air, and fire. Matter could easily be divided into smaller and smaller pieces. Roots of Atomic Theory

8. Greek philosopher First to propose that matter was NOT infinitely divisible. Atomos- Greek word meaning INDIVISIBLE. Matter is composed of atoms which move through empty space Size, shape, and movement of atoms determine the properties of matter Democritus (460 BC -370 BC)

9. Democritus – Atomic Model & Analogy ~ 400 BC Atomic ModelAnalogy Atoms are small, hard particles that are all made of the same material, but are formed into different shapes and sizes Legos

10. What holds the atoms together? THINK – PAIR - SHARE Why do you think it was hard for Democritus to defend his ideas? Lack of experimentation Ahead of his time Democritus – Criticism

11. One of the most influential Greek Philosophers Rejected Democritus’ notion of atoms because it contradicted his own ideas about nature. Because he was so influential, this led to Democritus’ atomic theory to be rejected. Aristotle (384 – 322 BC)

12. Aristotle – Atomic Model & Analogy ~ 300 BC – 1800’s Atomic ModelAnalogy All matter was made of only four elements & four Properties Fire, air, water, and earth Hot, cold, dry, and wet Death of Chemistry for 2000 years!!!

13. Marks the beginning of modern atomic theory Revived Democritus’ idea of atoms based on the results of his scientific research Studied numerous chemical reactions Determined the mass ratios of the elements involved in those reactions John Dalton (1766 – 1844)

14. Matter is composed of extremely small particles called atoms which are indivisible and indestructible. Atoms of a given element are identical in size, mass, and chemical properties. Atoms of a specific element are different from those of another element. Different atoms combine in simple whole-number ratios to form compounds. In a chemical reaction, atoms are separated, combined, or rearranged. Dalton’s Atomic Theory – 1803

15. Dalton – Atomic Model & Analogy 1803 Atomic ModelAnalogy All matter is made of atoms. All atoms of a given element are alike, atoms of different elements are different. Atoms combine in whole-number ratios. In chemical reactions, atoms are separated, combined, or rearranged. Billiard Ball

16. Law of conservation of Mass- mass is conserved in any process. Dalton’s Atomic Theory – explains the conservation of mass in chemical reactions as the result of SEPARATION, COMBINATION, or REARRANGEMENT of atoms. Atoms are NOT created, destroyed or divided in the process. Conservation of Mass

17. Which of these reactions show Dalton’s Theory? Dalton’s Theory : Practice Question

18. Which of these reactions show Dalton’s Theory? Dalton’s Theory : Practice Question Solution

19. Six atoms of Element A combine with eight atoms of Element B to produce six compound particles. How many atoms of Element A does each compound particle contain? How many atoms of element B does each compound particle contain? Are all of the atoms used to form compounds? Dalton’s Theory: Practice Question

20. Dalton’s Theory Practice Solution Element A (6) Element B (8) Compound (6 units) Only have 6 A Elements – can have up to 6 compound units Although you have 8 B Elements – not enough for 2 per compound Must have 2 leftover B Elements

21. Section 2: Defining the Atom

22. What is an atom? How can the subatomic particles be distinguished in terms of relative charge and mass? Where are the locations of the subatomic particles within the structure of the atom? Section 2: Essential Questions & Vocabulary Atom Cathode ray Electron Nucleus Proton Neutron Model Vocabulary

23. Section 2: Main Idea An atom is made up of a nucleus containing protons and neutrons; electrons move around the nucleus.

24. Smallest particle of an element that retains the properties of the element. How big is an atom? ~1.3 x 10 -10 m What is an Atom? If you could increase the size of an atom to make it as big as an orange. In this new scale, the orange would be as big as Earth!!

25. WHAT IS AN ATOM LIKE? Now that scientists were convinced on the existence of atoms, a new set of questions arises!!!

26. When an electric charge is applied, a ray of radiation travels from the cathode to the anode, called a cathode ray. Cathode rays are a stream of particles carrying a negative charge. Cathode Ray Tube Link to video on Cathode Ray Tubes by clicking on the image

27. Completed a series of cathode-ray tube experiments at Cambridge University. Studied mathematics and physics Won the Nobel Prize in 1906 J.J Thomson (1856-1940)

28. JJ Thomson: Discovery of the Electron 1897  Measured the effects of both magnetic and electric fields on the cathode ray to determine the charge to mass ratio.  Particles that compose cathode rays are negatively charged  Charge to mass ratio of the cathode-ray is always the same.  Concluded that all cathode rays are composed of identical negatively charged particles called ELECTRONS  Experiments revealed the electron has a very large charge for its tiny mass PROOF – THERE MUST BE A PARTICLE SMALLER THAN THE ATOM!!

29. JJ Thomson – Atomic Model & Analogy 1897 Atomic ModelAnalogy Discovered the presence of a negative particles in the atom. Atoms are made from a positively charged substance with negatively charged electrons scattered about, like raisins in a pudding or chocolate chips in a cookie. Plum Pudding Or Chocolate Chip Cookie

30. American Physicist Determined the charge of an electron Nobel Prize Winner in 1923 Robert Millikan (1868 – 1953)

31. Oil Droplets are dropped X-Ray knocks electrons from air which then attach to the oil droplets. Millikan could vary the electric field strength to makes the oil droplets move more slowly, rise, or become suspended. Could calculate the charge on the droplets based on their rate of fall. The magnitude of charge on the drops always changed by a discrete amount Smallest common denominator = 1.602 x 10 -19 C Oil Drop Experiment – Charge of an Electron ~1910

32. JJ Thomson - charge to mass ratio of the electron. Millikan - charge of the electron. CAN THOMSON’S INFORMATION ALONG WITH MILLIKAN’S BE USED TO GET THE MASS OF THE ELECTRON? The Electron – Mass??? Mass of electron = 9.1 x 10 -28 g

33. Recap of what we know: Matter is neutral! All matter is made up of atoms which have electrons (- charge) Electrons are much lighter than the lightest atom known! Questions: If electrons are part of all matter and they possess a negative charge, how can matter be neutral? If the mass of an electron is so small, what accounts for the mass of a typical atom? Plum Pudding Model New questions arise

34. JJ Thomson’s student Early work – discovered radioactive half-life Nobel Prize in Chemistry – 1908 Father of nuclear physics Studied how alpha particles interacted with matter Alpha particles – positively charged particles Ernest Rutherford (1871-1937)

35. Aimed a narrow beam of alpha particles at a thin sheet of gold foil. A flash of light is produced when the particle strikes the gold foil WHAT DID HE EXPECT? Positive charge is evenly distributed Path of α – particle should not be altered α – should continue in a straight path Rutherford’s Gold Foil Experiment

36. Rutherford’s Gold Foil Experiment What actually happened!!

37. Atoms consist of mostly empty space Almost all of the atom’s positive charge and mass is contained in a tiny dense region at the center of the atom He called this dense region – the NUCLEUS Electrons are held within the atom by their attraction to the positively charged nucleus Opposite charges attract Rutherford’s Gold Foil Experiment Conclusions

38. Rutherford – Atomic Model & Analogy 1911 Atomic ModelAnalogy Atoms have a small, dense, positively charged center that he called NUCLEUS Nucleus is tiny compared to the atom because the atom is mostly empty space Cherry with a Pit

39. Worked with Rutherford (after WWI) Proved the existence of neutrons Elementary charges devoid of any electrical charge. Nobel Prize in Physics – 1935 Manhattan Project – Development of Atomic Bomb James Chadwick (1891-1974) 1932 - Neutron

40. Bohr – Atomic Model & Analogy 1913 Atomic ModelAnalogy Theorized electrons move in definite orbits around the nucleus like planets circle the sun. Energy levels are located at certain distances from the nucleus. Solar System

41. Modern – Atomic Model & Analogy Late 20 th Century – 21 st Century Atomic ModelAnalogy Schrodinger, Heisenberg, Einstein & many other scientists Electrons move at high speeds in an electron could around the nucleus In the ELECTRON CLOUD, electrons orbit around the nucleus billions of times in one second Electron’s motion is dependent on the AMOUNT of ENERGY they contain Cotton Balls

42. Atomic Theory Ted Ed talk !!

43. All atoms have 3 fundamental subatomic particles Protons, neutrons, electrons Atoms are spherically shaped, with small, dense positively charged nucleus surrounded by negatively charged electrons. Most of an atom consists of fast moving electrons that are held within the atom by their attraction to the positively charged nucleus. Nucleus is composed of neutrons (neutral charge) and protons (positive charge) Scientists have determined that protons and neutrons are composed of “quarks” Scientists unsure if and how “quarks” affect chemical behavior. Chemical behavior can be explained by considering only the atom’s electrons. Completing the Model of the Atom

44. Section 3: How Atoms Differ

45. His the atomic number used to determine the identity of an atom? What is an isotope? Why are atomic masses not whole numbers? Given the mass number and the atomic number, how are the number of electrons, protons, and neutrons in an atom calculated? Section 3: Essential Questions & Vocabulary Vocabulary Atomic Number Isotope Mass Number Atomic Mass Unit (AMU) Atomic Mass

46. ATOMIC NUMBER – the number of protons in the nucleus of an atom the number of protons in an atom identifies it as an atom of a particular element. Atomic Number All atoms are neutral. So the number of protons must be equal to the number of electrons.

47. Atomic Number Practice

49. Atomic Number Practice SOLUTIONS Oxygen Sodium Chlorine Uranium 6 6 8 11 92 17

50. Sum of the number of protons and neutrons in the nucleus Always a whole number. Not on the periodic table. Mass Number © Addison-Wesley Publishing Company, Inc.

51. Mass Number : Practice

52. Isotopes Atoms of the same element with different mass numbers. Mass # Atomic #  Nuclear symbol:  Hyphen notation: carbon-12

53. If isotopes are the same element but have different mass number, what is different in each isotope? Number of Neutrons Isotope Question? © Addison-Wesley Publishing Company, Inc.

54. In nature, most elements are found as mixtures of isotopes. Usually, the relative abundance of each isotope is constant regardless of where the element is obtained. Natural Abundance of Isotopes

55. C. Johannesson Relative Atomic Mass 12 C atom = 1.992 × 10 -23 g  1 p= 1.007276 amu 1 n = 1.008665 amu 1 e - = 0.0005486 amu © Addison-Wesley Publishing Company, Inc.  atomic mass unit (amu)  1 amu= 1 / 12 the mass of a 12 C atom

56. Atomic Mass Weighted AVERAGE of all isotopes of that element Atomic mass found on the Periodic Table Round to 2 decimal places Write the relative abundance percentage as a decimal.

57. Average Atomic Mass : Practice EX: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16 O, 0.04% 17 O, and 0.20% 18 O.

58. Average Atomic Mass Practice Solution EX: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16 O, 0.04% 17 O, and 0.20% 18 O. 16 O = 16 x 0.9976 = 15.9616 17 O = 17 x 0.0004 = 0.0068 18 O = 18 x 0.0020 = 0.0360 + Average Atomic Mass for O = 16. 00 amu Add all isotope masses together!!

59. Average Atomic Mass : Practice EX: Find chlorine’s average atomic mass if approximately 8 of every 10 atoms are chlorine-35 and 2 are chlorine-37.

60. Average Atomic Mass : Practice Solution 35 Cl = 0.80 x 35 = 28.0 37 Cl = 0.20 x 37 = 7.4 + Average Atomic Mass = 35.4 amu