Atomic Theory of Matter and Modern View of Atomic Structure

1. AP Chapter 2 – Atoms, Molecules, and Ions Jennie L. Borders

2. Section 2.1 – The Atomic Theory of Matter • Democritus was the first to describe atoms as indivisible. • Dalton’s Atomic Theory • Each element is composed of extremely small particles called atoms. • All atoms of a given element are identical, but the atoms of one element are different from the atoms of all other elements. • Atoms of one element cannot be changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions. • Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kinds of atoms.

3. Laws for Atoms • The law of constant composition states that in a given compound, the relative number and kinds of atoms are constant. • The law of conservation of mass states that 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 states that if two elements A and B combine to from 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.

4. Section 2.2 – The Discovery of Atomic Structure • J.J. Thomson discovered electrons using the cathode ray tube. • Ernest Rutherford or Eugen Goldstein discovered protons. • James Chadwick discovered neutrons.

5. Rutherford’s Experiment • J.J. Thomson made the plum pudding model of the atom. • Rutherford tested the plum pudding model by shooting alpha particles (double positive particles) at gold foil. • Since some alpha particles bounced back and did not pass through the foil, Rutherford determined that there was a dense, positive nucleus in the center of the atom. • Since many alpha particles passed straight through the foil, Rutherford determined that most of the atom is empty space.

6. Section 2.3 – The Modern View of Atomic Structure • The charge on an electron is -1.602 x 10-19C, which we call a -1 charge. • The charge on a proton is +1.602 x 10-19C, which we call a +1 charge. • Neutrons have no charge. • Atoms have an equal number of protons and electrons.

7. Size of the Atom • Since atoms are so small, a unit called angstroms is commonly used for an atom’s diameter. 1 angstrom = 10-10m. • The atomic mass unit (amu) is used for the masses of atoms. 1 amu = 1.66054 x 10-24g. • Most of the space of an atom is the electron cloud, but most of the mass of the atom is in the nucleus (protons and neutrons), since electrons have almost 2000 times less mass than protons and neutrons.

8. Numbers for Atoms • Atoms are identified by the number of protons. • The atomic number is the number of protons and can be found at the top of the box on the periodic table. Atomic number is the bottom number when using shorthand. • Since atoms have no charge, they have the same number of protons and electrons.

9. Sample Exercise 2.1 • The diameter of a U.S. dime is 17.9 mm, and the diameter of a silver atom is 2.88 angstrom. How many silver atoms could be arranged side by side across the diameter of a dime?

10. Practice Exercise 1 • Which of the following factors determines the size of an atom? a. the volume of the nucleus b. the volume of the space occupied by the electrons in the atoms c. the volume of a single electron, multiplied by the number of electrons in the atom d. the total nuclear charge e. the total mass of the electrons surrounding the nucleus

11. Practice Exercise 2 • The diameter of a carbon atom is 1.54 angstrom. 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?

12. Mass Numbers • The mass number of an atom is the number of protons and neutrons in the nucleus of the atom. • To calculate the number of neutrons, you can do mass number minus atomic number. • Isotopes are atoms of the same element that have different numbers of neutrons. They have the same atomic number since they have the same number of protons but different mass numbers since the neutrons are not the same.

13. Sample Exercise 2.2 • How many protons, neutrons, and electrons are in an atom of a. 197Au b. strontium-90

14. Practice Exercise 1 • Which of these atoms has the largest number of neutrons? a. b. c. d. e. 148Eu 157Dy 149Nd 162Ho 159Gd

15. Practice Exercise 2 • How many protons, neutrons, and electrons are in an atom of a. 138Ba b. Phosphorus - 31

16. Sample Exercise 2.3 • Magnesium has three isotopes with mass numbers 24, 25, and 26. a. Write the complete chemical symbol (superscript and subscript) for each. b. How many neutrons are in an atom of each isotope?

17. Practice Exercise 1 • Which of the following is an incorrect representation for a neutral atom? a. b. c. d. e. 63Li 136C 6330Cu 3015P 10847Ag

18. Practice Exercise 2 • Give the complete chemical symbol for the atom that contains 82 protons, 82 electrons, and 126 neutrons.

19. Section 2.4 – Atomic Weights • The mass number is the number of protons and neutrons in an atom. • The atomic mass is the average mass of all of the isotopes of the element multiplied by their relative abundances.

20. Sample Exercise 2.4 • Naturally occurring chlorine is 75.78% 35Cl (atomic mass 34.969 amu) and 24.22% 37Cl (atomic mass 36.966 amu). Calculate the atomic weight of chlorine.

21. Practice Exercise 1 • There are two stable isotopes of copper found in nature, 63Cu and 65Cu. If the atomic weight of copper Cu is 63.546 amu, which of the following statements are true? a. b. c. All copper atoms have a mass of 63.546 amu. 65Cu contains two more protons than 63Cu. 63Cu must be more abundant than 65Cu.

22. Practice Exercise 2 • Three isotopes of silicon occur in nature: 28Si (92.23%), atomic mass 27.97693 amu; 29Si (4.68%), atomic mass 28.97649 amu; and 30Si (3.09%), atomic mass 29.97377 amu. Calculate the atomic weight of silicon.

23. Section 2.5 – The Periodic Table • The horizontal rows on the periodic table are called periods. • The vertical columns on the periodic table are called groups. • Elements in the same group have similar physical and chemical properties. • Metals are on the left of the periodic table and nonmetals are on the right of the periodic table.

24. Metals vs Nonmetals • Metals are on the left and middle of the periodic table. They are lustrous, good conductors, and solid (except mercury). • Nonmetals are on the right of the periodic table. They are brittle, mainly liquids and gases, and poor conductors (except carbon). • Metalloids surround the stair-step line and have properties of both metals and nonmetals.

25. Sample Exercise 2.5 • Which two of these elements would you expect to show the greatest similarity in chemical and physical properties: B, Ca, F, He, Mg, P?

26. Practice Exercise 1 • A biochemist who is studying the properties of certain sulfur (S) – containing compounds in the body wonders whether trace of another nonmetallic element might have similar behavior. To which element should she turn her attention? a. F b. As c. Se d. Cr e. P

27. Practice Exercise 2 • Locate Na (sodium) and Br (bromine) in the periodic table. Give the atomic number of each and classify each as metal, metalloid, or nonmetal.

28. Section 2.6 – Molecules and Molecular Compounds • The seven diatomic elements are H2, N2, O2, F2, Cl2, Br2, and I2. • Molecular compounds consist of nonmetals. • The molecular formula shows the exact number and types of atoms in a molecule. • The empirical formula is the lowest whole number ratio of atoms in a molecule.

29. Sample Exercise 2.6 • Write the empirical formulas for a. glucose, a substance also known as either blood sugar or dextrose – molecular formula C6H12O6. b. nitrous oxide, a substance used as an anesthetic and commonly called laughing gas – molecular formula N2O.

30. Practice Exercise 1 • Tetracarbon dioxide is an unstable oxide of carbon with the following molecular structure: • What are the molecular and empirical formulas of this substance? a. C2O2, CO2 b. C4O, CO c. CO2, CO2 d. C4O2, C2O e. C2O, CO2

31. Practice Exercise 2 • Give the empirical formula for decaborane, whose molecular formula is B10H14.

32. Section 2.7 – Ions and Ionic Compounds • An ion is an atom that has gained or lost electrons and has a charge. • A cation is a positive ion (lost electrons). • An anion is a negative ion (gains electrons).

33. Sample Exercise 2.7 • Give the chemical symbol, including the superscript indicating mass number for, a. the ion with 22 protons, 26 neutrons, and 19 electrons. b. the ion of sulfur that has 16 neutrons and 18 electrons.

34. Practice Exercise 1 • In which of the following species is the difference between the number of protons and the number of electrons largest? a. Ti2+ b. P3- c. Mn d. Se2- e. Ce4+

35. Practice Exercise 2 • How many protons, neutrons, and electrons does the 79Se2-ion possess?

36. Sample Exercise 2.8 • Predict the charge expected for the most stable ion of barium and the most stable ion of oxygen.

37. Practice Exercise 1 • Although it is helpful to know that many ions have the electron arrangement of a noble gas, many elements, especially among the metals, form ions that do not have a noble-gas electron arrangement. Use the periodic table to determine which of the following ions has a noble-gas electron arrangement, and which do not. For those that do, indicate the noble-gas arrangement they match: a. Ti4+ b. Mn2+ c. Pb2+ d. Te2- e. Zn2+

38. Practice Exercise 2 • Predict the charge expected for the most stable ion of a. Aluminum b. Fluorine

39. Ionic Compounds • The charges for an ion can be found by looking at which group the ion is in on the periodic table. • An ionic compound is made between a cation and anion held together by an ionic bond (the transfer of electrons). They are generally a metal and a nonmetal.

40. Sample Exercise 2.9 • Which of these compounds would you expect to be ionic: N2O, Na2O, CaCl2, SF4?

41. Practice Exercise 1 • Which of these compounds are molecular: CBr4, FeS, P4O6, PbF2?

42. Practice Exercise 2 • Give a reason why each of the following statements is a safe prediction: a. Every compound of Rb with a nonmetal is ionic in character. b. Every compound of nitrogen with a halogen element is a molecular compound. c. The compound MgKr2does not exist. d. Na and K are very similar in the compounds they form with nonmetals. e. If contained in an ionic compound, calcium (Ca) will be in the form of the doubly charged ion, Ca2+.

43. Sample Exercise 2.10 • Write the empirical formula of the compound formed by a. Al3+and Cl-ions b. Al3+and O2-ions c. Mg2+and NO3-ions

44. Practice Exercise 1 • Which of the following nonmetals will form an ionic compound with Sc3+that has a 1:1 ratio of cations to anions? a. Ne b. F c. O d. N

45. Practice Exercise 2 • Write the empirical formula for the compound formed by a. Na+and PO43- b. Zn2+and SO42- c. Fe3+and CO32-

46. Section 2.8 – Naming Inorganic Compounds • Cations have the same element name. • Anions have the ending –ide. • When naming ionic compounds that start with a regular metal (group 1, group 2, or Al), then name the cation and add –ide to the anion unless it is a polyatomic ion. • When naming ionic compounds that start with a transition metal, then name the cation and add a roman numeral to indicate the charge. Then add –ide to the anion unless it is a polyatomic ion.

47. Sample Exercise 2.11 • Based on the formula for the sulfate ion, predict the formula for a. The selenate ion b. The selenite ion

48. Practice Exercise 1 • Which of the following oxyanions is incorrectly named? a. ClO2-, chlorate b. IO4-, periodate c. SO32-, sulfite d. IO3-, iodate e. NO2-, nitrite

49. Practice Exercise 2 • The formula for the bromate ion is analogous to that for the chlorate ion. Write the formula for the hypobromite and bromite ions.

50. Sample Exercise 2.12 • Name the ionic compounds a. K2SO4 b. Ba(OH)2 c. FeCl3