The great variety of colors in a garden scene The texture of the fabric in your clothes

3. Look Around You • The great variety of colors in a garden scene • The texture of the fabric in your clothes • The solubility of sugar in a cup of coffee • The transparency of a window

4. How Do We Explain? • What makes diamonds transparent and hard, while table salt is brittle and dissolves in water? • Why does paper burn, and why does water quench fires? • Where does the beautiful colors of flowers come from? • The structure and behavior of atoms are key to understanding the properties of matter.

5. Variety of Elements • The diverse properties results from only about 100 different elements • How do atoms combine with one another? • What rules govern the ways in which atoms can combine? • How do the properties of a substance relate to the kinds of atoms it contains? • What is an atom like, and what make their difference?

6. 2.1 The Atomic Theory of Matter • History • Democritus and Greek philosophers (BC 400) • The material world mustbe made up of tiny indivisible particles • atomos: indivisible or uncuttable • Plato and Aristotle • There can be no ultimately indivisible particles • The “atomic” view of matter faded for many centuries • Newton (1642-1727) • Air is composed of something invisible and in constant motion • Still very different from thinking of atoms as the fundamental building blocks

7. 2.1 The Atomic Theory of Matter • Dalton’s Atomic Theory • Chemists learned to measure the amounts of elements that reacts with one another to form new substances • Dalton’s atomic theory • Introduced during the period from 1803 to 1807 • The theory was based on the four postulates given in the figure in the next page

9. 2.1 The Atomic Theory of Matter • Dalton’s Postulates • Dalton’s theory explains several simple laws of chemical combination. • The law of constant composition • In a given compound, the relative numbers and kinds of atoms are constant • The law of conservation of mass (matter) • The total mass of materials present after a chemical reaction is the same as the total mass present before the reaction • The law of multiple proportions • If two elements A and B combine to form more than one compound, the masses of B that can combine with a given mass of A are in the ratio of small whole numbers (H2O and H2O2)

10. 2.2 The Discovery of Atomic Structure • Atomic Images • Dalton had no direct evidence for the existence of atoms • Scientists have developed methods for more detailed probing of the nature of matter • Today, we can measure the properties of individual atoms and even provide images of them FIGURE 2.2 An image of the surface of silicon obtained byscanningtunneling microscopy (STM)

11. 2.2 The Discovery of Atomic Structure • Cathode Rays and Electrons • Thomson found that cathode rays are streams of negatively charged particles

12. 2.2 The Discovery of Atomic Structure • Cathode Rays and Electrons • The charge/mass ratio: 1.76 X 108 C/g

13. 2.2 The Discovery of Atomic Structure • Millikan’s Oil Drop Experiment • Robert Millikan (University of Chicago) determined the charge on the electron in 1909.

14. 2.2 The Discovery of Atomic Structure • Radioactivity • Radioactivity is the spontaneous emission of radiation by an atom. • It was first observed by Henri Becquerel. • Marie and Pierre Curie also studied it. • Three types of radiation were discovered by Ernest Rutherford: • a particles • b particles • g rays

15. 2.2 The Discovery of Atomic Structure • The Nuclear Model • The prevailing theory was that of the “plum pudding” model, put forward by Thomson. • It featured a positive sphere of matter with negative electrons imbedded in it.

16. 2.2 The Discovery of Atomic Structure • The Nuclear Model • Ernest Rutherford shot a particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.

17. 2.2 The Discovery of Atomic Structure • The Nuclear Model • Since some particles were deflected at large angles, Thompson’s model could not be correct. • 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. • Protons were discovered by Rutherford in 1919. • Neutrons were discovered by James Chadwick in 1932.

18. 2.3 The Modern View of Atomic Structure • Subatomic Particles • Protons and electrons are the only particles that have a charge. • Protons and neutrons have the same mass. • The mass of an electron is so small we ignore it. FIGURE 2.11 The structure of an atom.

19. 2.3 The Modern View of Atomic Structure • Subatomic Particles • Every atom has an equal number of electrons and protons, so atoms have no net charge. • Atomic mass unit, 1 amu = 1.66054 X 10-24 g • Atom’s size: 1-5 Å

20. 2.3 The Modern View of Atomic Structure Sample Exercise 2.1Atomic Size The diameter of a US dime is 19 mm, and the diameter of a silver atom is 2.88 Å. How many silver atoms could be arranged side by side across the diameter of a dime? Practice Exercise The diameter of a carbon atom is 1.54 Å. (a) Express this diameter in picometers. (b) How many carbon atoms could be aligned side by side across the width of a pencil line that is 0.20 mm wide?

21. 2.3 The Modern View of Atomic Structure • The Diameters of Atomic Nuclei • About 10-4Å • Density of nucleus: 1013 ~ 1014 g/cm3 • A match box full of nuclei would weigh over 2.5 billion tons!

22. 2.3 The Modern View of Atomic Structure • Atomic Numbers, Mass Numbers, & Isotopes • What makes the difference between carbon and oxygen? • The atoms of each element have a characteristic number of protons • Atomic number: the number of protons in the nucleus • Isotopes: atoms with identical atomic numbers but different mass numbers.

24. 2.5 The Periodic Table • The most significant tool that chemists use for organizing and remembering chemical facts

26. 2.4 Atomic Weights • The Atomic Mass Scale • The atomic mass unit (amu) • Defined by assigning a mass of exactly 12 amu to an atom of 12C • 1H: 1.0078 amu, 16O: 15.9949 amu • Most elements occur in nature as mixtures of isotopes. • 12C: 98.93%, 13C : 1.07% (0.9893)(12 amu) + (0.0107)(13.00335 amu) = 12.01 amu • Average Atomic Masses atomic weight

28. ▲FIGURE 2.12 A mass spectrometer. ▲FIGURE 2.12 Mass spectrum of atomic chlorine.

29. 2.5 The Periodic Table • Periodicity • Many elements show very strong similarities to one another • When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.

30. 2.5 The Periodic Table • The most significant tool that chemists use for organizing and remembering chemical facts

31. 2.5 The Periodic Table • It is a systematic catalog of the elements. • Elements are arranged in order of atomic number.

32. 2.5 The Periodic Table • The rows on the periodic chart are periods. • Columns are groups. • Elements in the same group have similar chemical properties.

33. 2.5 The Periodic Table • Groups • Many groups are known by their names. • “Coinage metals”: Group 11

34. 2.5 The Periodic Table • Metals, Nonmetals, and Metalloids • Nonmetals generally differ from the metals in appearance and in other physical properties. • A metalloid is a chemical element with properties that are in-between or a mixture of those of metals and nonmetals

36. 2.5 The Periodic Table • Isolated Pu • Identified the elements having atomic numbers 95 through 102 • Identified element number 106 • ACS proposed that element number 106 be named seaborgium ◄ FIGURE 2.17

37. 2.6 Molecules and Molecular Compounds • Molecules • Only the noble-gas elements are normally found in nature as isolated atoms. • A molecule is an assembly of two or more atoms tightly bound together. • Molecules behave in many ways as a single, distinct object a molecule atoms

38. 2.6 Molecules and Molecular Compounds • Molecules and Chemical Formulas • Many elements are found in nature in molecular form • Two different molecular forms of oxygen • O2: a diatomic molecule, essential for life, odorless • O3: a triatomic molecule, toxic, pungent smell • Diatomic molecules

39. 2.6 Molecules and Molecular Compounds • Molecules and Chemical Formulas • Molecular compounds are composed of two or more different atoms • Molecules vs Compounds • Most molecular substances that we will encounter contain only nonmetals.

40. 2.6 Molecules and Molecular Compounds • Molecular and Empirical Formulas • Empirical formulas • H2O • HO • CH2 • CH2O • Molecular formulas • H2O • H2O2 • C2H4 • C6H12O6 • Why do we need empirical formulas? • Certain common methods of analyzing substances lead to the empirical formula only

41. 2.6 Molecules and Molecular Compounds • Picturing Molecules • A structural formula shows which atoms are attached to which within the molecule. • A perspective drawing gives some sense of three dimensional shape • Ball-and-stick models show the accurate angles between bonds • A space-filling model shows the relative sizes of the atoms.

42. 2.7 Ions and Ionic Compounds • Cations and Anions • 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 (metal atoms). • Anions are negative and are formed by elements on the right side of the periodic chart (nonmetal atoms).

43. 2.7 Ions and Ionic Compounds • Predicting Ionic Charges • Many atoms gain or lose e- to make the same number of e- as the noble gas. Figure2.20. Predictable charges of some common ions

44. 2.7 Ions and Ionic Compounds • Ionic Compounds • A compound that contains both positively and negatively charged ions. • Generally combinations of metals and nonmetals

45. 2.7 Ions and Ionic Compounds • Writing Empirical Formulas • for Ionic Compounds • The ionic compound formed from Mg2+ and N3-

46. 97% of the mass of most organisms: O, C, H, N, P, and S • 70% of the mass of most cells: H2O • C is the most prevalent element in the solid components of cells

47. 2.8 Naming Inorganic Compounds • Names and Formulas of Ionic Compounds • Positive ions (Cations) • Cations formed from metal atoms • Ions of the same element that have different charges exhibit different properties • Metals that form only one cation - group 1A/2A, Al3+, Ag+, Zn2+ • Cations from nonmetals

48. 2.8 Naming Inorganic Compounds Fe(III) Fe(II)

49. 2.8 Naming Inorganic Compounds • Names and Formulas of Ionic Compounds • Negative ions (Anions) • Monatomic and simple polyatomic anions • Oxyanions

50. 2.8 Naming Inorganic Compounds • Names and Formulas of Ionic Compounds • Negative ions (Anions) • Anions containing H+ • Older method: HCO3- bicarbonate ion; HSO4- bisulfate ion (a) SeO42- (b) SeO32-