CHEM 120: Introduction to Inorganic Chemistry

1. CHEM 120: Introduction to Inorganic Chemistry Instructor: Upali Siriwardane (Ph.D., Ohio State University) CTH 311, Tele: 257-4941, e-mail: upali@chem.latech.edu Office hours: 10:00 to 12:00 Tu & Th ; 8:00-9:00 and 11:00-12:00 M,W,& F

2. Chapters Covered and Test dates • Tests will be given in regular class periods  from  9:30-10:45 a.m. on the following days: September 21,     2004 (Test 1): Chapters 1 & 2 • October 6,         2004(Test 2):  Chapters  3, & 4 • October 20,         2004 (Test 3): Chapter  5 & 6 • November 3,        2004 (Test 4): Chapter  7 & 8 • November 15,      2004 (Test 5): Chapter  9 & 10 • November 17,      2004 MAKE-UP: Comprehensive test (Covers all chapters • Grading: • [( Test 1 + Test 2 + Test3 + Test4 + Test5)] x.70 + [ Homework + quiz average] x 0.30 = Final Average •                               5

3. Chapter 2: The composition and structure of the atom

4. Matter and structure • Picture of matter has evolved (and is still evolving) over the years. • Democritus (Greek philosopher, 470-380 B.C.) --atomic theory

5. Composition of the atom • The smallest unit of an element that retains the properties of that element is called an ____ • Basic structural unit of an element is an ____ • An _____ is incredibly small.

6. Makeup of an atom for a chemist • An atom is composed of • The _________ (positively charged) and _______ (uncharged) are found in a very small, dense portion of the atom called the nucleus . • __________ (negatively charged) surround the nucleus in a very diffuse region and have a much smaller mass than the proton and neutron.

7. Name charge mass(amu) mass(g) • Electron (e) -1 5.4x10-4 9.1095x10-28 • Proton (p) +1 1.00 1.6725x10-24 • Neutron (n) 0 1.00 1.6750x10-24

8. Elemental Symbols • How are elements given symbols? • Chemical symbols can be one or two letters. The first letter is always a capital case and the second letter is always a small case. Some symbols are taken from the Latin or German names of elements. • Na = sodium, K = potassium, Fe = iron, Cu = copper, Ag = silver,  Sn = tin, Sb = antimony, W = tungsten, Au = gold, Hg = mercury, Pb = lead

9. Symbolic notation for element • Mass no. A c charge X Atomic no. Z • X is the chemical symbol for the element • The atomic no.(Z) is the • The mass no.(A) is the

10. If the mass number (A) = total no. of protons and neutrons in the nucleus, • The number is neutrons is? • Number of neutrons = • For neutral atoms, the no. of protons in the nucleus = no. of electrons outside of the nucleus and the overall charge is zero.

11. Isotopes (or are all atoms of a given element the same?) • Atoms with the same atomic number but different mass number (therefore diff. nos. of neutrons) are called isotopes • Most elements have two or more isotopes. • 11H 21H 31H

12. 33 S 16 • How many protons? • How many neutrons? • How many electrons? • Is it necessary to include the atomic no.?

13. What is the atomic number of bromine? • How many protons does a Br atom have? • How many neutrons does a Br atom with mass number 79 have? • How many electrons does a (neutral) Br atom have?

14. How many protons, neutrons, and electrons are in • 20984Po • 13656Ba

15. Average atomic mass • The atomic masses of an element are the weighted averages of all the isotopes of that element taking into account the relative abundance each isotope. • Average at.mass = %abundance isotope 1/100% x mass isotope 1 + %abundance isotope 2/100% x mass isotope 2 + etc…...

16. Isotopic mass problem • The atomic masses of the two stable isotopes of boron, 105B (19.78%) and 115B (80.22%) are 10.0129amu and 11.0093 amu respectively. What is the “average” atomic mass of boron?

17. Question 2.3 • The element neon has three naturally occurring isotopes. One of these has a mass of 19.99 amu and a natural abundance of 90.48%. A second isotope has a mass of 20.99 amu and a natural abundance of 0.27%. A third has a mass of 21.99 amu and a natural abundance of 9.25%. Calculate the atomic mass of neon.

18. Harder isotope problem • A hypothetical atom has two isotopes only. • Isotope one has a percent abundance of 60%. • If the average isotopic mass is 120.0 amu what are the masses of isotopes one and two?

19. Ions • Ions are charged particles that are a result of the atom • _______ have more electrons than protons and are negatively charged. The original atom has gained electron(s). • ________ have more protons than neutrons and are positively charged. The original atom has lost electron(s).

20. Charge on an ion = no. of protons - no. of electrons. • 3717Cl gains one electron g • 13856Ba loses two electrons g Like charges repel each other, opposite attract.

21. 2.24: Write the symbol for an isotope • that contains 92 protons and 146 neutrons • 2.30: Which are true? • An atom with an atomic number of 7 and a mass number of 14 is identical to an atom with an atomic number of 6 and a mass number of 14. • Neutral atoms have the same number of electrons as protons. • The mass of an atom is due to the sum of the no. of protons, neutrons and electrons.

22. 2.28 from end of chapter

23. Development of the atomic theory • Democritus (Greek philosopher) fifth century B.C. : matter consists of very small, indivisible particles--atomos (atoms--uncuttable) • Plato and Aristotle not accept this idea.

24. Dalton’s Atomic Theory (1808) • Marked beginning of modern era of chemistry

25. Dalton’s hypotheses • Elements are composed of extremely small particles, called atoms. • All atoms of an element are identical (same size, mass, chem. prop). • The atoms of one element are different from the atoms of all other elements. • An atom cannot be created, divided, destroyed or converted into any other type of atom. • Compounds are composed of atoms of more than one element in simple whole-number ratios. • A chemical reaction involves the separation, combination, or rearrangement of atoms, not their creation or destruction.

26. X2Y Law of definite proportions

27. Which of Dalton’s postulates are considered true today? • Elements are composed of extremely small particles, called atoms. • All atoms of an element are identical (same size, mass, chem. prop). • The atoms of one element are different from the atoms of all other elements. • An atom cannot be created, divided, destroyed or converted into any other type of atom. • Compounds are composed of atoms of more than one element in simple whole-number ratios. • A chemical reaction involves the separation, combination, or rearrangement of atoms, not their creation or destruction.

28. Subatomic particles • Is the atom really indestructible, or does it consist of even smaller particles?

29. Electrons • 1890’s: cathode ray tube experiments (tube sealed with metal electrodes in it and evacuated of air) • Apply high voltage source, invisible ray produced (see effect by fluorescence when ray strikes coated surface).

30. Electrons (fig 2.3) A. Magnetic Field C. Electric field B. No field or balanced fields. Thomson determined charge/mass ratio of -1.76x108C/g)

31. Find rays have same properties regardless of metal used in constructing the cathode. • Experiments show that cathode rays are made of charged particles that interact with electric and magnetic field when moving. • Particles are negatively charged (repelled by the negative plate, attracted toward the positive plate). • These negative particles are fundamental particles of matter. Called electrons. (1897 Thomson)

32. Are there other particles that make up the atom? • Atoms are neutral and contain electrons which are negatively charged. • Therefore there must be something positive present also.

33. Protons, which are positively charged, were discovered by Goldstein Protons have mass of 1.67262 x 10-24 g (1840 times greater than electron mass) and charge equal but opposite in sign from e-.

34. Atomic strucure model: or where might the positive charge be? Thomson’s plum pudding model is a uniform positive sphere where negative electrons embedded in it.

35. Natural radioactivity • Spontaneous emission of particles and/or radiation. a rays = positively charged helium nuclei b rays = electrons (negatively charged particles) g rays = high-energy radiation (photons), with no charge

36. Scattering of a particles by gold foil Experiment of Geiger Majority of a particles undeflected or slightly deflected but some displayed large deflections.

37. Rutherford picture of atom from scattering expt (1910) • Atom is mostly empty space with positive charge (protons) concentrated in a dense central core called the nucleus. (Neutrons are also in the nucleus--Chadwick experimental evidence in 1932 for neutron) • Positive core of atom deflects  particles strongly. • ‘It was almost as incredible as if you fired a 15-inch shell at a piece of tissue and it came back and hit you.”

38. 2.4: Relationship between light and atomic structure

39. Light and atomic structure • Rutherford atom doesn’t say much about electron location except that they’re in a region outside of the nucleus that is mostly empty space. • Let’s look at electromagnetic spectrum of light. (energy and wavelength)

40. roygbiv High energy Short wavelength Low energy Long wavelength Electromagnetic spectrum 7.1

41. Pass ordinary light through a prism get continuous spectrum of all wavelengths • But if look at emitted light from a tube containing hydrogen or another gas get an emission spectrum like

42. Where did all the colors go?

43. Bohr atom • Electrons orbit nucleus like a planet around the sun in circular orbits (held electrostatically). • Hydrogen atom consists of 1 electron orbiting 1 proton

44. Electron can only be located in certain stable orbits. Bohr assumed that the energy of the e-’s orbit and its radius are ________. Not all energies or radii are allowed. • Moving from one orbit (quantum level) to another causes atom to absorb or emit a photon (particle of light) of energy

45. Color and energy of line in spectrum correspond to difference in energy btn 2 levels

46. Summary of results of Bohr’s theory (p 47) • Light energy can be absorbed and emitted by promotion and relaxation of electrons from one energy level to another--see a line in spectrum corresponding to energy difference (photon emission) btn levels. • Don’t see all colors, just those that correspond to energy difference btn levels

47. Electrons can be found only in certain allowed energy levels (orbits). Not all energies or radii of orbits are allowed--quantized. • As orbits get further from nucleus, energy of orbit increases. • According to Bohr one can know the location and energy of an electron in an atom with certainty. • Read the summary of Bohr theory on p 46

48. Modern atomic theory • Bohr model only good for one electron atoms and quantization assumed. • Later developments: • deBroglie noted that electrons had both wave and particle properties: wave-particle duality of matter. Need both concepts to describe electrons.

49. Heisenberg Uncertainty Principle: It is impossible to know simultaneously how fast an electron is moving and its position with certainty.

50. This leads to: • Electrons do not move around the nucleus in well-defined orbits but are located in orbitals which are regions in space where there is a large probability of finding an electron (electron cloud). • These electron clouds are denser in some regions than others. The electron density is proportional to the probability of finding the electron at any point in time.