Atoms, Molecules, and Ions

1. Atoms, Molecules, and Ions • Classification and Compositions of Matter • Scientific Laws and Atomic Theories: • Dalton’s Atomic Theory • Atomic Compositions and Structures • Isotopes • Periodic Table • Molecules and Ions • Types of Compounds • Naming Compounds

2. Classification of Matter

3. Matter According to Ancient Philosophy • Matter is composed of four basic elements: Earth, water, wind, and fire • Two schools of thought emerged during Greek Civilization: • Aristotle and his followers believed that matter to be infinite – not composed of discrete unit. • Democritus and Leucippus believed that matter is made of discrete units called “atomos” that is indivisible.

4. The “Development” of Theory • Early 18th Century, chemists attempted to explain the combustion process. • Why do some materials burn, while others don’t? • What they observed: “When a piece of wood was burned, what was left was ash that had mass much less than the original wood.” What happen to the rest of the wood?

5. The Phlogiston Theory? • Georg Stahl (1659-1735) proposed the phlogiston theoryto explain the principle of combustion: • Combustible materials contain phlogiston; • When a substance burns it loses phlogiston, so that the mass decreases; • Non-combustible substances do not contain phlogiston.

6. Experimental Science versus Philosophy • Antoine Lavoisier (1743-1794) performed quantitative experiments to study chemical reactions, including combustion process; • He noted that when a substance is burned, it produced a product with a higher mass. He concluded that: • The Phlogiston theory was incorrect; • Combustion involves reaction with oxygen gas and formed a product with more mass, because oxygen is added to form the product; • Mass is conserved during chemical reactions.

7. Antoine Lavoisier & the Law of Conservation of Mass • The total mass of substances is the same before and after the reaction; • 5.00 g Zn + 5.00 g S  7.45 g of ZnS + 2.55 g of S; • total mass before reaction = 5.00 g + 5.00 g = 10.00 g; • total mass after reaction = 7.45 g + 2.55 g = 10.00 g; • Mass is conserved during reaction. • (Mass is neither created or destroyed during chemical reactions.)

8. Experimental Science versus Philosophy • Joseph Proust (1754-1826) performed experiments that analyzed compositions of various compounds, and found that: A given compound always has the same composition of the element (in mass %) regardless of its origin or sample size.

9. The Law of Constant Composition Examples: • Copper carbonate is composed of 51.4% Cu, 9.7% C, and 38.9% O, by mass. (b) Sodium chloride is composed of 39.34% Na and 60.66% Cl, by mass. (c) Sugar is composed 42.1% C, 6.48% H, and 51.4% O, by mass.

10. Two Important Fundamental Laws • The Law of Conservation of Mass: (Mass is neither created nor destroyed.) In chemical reactions, the total mass of substances is conserved. • The Law of Definite Proportions: A given compound contains the same compositions of the elements it is made up of, regardless of the origin.

11. The Need for a Theory • Scientific laws are statements that generalize properties of matter and their reactions; • Theories are needed to explain these natural laws, like: • How do we explain that mass is not gained or lost during chemical reactions? • How do we explain that calcium carbonate always contains calcium, carbon and oxygen in the same exact proportions?

12. John Dalton (1766-1844) • Born in Eaglesfield, England; he was color blind; • Starting at age 15, he helped older brother run Quaker boarding school, and became school principal at 19; • Taught himself science and mathematics; • At 26, he became teacher of mathematics and natural philosophy at Manchester’s new College; • In 1903 he published a scientific paper on the law of partial pressures of gas, known as Dalton’s law of partial pressures; • In 1905 he published the Dalton’s Atomic Theory;

13. Dalton’s Atomic Theory (1805) 1. Elements are made up of discrete particles called atoms; 2. Atoms of the same element are identical; atoms of different elements are different; 3. Atoms are not created or destroyed during chemical reactions; they are only re-grouped. 4.A compound contains atoms of at least two different elements; they are combined in a simple whole number ratio; 5. A given compound always contains the same number and type of atoms;

14. The Law of Multiple Proportion • Dalton also noted that: When two elements combine (react) to form more that one compounds, there is a simple ratio in the masses of one of the elements that combined with a fixed mass of the second element in these compounds.

15. The Law of Multiple Proportion • Example-1: • Carbon reacts with oxygen to form two compounds: X and Y. In compound X, there are 1.33 g of oxygen for every gram of carbon; in compound Y, there are 2.66 g of oxygen for every gram of carbon. Fixing the mass of carbon in X and Y, we noted the masses of oxygen in X and Y show a simple whole number ratio of 1:2. From this mass ratio, the formula of X and Y were derived as X = CO, and Y = CO2

16. The Law of Multiple Proportion • Example-2: Sulfur reacts with oxygen to form two compounds, A and B, where compound A contains about 1.00 g of oxygen for every gram of sulfur, and compound B contains 1.50 g of oxygen for every gram of sulfur. Thus, for a fixed mass of sulfur in A and B, the ratio of the mass of oxygen in A to that in B is 2:3. Their formulas are: A = SO2 and B = SO3, which reflex the

17. Law of Multiple Proportions Exercise #1: • Sulfur reacts with fluorine gas to form three different compounds, A, B and C. Compound A is found to contain 1.000 g of sulfur and 1.185 g of fluorine; compound B contains 1.000 g of sulfur and 2.370 g of fluorine; while compound C contains 1.000 g of sulfur and 3.556 g of fluorine. Show that these data illustrate the law of multiple proportions. (a) If A = SF, what are the formula of B and C? (b) If A = SF2, what are the formula of B and C?

18. Law of Multiple Proportions Exercise #1: Answer: Mass ratios of fluorine in A, B, and C: 1.185 g-to-2.370 g-to-3.556 g = 1:2:3 Mass ratios for fluorine in A, B & C shows simple whole number ratios – implies Law of Multiple proportions. Formulas: (a) If A = SF, B = SF2; and C = SF3; (b) If A = SF2; B = SF4; and C = SF6;

19. Gay-Lussac Interpretation of Combining Volume

20. Principle of Chemical Combination • Gay-Lussac’s Law of Combining Volumes: In reactions that involve gaseous reactants and products, there exist “simple ratios” of their volumes measured under the same temperature and pressure. Examples: • 1 L of hydrogen gas reacts with 1 L of chlorine gas to form 2 L of hydrogen chloride gas; • 2 L of hydrogen gas reacts with 1 L of oxygen to form 2 L of water vapor.

21. Interpretation of Gay-Lussac’s Experiments • According to Avogadro’s law: “equal volumes of gases at the same temperature and pressure contain the same number of molecules” • 1 L of hydrogen + 1 L of chlorine  2 L of hydrogen chloride • This implies: 1 H-molecule + 1 Cl-molecule  2 HCl molecules. Hydrogen and chlorine molecules must be diatomic (2 atoms per molecule), and the reaction may be written as follows: H2(g) + Cl2(g) 2 HCl(g)

22. Interpretation of Gay-Lussac’s Experiments • 2 L of hydrogen + 1 L of oxygen  2 L of water vapor implies: 2 H-molecules + 1 O-molecule  2 water molecules. a) Hydrogen and oxygen gases contains diatomic molecules (H2 and O2), and water has the formula H2O. b) The above reaction can be represented by the equation: 2H2(g) + O2(g) 2 H2O(g)

23. J.J. Thomson (1856 – 1940) • Born December 18, 1856, Cheetham Hill, England; • Attended Owens College, Manchester in 1870; Trinity College, Cambridge in 1876 for graduate study; • Became fellow at Cavendish Lab under Lord Raleigh in 1880; lecturer in 1883 and Professor of Physics in 1884; • Started “cathode ray experiments” in 1894 that led to the discovery of electron and “plum-pudding model”; • Received Nobel prize in physics in 1906; • Knighted by King Edward VII in 1908;

24. Cathode Rays and Discovery of Electrons • In 1895, J.J. Thomson discovered cathode-ray while studying the flow of electric current through a vacuum.

25. Cathode-ray Tube used by J.J. Thomson

26. Cathode Ray

27. Characteristics of Cathode Ray • Rays originates from the cathode plate; • It contains negatively charged particles - it bends in electric and magnetic fields in the direction that indicates negatively charged particles; • The charge-to-mass ratio of cathode ray particles is constant at -1.76 x 108 C/g, regardless of the materials used as a cathode; • Conclusion: cathode ray is a beam of negatively charged particles now called electrons.

28. Modern Version of Cathode-ray Tube

29. Thomson’s Model of an Atom • J.J. Thomson proposed the “Plum-pudding” model: (a) Atom is composed of a diffused mass of matter (like a cotton ball) with positive charges, with electrons loosely embedded on its surface; (b) The number of electrons present in an atom must yield a total negative charge that is equal to the magnitude of positive charge in the atom.

30. Plum-Pudding Model

31. Ernest Rutherford (1871-1937) • Born and received early education in New Zealand; • Attended Canterbury College for B.Sc., then moved to Cambridge in 1894 to work under J.J. Thomson; • Become Chair of Physics at McGill University, Montreal, Canada in 1898; then became Physics Professor at University of Manchester in 1907; • Started “alpha particles scattering experiments” with Han Geiger and Ernest Marsden; • Discovered atomic nucleus and “nuclear model” in 1910; • Nobel price in chemistry in 1908; knighted in 1914.

32. Alpha Particles Scattering Experiment

33. Rutherford’s Nuclear Model

34. Rutherford’s Nuclear Atomic Model 1. Atom contains nucleus, composed of protons and neutrons; 2. The nucleus is much, much smaller than the atom; 3. Electrons occupy the vast “empty space” surrounding the nucleus; 4. The mass of atom is concentrated in the nucleus; 5. Proton or neutron is almost 2000 times larger and more massive than electron;

35. The Atomic Structure & Composition

36. A Version of Nuclear Model of Atom

37. Robert Millikan (1868 – 1953) • Born in Morrison, Illinois; • Went to Oberlin College, Ohio, in 1886, and graduated with a bachelor degree in 1891; • Received his Ph.D. in Physics from Colombia U in 1895; • Became Professor at University of Chicago 1910-1921; • Determined the correct charge on an electron (1910) • In 1921, became Director of the Norman Bridge Laboratory of Physics at Cal Tech, Pasadena; • Received Nobel prize in physics in 1923;

38. What is the Charge of an Electron?Millikan’s Oil-Drop Experiments

39. Relative and Absolute Masses Relative MassAbsolute Mass • Proton: 1.007276 amu; 1.673 x 10-27 kg. • Neutron: 1.008665 amu; 1.675 x 10-27 kg. • Electron: 0.000549 amu; 9.109 x 10-31 kg.

40. Relative and Absolute Charges RelativeAbsolute • Proton = +1; +1.602 x 10-19 C; • Neutron = 0; • Electron = -1; -1.602 x 10-19 C;

41. Atoms and Isotopes • Atoms of the same element can have different masses; they are called Isotopes • Isotopes: Atoms having same number of protons but different number of neutrons; • Element is identified by chemical symbol X or atomic numberZ; • Isotope isidentified by chemical symbolX and mass numberA; Atomic number (Z) = # of protons; Mass number (A) = # of protons + # of neutrons; # neutrons = (A – Z)

42. Exercise #2: Isotope Symbols Write the symbols of isotopes that contain the following: (a) 10 protons, 10 neutrons, and 10 electrons. (b) 12 protons, 13 neutrons, and 10 electrons. (c) 15 protons, 16 neutrons, and 15 electrons. (d) 17 protons, 18 neutrons, and 18 electrons. (e) 24 protons, 28 neutrons, and 21 electrons. (a) 20Ne; (b) 25Mg2 +; (c) 31P; (d) 35Cl-; (e) 52Cr3+

43. Exercise #3: Isotopes Indicate the number of protons, neutrons, and electrons in each isotope with the following symbols. (a) 60Ni (b) 239Pu4+ (c) 79Se2- Answers: (a) 28 protons, 32 neutrons, and 28 electrons; (b) 94 protons, 145 neutrons, and 90 electrons; (c) 34 protons, 45 neutrons, and 36 electrons.

44. Molecules and Ions • Molecule: A neutral particle that contains two or more atoms bound together by covalent bonds. • Ion = electrically charged particle; • Cation = positive ion; • Anion = negative ion; Atoms may lose electrons to form cations, or gain electrons to form anions.

46. Periodic Table • The modern Periodic Table is divided into 18 columns (groups) and 7 rows (periods). • Groups are numbered 1 – 18 in the IUPAC configuration, or 1A – 8A and 1B – 8B in the ACS configuration. • In each period, elements are arranged left-to-right in increasing atomic number; • Within each group, elements share similar chemical and physical characteristics.

47. Major Classifications of Elements • Main group elements: • Group 1A (1): the alkali metals; • Group 2A (2): the alkaline Earth metals; • Groups 3A (13), 4A (14), 5A (15), and 6A (16), • Group 7A (17): the halogens, and • Group 8A (18): the noble gases. • Transition metals: Groups 3B (3) – 2B (12) ; contains heavy metals. • Metalloids (semi-metals): B, Si, Ge, As, Sb, Te, Po, At

48. Classification of Elements in The Periodic Table

49. Characteristics of Metals, Nonmetals & Metalloids • Metals: • Solid, except mercury; have shiny appearance; • Good conductors of heat and electricity; • Malleable and ductile; • Only react with nonmetals; • Metals lose electrons and become cations; • Metals cannot react with one another. • Compounds of metals are primarily ionic;

50. Characteristics of Metals, Nonmetals & Metalloids • Nonmetals: • Mainly gases; bromine is liquid; few are solids; • Poor conductors of electricity; • Solids are brittle and not lustrous. • In reactions with metals, nonmetals gain electrons and become anions; • Nonmetals also react with one another or with metalloids to form molecular compounds.