2. V.Montgomery & R.Smith1 Dalton’s Atomic Theory (1808) 1.All matter is made of tiny indivisible particles called atoms. 2.Atoms of the same element are identical. The atoms of any one element are different from those of any other element. 3.“Atoms are like billiard balls, spherical and solid inside.”

3. V.Montgomery & R.Smith2 Dalton’s Atomic Theory (1808) 4.Atoms of different elements can combine with one another in simple whole number ratios to form compounds. 5.Chemical reactions occur when atoms are separated, joined, or rearranged;however, atoms of one element are not changed into atoms of another by a chemical reaction.

4. Summary of the Atom atoms are the smallest particles that can be uniquely associated with an element each element has unique atoms atoms are composed of e -, p and n atoms are electrically neutral (# of e - = # of p) for a single element, isotopes differ only in number of n (neutrons) atoms have characteristic masses (atomic weights) atoms combine with one another in definite, whole number proportions to make compounds

5. ~ 10 -10 m nucleus Mass 9 x 10 -31 kg Mass > 10 -26 kg electron

6. ~ 1 – 7 x 10 -15 m (1 – 7 fermi)

8. atoms are dominantly empty space: If an oxygen atom had a total radius of 100 km, the nucleus would be a ~1 m diameter sphere in the middle. electron orbits The Spacious Atom

9. In a simplistic model, electrons float around the nucleus in energy levels called shells. As the number of electrons increases, they start to fill shells farther out from the nucleus. In most cases, electrons are lost or gained only from the outermost shell. electron orbits Electrons in Orbit

10. Atom Nucleus

11. The Nuclear Model of the atom

12. Mass of an electron is approximately 1/1840 th of a proton or neutron. Mass of a neutron is very close to the mass of a proton. 1 atomic mass unit (amu or u)= 1.66054x10 -24 g 1 g= 6.022 x10 23 amu 1amu is defined as 1/12 th the mass of an atom of carbon-12.

13. Subatomic Particles ParticleSymbol Charge Relative Mass Electron e - 1- 0 Proton p + 1+ 1 Neutron n0 1

14. Atomic Number(nuclear charge) 11 Na Atomic Number Symbol

15. All atoms of an element have the same number of protons 11 Na 11 protons Sodium

16. Number of Electrons If an atom is neutral ; The net charge is zero Number of protons = Number of electrons Atomic number = Number of electrons in a neutral atom

17. Ions Have a net electrical charge since the total number of electrons isn’t equal to the number of protons. Can be anions (e - >p + ;formed as a result of gaining electrons; negatively charged), cations (p + >e - ;formed as a result of losing electrons; positively charged). e - + q(charge of the ion) = proton number

18. In chemical rxns, atoms never gain or lose protons. It’s the interaction of electrons.

19. Mass Number(nucleon number) Counts the number of protons and neutrons in an atom A= p + + n 0

20. Atomic Symbols Show the mass number and atomic number Give the symbol of the element 23 Na 11 (Z)atomic number (A)mass number sodium-23

21. Notation for Atoms 12 C 13 C C only one isotope of carbon all isotopes of carbon only one isotope of carbon

22. Basic Definitions “atomic number” = number of protons in the nucleus; “mass number” = sum of protons + neutrons in the nucleus “isotopic mass” = mass of a single isotope

23. More Atomic Symbols 163165 O P Zn 81530 8 p + 15 p + 30 p + 8 n16 n35 n 8 e - 15 e - 30 e -

24. Isotopes Atoms with the same number of protons, but different numbers of neutrons. Atoms of the same element (same atomic number) with different mass numbers Isotopes of chlorine 35 Cl 37 Cl17 chlorine - 35 chlorine - 37

25. Since both isotopes have the same number of protons and electrons, they have identical chemical properties. Since physical properties depend on the mass of particles as well, isotopes will often have slightly different physical properties such as density, mass, rate of diffusion etc.

26. Natural abundances of isotopes Natural chlorine contains : 75 % 35 Cland 25 % 37 Cl. 17 17 These percentages are known as the natural abundances of the isotopes and determined by “mass spectrometry.”

27. Relative atomic mass Is the weighted mean. Relative atomic mass has no unit since the atomic mass unit, amu is cancelled out in calculation w/ respect to C-12.

28. Exercise 1 The molar mass of iridium is 192.2 g / mol. What are the naturally occurring percentages of the two isotopes of Ir-191 and Ir-193?

29. solution Iridium is a mixture of 40% 191 Ir and 60 % 193 Ir.

30. Isotopes of Hydrogen element has the biggest abundance in nature. Tritium Deuterium

31. Radioactive isotopes Are produced by exposing the natural element to a flux(flow) of neutrons in a nuclear reactor. The nucleus of an atom captures an additional neutron and forms radioisotope.

32. Usages of radioisotopes 1.The rate of radioactive decay is used to date objects (C-14).  Naturally occurring C has a fixed proportion of C-14 due to exchange w/ C in the atmosphere. When the plant is dead, the exchange stops & the proportion of C- 14 starts to decrease in the plant due to radioactive decay. This decayed amount is used to date the plant. After about 5,700 yrs, the proportion of C-14 falls to about half its initial value.

33. 2. as tracers.  Radioactive isotope reacts chemically & biologically. For example the activity of thyroid gland can be measured w/ monitoring the increase in radioactivity of the gland after taking a drink including iodine radioisotopes (I-125 and I-131) since the thyroid gland absorbs the radioactive iodine when it works. Usages of radioisotopes

34. 3. Source of gamma rays and therefore, source intense radioactivity.  Cobalt-60 is an example of such a radioactivity source. It’s used in radiation treatment for cancer and industrially as well Usages of radioisotopes

35. A mass spectrometer is an instrument which separates particles according to their masses, records the relative proportions of these, and determine natural abundances of the isotopes of an element. Therefore, it also allows us to calculate the atomic mass of an element. The most accurate way for determining atomic and molecular weights is provided by mass spectrometer. Mass spectrometry

36. mass spectrometer, invented by the English physicist Francis William Aston (1877-1945) when he was working in Cambridge with J. J. Thomson. It was in his use of this instrument that the existence of isotopes of elements was discovered. Aston eventually discovered many of the naturally occurring isotopes of non-radioactive elements. He was awarded the Nobel Prize for Chemistry in 1922. Mass spectrometry

37. A B C D E accelerating F

38. A: a gaseous sample is very slowly introduced to the mass spectrometer. B: atoms/molecules are bombarded by a stream of high energy electrons to produce positive ions, mostly w/ a 1+ charge. These electrons collide w/ electrons in the particle knocking them out and leaving a positive ion. C: positively charged ions are accelerated high enough to make the particles pass through the slits and magnetic field by high electrical voltage on the negatively charged grid. With the slits, the ions were made a beam of ions.

39. D: Fast moving ions enter a magnetic field produced by an electromagnet. Ions are deflected by a magnetic field into a curved path. The deflection of the ions depends on “charge to mass ratio(q/m). ”The more massive the ion, the less the deflection. The ions w/ equal mass and charge will deflect the same. E: By changing the strength of the magnetic field or the accelerating voltage on the negatively charged grid, ions of varying masses can be made to enter the detector at the end of the instrument.

40. On the detector, ions are collected on a metal plate and the current flows through the metal plate to neutralise the ions and this current is recorded. In this way, the relative abundances of ions of different masses in the sample can be determined and put into a graph called “mass spectrum.”

41. F: The mass spectrometer must be at a high vacuum for its correct operation and its correct operation depends on particles being able to pass through it w/o colliding with any other particles.

42. A: vapourised sample introduced B: ionization by electron bombardment C: Positive ions accelerated by electrical field D: ions deflected by a magnetic field E: detector records ions of a particular mass F: vacuum prevents molecules colliding

43. Mass spectrometer is used to identify the molecular structure of a compound or analyze mixtures of substances. When a molecule loses an electron, it falls apart, forming fragments. Mass spectrometer measures the mass of these fragments, producing a chemical “fingerprint” of the molecule and providing clues about how the atoms were connected together in the molecule.

44. Atomic weight measurements How was the atomic weight measured? By mass spectrometry –This also measures % natural abundance for a given isotope -The graph is called as “mass spectrum.”

45. P.53  check out the graph.

46. Atomic weight calculation There are three naturally occuring isotopes of neon (Ne): 20 Ne isotopic mass = 19.99244018 amu 21 Ne isotopic mass = 20.9938467 amu 22 Ne isotopic mass = 21.9913855 amu the atomic weight is reported in text as: 20.1797 amu

47. Learning Check 1 Naturally occurring carbon consists of three isotopes, 12 C, 13 C, and 14 C. State the number of protons, neutrons, and electrons in each of these carbon atoms. 12 C 13 C 14 C 6 6 6 #P _______ _______ _______ #N _______ _______ _______ #E _______ _______ _______

48. Solution 12 C 13 C 14 C 6 6 6 #P __6___ _ 6___ ___6___ #N __6___ _ _7___ ___8___ #E __6___ _ 6___ ___6___

49. Learning Check 2 An atom of zinc has a mass number of 65. A.Number of protons in the zinc atom 1) 302) 353) 65 B.Number of neutrons in the zinc atom 1) 302) 353) 65 C. What is the mass number of a zinc isotope with 37 neutrons? 1) 372) 653) 67

50. Solution An atom of zinc has a mass number of 65. A.Number of protons in the zinc atom 1) 30 B.Number of neutrons in the zinc atom 2) 35 C. What is the mass number of a zinc isotope with 37 neutrons? 3) 67

51. Learning Check 3 Write the atomic symbols for atoms with the following: A. 8 p +, 8 n, 8 e - ___________ B.17p +, 20n, 17e - ___________ C. 47p +, 60 n, 47 e - ___________

52. Solution 16 O A. 8 p +, 8 n, 8 e - 8 B.17p +, 20n, 17e - 37 Cl 17 C. 47p +, 60 n, 47 e - 107 Ag 47

53. Learning Check 4 An atom has 14 protons and 20 neutrons. A.Its atomic number is 1) 142) 163) 34 B. Its mass number is 1) 142) 163) 34 C. The element is 1) Si2) Ca3) Se D.Another isotope of this element is 1) 34 X 2) 34 X 3) 36 X 16 14 14

54. Solution An atom has 14 protons and 20 neutrons. A.It has atomic number 1) 14 B. It has a mass number of 3) 34 C. The element is 1) Si D.Another isotope of this element would be 3) 36 X 14

55. Masses of Atoms A scale designed for atoms gives their small atomic masses in atomic mass units (amu) An atom of 12 C was assigned an exact mass of 12.00 amu Relative masses of all other atoms was determined by comparing each to the mass of 12 C An atom twice as heavy has a mass of 24.00 amu. An atom half as heavy is 6.00 amu.

56. Average atomic mass(atomic weight) “atomic weight or mass” = average mass of an atom calculated from the masses and natural abundances of all isotopes (use atomic weights to calculate the molecular weights of compounds from their constituent elements!)

57. Atomic Mass Average atomic mass is based on all the isotopes and their abundance % Atomic mass is not a whole number Na 22.99

58. Calculating Atomic Weight or Mass Percent(%) abundance of isotopes Mass of each isotope of that element Weighted average = mass isotope 1 (%) + mass isotope 2 (%) + … 100 100

59. Naturally occurring C is composed of 98.93 % 12 C and 1.07 % 13 C. The masses of these nuclides are 12 amu (exactly) and 13.00335 amu, respectively. Average atomic mass(atomic mass) or atomic weight

60. Atomic Mass of Magnesium Isotopes Mass of Isotope Abundance 24 Mg =24.0 amu 78.70% 25 Mg = 25.0 amu 10.13% 26 Mg = 26.0 amu 11.17% Atomic mass (average mass) Mg = 24.3 amu Mg 24.3

61. Atomic mass calculation How was the atomic mass calculated? multiply each isotopic mass by the reported natural abundance for the isotope, then: add these individual contributions for each isotope to get the average atomic mass for the element

62. Atomic mass calculation There are three naturally occuring isotopes of neon (Ne): 20 Ne mass # = 19.99244018 amu (90.51%) 21 Ne mass # = 20.9938467 amu (0.27%) 22 Ne mass # = 21.9913855 amu (9.22%) the atomic mass is reported in text as: 20.1797 amu 18.10 + 0.057 + 2.03 = 20.19 amu

63. Learning Check 5 Gallium is a metallic element found in small lasers used in compact disc players. In a sample of gallium, there is 60.2% of gallium-69 (68.9 amu) atoms and 39.8% of gallium-71 (70.9 amu) atoms. What is the atomic mass of gallium ?

64. Solution Ga-69 68.9 amu x 60.2 = 41.5 amu for 69 Ga 100 Ga-71 (%/100) 70.9 amu x 39.8 = 28.2 amu for 71 Ga 100 Atomic mass Ga = 69.7 amu

65. Finding An Isotopic Mass A sample of boron consists of 10 B (mass 10.0 amu) and 11 B (mass 11.0 amu). If the average atomic mass of B is 10.8 amu, what is the % abundance of each boron isotope?

66. Assign X and Y values: X = % 10 B Y = % 11 B Determine Y in terms of X X + Y = 100 Y = 100 - X Solve for X: X (10.0) + (100 - X )(11.0) = 10.8 100 100 Multiply through by 100 10.0 X + 1100 - 11.0X = 1080

67. Collect X terms 10.0 X - 11.0 X = 1080 - 1100 - 1.0 X = -20 X = -20 = 20 % 10 B - 1.0 Y = 100 - X % 11 B = 100 - 20% = 80% 11 B

68. Learning Check 6 Copper has two isotopes 63 Cu (62.9 amu) and 65 Cu (64.9 amu). What is the % abundance of each isotope? (Hint: Check Zumdahl or any other chemistry text for atomic mass) 1) 30%2) 70%3) 100%

69. Solution 2) 70% Solution 62.9X + 6490 = 64.9X = 6350 -2.0 X = -140 X = 70%

70. Atomic Masses 13 C 12 C 13.00335 amu (1.11%) 12.0000 amu (98.89%) atomic weight of C = 12.01115 amu WHY?

71. Calculating masses of atoms relative to 12 C (mass of 12 C atom) * 1.58320 = mass of F atom = 18.99840 reported atomic weight of F = 18.9984

72. Charged Atoms: Ions Left to their own devices, atoms are electrically neutral. That means that they have an equal number of protons and electrons. During the course of most natural events, protons are not gained or lost, but electrons may be. Atoms with more or fewer electrons than protons are electrically charged. They are called ions: an atom that loses electrons takes on a positive charge (cation); an atom that gains electrons takes on a negative charge (anion). Complex cations and anions can also occur: (NH 4 ) +1, (SO 4 ) -2

73. An ISOTOPE is one of a set of nuclides with the same Z and consequently different A. (ie isotopes are the same chemical element but different masses). e.g. An ISOBAR is one of a set of nuclides with the same A but different N and Z. e.g An ISOTONE is one of a set of nuclides with the same N and consequently different A. e.g.

74. More on atomic notation, which is based on the nuclear structure: –Isotope: same Z, different A and N –Isobar: same A, different Z and N –Isotone: same N, different Z and A Example: From the following list of atoms, which are isotopes, isobars, and isotones? Component Atom AZN Xe I Cs I