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Chapter 2 Atoms & Elements

Chapter 2 Atoms & Elements

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Chapter 2 Atoms & Elements

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  1. Chapter 2Atoms & Elements General Chemistry I T. Ara

  2. A. Atomic Structure • Although much of Dalton’s Atomic Theory still holds true, experiments carried out in the beginning of the 20th century suggested that atoms are made up of even smaller subatomic particles. • The first two subatomic particles to be detected were the electron and the proton.

  3. 1. Electrons & Protons • Electrons are small (9.1094x10-28 g) and negatively charged (–1). • Protons are larger (1.6726x10-24 g – almost 2000 times the mass of an electron) and positively charged (+1). • Every neutral atom must have an equal number of electrons (–1) and protons (+1).

  4. One More Problem. . . • Most neutral atoms have masses greater than the sum of the masses of their electrons and protons. • There must be another subatomic particle that has mass, but no charge! 2. Neutrons are similar in mass to protons (1.6749x10-24 g), but they are uncharged.

  5. A. Atomic Structure (cont.) • If atoms are composed of subatomic particles like electrons, protons, and neutrons, how are these particles arranged? • In 1910, Ernest Rutherford carried out an experiment that gave scientists important information about the structure of the atom…

  6. Rutherford’s Gold Leaf Experiment Conclusion: All of the positive charge and most of an atom’s mass is concentrated in a very small region (the nucleus), and only the electrons occupy the space outside of the nucleus. Atoms are mostly vacant space!

  7. 3. Nucleus • In an atom, the protons and neutrons make up the nucleus, and the electrons occupy the space around the nucleus (more on electrons later). • The electron cloud of an atom is approximately 10,000 times larger than its nucleus.

  8. 3. Nucleus • The smallest unit of an element is an atom. • The number of protons in an atom’s nucleus determines what type of element it is. • All atoms of the same element have the same number of protons in their nuclei. a) Atomic Number (Z): number of protons in a given nucleus

  9. All atoms of the same element must have the same number of protons, but they can have different numbers of neutrons. • A complete description of an atom’s nuclear composition must therefore include the mass numberin addition to the atomic number. Mass number b) Mass Number: the sum of the protons and neutrons in the nucleus of an atom Why is it called the mass number? A X Z Atomic number Atomic symbol

  10. c) Atomic Mass Unit (amu) • Atoms are so small that it is not practical to describe atomic masses in terms of grams. • For simplicity, a smaller unit (atomic mass unit) is used in which atomic masses are defined relative to a standard. • The carbon atom (which has 6 protons and 6 neutrons) is defined as having a mass of exactly 12 amu.

  11. i) Mass of Subatomic Particles • As a result, the masses of the three subatomic particles in atomic mass units are: • Proton: 1.0073 amu • Neutron: 1.0087 amu • Electron: 0.00055 amu Notice that protons and neutrons each have masses very close to 1 amu while electrons are much smaller (negligible).

  12. This makes it easy to estimate the atomic mass of an atom in atomic mass units: Atomic Mass (in amu) ≈ # protons + # neutrons = Mass Number (A) • The Mass Number (A) is an approximation of the mass of an individual atom.

  13. Draw the atomic symbol for a copper atom with 29 protons & 34 neutrons.

  14. Draw the atomic symbol for a copper atom with 29 protons & 34 neutrons. • Atomic number (Z) = 29 • Mass number (A) = 29 + 34 = 63 29Cu 63

  15. 4. Atomic Mass & Weight • The mass number is an approximate atomic mass in atomic mass units. • The exact mass of an atom can be measured using a mass spectrometer.

  16. 4. Atomic Mass & Weight Mass Spectrum of Neon Why does the mass spectrum of neon show atoms with three different masses? We know that all neon atoms must have the same number of protons (10). These different masses must result from atoms with different numbers of neutrons.

  17. a) Isotopes Isotopes: atoms with the same atomic number (Z) but different mass numbers (A). eg. Two isotopes of chlorine (Z = 17) with different numbers of neutrons

  18. Fill in the table with the correct data. • protons electrons neutrons 13 C 6 23 Na 11 7 Li 3 18 O 8 16 O 8

  19. Fill in the table with the correct data. • protons electrons neutrons 13 C 6 6 23 Na 11 7 Li 3 18 O 8 16 O 8

  20. Fill in the table with the correct data. • protons electrons neutrons 13 C 6 6 6 23 Na 11 7 Li 3 18 O 8 16 O 8

  21. Fill in the table with the correct data. • protons electrons neutrons 13 C 6 6 7 6 23 Na 11 7 Li 3 18 O 8 16 O 8

  22. Fill in the table with the correct data. • protons electrons neutrons 13 C 6 6 7 6 23 Na 11 11 7 Li 3 18 O 8 16 O 8

  23. Fill in the table with the correct data. • protons electrons neutrons 13 C 6 6 7 6 23 Na 11 11 11 7 Li 3 18 O 8 16 O 8

  24. Fill in the table with the correct data. • protons electrons neutrons 13 C 6 6 7 6 23 Na 11 11 12 11 7 Li 3 18 O 8 16 O 8

  25. Fill in the table with the correct data. • protons electrons neutrons 13 C 6 6 7 6 23 Na 11 11 12 11 7 Li 3 3 18 O 8 16 O 8

  26. Fill in the table with the correct data. • protons electrons neutrons 13 C 6 6 7 6 23 Na 11 11 12 11 7 Li 3 3 3 18 O 8 16 O 8

  27. Fill in the table with the correct data. protons electrons neutrons 13 C 6 6 7 6 23 Na 11 11 12 11 7 Li 3 3 4 3 18 O 8 16 O 8

  28. b) Percent Abundance • A naturally occurring sample of an element is generally a mixture of isotopes. • For example, a sample of boron consists of two isotopes: 10B and 11B. Percent Abundance: the proportion of atoms of each isotope in a natural sample of an element; the percentage of atoms of a particular isotope

  29. c) Atomic Weight • The atomic weightisthe average atomic mass for the naturally occurring element, expressed in atomic mass units. • The atomic weight of an element is NOT equal to the mass of any of the individual isotopes.

  30. d) Calculating Atomic Weight • The atomic weight of an element can be calculated using the percent abundance and isotopic mass for each isotope – think of it as a weighted average. eg. If an element is composed of two isotopes (A & B): Atomic Weight = (mass A)(%ab A) + (mass B)(%ab B)

  31. d) Calculating Atomic Weight Calculate the atomic weight of lithium, Li, from the following data: Isotope Isotopic Mass % Abundance Li-6 6.015121 amu 7.500% Li-7 7.016003 amu 92.50%

  32. d) Calculating Atomic Weight Calculate the atomic weight of lithium, Li, from the following data: Isotope Isotopic Mass % Abundance Li-6 6.015121 amu 7.500% Li-7 7.016003 amu 92.50% Li-6: (6.015121 amu)(0.07500) = 0.45113 amu Li-7: (7.016003 amu)(0.9250) = 6.4898 amu Total: 0.45113 amu + 6.4898 amu = 6.941 amu (We’ll talk about significant figures later.)

  33. e) Calculating Percent Abundance • The same equation can be used to calculate percent abundance if the atomic weight is known. Atomic Weight = (mass A)(%ab A) + (mass B) (%ab B) • What other mathematical relationship do we need to use to do this calculation?

  34. e) Calculating Percent Abundance Silver has two naturally occurring isotopes, one of mass 106.91 amu (A), and the other of mass 108.90 amu (B). The atomic weight is 107.87 amu. Calculate the percent abundance for these isotopes.

  35. e) Calculating Percent Abundance Silver has two naturally occurring isotopes, one of mass 106.91 amu (A), and the other of mass 108.90 amu (B). The atomic weight is 107.87 amu. Calculate the percent abundance for these isotopes. n = % abundance of A 1-n = % abundance of B [sum for A and B must equal 1 (100%)] 106.91n + (108.90) (1-n) = 107.87 106.91n + 108.90 – 108.90n = 107.87 -1.99n = -1.03 n = .518 % Abundance A = 51.8% % Abundance B = 48.2%

  36. 5. Ions • Ion:An electrically charged particle obtained from an atom, or group of atoms, by adding or removing electrons. -When an atom loses an electron, it becomes a positively charged cation. -When an atom gains an electron, it becomes a negatively charged anion.

  37. 5. Ions

  38. 5. Ions • protons electrons neutrons 40 Ca2+ 20 23 Na+ 11 18 O2- 8

  39. 5. Ions • protons electrons neutrons 40 Ca2+ 20 20 23 Na+ 11 18 O2- 8

  40. 5. Ions • protons electrons neutrons 40 Ca2+ 20 18 20 23 Na+ 11 18 O2- 8

  41. 5. Ions • protons electrons neutrons 40 Ca2+ 20 18 20 20 23 Na+ 11 18 O2- 8

  42. 5. Ions • protons electrons neutrons 40 Ca2+ 20 18 20 20 23 Na+ 11 11 18 O2- 8

  43. 5. Ions • protons electrons neutrons 40 Ca2+ 20 18 20 20 23 Na+ 11 10 11 18 O2- 8

  44. 5. Ions • protons electrons neutrons 40 Ca2+ 20 18 20 20 23 Na+ 11 10 12 11 18 O2- 8

  45. 5. Ions • protons electrons neutrons 40 Ca2+ 20 18 20 20 23 Na+ 11 10 12 11 18 O2- 8 8

  46. 5. Ions • protons electrons neutrons 40 Ca2+ 20 18 20 20 23 Na+ 11 10 12 11 18 O2- 8 10 8

  47. 5. Ions • protons electrons neutrons 40 Ca2+ 20 18 20 20 23 Na+ 11 10 12 11 18 O2- 8 10 10 8

  48. B. The Periodic Table:A VERY Brief Introduction • By arranging the elements in order of increasing atomic number, one can observe the periodicity of the properties of the elements. • The periodic table is a tabular arrangement of elements in rows (periods) and columns (groups), highlighting the regular repetition of the properties of the elements.

  49. a) Atomic Symbol • Each element in the periodic table is represented by an atomic symbol. Atomic Symbol: One- or two-letter symbol used to represent an atom corresponding to a particular element (Symbols frequently contain the first letter of the element, capitalized, sometimes with an additional letter. Careful, some symbols are derived from the name of the element in another language!)

  50. B. The Periodic Table • b) Groups/Periods: • The periodic table organizes elements in such a way that elements with similar properties occur invertical columns called • groups. • The horizontal rows • in the table are called • periods.