2.1 Atoms and Their Structure

1. 2.1 Atoms and Their Structure Objective Relate historic experiments to the development of the modern model of the atom. Illustrate the modern model of the atom. Interpret the information available in an element block of the periodic table.

2. Early Ideas About Matter • Our current model of the composition of matter is based on hundreds of years of work. • About 2500 years ago, the Greek philosophers thought about the nature of matter and its composition. • They proposed that matter was composed of four fundamental elements…

3. And these were… • Air • Water • Fire • Earth

4. Greek Philosophers • They also debated whether matter could be divided into smaller and smaller pieces endlessly or whether there was an ultimate smallest particle that could no longer be divided. • They were excellent observers, but they did not test their theories with experiments.

5. Democritus, 460-370 B.C. • Proposed that the world is made up a empty space and tiny particles called atoms. • He thought that atoms were the smallest particles of matter and that different types of atoms exist for every type of matter. • Atomic Theory- matter is made up of fundamental particles called atoms.

6. Antoine Lavoisier (1743-1794) • French chemist • In 1782 he made measurements of a chemical change in a sealed container.

7. Lavoisier’s Conclusion • He observed that the mass of the reactants before the chemical reaction were equal to the mass of the products after the reaction. • In a sealed container 2.0 g of hydrogen gas reacted with 16.0 g of oxygen gas to produce 18.0 g of water • Matter was neither created nor destroyed during a chemical reaction, but changed. • This became known as the law of conservation of matter

8. Law of Conservation of Mass

9. Joseph Proust • Joseph Proust, French chemist • September 26, 1754 – July 5, 1826

10. Proust’s Contribution • Observed that the composition of water is always 11% hydrogen and 89% oxygen by mass. Regardless of its source • He observed many other compounds and observed that the elements that composed the compounds were always in a certain proportion by mass. This is referred to as law of definite proportions.

11. John Dalton

12. Dalton’s Atomic Theory • John Dalton (1766-1844), an English schoolteacher and chemist • Studied the experiments of Lavoisier, Proust, and many others. • Based on this he developed his atomic theory.

13. Dalton’s Atomic Theory • All matter is made up of atoms. • Atoms are indestructible and cannot be divided into smaller particles(Atoms are indivisible) • All atoms of one element are exactly alike, but they are different from atoms of other elements.

14. Dalton’s Atomic Theory • This gave chemist a model of the particle nature of matter, but it also raised a lot of questions. • If all elements are made up of atoms then why are there so many different types of elements? • What makes one atom different from another?

15. Dalton’s Atomic Theory • Experiments performed late in the 19th century began to explain the properties and behavior of substances • This was done by the discovery of three smaller particles. • Protons, electrons and neutrons

16. Atomic Theory, Conservation of Matter and Recycling • What happens to the stuff you throw away? What happens to the atoms? • As you have already learned matter is neither created nor destroyed, so what happens.

17. What happens? • When waste is incinerated or buried the atoms may combine with oxygen and other substances to form new compounds. • The atoms are “recycled”

18. Recycling • Recycling has become a part of life now. • Much of what you buy is either recycled or can be recycled. • We have found the advantage of recycling materials and therefore atoms. We mimic what nature does and conserve our natural resources.

19. Hypotheses • Hypothesis – testable prediction to explain observations. • Hypotheses are based on observations. • They can be proven correct or incorrect by the experiments that are designed to test them.

20. Hypothesis

21. Theories • Theory- an explanation based on many observations and supported by the results of many experiments. • As scientists gather new information, a theory may be revised or replaced.

22. Theory

23. Laws • Scientific Law – a fact of nature that is observed so often that it is accepted as the truth. • Examples: • Sun rises in the east • A law can generally be used to make predictions but does not explain why something happens. • Theories explain laws. • One part of Dalton’s atomic theory explains why the law of conservation of matter is true.

24. Review • What is the difference between a theory and a hypothesis? • What is the Law of Conservation of Mass?

25. The Discovery of Atomic Structure • Dalton’s atomic theory was almost true. • He assumed that atoms are the ultimate particles of matter and can’t be broken up into smaller particles and that all atoms of the same element are identical. • His theory needed to be modified with the discovery of electrons, protons and neutrons.

26. The Electrons • In 1897, JJ Thomson a British physicist discovered that the solid ball of model was not accurate. • His experiments involved the use of a cathode ray.

27. Cathode Ray • Composed of a vacuum tube • At the end of the tube is a piece of metal called on electrode which is connected to a metal terminal on the outside of the tube. • The electrode becomes electrically charged when connected to a high voltage. • When the electrode is charged rays travel in the tube from the negative electrode(the cathode) to the positive electrode (the anode)

28. Cathode Ray

29. Cathode Ray

30. Thomson’s Discovery • What he found that the ray would bend towards a positive charged plate and away from a negatively charged plate. • He concluded that the rays are composed of invisible negatively charged particles he called electrons • Electron – negatively-charged particle

31. Early Atomic Model • The early atomic model was referred to as the “plum pudding model”; you could more closely relate it to a chocolate chip cookie • Scientists believed that atoms were balls of positive charge with the negatively charged particles embedded in them.

32. Changes in the Atomic Model Plum Pudding Model Solid Ball Model

33. Ernest Rutherford • In 1909 he carried out the first of the experiments that would reveal an arrangement far different from the “plum pudding” model

34. Rutherford’s Experiment • Gold Foil Experiment • He set up a lead-shielded box containing polonium, which emitted a positively charged beam of alpha particles. • When the beam struck a sheet of gold foil, most of the particles passed straight through the foil • However, some of the particles from the beam were deflected. Some were only slightly deflected and some bounced straight back.

35. Gold Foil Experiment

36. Gold Foil Experiment

37. The Nuclear Model of the Atom • Based on Rutherford’s work the team devised a new model for the atom. • Because some of the particles bounced straight back they concluded that atoms must have a dense central core called the nucleus • Nucleus – a small, dense, positively charged central core of an atom

38. Nuclear Model

39. Changes in the Atomic Model Plum Pudding Model Nuclear Model Solid Ball Model

40. Atomic Model • It was hard for people to grasp that atoms contained a lot of empty space. • When they looked at a rock, it was very difficult to see how most of this object could be empty space. • If you were to enlarge an atom of hydrogen so the nucleus was the size of a golf ball, the electron would be a mile a way. • In one drop of water there is 6,500,000,000,000,000,000,000 (6.5 sextillion) atoms

41. Atomic Numbers and Masses • The nucleus of an atom is composed of protons and neutrons. Electrons move in the space around the nucleus. • Atomic Number – the number of protons in the nucleus of an atom. • Every element has a unique atomic number. Therefore, it is the number of protons that determines the identity of an element.

42. Atoms • Protons - positively charged subatomic particles. • Neutrons – subatomic particle that does not have a charge, it is neutral • Atoms have no overall charge. • Therefore the number of protons and electrons must be equal. • If you elements atomic number is 2 • 2 protons • 2 electrons

43. Mass Number • Mass number - the number of protons and neutrons in the nucleus of an atom. • Isotope – atom of an element with a different number of neutrons and therefore a different mass number • Ne-20 = neon 20(10 protons, 10 neutrons) • Ne-21 = neon 21(10 protons, 11 neutrons)

44. Neon

45. Composition of Atoms • You can determine the composition of any element if you know the • Atomic number • Mass number

46. Atomic Mass • Since elements have different isotopes the atomic mass is the weighted average. • Located under the symbol on the periodic table. • You can use the atomic mass to determine the mass number of a element • You take the atomic mass rounded up

47. Determining Atomic Mass 62.930 X 0.6917 = 43.529 64.928 X 0.3083 = 20.017 43.529 + 20.017 = 63.546

48. Review • What information does the atomic number give us? • How do we determine the number of neutrons of any element? • Who discovered the electron? • What discovery did the Gold Foil Experiment led to?