1. Atoms, Molecules, and Ions

2. History Highlights  600 B.C. group of Greek philosophers became dissatisfied with myths  600 B.C. - 400 B.C. Golden Age of Philosophy  Answer ??? with logical possibilities  Some believed single substance that was basis for everything else  Water - Thales  Air - Anaximenes

3. The Golden Age continues  Universe composed of 4 elements earth, air, fire and water  Empedocles – 4 elements combine in different proportions to make up all objects in the universe  Leucippus and Democritus trying to determine smallest particle of matter (idea of atom)

4. Atomic Theory of Matter The theory that atoms are the fundamental building blocks of matter reemerged in the early 19th century, championed by John Dalton.

5. Dalton’s Postulates  Each element composed of atoms  All atoms of a given element are identical  Atoms of an element are not changed during a chemical reaction (the atoms just move from one chemical species to another)  Compounds are formed when atoms of more than one element combine

6.  Law of conservation of mass The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place.  Law of constant composition Also known as the law of definite proportions. The elemental composition of a pure substance never varies.  Law of multiple proportions When two elements form different compounds, the mass ratio of the elements in one compound is related to the mass ratio in the other by a small whole number. Atoms

7. Discovery and Properties of Electrons  Humphrey Davy (early 1800’s) - passed electricity through compounds compounds decomposed into elements compounds are held together by electrical forces  Michael Faraday - (1832-1833) - amount of reaction that occurs during electrolysis is proportional to current passed through compounds Matter (atoms) is electrical in nature.

8. Discovery and Properties of Electrons  Cathode Ray Tubes - (late 1800’s & early 1900’s)  2 electrodes in a glass tube with a gas at low pressure  voltage applied to tube causing a glow discharge  “rays” emitted from cathode (- end) to anode (+ end)  Cathode Rays must be negatively charged!

9. Discovery and Properties of Electrons  J.J. Thomson - (1897) - changed cathode ray tube experiments by adding two adjustable voltage electrodes into the experiment

10. The Electron  Streams of negatively charged particles were found to emanate from cathode tubes.  J. J. Thomson is credited with their discovery (1897).  Thomson measured the charge/mass ratio of the electron to be 1.76  10 8 coulombs/g of e-.

11. Cathode Ray Tube

12. Millikan Oil Drop Experiment Robert A. Millikan University of Chicago 1 st American Nobel Laureate  determined the charge and mass of the electron (1909)  oil drop experiment

13. Millikan Oil Drop Experiment  Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would yield the other.  charge on 1 electron = -1.60219 x 10 -19 coulomb  using Thomson’s charge to mass ratio we get that the mass of 1 electron is 9.11 x 10 -28 g

14. Canal Rays and Protons  Goldstein (1886) - “Canal Rays” streams of positively charged particles in cathode rays flow in opposite direction of cathode rays  must be positive postulated existence of “proton”

15. Radioactivity  Three types of radiation were discovered by Ernest Rutherford:   particles   particles   rays

16. Spontaneous emission of high energy radiation A radioactive substance is placed in a shield containing a small hole so that a beam of radiation is emitted from the hole. The radiation is passed between two electrically charged plates and detected. Three spots are noted on the detector: a spot in the direction of the positive plate, a spot which is not affected by the electric field, a spot in the direction of the negative plate. Radioactivity

17. A high deflection towards the positive plate corresponds to radiation which is negatively charged and of low mass. This is called  -radiation (consists of electrons). No deflection corresponds to neutral radiation. This is called  -radiation. Small deflection towards the negatively charged plate corresponds to high mass, positively charged radiation. This is called  -radiation. Radioactivity

18. The Atom, circa 1900:  “Plum pudding” model, put forward by Thomson.  Positive sphere of matter with negative electrons imbedded in it.

19. Discovery of the Nucleus  Ernest Rutherford - 1910 - basic picture of atom  Geiger & Marsden’s experiment on  - particle scattering from thin Au foils  Ernest Rutherford shot  particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.

20. In order to get the majority of  -particles through a piece of foil to be undeflected, the majority of the atom must consist of a low mass, diffuse negative charge - the electron. To account for the small number of high deflections of the  -particles, the center or nucleus of the atom must consist of a dense positive charge. Rutherford and the Nuclear Atom

21. The Nuclear Atom Since some particles were deflected at large angles, Thomson’s model could not be correct.

22. The Nuclear Atom  Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom.  Most of the volume of the atom is empty space.  nearly all of atom’s mass in nucleus  nuclear diameter is 1/10,000 to 1/100,000 times less than atom’s radius  nuclear density is 10 15 g/mL- equivalent to 3.72 x 10 9 tons/in 3

23. The atom consists of positive, negative, and neutral entities (protons, electrons, and neutrons). Protons and neutrons are located in the nucleus of the atom, which is small. Most of the mass of the atom is due to the nucleus. Electrons are located outside of the nucleus. Most of the volume of the atom is due to electrons. The Modern View of Atomic Structure

25. Subatomic Particles  Protons were discovered by Rutherford in 1919.  Neutrons were discovered by James Chadwick in 1932  Protons and electrons are the only particles that have a charge.  Protons and neutrons have essentially the same mass.  The mass of an electron is so small we ignore it.

26. Symbols of Elements Elements are symbolized by one or two letters.

27. Atomic Number All atoms of the same element have the same number of protons: The atomic number (Z)

28. Atomic Mass The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.

29. Isotopes:  Atoms of the same element with different masses.  Isotopes have different numbers of neutrons. 11 6 C 12 6 C 13 6 C 14 6 C

30. Isotopes of Hydrogen

31. Atomic Mass Atomic and molecular masses can be measured with great accuracy with a mass spectrometer.

32. Average Mass  Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations.  Average mass is calculated from the isotopes of an element weighted by their relative abundances.

33. Periodic Table:  A systematic catalog of elements.  Elements are arranged in order of atomic number.

34. Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.

35. Molecular Compounds Molecular compounds are composed of molecules and almost always contain only nonmetals.

36. Diatomic Molecules These seven elements occur naturally as molecules containing two atoms.

37. Types of Formulas  Empirical formulas give the lowest whole-number ratio of atoms of each element in a compound.  Molecular formulas give the exact number of atoms of each element in a compound.

38. Types of Formulas  Structural formulas show the order in which atoms are bonded.  Perspective drawings also show the three-dimensional array of atoms in a compound.

39. Ions  When atoms lose or gain electrons, they become ions.  Cations are positive and are formed by elements on the left side of the periodic chart.  Anions are negative and are formed by elements on the right side of the periodic chart.

40. Ionic Bonds Ionic compounds (such as NaCl) are generally formed between metals and nonmetals.

41. Writing Formulas  Because compounds are electrically neutral, one can determine the formula of a compound this way:  The charge on the cation becomes the subscript on the anion.  The charge on the anion becomes the subscript on the cation.  If these subscripts are not in the lowest whole-number ratio, divide them by the greatest common factor.

42. Common Cations

43. Common Anions

44. Inorganic Nomenclature  Write the name of the cation.  If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion.  If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses.

45. Inorganic Nomenclature

46. Patterns in Oxyanion Nomenclature  When there are two oxyanions involving the same element:  The one with fewer oxygens ends in -ite  NO 2 − : nitrite ; SO 3 2− : sulfite  The one with more oxygens ends in -ate  NO 3 − : nitrate; SO 4 2− : sulfate

47. Patterns in Oxyanion Nomenclature  The one with the second fewest oxygens ends in -ite  ClO 2 − : chlorite  The one with the second most oxygens ends in -ate  ClO 3 − : chlorate

48. Patterns in Oxyanion Nomenclature  The one with the fewest oxygens has the prefix hypo- and ends in -ite  ClO − : hypochlorite  The one with the most oxygens has the prefix per- and ends in -ate  ClO 4 − : perchlorate

49. Acid Nomenclature  If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro- :  HCl: hydrochloric acid  HBr: hydrobromic acid  HI: hydroiodic acid

50. Acid Nomenclature  If the anion in the acid ends in -ite, change the ending to -ous acid:  HClO: hypochlorous acid  HClO 2 : chlorous acid

51. Acid Nomenclature  If the anion in the acid ends in -ate, change the ending to -ic acid:  HClO 3 : chloric acid  HClO 4 : perchloric acid

52. Nomenclature of Binary Compounds  The less electronegative atom is usually listed first.  A prefix is used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however.)

53. Nomenclature of Binary Compounds  The ending on the more electronegative element is changed to -ide.  CO 2 : carbon dioxide  CCl 4 : carbon tetrachloride

54. Nomenclature of Binary Compounds If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one: N 2 O 5 : dinitrogen pentoxide