1. Topic 2.1: Atomic Structure Honors Chemistry 2014-15 Mrs. Peters 1

2. Atomic Structure 2.1: The nuclear atom EI: The mass of the atoms is concentrated in its minute, positively charged nucleus. NOS: 1.Evidence and improvements in instrumentation – alpha particles were used in the development of the nuclear model of the atom that was first proposed by Rutherford. (1.8) 2.Paradigm shifts- the subatomic particle theory of matter represents a paradigm shift in science that occurred in the late 1800s (2.3) 2

3. Atomic Structure 2.1: The nuclear atom Understandings: 1.Atoms contain a positively charged dense nucleus composed of protons and neutrons (nucleons) 2.Negatively charged electrons occupy the space outside the nucleus 3.The mass spectrometer is used to determine the relative atomic mass of an element from its isotopic composition. 3

4. Atomic Structure 2.1: The nuclear atom Applications and Skills: 1.Use of the nuclear symbol notation A Z X to deduce the number of protons, neutrons, and electrons in atoms and ions. 2.Calculations involving non-integer relative atomic masses and abundance of isotopes from given data, including mass spectra. 4

5. NOS: Paradigm Shift History behind Atomic Theory Democritus ( 420 BCE ) first proposed the idea that matter may be made up of small, indivisible particles called atoms. Aristotle (384-322 BCE) Greek philosopher; matter composed of earth, air, fire, water. This view dominated thought until 17th century 5

6. NOS: Paradigm Shift History behind Atomic Theory Atomism developed in Chinese & Arabic cultures during the Dark Ages in Europe. John Dalton (1766-1844) was the first to base atomic theory on scientific evidence. 6

7. NOS: Paradigm Shift Dalton’s Atomic Theory Elements are made of tiny particles called atoms. All atoms of a given element are identical. The atoms of a given element are different from those of any other element. 7

8. NOS: Paradigm Shift Dalton’s Atomic Theory Atoms of one element can combine with atoms of other elements to form compounds. A given compound always has the same relative number of types of atoms. Atoms cannot be created, nor divided into smaller particles, nor destroyed in the chemical process. A chemical reaction simply changes the way atoms are grouped together. 8

9. NOS: Paradigm Shift Evidence for sub-atomic particles 1897: J.J. Thomsen: Cathode Ray Tube Evidence for electrons: Bent a stream of rays originating from the negative electrode (cathode). Stream of particles with mass & negative charge. 9

10. NOS: Paradigm Shift Evidence for sub-atomic particles 1909: Ernest Rutherford: Gold Foil Evidence for protons & nucleus: Alpha particles deflected passing through gold foil 10

11. NOS: Paradigm Shift Evidence for sub-atomic particles 1932: James Chadwick: Beryllium Evidence for neutrons: Alpha particles caused beryllium to emit rays that could pass through lead but not be deflected, 11

12. U1. and U2. Atomic Structure Sub-Atomic Particles: Proton: Located in the nucleus Relative charge of +1 Relative mass of 1 amu Neutron: Located in the nucleus Relative charge of 0 Relative mass of 1 amu 12 www.green-planet-solar-energy.comwww.green-planet-solar-energy.com, 3.bp.blogspot.com

13. U1. and U2. Atomic Structure Sub-Atomic Particles Electron: Located in cloud surrounding the nucleus Relative charge of –1 Relative mass of 0.0005 amu 13 www.green-planet-solar-energy.comwww.green-planet-solar-energy.com, 3.bp.blogspot.com

14. U1. and U2. Atomic Structure Nucleus consists of protons and neutrons with the electrons surrounding the nucleus. In a neutral atom, the #protons = # electrons. 14

15. A1. Nuclear Symbol Notation Atomic Number (Z) The atomic number is the number of protons in the nucleus. It determines the identity of an atom. All oxygen atoms have 8 protons in the nucleus All lead atoms have 82 protons in the nucleus 15

16. A1. Nuclear Symbol Notation Atomic Number (Z) It also tells us the number of electrons in a neutral atom A neutral sodium atom contains 11 protons and 11 electrons A neutral bromine atom contains 35 protons and 35 electrons 16

17. A1. Nuclear Symbol Notation Mass Number (A) It is not practical to measure the masses of atoms in grams due to their small size. Scientists devised a measurement called atomic mass units (amu). 17

18. A1. Nuclear Symbol Notation Mass Number (A) Protons have a mass of 1 amu Neutrons have mass of 1 amu Electrons have mass of 0 amu. Mass Number of atoms = # protons + # neutrons 18

19. A1. Nuclear Symbol Notation Mass Number (A) Mass Number of atoms = # protons + # neutrons **Round the Relative Atomic Mass to a whole number to find the Mass Number Ex: Lithium = 6.94 Mass Number is 7 Magnesium = 24.31 Mass Number is 24 19

20. How to read the Periodic Table 20 Lithium 3 Li 6.94 Element Name Atomic Number Element Symbol Relative Atomic Mass

21. A1. Nuclear Symbol Notation Atomic Name: Element Name - A (mass number) Ex: Carbon-12 Nuclear Symbol: 21 X A Z Element Symbol Mass Number (Protons + Neutrons) Atomic Number

22. Important Terms Isotope Atoms of the same element can have different numbers of neutrons, thus they will have different atomic masses. These are called isotopes of the element. These are the same element, just different numbers of neutrons and mass. 22

23. Important Terms Isotope Example There are three isotopes of hydrogen: Hydrogen- 1 has 1 proton, 1 electron, 0 neutrons Hydrogen- 2 has 1 proton, 1 electron, 1 neutron Hydrogen- 3 has 1 proton, 1 electron, 2 neutrons 23

24. A1. Deduce the symbol given its mass number and atomic number Consider an atom that has an atomic number of 29 and a mass number of 63. What is its name and symbol? Name: Symbol: 24

25. A. Deduce the symbol given its mass number and atomic number Consider an atom that has an atomic number of 29 and a mass number of 63. What is its name and symbol? atomic number of 29 identifies it as copper Name: Copper-63 Symbol: 63 Cu 29 25

26. A1. Deduce the symbol given its mass number and atomic number Consider an atom that has A=32 and Z=16. What is its name and symbol? Name: Symbol: 26

27. A1. Deduce the symbol given its mass number and atomic number Consider an atom that has A=32 and Z=16. What is its name and symbol? Z=16 identifies it as sulfur Name: Sulfur-32 Symbol: 32 S 16 27

28. A1. Deduce the symbol given its mass number and atomic number Consider an atom that has an atomic number of 74 and a mass number of 185. What is its name and symbol? Consider an atom that has A=127 and Z=53. What is its name and symbol? 28

29. A1. Deduce the symbol given its mass number and atomic number Consider an atom that has an atomic number of 74 and a mass number of 185. What is its name and symbol? atomic number of 74 identifies it as tungsten Name: Tungsten-185 Symbol: 185 W 74 Consider an atom that has A=127 and Z=53. What is its name and symbol? Z=53 identifies it as iodine Name: Iodine-127 Symbol: 127 I 53 29

30. A1. Deduce protons, neutrons, and electrons in atoms and ions from the A, Z, and charge Consider the neutral carbon-12 atom. Find the A, Z, protons, neutrons, electrons, and symbol Name is Carbon-12 Atomic mass (A) = 12 Atomic number (Z) = 6 Protons = 6 (atomic number) Neutrons = 6 (mass – protons) Electrons = 6 (neutral atom so same as protons) Symbol is 12 C 6 30

31. A1. Deduce protons, neutrons, and electrons in atoms and ions from the A, Z, and charge Consider an atom that has 9 protons, 9 electrons, and 10 neutrons. What is its atomic number, atomic mass, name, and symbol? Z=9 (atomic number = # protons) A=19 (atomic mass = protons + neutrons) Fluorine-19 (name and mass) 19 F (neutral because protons = electrons) 9 31

32. A1. Deduce protons, neutrons, and electrons in atoms and ions from the A, Z, and charge Consider a neutral atom with A=75 and Z=33. How many protons, neutrons, and electrons are in the atom. What is the name and symbol? Consider a neutral atom with A=77 and Z=33. How many protons, neutrons, and electrons are in the atom. What is the name and symbol? 32

33. A1. Deduce protons, neutrons, and electrons in atoms and ions from A, Z, and charge Consider a neutral atom with A=75 and Z=33. How many protons, neutrons, and electrons are in the atom. What is the name and symbol? Protons = 33Neutrons = 42Electrons = 33 Name: Arsenic-75 Symbol: 75 As 33 Consider a neutral atom with A=77 and Z=33. How many protons, neutrons, and electrons are in the atom. What is the name and symbol? Protons = 33Neutrons = 44Electrons = 33 Name: Arsenic-77 Symbol: 77 As 33

34. Important Terms Ions are charged particles formed when atoms gain or lose electrons resulting in unequal numbers of protons and electrons Cations: Atoms that lose electrons become positively charged Anions: Atoms that gain electrons become negatively charged 34

35. A1. Use of Nuclear Symbol Notation Nuclear Symbol: 35 X A Z Element Symbol Mass Number (Protons + Neutrons) Atomic Number + Charge (+, - or nothing) Determined by electrons

36. Important Terms How many protons, neutrons, and electrons are in an ion of K-39 that has lost one electron? What is the charge of the ion? What is its symbol? Protons = 19 Neutrons = 20 Electrons = 18 Charge = 1+ or +1Symbol is 39 K 1+ 19 36

37. A1. Deduce from nuclear symbol notation The symbol of an anion is 31 P 3-. Calculate the number of 15 protons, neutrons, and electrons. What is Z and what is A? What is the symbol of a species containing 26 protons, 30 neutrons, and 23 electrons? What is the symbol of a species with A=56, Z=26, and 24 electrons? 37

38. A1. Deduce the number of protons, neutrons, and electrons in atoms and ions from A, Z, and charge The symbol of an anion is 31 P 3-. Calculate the number 15 protons, neutrons, and electrons. What is Z and what is A? #P = 15; #N = 16; #E = 18; Z= 15; A = 31 What is the symbol of a species containing 26 protons, 30 neutrons, and 23 electrons? 56 Fe 3+ 26 What is the symbol of a species with A=56, Z=26, and 24 electrons? 56 Fe 2+ 26 38

39. A2. Discuss the use of radioisotopes. Radioisotopes: isotopes of elements that have become radioactive because the nucleus is unstable and breaks down spontaneously emitting radiation. Radioisotopes can occur naturally or be created artificially Examples: Carbon-14; Iodine-125; Strontium-90; Cobalt-60, Iodine-131 39

40. A2. Discuss the use of radioisotopes. Uses of Radioisotopes Nuclear power generation Sterilization of surgical instruments Crime detection Food preservation Dating artifacts Treating and diagnosing disease 40

41. U3. The Mass Spectrometer Mass Spectrometers: Instruments that measure charge-to-mass ratio of charged particles. Used to measure masses of isotopes as well as isotopic abundance 41

42. U3. The Mass Spectrometer How a Mass Spec Works: 1. Vaporization: sample is heated to gas state 2. Ionization: sample gas is turned into ions by blasting free electrons to knock electrons off from the gas atoms, creating positive ions 3. Acceleration: increases the speed of particles, using an electric field 42

43. U3. The Mass Spectrometer How a Mass Spec Works: 4. Deflection: using an electromagnet to create a magnetic field, amount of deflection depends on mass and charge of the ion (think of cars going around a corner) 5. Detection: measures both mass and relative amounts (abundance) of all the ions present 43

44. U3. The Mass Spectrometer Mass Spectrometer Video http://www.youtube.com/watch?feature=fv wp&v=lxAfw1rftIA&NR=1 http://www.youtube.com/watch?feature=fv wp&v=lxAfw1rftIA&NR=1 44

45. U3. Relative Atomic Mass Relative Atomic Mass Mass numbers (atomic mass) on the periodic table are weighted averages of the isotopes. Based on 12 C. Has 6 protons, 6 neutrons, and 6 electrons Has a relative atomic mass of exactly 12.000 45

46. U3. Relative Atomic Mass Relative Atomic Mass One amu is exactly 1/12 of the mass of a carbon-12 atom. All other isotopes are measured compared to this value. 46

47. U3. Relative Atomic Mass Average relative atomic mass: the weighted average for all of the isotopes of a given element, based on the percent abundance of each 47

48. U3. Relative Atomic Mass To determine Average Relative Atomic Mass: Need masses of each isotopes Need abundance (percentage) of each isotope o The mass spec is used to determine these values This is the value shown on the periodic table 48

49. U3. Relative Atomic Mass A sample of neon is placed in the mass spectrometer 49

50. U3. Relative Atomic Mass A sample of neon is placed in the mass spectrometer The results show the abundance for each isotope of an element o 90.92% is neon-20 o 0.26% is neon-21 o 8.82% is neon-22 50

51. A2. Calculate non-integer relative atomic masses and abundance of isotopes from given data. Example 1.95.5% =.955.5%=0.005 2.24.5 x.955= 23.4 and 23.7 x.005 =.119 3.23.4 +.119 = 23.519 4.23.5 amu How to Determine Relative Atomic Mass 1. Convert the percent abundance for each isotope into decimal 2. Multiply the mass for each isotope by the abundance 3. Add all product values from step 2. 4. Include amu for the units of the value. 51

52. A2. Calculate non-integer relative atomic masses and abundance of isotopes from given data. Three isotopes of magnesium occur in nature. Their abundances and masses, determined by mass spectrometry, are listed in the table on the right. Use this information to calculate the atomic weight of magnesium. Three isotopes: 24, 25, 26 Percentage of each isotope: Given Multiply the percent of each isotope by its mass 23.98504 x.7899 = 18.95 amu 24.98584 x.1000 = 2.499 amu 25.98259 x.1101 = 2.861 amu Add these values = 24.31 amu Isotope% AbundanceMass (amu) 24 Mg78.9923.98504 25 Mg10.0024.98584 26 Mg11.0125.98259 52

53. A2. Calculate non-integer relative atomic masses and abundance of isotopes from given data. Calculate the atomic weight of chromium using the following data for the percent natural abundance and mass of each isotope: 4.35% 50 Cr (49.9461 amu); 83.79% 52 Cr (51.9405 amu); 9.50% 53 Cr (52.9406 amu); 2.36% 54 Cr (53.9389 amu) 53

54. A2. Calculate non-integer relative atomic masses and abundance of isotopes from given data. Calculate the atomic weight of chromium using the following data for the percent natural abundance and mass of each isotope: 4.35% 50 Cr (49.9461 amu); 83.79% 52 Cr (51.9405 amu); 9.50% 53 Cr (52.9406 amu); 2.36% 54 Cr (53.9389 amu) 49.9461 x.0435 = 2.17 amu 51.9405 x.8379 = 43.52 amu 52.9406 x.0950 = 5.03 amu 53.9389 x.0236 =+ 1.27 amu 51.99 amu 54

55. A2. Calculate non-integer relative atomic masses and abundance of isotopes from given data. Determine the atomic weight of lead using the data from the mass spectrum of lead Four isotopes: 204, 206, 207, 208 Percentage of each isotope: Total # isotopes is 10 (1+2+2+5) 204: 1/10 = 10% 206: 2/10 = 20% 207: 2/10 = 20% 208: 5/10 = 50 % Multiply the percent of each isotope by its mass 204 x.1 = 20.4 206 x.2 = 41.2 207 x.2 = 41.4 208 x.5 = 104 Add these values 20.4 + 41.2 + 41.4 + 104 = 207 55 Mass Spectrum of Lead

56. A2. Calculate non-integer relative atomic masses and abundance of isotopes from given data. The atomic weight of gallium is 69.72 amu. The masses of the naturally occurring isotopes are 68.9257 amu for 69 Ga and 70.9249 amu for 71 Ga. Calculate the percent abundance of each isotope. Let x = % abundance of 69 Ga. Then 1-x = % abundance of 71 Ga. 68.9257x + 70.9249(1-x) = 69.72 amu 68.9257x + 70.9249 – 70.9249x = 69.72 -1.9992x = -1.20 x = 0.600 = decimal value of 69 Ga so 60.0% 69 Ga 1-x = 0.400 = decimal value 71 Ga so 40.0 % 71 Ga 56

57. A2. Calculate non-integer relative atomic masses and abundance of isotopes from given data. The atomic weight of copper is 63.546 amu. The masses of the two naturally occurring isotopes are 62.9298 amu for 63 Cu and 64.9278 amu for 65 Cu. Calculate the percent of 63 Cu in naturally occurring copper. 57

58. A2. Calculate non-integer relative atomic masses and abundance of isotopes from given data. The atomic weight of copper is 63.546 amu. The masses of the two naturally occurring isotopes are 62.9298 amu for 63 Cu and 64.9278 amu for 65 Cu. Calculate the percent of 63 Cu in naturally occurring copper. Let x = % abundance of 63 Cu. Then 1-x = % abundance of 65 Cu. 62.9298x + 64.9278(1-x) = 63.546 amu 62.9298x + 64.9278 – 64.9278x = 63.546 -1.998x = -1.382 x = 0.6917 = decimal of 63 Cu so 69.17% 63 Cu 58