Principles of Chemistry, Chapt . 2: Atomic Structure and The Elements The Structure of Atoms

1. Principles of Chemistry, Chapt. 2: Atomic Structure and The Elements • The Structure of Atoms • protons, neutrons, and electrons • Atomic Structure and Properties—the Elements • atomic mass, atomic number, isotopes • The Mole Concept: 6.02 x 1023 • The Periodic Table Homework: Chapt. 2 Problems 26, 29, 37, 43, 75

2. Atomic Theory in a single Slide ~ 10-10 meters = 1 angstrom (Å) Smeared out electron charge cloud _ Most of the Chemistry is here 10-14 m Most of the mass is here + Protons and neutrons + + + + + + +

3. STM Image: Oxygen atoms at the surface of Al2O3/Ni3Al(111) Electronic charge cloud surrounding the nucleus S. Addepalli, et al. Surf. Sci. 442 (1999) 3464

4. What’s inside the nucleus: ParticleMass (amu)Charge Proton (p+) 1.007 amu +1 Neutron (n0) 1.009 amu 0 What’s outside the nucleus: Electron (e-) .00055 amu -1 Note: mass ratio of electron/proton (Mp+/Me-) = 1836 For any atom: # of electrons = # of protons: Why?

5. Atomic Theory: Late 19th Century Atomic theory—everything is made of atoms—generally accepted (thanks to Ludwig Boltzman, and others). Mendeleev/periodic table—accepted, but the basis for periodic behavior not understood What are atoms made of? How are they held together?

6. Radioactive material β-particles (“–”) + – Beam of , , and  Gamma ray (γ) No charge, no deflection Electrically charged plates α-particle (“+” ) Heavier, deflected less than β Radioactivity Electrical behavior: “+” attracts “-” but like charges repel Atoms must contain smaller sub-units. Alpha particle  2 n0 + 2p+ Beta particel  electron (e-) Gamma photon

7. – high voltage + Electrons fluorescent screen • Thomson (1897) discovered the e-: cathode ray • “Cathode rays” • Travel from cathode (-) to anode (+). • Negative charge (e−). • Emitted by cathode metal atoms. • Electric and magnetic fields deflect the beam. • Gives mass/charge of e- = −5.60 x 10-9 g/C • Coulomb (C) = SI unit of charge

8. Essence of the Thompson Experiement (and old fashioned TV’s) Electric field exerts Force + plate repels +charged particles - Plate repels – charged particles y + x + d _ _ Phosphor screen F = Eq = ma d = displacement = ½ at2 = Eq/m (t = L/Vx) Therefore, the greater the displacement, the lower the mass of the particle

9. Electrons • Millikan (1911) studied electrically-charged oil drops.

10. mass charge me = charge x Electron mass Charge on each droplet was: n (−1.60 x 10-19 C) with n = 1, 2, 3,… n (e- charge) • These experiments give: Modern value = −1.60217653 x 10-19 C. = −1 “atomic units”. = (-1.60 x 10-19 C)(-5.60 x 10-9 g/C) = 8.96 x 10-28 g Modern value = 9.1093826 x 10-28 g

11. Protons • Atoms gain a positive charge when e- are lost. Implies a positive fundamental particle. • Hydrogen ions had the lowest mass. • Hydrogen nuclei assumed to have “unit mass” • Called protons. Modern science: mp = 1.67262129 x 10-24 g mp ≈ 1800 x me. Charge = -1 x (e- charge). = +1.602176462 x 10-19 C = +1 atomic units

12. The Nuclear Atom • How were p+ and e- arranged? • Thompson: • Ball of uniform positive charge, with small negative dots (e-) stuck in it. • The “plum-pudding” model. 1910Rutherford (former Thompson graduate student) fired α-particles at thin metal foils. Expected them to pass through with minor deflections.

13. Rutherford Scattering Experiment

14. Different Models of the Atom: different scattering results • “Plum pudding model” • + and – charges evenly distrubted • low, uniform density of matter • No back scattering α particles Rutherford’s explanation of results: Small regions of very high density + charge in the dense regions - Charges in region around it α particles From wikipedia

15. α particles The Nuclear Atom Some Large Deflections were osbserved Rutherford “It was about as credible as if you had fired a 15-inch shell at a piece of paper and it came back and hit you.”

16. α particles The Nucleus Most of the mass and all “+” charge is concentrated in a small core, the nucleus. ≈10,000 times smaller diameter than the entire atom. e- occupy the remaining space.

17. The Atom is mostly ‘empty’ Charge cloud Diameter ~ 1 Å Mass ~ 10-30 kg Nucleus diameter~ 10-4 Å = 10-14 meters Mass ~ 10-27 Kg

18. Most Chemistry involves rearrangement of outermost electrons, not nuclei Example: H  1p+ , 1 e- H + H  H2 +

19. How to “see” atoms:Scanning Tunneling Microscopy

20. 7 Å Epitaxial Al2O3(111) film on Ni3Al(111) (Kelber group): • Grown in UHV • Uniform • No Charging Start with ordered films growth studies Proceed to amorphous films on Si(100) STM Surface terminated by hexagonal array of O anions S. Addepalli, et al. Surf. Sci. 442 (1999) 3464

21. Neutrons • Atomic mass > mass of all p+ and e- in an atom. • Rutherford proposed a neutral particle. Chadwick (1932) fired -particles at Be atoms. Neutral particles, neutrons, were ejected: mn ≈ mp (0.1% larger). mn = 1.67492728 x 10-24 g. Present in all atoms (except ‘normal’ H).

22. ~ 10-10 meters = 1 angstrom (Å) Smeared out electron charge cloud _ 10-14 m + Protons and neutrons + + + + + + +

23. The Nuclear Atom Nucleus • Contains p+ and n0 • Most of the atomic mass. • Small (~10,000x smaller diameter than the atom). • Positive (each p+ has +1 charge). Electrons • Small light particles surrounding the nucleus. • Occupy most of the volume. • Charge = -1. Atoms are neutral. Number of e− = Number of p+

24. A neutron can decay into a proton and electron: n0 p+ + e- This can cause decay of a radioactive element, e.g., # of p+ + n0 14 14 C C Elemental symbol (carbon) 6 6 Atomic No. (# e- = # p+ Carbon with 6 protons and 8 neutrons is unstable (radioactive) Carbon with 6 protons and 6 neutrons is stable (non-radioactive 12 C 6 radioctive stable

25. An atom of 14C can undergo decay to N as a neutron turns into a proton + an emitted electron 14 14 N C + e- 7 6 1 p+ 1 n0 + an electron (emitted)

26. 1 12 Atomic mass unit (amu) = (mass of C atom) that contains 6 p+ and 6 n0. Atomic Numbers & Mass Numbers • Same element - same number of p+ Atomic number (Z) = number of p+ 1 amu = 1.66054 x 10-24 g Particle Mass Mass Charge (g) (amu) (atomic units) e− 9.1093826 x 10-28 0.000548579 −1 p+ 1.67262129 x 10-24 1.00728 +1 n0 1.67492728 x 10-24 1.00866 0

27. 1 1 Hydrogen isotopes: H 1 p+, 0 n0 2 1 3 1 H 1 p+, 1 n0 H 1 p+, 2 n0 Isotopes and Atomic Weight • Isotopes • Atoms of the same element with different A. • equal numbers of p+ • different numbers of n0 deuterium (D) tritium (T)

28. ISOTOPES: SAME Element, Different numbers of neutrons Carbon: atomic no. = 6  6 protons in the nucleus+ 6 electrons 12 14 C C Atomic mass = 12 amu (12 gr/mole) Therefore , 6 protons + 6 neutrons Atomic mass = 14 amu Therefore, 6 protons+ 8 neutrons

29. Isotopes Display the same chemical reactivities (which depend mainly on the outer arrangement of the electrons) 12C + O2 CO2 14C +O2  CO2 12 14 C C Isotopes display different nuclear properties stable Radioactive: spontaneously emits electrons. Half-life ~ 5730 years

30. Isotopes and Moles (more on this later) and isotope abundance: 1 mole = 6.02 x 1023of anything! 1 mole of atoms = 6.02 x 1023 atoms Molar Mass (in grams) = average atomic mass (in amu) 1 mole of H atoms = 1.008 gr. Why not 1.000 gr??  most atoms are 1H, but some are 2H (deuterium)

31. Average atomic mass of H = 1.008 amu 100 atoms have a mass of 100.8 amu # of 2H atoms = n # of 1H atoms = 100 –n (assume these are the only two isotopes that matter) Mass of 100 atoms = n x 2.000 +(100-n) x 1.000 = 100.8 amu 2n + 100-n = 100.8 n = 0.8 So, out of every 100 atoms , have 0.8 2H atoms Out of every 1000 atoms, have 8 2H atoms Natural abundance of “heavy hydrogen (deuterium) is then 0.8%

32. 24Mg 25Mg 26Mg number of p+ 12 12 12 number of n0 12 13 14 mass / amu 23.985 24.986 25.982 Isotopes and Atomic Weight • Most elements occur as a mixture of isotopes. Magnesium is a mixture of:

33. Isotopes and Atomic Weight • For most elements, the percent abundance of its isotopes are constant (everywhere on earth). • The periodic table lists an average atomic weight. Example Boron occurs as a mixture of 2 isotopes, 10B and 11B. The abundance of 10B is 19.91%. Calculate the atomic weight of boron.

34. 19.91 100 (10.0129 amu) = 1.994 amu 10B 80.09 100 (11.0093 amu) = 8.817 amu 11B Isotopes and Atomic Weights Boron occurs as a mixture of 2 isotopes, 10B and 11B. The abundance of 10B is 19.91%. Calculate the atomic weight of boron. Atomic mass = Σ(fractional abundance)(isotope mass) % abundance of 11B = 100% - 19.91% = 80.09% Atomic weight for B = 1.994 + 8.817 amu = 10.811 amu

35. Isotopes and Atomic Weight Periodic table: 5 Atomic number (Z) B Symbol Boron Name 10.811 Atomic weight

36. Amounts of Substances: The Mole • A counting unit – a familiar counting unit is a “dozen”: 1 dozen eggs = 12 eggs 1 dozen donuts = 12 donuts 1 dozen apples = 12 apples • 1 mole (mol) = Number of atoms in 12 g of 12C • Latin for “heap” or “pile” • 1 mol = 6.02214199 x 1023 “units” • Avogadro’s number

37. 6.0 x 1018 ft3 8.4 x 1013 ft2 =7.1 x 104 ft = 14 miles ! Amounts of Substances: The Mole • A green pea has a ¼-inch diameter. 48 peas/foot. • (48)3 / ft3 ≈ 1 x 105 peas/ft3. • V of 1 mol ≈ (6.0 x 1023 peas)/(1x 105 peas/ft3) • ≈ 6.0 x 1018 ft3 U.S. surface area = 3.0 x 106 mi2 = 8.4 x 1013 ft2 height = V / area, 1 mol would cover the U.S. to:

38. Amounts of Substances: The Mole • 1 mole of an atom = atomic weight in grams. 1 Xe atom has mass = 131.29 amu 1 mol of Xe atoms has mass = 131.29 g 1 He atom has mass = 4.0026 amu 1 mol of He has mass = 4.0026 g There are 6.022 x 1023 atoms in 1 mol of He and 1 mol of Xe – but they have different masses. … 1 dozen eggs is much heavier than 1 dozen peas!

39. 1 mol Cu 63.546 g Conversion factor: = 1 1 mol Cu 63.546 g Molar Mass and Problem Solving Example How many moles of copper are in a 320.0 g sample? Cu-atom mass = 63.546 g/mol (periodic table) nCu = 320.0 g x = 5.036 mol Cu n = number of moles

40. 1 mol B 10.81 g nB = (1.000 g) = 0.092507 mol B Molar Mass and Problem Solving • Calculate the number of atoms in a 1.000 g sample of boron. B atoms = (0.092507 mol B)(6.022  1023 atoms/mol) = 5.571  1022 B atoms

41. Dimensional Analysis and Problem Solving Special Homework Problem: Due Tues. Recitation Density = mass/volume Problem: Assume that a hydrogen atom has a spherical diameter of 1 angstrom Assume that the nucleus (1 proton) has a diameter of 10-4 angstrom Calculate the densities of the nucleus, and of the electron charge cloud in kg/m3 Calculate the ratio of the two densities: R = dnucleus/delectron cloud Mass of proton = 1.67 x 10-27 kg Mass of electron = 9 x 10-31 kg

42. The Periodic Table • Summarizes • Atomic numbers. • Atomic weights. • Physical state (solid/liquid/gas). • Type (metal/non-metal/metalloid). • Periodicity • Elements with similar properties are arranged in vertical groups.

43. The Periodic Table In the USA, “A” denotes a main group element… International system uses 1 … 18. …”B” indicates a transition element.

44. The Periodic Table Main group metal Transition metal Metalloid Nonmetal

45. The Periodic Table Period number A period is a horizontal row

46. The Periodic Table Group 8A Noble gases Group 1A Alkali metals (not H) A group is a vertical column Group 7A Halogens Group 2A Alkaline earth metals

47. The Alkali Metals and Alkaline Earth Metals • Alkali metals (group 1A; 1) • Alkaline earth metals (group 2A; 2) • Grey … silvery white colored. • Highly reactive. • Never found as native metals. • Form alkaline solutions.

48. Transition Elements, Lanthanides & Actinides • Transition Elements (groups 1B – 8B) • Also called transition metals. • Middle of table, periods 4 – 7. • Includes the lanthanides & actinides. • Lanthanides and Actinides • Listed separately at the bottom. • Chemically very similar. • Relatively rare on earth. • (old name: rare earth elements)

49. Groups 3A to 8A • Groups 3A to 6A • Most abundant elements in the Earth’s crust and atmosphere. • Most important elements for living organisms. • Halogens (group 7A; 17) • Very reactive non metals. • Form salts with metals. • Colored elements. • Noble gases (8A; 18) • Very low reactivity. • Colorless, odorless gases.

50. Sizes of Atoms and Units • Atoms are very small. • 1 tsp of water contains 3x as many atoms as there are tsp of water in the Atlantic Ocean! Impractical to use pounds and inches... • Need a universal unit system • The metric system. • The SI system (Systeme International) - derived from the metric system.