Chapter 3

1. Chapter 3 Atoms: The Building Blocks of Matter

2. Chapter 3 The Atom: From Philosophical Idea To Scientific Theory

3. The Atom: From Philosophical Idea To Scientific Theory • Remember the idea that matter is made of atoms dates back to 400 BC. • Not proven experimentally until 1700’s

4. ~ ~ The GreeksHistory of the Atom • In 400 B.C the Greeks tried to understand matter (chemicals) and broke them down into earth, wind, fire, and air.

5. Greek Model • Greek philosopher • Idea of ‘democracy’ • Idea of ‘atomos’ • Atomos = ‘indivisible’ • ‘Atom’ is derived • No experiments to support idea Democritus Democritus’s model of atom No protons, electrons, or neutrons Solid and INDESTRUCTABLE

6. Alchemy • After that chemistry was ruled by alchemy. • They believed that that could take any cheap metals and turn them into gold. • Alchemists were almost like magicians. • elixirs, physical immortality

7. GOLD SILVER COPPER IRON SAND . . . . . . . . . . . . . . . Alchemy Alchemical symbols for substances… transmutation: changing one substance into another D In ordinary chemistry, we cannot transmute elements.

8. Contributions of alchemists: • Information about elements • - the elements mercury, sulfur, and antimony were discovered • - properties of some elements • Develop lab apparatus / procedures / experimental techniques • - alchemists learned how to prepare acids. • - developed several alloys • - new glassware

9. ALCHEMY Timeline Greeks (Democratus ~450 BC) Discontinuous theory of matter Issac Newton (1642 - 1727) 400 BC 300 AD 1000 2000 American Independence (1776)

10. Foundations of Atomic Theory Observations and chemical reactions led to the following scientific laws that describe how compounds are formed.

11. Foundations of Atomic Theory • Law of Conservation of Mass (Lavoisier) • Mass is neither destroyed nor created during ordinary chemical reactions or physical changes. • Law of Definite Proportions (Proust) • The fact that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound. • Law of Multiple Proportions (Dalton) • If two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a fixed mass of the first element is always a ratio of small whole numbers.

12. H H O O H2 H2O H H O2 + O H H H2 O H H2O H Conservation of Mass 2 H2 + O2 2 H2O 4 atoms hydrogen 2 atoms oxygen 4 atoms hydrogen 2 atoms oxygen

13. H H O O H2 H2O H H O2 + O H H H2 O H H2O H Legos are Similar to Atoms Legos can be taken apart and built into many different things. Atoms can be rearranged into different substances.

14. High voltage High voltage After reaction O2 H2O 0 g H2 40 g O2 300 g (mass of chamber) + 385 g total Conservation of Mass electrodes Before reaction glass chamber 5.0 g H2 H2 O2 80 g O2 45 g H2O ? g H2O 300 g (mass of chamber) + 385 g total Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 204

15. Law of Definite ProportionsJoseph Louis Proust (1754 – 1826) • Each compound has a specific ratio of elements • It is a ratio by mass • Water is always 16 grams of oxygen for every 2 grams of hydrogen

16. Law of Definite Proportions Whether synthesized in the laboratory or obtained from various natural sources, copper carbonate always has the same composition. Analysis of this compound led Proust to formulate the law of definite proportions. + + 53 g of copper 103 g of copper carbonate 40 g of oxygen 10 g of carbon

17. Law of Multiple ProportionsJohn Dalton (1766 – 1844) If two elements form more than one compound, the ratio of the second element that combines with the first element in each is a simple whole number. H2O H2O2 water hydrogen peroxide Ratio of oxygen is 1:2 (an exact ratio)

18. Dalton’s Atomic Theory • English chemist in the early 1800’s • Dalton stated that elements consisted of tiny particles called atoms • He also called the elements pure substances because all atoms of an element were identical and that in particular they had the same mass.

19. Dalton’s Atomic Theory • All matter consists of extremely small particles that are indivisible and indestructiblecalled atoms. • Atoms of a given element are identical in their physical and chemical properties. Atoms of different elements have different physical and chemical properties. • Atoms of different elements combine in simple, whole number ratios to form chemical compounds. • Atoms cannot be subdivided, created or destroyed. • In chemical reactions, atoms are combined, separated, or arranged. • Although some exceptions have been discovered, the theory still stands today, and has been expanded and modified.

20. Carbon dioxide, CO2 Methane, CH4 Water, H2O Daltons’ Models of Atoms

21. Structure of Atoms • Scientist began to wonder what an atom was like. • Was it solid throughout with no internal structure or was it made up of smaller, subatomic particles? • It was not until the late 1800’s that evidence became available that atoms were composed of smaller parts.

22. Chapter 3 The Structure of the Atom

23. Structure of Atoms • The cathode ray tube led to the discovery of electrons, small, negatively charged particles at the turn of the century. • J. J. Thomson - English physicist • Made a piece of equipment called a cathode ray tube. • It is a vacuum tube - all the air has been pumped out.

24. A Cathode Ray Tube Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 58

25. Thomson’s Experiment voltage source - + vacuum tube metal disks

26. Thomson’s Experiment voltage source ON - OFF + Passing an electric current makes a beam appear to move from the negative to the positive end

27. Thomson’s Experiment voltage source ON - OFF +

28. + - Thomson’s Experiment voltage source ON - OFF + By adding a magnetic field… he found that the moving pieces were negative.

29. J.J. Thomson • He proved that ALL atoms of any element must contain these negative particles. • He knew that atoms did not have a net negative charge and so there must be balancing the negative charge. J.J. Thomson

30. Spherical cloud of Positive charge Electrons Plum Pudding Model • In 1910 proposed the Plum Pudding model • Negative electrons were embedded into a positively charged spherical cloud. Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 56

31. Fluorescent screen with phosphor dots Red beam Green beam Blue beam Shadow mask Television Picture Tube Blue beam Green beam Red beam Glass window Shadow mask Fluorescent screen Electron gun Electron beam Deflecting electromagnets

32. Ernest Rutherford (1871-1937) • Learned physics in J.J. Thomson’ lab. • Noticed that ‘alpha’ particles were sometime deflected by something in the air. • Gold-foil experiment

33. Lead collimator Gold foil a particle source q Rutherford ‘Scattering’ • In 1909 Rutherford undertook a series of experiments • He fired a (alpha) particles at a very thin sample of gold foil • According to the Thomson model the a particles would only be slightly deflected • Rutherford discovered that they were deflected through large angles and could even be reflected straight back to the source

34. Rutherford’s Apparatus beam of alpha particles radioactive substance circular ZnS - coated fluorescent screen gold foil Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 120

35. What He Expected • The alpha particles to pass through without changing direction (very much). • Because… • The positive charges were spread out evenly. Alone they were not enough to stop the alpha particles California WEB

36. What he expected… California WEB

37. What he got…

38. Density and the Atom • Since most of the particles went through, the atom was mostly empty. • Because the alpha rays were deflected so much, the positive pieces it was striking were heavy. • Small volume and big mass = big density • This small dense positive area is the nucleus California WEB

39. Rutherford’s Experiment Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 56

40. The Rutherford Atom n + Zumdahl, Zumdahl, DeCoste, World of Chemistry2002, page 57

41. Size of an atom • Atoms are incredibly tiny. • Measured in picometers (10-12 meters) • Hydrogen atom, 32 pm radius • Nucleus tiny compared to atom • Radius of the nucleus near 10-15 m. • Density near 1014 g/cm • If the atom was the size of a stadium, the nucleus would be the size of a marble. California WEB

42. Atoms All atoms have similar structure • protons and neutrons cluster together to form a nucleus, or central core • electrons orbit the space surrounding the nucleus • then things change depending on the element

43. Chapter 3 Counting Atoms

44. Atoms Each element has a characteristic number of protons • Si- silicon 14 protons • H- hydrogen 1 ptoton • Ag- silver 47 protons ↓ Atomic number

45. Atomic Particles • Protons are constant for an element, but electrons and neutrons can vary • When electrons vary, the charge of the atom changes • When neutrons vary, you have a different isotope of the atom • Isotope- one of 2 or more atoms having the same number of protons but different numbers of neutrons

46. Isotopes Hydrogen 3 isotopes protium deuterium tritium 1 proton 1 proton 1 proton 0 neutrons 1 neutron 2 neutrons • isotopes have very similar chemical properties Deuterium or hydrogen-2

47. X Mass number # protons Atomic number # protons + # neutrons mass number Isotopes Contain the symbol of the element, the mass number and the atomic number

48. Isotopes Atomic Number = number of protons # of protons determines kind of atom/element Atomic Number = number of electrons in a neutral atom Mass Number = the number of protons + neutrons California WEB

49. + + + + + + Mass Number • mass # = protons + neutrons • always a whole number • NOT on the Periodic Table! Neutron Electrons Nucleus Proton Nucleus Carbon-12 Neutrons 6 Protons 6 Electrons 6

50. + F 19 9 Isotopes • Find the • number of protons • number of neutrons • number of electrons • Atomic number • Mass number = 9 = 10 = 9 = 9 = 19