1. Warm-UP 9/17 (pick up all three papers from the side table when you come in) On a clean sheet of paper, answer the following questions: 1. 1. Who discovered the nucleus of an atom? What experiment led to this discovery? 2. 2. What scientist discovered the electron? What experiment led to this discovery? 3. 3. What is the atomic number and symbol of the following elements: a. tinc. silicon b. mercuryd. sodium

2. Warm-UP 9/2 pre-AP (It’s FRIDAY!) Answer the following on a clean sheet of paper. You will have 15 minutes to complete the warm-up. 1. Use atomic number and mass number to determine the # of subatomic particles. a. Phosphorus (atomic # 15) with a mass number of 32 b. The isotope of rubidium that contains 47 neutrons 2. Write hyphen and nuclear symbol notation for isotopes. a. Sodium isotope with 13 neutrons b. An atom that has 12 protons and 13 neutrons c. The isotope of iron that has a mass number of 56

3. Warm-UP 9/2 (It’s FRIDAY!) Answer the following on a clean sheet of paper. You will have 15 minutes to complete the warm-up. 1. Use atomic number and mass number to determine the # of subatomic particles (protons, electrons, and neutrons). a. Phosphorus (atomic # 15) with a mass number of 32 b. The isotope of rubidium that contains 47 neutrons 2. Write hyphen notation for the following isotopes. a. Sodium isotope with 13 neutrons b. The isotope of iron that has a mass number of 56

4. Warm-Up 9/6/2011 Where is the most dense part of an atom located? What did Rutherford’s experiment determine? Convert the following hyphen notations for an isotope into nuclear symbol notation: C-14 H-3 Na-23 Mg-78

5. Warm-Up 9/7 On your desk is an “isotope and average atomic mass” worksheet. You need to start working on the front part. Please turn in your Atomic Number, Mass Number, and Isotopes homework at the beginning of class. If you haven’t given me your calculator contract, please do so today.

6. Warm-UP 9/8 pre-AP Answer the following on a clean sheet of notebook paper: Verify that the atomic mass of magnesium is 24.31, given the following : 24Mg= 23.985042 amu, 78.99% 25Mg= 24.985837 amu, 10.00% 26Mg= 25.982593 amu, 11.01% * When done, work on your “Isotopes and Average Atomic Mass” assignment from Tuesday. This will be turned in before the test.

7. Chapter 4-The Structure of the Atom 4.1-The Atom: Early Theories of Matter 4.2-Subatomic Particles and the Nuclear Atom 4.3-How atoms differ 4.4- Unstable Nuclei and Radioactive Decay

8. 4.1-The Atom: Early Theories of Matter Pages 87-91

9. The Philosophers Democritus- (470-370 B.C.) Aristotle (384-322 B.C.) John Dalton (1766-1844)

10. Democritus:  400 BC  Developed the idea that matter was divisible into small particles, but not infinitely small particles, called atoms (atom means indivisible in Greek)

11. Aristotle: Lived in the same era as Democritus, but did not believe in atoms; he believed that all matter was continuous—his theory lasted until 1700s

12. In general, scientists in the 1700s believed:  Elements could not be broken down by ordinary means  Compounds formed from elements have different properties than the elements that formed them  Chemical reactions resulted in the transformation of a substance(s) into a new substance

13. Law of Conservation of Mass Mass is neither created nor destroyed during chemical/physical changes

14. Dalton’s Atomic Theory (1808)  Matter is made of atoms  Atoms of one element are identical; atoms of different elements differ in size, mass, etc.  Atoms cannot be subdivided, created or destroyed  Atoms of different elements combine in whole number ratios to form compounds  Atoms are combined, separated, or rearranged during chemical reactions

15. Modern Atomic Theory  All matter is composed of atoms  Atoms of one element differ in properties from atoms of another element

16. 4.2-Subatomic Particles and the Nuclear Atom Pages 92-97

17. Discovery of the Electron  Discovered with cathode-rays  Objects inside the cathode-ray tube cast shadows  A paddle wheel inside the cathode-ray tube spun and moved from cathode to anode  Cathode rays were deflected away from a negative charge by magnetic field

18.  Deduced that cathode rays were made of negative charged particles  Called them electrons J.J. Thomson (1897)

19. Other Inferences Made Based on Cathode Ray Experiments  Atoms are divisible  All atoms contain negative electrons  Positive charge must be present to balance the negative electrons  Electrons are so small (9.1 x 10 -31 kg) that other subatomic particles must account for most of the atom’s mass

20. Ernest Rutherford (1911) Discovery of the Atomic Nucleus  Gold foil bombarded with alpha particles experiment proved that a powerful force had to exist inside the atom  The force had to occupy a very small space, was very dense, and positively charged  Called it the nucleus  Suggested that electrons “orbited” the nucleus

21. Rutherford’s Gold Foil Experiment

22.  Two types of particles  protons (p+) 1.673 x 10 -27 kg  neutrons (n) 1.675 x 10 -27 kg  Same number of protons and electrons in a neutral atom  Number of protons differ for each element  Strong nuclear forces hold the atom together Composition of the Nucleus

23. Size of Atoms  Atomic radius-the distance from the center of the nucleus to the outside of the electron cloud  Expressed in picometers (pm)  1 pm=10 -12 m

24. 9/17 Pre-AP. Today you will need… Take out the paper you were given yesterday Take out your periodic table You will need a calculator today We will talk about basic atom stuff and then calculate average atomic mass.

25. Be prepared to answer the following questions: 1. 1. Who discovered the nucleus of an atom? What experiment led to this discovery? 2. 2. What scientist discovered the electron? What experiment led to this discovery? 3. 3. What is the atomic number, number of electrons, number of protons, and symbol of the following elements: a. tinc. silicon b. mercuryd. sodium

26. Atomic Structure

27. Parts of an Atom   Smallest particle of an element that retains its chemical properties   Composed of two regions   Nucleus--dense core containing protons and neutrons   Electron cloud—large area containing the electrons   Three primary subatomic particles are protons (p+), neutrons (n) and electrons (e-)

28. Atomic Number All atoms made of same basic particles: Protons, electrons, neutrons All atoms made of same basic particles: Protons, electrons, neutrons Differ in number of each particle Differ in number of each particle Atomic number-number of protons in the nucleus of an atom Atomic number-number of protons in the nucleus of an atom Determines the identity of the atom Determines the identity of the atom Where do you find atomic number? Where do you find atomic number? Above the symbol for the element on the periodic table Above the symbol for the element on the periodic table If an atom is neutral, what else does atomic # tell you? If an atom is neutral, what else does atomic # tell you? Number of electrons Number of electrons

29. So…. Atomic number=number of protons=number electrons (in a neutral atom) Atomic number=number of protons=number electrons (in a neutral atom) ALL THE SAME NUMBER!!!

30. Isotopes All atoms of the same element must have same number of protons (atomic #) and electrons (if neutral) All atoms of the same element must have same number of protons (atomic #) and electrons (if neutral) Can have different number of neutrons. Can have different number of neutrons. This gives them different masses. This gives them different masses. Called isotopes of that element (or nuclides) Called isotopes of that element (or nuclides) Example: hydrogen Example: hydrogen

31. Hydrogen How many protons? How many protons? 1 How many electrons? How many electrons? 1 3 different isotopes of hydrogen: 3 different isotopes of hydrogen: Protium (0 neutrons) Protium (0 neutrons) Deuterium (1 neutron) Deuterium (1 neutron) Tritium (2 neutrons) Tritium (2 neutrons)

32. Mass number=Mass of NUCLEUS Sum of the protons and neutrons in the nucleus of an atom (masss #=p + n). NOT ON PERIODIC TABLE! Can round atomic mass to estimate. Sum of the protons and neutrons in the nucleus of an atom (masss #=p + n). NOT ON PERIODIC TABLE! Can round atomic mass to estimate. Used to identify an isotope Used to identify an isotope Example Example An isotope of what element has an atomic number of 17 and a mass number of 40? An isotope of what element has an atomic number of 17 and a mass number of 40? Chlorine Chlorine How many of each subatomic particle is found in this atom? How many of each subatomic particle is found in this atom? 17 protons, 17 electrons, 23 neutrons (40-17) 17 protons, 17 electrons, 23 neutrons (40-17)

33. What information does the atomic number tell you about an atom of a specific element? How many protons & electrons would a neutral phosphorus atoms have? What information does the mass number of an atom tell you? If a nitrogen isotope has a mass number of 14, how many of each subatomic particle would the atom have? Using the periodic table, answer the following questions:

34. Designating Isotopes: Hyphen Notation Write the mass number after the name of the element. Protium=hydrogen-1 (H-1) What is the hyphen notation for the following elements? Lithium that has 4 neutrons Lithium-7 (Li-7) A neutral element with 6 electrons and 8 neutrons. Carbon-14 (C-14)

35. Designating Isotopes: Nuclear Symbol Uses both the atomic number and the mass number with the symbol of the element Place the mass number as a superscript and the atomic number as a subscript to the left of the symbol. Protium: Hint: When you subtract the top from the bottom, you get the # of neutrons. Should always be positive!

36. Nuclear Symbol Practice Nickel isotope that has a mass number of 50. Calcium isotope that has 10 neutrons. Isotope of an element with 4 protons, 4 electrons, and 2 neutrons.

37. Using the periodic table, answer the following questions: If a nitrogen isotope has a mass number of 14, how many of each subatomic particle would the atom have? Write the hyphen and nuclear symbol notation for this nitrogen isotope.

38.  Carbon-12 was arbitrarily chosen as standard by which other atoms are compared  Atomic mass unit is equal to 1/12 the mass of a carbon-12 atom  Mass of a magnesium atom is about twice carbon-12, so its mass is 23.985 amu  Protons and neutrons have a mass ~1 amu Relative Atomic Mass

39. Weighted average of atomic masses of naturally occurring isotopes EXAMPLE: You have 100 marbles, 25% have a mass of 2 g and 75% have a mass of 3g CALCULATE: (.25 x 2g) + (.75 x 3g) = 2.75g, which is the average mass of all the marbles Average Atomic Mass

40. Example: Grades are weighted averages… You are in a level chemistry class and you currently have a 85 for your summative average (60%) and a 54 for your formative average (40%). What is your grade? You are in a level chemistry class and you currently have a 85 for your summative average (60%) and a 54 for your formative average (40%). What is your grade?

41. Example: Grades are weighted averages… You are in a pre-AP chemistry class and you currently have a 85 for your summative average (70%) and a 54 for your formative average (30%). What is your grade? You are in a pre-AP chemistry class and you currently have a 85 for your summative average (70%) and a 54 for your formative average (30%). What is your grade?

42. Average Atomic Mass Example Oxygen has 3 naturally occurring isotopes. Oxygen-16 (15.99 amu) accounts for 99.762%, oxygen-17 (17.00 amu) for 0.038%, and oxygen-18 (18.00 amu) for 0.200 %. Calculate the average atomic mass for oxygen. Avg. AM=(15.99 amu)(0.99762) + (17.00 amu)(0.00038) + (18.00 amu)(0.00200)=15.9992 amu Round to 3 places past decimal  15.999 amu

43. Example… Look at number one on your worksheet and try to calculate average atomic mass for the isotopes on your own. Look at number one on your worksheet and try to calculate average atomic mass for the isotopes on your own.