Notes for chapter 4

1. Atomic Structure Defining the Atom

2. Opening What are we going to learn today? Essential Questions: How did Democritus describe atoms? How did John Dalton further Democritus’s ideas on atoms? What instruments are used to observe individual atoms? GPS Standards: SCSh1b – Recognize that different explanations can be given for the same information SCSh7c – Understand how major shifts in scientific knowledge occur. SCSh7d – Hypotheses often cause scientists to develop new experiments that produce additional data. SCSh7e – Testing, revising and occasionally rejecting new and old theories never ends.

3. Why are we doing this? (logbook) What are some objects that require experimental data in order to “picture” them, either because they are small or inaccessible? Here’s how. (agenda) Discuss early models of the atom Begin gathering information for an atomic theory timeline. Complete Section Review WS Wrap-up Evaluate and criticize the following statements: “All atoms are identical.” Chemical reactions occur when atoms of one element change into atoms of another element.”

4. Early Models of the atom atom Smallest particle of an element that retains its identity in a chemical reaction Democritus (460BC – 370BC) Greek philosopher One of the first to suggest the existence of atoms Believed that atoms were indivisible and indestructible No attempt to explain chemical behavior No experimental support

5. John Dalton (1766 – 1844) English chemist and schoolteacher Dalton’s atomic theory All elements are composed of tiny indivisible particles called atoms. Atoms of the same element are identical. The atoms of any one element are different from those of any other element. Atoms of different elements can physically mix together or can chemically combine in simple whole-number ratios to form compounds. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction.

6. Sizing up atoms Pure copper penny contains about 2.4x1022 atoms of copper 6x109 people on Earth Radius of most atoms is between 5x10-11m and 2x10-10m. Individual atoms can only be observed by scanning tunneling electron microscope

7. Scanning tunneling electron microscope

8. SCE Microscope depiction of electron cloud

9. Diagram of scanning tunneling electron microscope

10. Quantum forest

11. Opening – 8/29/11 What are we going to learn today? GPS Standards SCSh1b – Recognize that different explanations can be given for the same information SCSh1c – Explain that further understanding of scientific problems relies on the design and execution of new experiments which may reinforce or weaken opposing explanations. SCSh7c – Understand how major shifts in scientific knowledge occur. SCSh7d – Hypotheses often cause scientists to develop new experiments that produce additional data. SCSh7e – Testing, revising and occasionally rejecting new and old theories never ends. SC3a – Discriminate between the relative size, charge, and position of protons, neutrons, and electrons in the atom.

12. Why are we doing this? Essential Questions What are 3 kinds of subatomic particles? How can you describe the structure of the nuclear atom? Here’s how. (agenda) Notes/discussion about the discovery of protons, neutrons, and electrons Students work in pairs to complete a chart about the three subatomic particles. Complete Section 4.2 Review Sheet

13. Subatomic particles Three kinds of subatomic particles Proton Positive charge Neutron No charge Electron Negative charge

14. Discovery of the electron J.J. Thomson (1856 – 1940) Discovered the electron in 1897 Cathode ray tube experiments Hypothesized that cathode rays are tiny negatively charged particles moving at high speed (electrons) Measured the charge to mass ratio of the electron Plum Pudding Model

15. Cathode ray deflected by a magnet Plum Pudding Model

16. Robert Millikan (1868 – 1953) determined the quantity of charge on an electron Used Thomson’s charge-mass ratio to calculate the mass of the electron (1916) Oil drop experiments

17. Millikan oil drop experiment

18. Discovery of the proton Eugen Goldstein (1850 – 1930) Found rays traveling in the direction opposite to that of the cathode rays in a cathode ray tube Called these rays canal rays (later renamed protons)

19. Discovery of the Neutron James Chadwick (1891 – 1974) Discovered the neutron (1932)

20. Properties of subatomic particles

21. Discovery of the Nucleus Ernest Rutherford (1871 – 1937) Gold foil experiments (1911) Findings Atom is mostly empty space Small positively charged nucleus Electrons move around outside the nucleus Nuclear model

22. Explanation of results of gold foil experiment

23. Comparison of Thomson’s plum pudding model (top) and Rutherford’s nuclear model (bottom) Notice that the nucleus in this model is solid. Protons and neutrons had not been discovered.

24. Ticket out the door Write a paragraph explaining how Rutherford’s gold foil experiment yielded new evidence about atomic structure. Hint: First describe the setup of the experiment. The explain how Rutherford interpreted his experimental data.

25. Opening Essential Questions: What makes one element different from another? How do you find the number of neutrons in an atom? How do isotopes of an element differ? How do you calculate the atomic mass of an element? Why is the periodic table useful? GPS Standards: SC3c – Explain the relationship of the proton number to the element’s identity. SC3d – Explain the relationship of isotopes to the relative abundance of atoms of a particular element.

26. Atomic Number The number of protons in an atom identifies the element. Atomic number the number of protons in the nucleus of an atom Each element has a unique atomic number Because atoms are neutral, the number of electrons(-1) must equal the number of protons(+1).

27. Sample Problem p. 111 How many protons and electron are in each of the following atoms? Fluorine Calcium Aluminum

28. Mass Number Mass number Total number of protons and neutrons in the nucleus of the atom # neutrons = mass number – atomic number

29. Hyphen notation Name of element followed by a hyphen and the mass number Examples Carbon – 12 Carbon – 14 Oxygen – 18

30. Nuclear Notation The symbol of the element Mass number as a superscript before the symbol Atomic number as a subscript before the symbol

31. isotopes Isotopes atoms of the same element that have different masses Atoms that have the same number of protons but different numbers of neutrons Atoms that have the same atomic number but different mass numbers

32. Isotopes of hydrogen Protium Hydrogen – 1 Deuterium Hydrogen – 2 Tritium Hydrogen – 3

33. Atomic mass Atomic mass unit (amu) 1/12 the mass of a carbon-12 atom Mass of a single proton or neutron is approximately 1amu Atomic mass weighted average mass (in amu) of the atoms in a naturally occurring sample of an element Mass shown on the periodic table

34. Calculating Atomic mass Atomic mass = [(relative abundance)(atomic mass of the isotope)] for each naturally occurring isotope Multiply the relative abundances (expressed as a decimal) times the mass of each isotope then add the results

35. Sample problem, p. 117 Atomic mass = 1.993 amu + 8.817 amu = 10.81 amu

36. Periodic Table Preview Periodic table – an arrangement of elements in which the elements are separated into groups based on a set of repeating properties Period Horizontal row on the Periodic Table 7 periods Properties vary as you move across a period Group or family Vertical column of the Periodic Table 18 groups Elements within a group have similar properties