1. William L Masterton Cecile N. Hurley http://academic.cengage.com/chemistry/masterton Edward J. Neth University of Connecticut Chapter 2 Atoms, Molecules, and Ions

2. Learning a Language (NOT IN NOTES) When learning a new language: Start with the alphabet Then, form words Finally, form more complex structures such as sentences Chemistry has an alphabet and a language; in this chapter, the fundamentals of the language of chemistry will be introduced

3. Outline (NOT IN NOTES) Atoms and Atomic Theory Components of the Atom Introduction to the Periodic Table Molecules and Ions Formulas of Ionic Compounds Names of Compounds

4. The Language of Chemistry (NOT in NOTES) This chapter introduces the fundamental language of chemistry Atoms, molecules and ions Formulas Names

5. The Structure of Matter – notes start here Atoms Composed of electrons, protons and neutrons Molecules Combinations of atoms Ions Charged particles

6. Greeks: Empedocles and Democritus Suggested the concept of atoms but were not taken seriously or credited with an atomic theory

7. John Dalton: credited with the first atomic model

8. Figure 2.1 - John Dalton and Atomic Theory

9. Atomic Theory 1.An element is composed of tiny particles called atoms 2.All atoms of the same element have the same chemical properties 3.In an ordinary chemical reaction, atoms rearrange their bonds but atoms are not created or destroyed 4.Compounds are formed when two or more atoms of different element combine

10. Fundamental Laws of Matter There are three fundamental laws of matter 1.Law of conservation of mass 2.Law of constant composition 3.Law of multiple proportions

11. Law of Conservation of Mass Matter is conserved in chemical reactions This applies to all chemical reactions but DOES NOT include nuclear reactions

12. Law of Constant Composition Compound always contains the same elements in the same proportions by mass. Pure water has the same composition everywhere.

13. Law of Multiple Proportions The masses of one element that combine with a fixed mass of the second element are in a ratio of small whole numbers. Compare CO and CO 2

14. Figure A – The Law of Multiple Proportions Two different oxides of chromium

15. Components of the Atom Atomic theory raised more questions than it answered Could atoms be broken down into smaller particles 100 years after atomic theory was proposed, the answers were provided by experiment Finding the Electrons: Protons: Neutrons:

16. Fundamental Experiments J.J. Thomson: Cavendish Laboratories, Cambridge, England

17. Figure 2.2 – J.J. Thomson and Ernest Rutherford

18. Figure 2.3 – Cathode Ray Apparatus

19. Electrons First evidence for subatomic particles came from the study of the conduction of electricity by gases at low pressures J.J. Thomson, 1897 Rays emitted were called cathode rays Rays are composed of negatively charged particles called electrons Electrons carry unit negative charge (-1) and have a very small mass (1/2000 the lightest atomic mass)

20. J.J. Thomson’s Model Every atom has at least one electron Atoms are known that have one hundred or more electrons There is one electron for each positive charge in an atom Electrical neutrality is maintained

21. Ernest Rutherford: McGill University, Canada and Manchester and Cambridge Universities, England Bombardment of gold foil with α particles (helium atoms minus their electrons Expected to see the particles pass through the foil Found that some of the alpha particles were deflected by the foil Led to the discovery of a region of heavy mass at the center of the atom

22. Gold Foil Experiment: Bombardment of gold foil with α particles (helium atoms minus their electrons) Expected to see the particles pass through the foil Found that some of the alpha particles were deflected by the foil Led to the discovery of a region of heavy mass at the center of the atom = nucleus

23. Figure 2.4 – Rutherford Backscattering

24. Nuclear Particles 1. Protons Mass nearly equal to the H atom Positive charge 2. Neutrons Mass slightly greater than that of the proton No charge

25. Atomic Mass The average mass of all of the isotopes of an element accounting for their relative abundances

26. Table 2.1 – Subatomic Particles

27. Terminology Atomic number, Z Number of protons in the atom Mass number, A Number of protons plus number of neutrons Mass # = p + + n 0

28. Nuclear symbolism A is the mass number Z is the atomic number X is the chemical symbol

29. Isotopes Isotopes are two atoms of the same element Same atomic number Different mass numbers n 0 = mass # - p + (n 0 = A – Z) Number of neutrons differs between isotopes

30. Isotopes of hydrogen (NOT IN NOTES) 1 H, 2 H, 3 H Hydrogen, deuterium, tritium Different masses Note that some of the ice is at the bottom of the glass – this is 2 H 2 O

31. Example 2.1

32. Radioactivity Radioactive isotopes are unstable (Radioactive decay is not a chemical process) 1. These isotopes decay over time 2. Emit other particles and are transformed into other elements Particles emitted 1. Beta (β) particles: High speed electrons 2. Alpha (α) particles: helium nuclei 3. Gamma (γ) rays: high energy light

33. Nuclear Stability depends on the neutron/proton ratio For light elements, n/p is approximately 1/1 For heavier elements, n/p is approximately 1.4/1

34. Figure 2.5 – The Nuclear Belt of Stability

35. 2.3 Introduction to the Periodic Table Dmitri Mendeleev: 1836-1907 Arranged elements by chemical properties Left space for elements unknown at the time Predicted detailed properties for several undiscovered elements: Sc, Ga, Ge By 1886, all these elements had been discovered, and with properties similar to those he predicted

36. Mendeleev’s P.T.

37. Introduction to the Periodic Table

38. Modern Periodic Table Period – a horizontal row on the periodic table Group – a vertical column on the periodic table Blocks – sections of elements with common properties Families – another name for group; emphasizes the similarity in properties within a group

39. Blocks in the Periodic Table Main group elements 1-2, 13-18 Transition Metals 3-12 Post-transition metals Elements in groups 13-15 to the right of the transition metals and under the staircase Ga, In, Tl, Sn, Pb, Bi

40. Families with Common Names (label on PT) Alkali Metals, Group 1(I) Alkaline Earth Metals, Group 2 (II) Halogens, Group 17 (VII) Noble Gases, Group 18 (VIII)

41. A Look at the Sulfur Group Sulfur (nonmetal), antimony (metalloid) and silver (metal)

42. Biological View of the Periodic Table – NOT IN YOUR NOTES, JUST AN FYI “Good guys” Essential to life Carbon, hydrogen, oxygen, sulfur and others “Bad guys” Toxic or lethal Some elements are essential but become toxic at higher concentrations Selenium

43. Figure 2.8 – Biologically Important and Toxic Elements

44. Example 2.3

45. 2.4 Molecules and Ions Molecule: Two or more atoms chemically combined 1. Atoms involved are often nonmetals 2. Covalent bonds are strong forces that hold the atoms together Molecular formulas: Number of each atom is indicated by a subscript Examples Water, H 2 O Ammonia, NH 3

46. Structural Formulas Structural formulas: a formulas that shows the bonding patterns within the molecule

47. Ions A charged particle that is the result of the loss or gain of electrons Cation – a positive ion (loss) Anion – a negative ion (gain) Examples: Na → Na + + e - O + 2e - → O 2-

48. Ionic Compounds Compounds formed from the electrostatic attraction of oppositely charged particles Sodium chloride (NaCl): Sodium cations and chloride anions associate into a continuous network

49. Forces: Ionic compounds are held together by strong forces Compounds are usually solids at room temperature High melting points often water-soluble

50. Solutions: When an ionic compound dissolves in water, the ions are released from each other conductivity – the ions in a solution support the transmission of an electric current Strong electrolytes – solutions that are very good conductors Weak electrolytes – solutions that are poor conductors Nonelectrolytes – solutions that do NOT conduct

51. Figure 2.12 – Electrical Conductivity

52. Formulas for Ionic Compounds Charge balance Each positive charge must have a negative charge to balance it Calcium chloride, CaCl 2 Ca 2+ Two Cl - ions are required for charge balance

53. Cations of Transition and Post-Transition Metals Polyvalent – exhibit multiple positive charges depending on conditions Iron Commonly forms Fe 2+ and Fe 3+ Lead Commonly forms Pb 2+ and Pb 4+

54. Polyatomic Ions Groups of atoms may carry a charge; these are the polyatomic ions OH - NH 4 +

55. Noble Gas Connections Atoms that are close to a noble gas (group 18 or VIII) form ions that contain the same number of electrons as the neighboring noble gas atom +1, +2, +3 skip -3, -2, -1 Noble Gases

56. Example 2.5

57. 2.6 Naming of Compounds Cations: element name Na +, sodium If polyvalent, a Roman numeral is used to denote the charge Fe 2+ iron(II)

58. Names of Compounds - Anions Monatomic anions are named by adding –ide to the element name Oxygen becomes oxide, O 2- Polyatomic ions keep their names

59. To name an ionic compound: name the cation first, then, name the anion (with the word 'ion' omitted). It is not necessary to indicate the number of cations and anions in the compound because it is understood that the total positive charges carried by the cations must equal the total negative charges carried by the anions.

60. KIpotassium ion + iodide ion = potassium iodide CoCl 2 cobalt(II) ion + two chloride ions = cobalt(II) chloride CoCl 3 cobalt(III) ion + three chloride ions = cobalt(III) chloride Hg 2 Cl 2 mercury(I) ion + two chloride ions = mercury(I) chloride AgNO 3 silver ion + nitrate ion = silver nitrate

61. Oxoanions Per _________ate ___________ate ___________ite Hypo_________ite

62. Table 2.3 – Oxoanions of Nitrogen, Sulfur and Chlorine

63. Binary Molecular Compounds Made of 2 nonmetal elements Never reduce subscripts Covalently bonded

64. Mono-1 di-2 tri-3 tetra-4 penta-5 hexa-6 hepta-7 octa-8 nona-9 Systematic naming 1. First name is the first element, with prefix to for number of atoms (EXCEPT NO MONO) 2. Second name is prefix with element name changed to –ide (INCLUDE MONO)

65. Some Examples Diphosphorus pentaoxide Sulfur dioxide Dinitrogen tetraoxide Hydrogen dioxide Carbon monoxide Phosphorus trichloride

66. Acids Ionic compounds with Hydrogen as the cation Naming: Common: (strong acids) HBrHIHCl H 2 SO 4 HNO 3 HClO 4 Br I Cl SO NO ClO 434

67. Oxyacids or Oxoacids: Acids with and oxoanion as the anion

68. Acids of Chlorine (example):

69. Examples: Hydrogen chloride (hydrochloric acid) Nitric acid Sulfuric acid Hypobromous acid Nitrous acid Phosphoric acid