Atomic Theory

1. Atomic Theory 2.1.1 State the position of protons, electrons and neutrons in the atom 2.1.2 State the relative masses and relative charges of protons, neutrons and electrons 2.1.3 Define the terms mass number (A), atomic number (Z) and isotopes of an element 2.1.4 Deduce the symbol for an isotope given its mass number and atomic number 2.1.5 Calculate the number of protons, neutrons and electrons in atoms and ions from the mass number, atomic number and charge. 2.1.6 Compare the properties of the isotopes of an element 2.1.7 Discuss the uses of radioisotopes

2. History of the atom Democritus (400 BC) suggested that the material world was made up of tiny, indivisible particles atomos, Greek for “uncuttable” Aristotle believed that all matter was made up of 4 elements, combined in different proportions Fire - Hot Earth - Cool, heavy Water - Wet Air - Light The “atomic” view of matter faded for centuries, until early scientists attempted to explain the properties of gases

3. Re-emergence of Atomic Theory John Dalton postulated that: • All matter is composed of extremely small, indivisible particles called atoms • All atoms of a given element are identical (same properties); the atoms of different elements are different

4. 3. Atoms are neither created nor destroyed in chemical reactions, only rearranged 4. Compounds are formed when atoms of more than one element combine • A given compound always has the same relative number and kind of atoms

5. Atoms are divisible! • By the 1850s, scientists began to realize that the atom was made up of subatomic particles • Thought to be positive and negative • How would we know this if we can’t see it or touch it?

6. Cathode Rays and Electrons Mid-1800’s scientists began to study electrical discharge through cathode-ray tubes. Ex: neon signs Partially evacuated tube in which a current passes through Forms a beam of electrons which move from cathode to anode Electrons themselves can’t be seen, but certain materials fluoresce (give off light) when energised

7. JJ Thompson observed that when a magnetic or electric field are placed near the electron beam, they influence the direction of flow opposite charges attract each other, and like charges repel. The beam is negatively charged so it was repelled by the negative end of the magnet Oh there you are!

8. http://www.chem.uiuc.edu/clcwebsite/video/Cath.mov • Magnetic field forces the beam to bend depending on orientation • Thompson concluded that: • Cathode rays consist of beams of particles • The particles have a negative charge

9. Thompson understood that all matter was inherently neutral, so there must be a counter A positively charged particle, but where to put it It was suggested that the negative charges were balanced by a positive umbrella-charge “Plum pudding model” “chocolate chip cookie model”

10. Rutherford and the Nucleus • This theory was replaced with another, more modern one • Ernest Rutherford (1910) studied angles at which a particles (nucleus of helium) were scattered as they passed through a thin gold foil • http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/ruther14.swf

12. Rutherford expected … Rutherford believed that the mass and positive charge was evenly distributed throughout the atom, allowing the a particles to pass through unhindered a particles

13. + Rutherford explained … Atom is mostly empty space Small, dense, and positive at the center Alpha particles were deflected if they got close enough a particles

14. The modern atom is composed of two regions: • Nucleus: Containing protons and neutrons, it is the bulk of the atom and has a positive charge associated with it • Electron cloud: Responsible for the majority of the volume of the atom, it is here that the electrons can be found orbiting the nucleus (extranuclear)

15. Major Subatomic Particles Atoms are measured in picometers, 10-12 meters Hydrogen atom, 32 pm radius Nucleus tiny compared to atom If the atom were a stadium, the nucleus would be a marble Radius of the nucleus is on the order of 10-15 m Density within the atom is near 1014 g/cm3

16. Elemental Classification Atomic Number (Z) = number of protons (p+) in the nucleus Determines the type of atom Li atoms always have 3 protons in the nucleus, Hg always 80 Mass Number (A) = number of protons + neutrons [Sum of p+ and nº] Electrons have a negligible contribution to overall mass In a neutral atom there is the same number of electrons (e-) and protons (atomic number)

17. E mass number A elemental symbol Z atomic number Nuclear Symbols Every element is given a corresponding symbol which is composed of 1 or 2 letters (first letter upper case, second lower), as well as the mass number and atomic number

18. 184 W 74 19 F 9 80 Br 35 Find the number of protons number of neutrons number of electrons atomic number mass number

19. Ions • Cation is a positively charged particle. Electrons have been removed from the element to form the + charge. ex: Na has 11 e-, Na+ has 10 e- • Anion is a negatively charged particle. Electrons have been added to the atom to form the – charge. ex: F has 9 e-, F- has 10 e-

20. 1 2 3 H H H 1 1 1 Hydrogen-1 Hydrogen-2 Hydrogen-3 Isotopes Atoms of the same element can have different numbers of neutrons and therefore have different mass numbers The atoms of the same element that differ in the number of neutrons are called isotopes of that element When naming, write the mass number after the name of the element

21. How heavy is an atom of oxygen? There are different kinds of oxygen atoms (different isotopes) 16O, 17O, 18O We are more concerned with average atomic masses, rather than exact ones Based on abundance of each isotope found in nature We can’t use grams as the unit of measure because the numbers would be too small Instead we use Atomic Mass Units (u) Standard u is 1/12 the mass of a carbon-12 atom Each isotope has its own atomic mass

22. Calculating Averages Average = (% as decimal) x (mass1) + (% as decimal) x (mass2) + (% as decimal) x (mass3) + … Problem: Silver has two naturally occurring isotopes, 107Ag with a mass of 106.90509 u and abundance of 51.84 % ,and 109Ag with a mass of 108.90476 u and abundance of 48.16 % What is the average atomic mass? Average = (0.5184)(106.90509 u) + (0.4816)(108.90476 u) = 107.87 amu

23. Average Atomic Masses • If not told otherwise, the mass of the isotope is the mass number in amu • The average atomic masses are not whole numbers because they are an average mass value • Remember, the atomic masses are the decimal numbers on the periodic table

24. Properties of Isotopes • Chemical properties are primarily determined by the number of electrons • All isotopes has the same number of electrons, so they have nearly identical chemical properties even though they have different masses. • Physical properties often depend on the mass of the particle, so among isotopes they will have slightly different physical properties such as density, rate of diffusion, boiling point… • The isotopes of an element with fewer neutrons will have: • Lower masses • faster rate of diffusion • Lower densities • lower melting and boiling points

25. More Practice Calculating Averages • Calculate the atomic mass of copper if copper has two isotopes • 69.1% has a mass of 62.93 amu • The rest (30.9%) has a mass of 64.93 amu • Magnesium has three isotopes • 78.99% magnesium 24 with a mass of 23.9850 amu • 10.00% magnesium 25 with a mass of 24.9858 amu • The rest magnesium 26 with a mass of 25.9826 amu • What is the atomic mass of magnesium?

26. Radioisotopes • Isotopes of atoms that have had an extra neutron attached to their nucleus. • Carbon-14 radioactive decay is used to measures the date of objects. • After 5700 years the amount of 14C will be half its original value. • Iodine-125 or 131 is used to monitor the activity of the thyroid gland (b/c the thyroid tends to absorb iodine)

27. Cobalt-60 produces gamma rays (intense radioactivity) and is used in radiation treatment of cancer. • Note: gamma rays are the shortest wavelength on the electromagnetic spectrum. They are the most dangerous and difficult to shield from.

28. 2.2 The Mass Spectrometer

29. Mass Spectrometer • The mass spectrometer is an instrument used: • To measure the relative masses of isotopes • To find the relative abundance of the isotopes in a sample of an element When charged particles pass through a magnetic field, the particles are deflected by the magnetic field, and the amount of deflection depends upon the mass/charge ratio of the charged particle.

30. Mass Spectrometer – 5 Stages • Once the sample of an element has been placed in the mass spectrometer, it undergoes five stages. • Vaporisation – the sample has to be in gaseous form. If the sample is a solid or liquid, a heater is used to vaporise some of the sample. X (s) X (g) or X (l)  X (g)

31. Mass Spectrometer – 5 Stages • Ionization – sample is bombarded by a stream of high-energy electrons from an electron gun, which ‘knock’ an electron from an atom. This produces a positive ion: X (g) X +(g) + e- • Acceleration – an electric field is used to accelerate the positive ions towards the magnetic field. The accelerated ions are focused and passed through a slit: this produces a narrow beam of ions.

32. Mass Spectrometer – 5 Stages • Deflection – The accelerated ions are deflected into the magnetic field. The amount of deflection is greater when: • the mass of the positive ion is less • the charge on the positive ion is greater • the velocity of the positive ion is less • the strength of the magnetic field is greater

33. Mass Spectrometer • If all the ions are travelling at the same velocity and carry the same charge, the amount of deflection in a given magnetic field depends upon the mass of the ion. • For a given magnetic field, only ions with a particular relative mass (m) to charge (z) ration – the m/z value– are deflected sufficiently to reach the detector.

34. Mass Spectrometer • Detection – ions that reach the detector cause electrons to be released in an ion-current detector • The number of electrons released, hence the current produced is proportional to the number of ions striking the detector. • The detector is linked to an amplifier and then to a recorder: this converts the current into a peak which is shown in the mass spectrum.

35. Atomic Structure – Mass Spectrometer • Isotopes of boron Ar of boron = (11 x 18.7) + (10 x 81.3) (18.7 + 81.3) = 205.7 + 813 100 = 1018.7 = 10.2 100

36. Mass Spectrometer – Questions • A mass spec chart for a sample of neon shows that it contains: • 90.9% 20Ne • 0.17% 21Ne • 8.93% 22Ne Calculate the relative atomic mass of neon You must show all your work!

37. (90.9 x 20u) + (0.17 x 21u) + (8.93 x 22u) • 100 Mass Spectrometer – Questions • 90.9% 20Ne • 0.17% 21Ne • 8.93% 22Ne Ar= 20.18u

38. 52.3 23.6 22.6 1.5 m/e 204 206 207 208 Mass Spectrometer – Questions Calculate the relative atomic mass of lead You must show all your work!

39. 20724.2 • 100 • (1.5 x 204) + (23.6 x 206) + (22.6 x 207)+(52.3 x 208) • 100 • 306 + 4861.6 + 4678.2 + 10878.4 • 100 Mass Spectrometer – Questions • 1.5% 204Pb • 23.6% 206Pb • 22.6% 207Pb • 52.3% 208Pb = Ar= 207.24

40. 2.3 Electron Arrangement 2.3.1 Describe the electromagnetic spectrum 2.3.2 Distinguish between a continuous spectrum and a line spectrum 2.3.3 Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels 2.3.4 Deduce the electron arrangement for atoms and ions up to Z=20

41. Electromagnetic radiation.

42. Electromagnetic Radiation • Most subatomic particles behave as PARTICLES and obey the physics of waves.

43. wavelength Visible light Amplitude wavelength Node Ultaviolet radiation Electromagnetic Radiation

44. Wavelengths and energy • Understand that different wavelengths of electromagnetic radiation have different energies. • Waves have a frequency • c=νλ • c=velocity of wave (2.998 x 108 m/s) • ν=(nu) frequency of wave, units are “cycles per sec” • λ=(lambda) wavelength

45. Electromagnetic Spectrum In increasing energy, ROYGBIV

46. increasing frequency increasing wavelength Electromagnetic Spectrum Long wavelength --> small frequency Short wavelength --> high frequency

47. Bohr’s Model • Why don’t the electrons fall into the nucleus? • Move like planets around the sun. • In circular orbits at different levels. • Amounts of energy separate one level from another.

48. Bohr postulated that: • Fixed energy related to the orbit • Electrons cannot exist between orbits • The higher the energy level, the further it is away from the nucleus • An atom with maximum number of electrons in the outermost orbital energy level is stable (unreactive) • Think of Noble gases

49. Those who are not shocked when they first come across quantum theory cannot possibly have understood it.(Niels Bohr on Quantum Physics)

50. Atomic Line Emission Spectra and Niels Bohr Bohr’s greatest contribution to science was in building a simple model of the atom. It was based on an understanding of theLINE EMISSION SPECTRAof excited atoms. • Problem is that the model only works for Hydrogen Niels Bohr (1885-1962)