1. The Atom and The Periodic Table

2. James Chadwick 1932 James Chadwick British Found an electrically neutral particle which resides in the nucleus and has almost the same mass as the proton. He named this “neutral proton” the neutron.

3. Discovery of the Neutron James Chadwick bombarded beryllium-9 with alpha particles, carbon-12 atoms were formed, and neutrons were emitted. + +

4. Atom: smallest particle of an element that retains the chemical properties of the element. -can be broken down, however they would lose their chemical identity The Atom

5. One of Dalton’s postulates said that, “All atoms of the same element are identical.” (i.e. same mass and properties) IS THIS TRUE?NO!!! Example: Boron is mined in Death Valley, CA. There are two “types” of Boron. Both have 5 p + but one has 5 n o while the other has 6n o. They are exactly alike chemically, but different in mass.

6. ISOTOPES Isotopes: same number of p +, but different number of n o -most elements have 2 stable isotopes -EXCEPTIONS who only have 1 = Al, F, P -Sn has 10!!!! -Refer to an isotope by its mass number (p + + n o ) for example: uranium-238 or 238 U

7. The number of p + = ATOMIC NUMBER represented by “Z” if Z = 5, then the element is Boron In the atom, what is the same as the number of p + ? WHY? p + = e - Atoms are electrically neutral. Hence, the number of p + determines the identity of the element and the number of n o determines the isotope of the element.    Atomic Number (Z) Nucleus - Proton (+1) Neutrons (0) If two elements have the same atomic number they are the same element. If two elements have different atomic numbers they are different.

8. Dalton’s postulate now reads: “All atoms of an element contain the same number of p + but can contain different numbers of n o.” A particular type of atom is called a NUCLIDE (another name for isotope). Protium, Deuterium, and Tritium are all nuclides of hydrogen. Particles that make up the nucleus are called NUCLEONS. The proton and neutron are nucleons.

9. # protons + # neutrons mass number The total number of nucleons = p + + n o = mass number (“M”) Isotopic notation: X M Z Isotopic Notation X = chemical symbol M = mass # Z = atomic #

10. Isotopic Notation

12. Symbols Find the –number of protons –number of neutrons –number of electrons –Atomic number –Mass number Br 80 35 = 35 = 45 = 35 = 80

13. Symbols If an element has 60 protons and 84 neutrons what is the –Atomic number –Mass number –number of electrons –Complete symbol Nd 144 60 = 60 = 144 = 60

14. Symbols If a neutral atom of an element has 78 electrons and 117 neutrons what is the –Atomic number –Mass number –number of protons –Complete symbol Pt 195 78 = 78 = 195 = 78

15. Symbols Find the number of protons number of neutrons number of electrons Atomic number Mass number Na 23 11 1+ Sodium ion = 11 = 12 = 10 = 11 = 23

16. Symbols If an element has an atomic number of 23 and a mass number of 51 what is the –number of protons –number of neutrons –number of electrons V 51 23 = 23 = 28 = 23

17.    Metals lose electrons to form positive ions (cations): Li ---> Li + + e - 3 p + 3 p + 3 e - 2 e - 4 n 0 4 n 0    Nucleus - Proton (+1) Neutrons (0) If # of protons > # of electrons it has a positive charge and we call it a cation.

18. Nonmetals gain electrons to form negative ions (anions): F + e - ---> F - 9 p + 9 p + 9 e - 10 e - 10 n 0 10 n 0  Nucleus - Proton (+1) Neutrons (0)                  If # of protons < # of electrons it has a negative charge and we call it an anion.

19. Protons Isotopic Notation 92 NeutronsElectrons 34 11 146 45 12 92 36 10 59 27 3+ Co 37 17 1– Cl 557+ Mn 25 238 92 U 23 11 1+ Na 79 34 2– Se 201817 301825 322427

20. The p + and n o are essentially equal in mass. As you can see, the electron has very little mass when compared to both the proton and the neutron. 99.95% of an atom’s mass is found in the nucleus. Protons are over 1800 times larger than electrons. So, chemists say electrons have no mass. This is not exactly true, it’s more like they have negligible mass. Atomic Mass

21. Subatomic particles Electron Proton Neutron NameSymbolCharge Relative mass Actual mass (g) e-e- p+p+ nono +1 0 1/1840 = 0 1 1 9.109389 x 10 -28 1.6762623 x 10 -24 1.6749286 x 10 -24

22. The mass of atoms is measured in amu or (u). Amu = 1/12 the mass of the carbon-12 nuclide 1 u = 1.66054 x 10 -24 g Carbon-12 is the standard. One C-12 atom has a mass of 12 u. The mass of the other elements is relative to this mass. carbon atom (12 amu) (1 amu)

23. For example….Methane For carbon 1 in approximately 90 atoms are carbon-13 The rest are carbon-12 the isotope that is 98.9% abundant. So, for approximately 90 methane molecules…1 carbon is carbon-13

24. Where’s Waldo? C-13

25. Subatomic Particles Quarks component of protons & neutrons 6 types –3 quarks = 1 proton or 1 neutron He

26. A mass spectrometer

27. Mass Spectrometry A mass spectrometer is a device that separates positive gaseous ions according to their mass-to- charge ratios. If a stream of positive ions having equal velocities is brought into a magnetic field: All the ions are deflected from their straight line paths into circular paths The lightest ions are deflected the most making a tighter circle Conversely, the heaviest ions are deflected the least A record of the separation of ions is called a mass spectrum.

28. Diagram of a simple mass spectrometer, showing the separation of neon isotopes

29. The mass spectrum of neon

30. Mass Spectrometry - - Photographic plate 196 199 201 204 198 200 202 Mass spectrum of mercury vapor Stream of positive ions +

31. Mass Spectrum for Mercury 196 197 198 199 200 201 202 203 204 Mass number Relative number of atoms 30 25 20 15 10 5 196 199 201 204 198 200 202 Mass spectrum of mercury vapor The percent natural abundances for mercury isotopes are: Hg-196 0.146% Hg-196 0.146% Hg-198 10.02% Hg-198 10.02% Hg-199 16.84% Hg-199 16.84% Hg-200 23.13% Hg-200 23.13% Hg-201 13.22% Hg-201 13.22% Hg-202 29.80% Hg-202 29.80% Hg-204 6.85% Hg-204 6.85% (The photographic record has been converted to a scale of relative number of atoms)

32. Mass spectrums reflect the abundance of naturally occurring isotopes. Hydrogen Carbon Nitrogen Oxygen Sulfur Chlorine Bromine 1 H = 99.985% 2 H = 0.015% 12 C = 98.90% 13 C = 1.10% 14 N = 99.63% 15 N = 0.37% 16 O = 99.762% 17 O = 0.038% 18 O = 0.200% 32 S = 95.02% 33 S = 0.75% 34 S = 4.21% 36 S = 0.02% 35 Cl = 75.77% 37 Cl = 24.23% 79 Br = 50.69% 81 Br = 49.31% Natural Abundance of Common Elements

33. Atomic Mass How heavy is an atom of oxygen? There are different kinds of oxygen atoms. More concerned with average atomic mass. Based on abundance of each element in nature. Don’t use grams because the numbers would be too small

34. Abundance = Percent (% ) = (part/whole) = mass individual /mass whole Average = Abundance 1 (mass 1 ) + Abundance 2 (mass 2 ) + etc AVERAGE ATOMIC MASS

35. Steps: Multiply the mass of each isotope by its Abundance. Add up all of the products from step 1.

36. (.2000)(10.0129)+(.8000)(11.00931)= 2.00250+8.807448=10.810028 =10.81003 u Isotope% AbundanceAtomic mass B-1020.00 %10.0129 u B-1180.00 %11.00931 u Example:

37. Mass Spectrum for Mercury 196 197 198 199 200 201 202 203 204 Mass number Relative number of atoms 30 25 20 15 10 5 196 199 201 204 198 200 202 Mass spectrum of mercury vapor The percent natural abundances for mercury isotopes are: Hg-196 0.146% Hg-196 0.146% Hg-198 10.02% Hg-198 10.02% Hg-199 16.84% Hg-199 16.84% Hg-200 23.13% Hg-200 23.13% Hg-201 13.22% Hg-201 13.22% Hg-202 29.80% Hg-202 29.80% Hg-204 6.85% Hg-204 6.85% (The photographic record has been converted to a scale of relative number of atoms)

38. The percent natural abundances for mercury isotopes are: Hg-196 0.146% Hg-196 0.146% Hg-198 10.02% Hg-198 10.02% Hg-199 16.84% Hg-199 16.84% Hg-200 23.13% Hg-200 23.13% Hg-201 13.22% Hg-201 13.22% Hg-202 29.80% Hg-202 29.80% Hg-204 6.85% Hg-204 6.85% (0.00146)(196) + (0.1002)(198) + (0.1684)(199) + (0.2313)(200) + (0.1322)(201) + (0.2980)(202) + (0.0685)(204) = x 0.28616 + 19.8396 + 33.5116 + 46.2600 + 26.5722 + 60.1960 + 13.974 = x x = 200.63956 amu Hg 200.6 80 (% "A")(mass "A") + (% "B")(mass "B") + (% "C")(mass "C") + (% "D")(mass "D") + (% "E")(mass "E") + (% F)(mass F) + (% G)(mass G) = AAM ABCDEFGABCDEFG

39. Example Chlorine has two naturally occurring isotopes, 35 Cl (34.9689 amu) and 37 Cl (36.9659 amu). If chlorine has an atomic mass of 35.4527 amu, what is the percent abundance of each chlorine isotope? Isotope% AbundanceAtomic mass Cl-35x34.9689 u Cl -371 – x36.9659 u Average Atomic mass35.4527

40. Isotope% AbundanceAtomic mass Cl-35x34.9689 u Cl -371 – x36.9659 u Average Atomic mass35.4527 x(34.9689)+(1 - x)(36.9659)=35.4527 34.9689 x+36.9659 - 36.9659 x=35.4527 36.9659-1.9970 x=35.4527 1.5132=1.9970 x.7577366049=x Abundance 1 (mass 1 ) + Abundance 2 (mass 2 ) + etc = Average Atomic mass