1. Chapter 2 Atoms, Molecules and Ions The Early History of Chemistry Before 16 th Century Greeks were the first to attempt to explain why chemical changes occur. Alchemy: Attempts to change cheap metals into gold. They invented the idea of atoms, that matter is not continuous. They discovered several elements and learned to prepare mineral acids.

2. The Early History of Chemistry 16 th Century German develop the systematic metallurgy (extraction of metal from ores) Swiss develop the medicinal application of minerals 17 th Century Robert Boyle: First chemist to perform quantitative experiments

3. Fundamental Chemical Laws Law of Conservation of Mass (Antoine Lavoisier, 18 th Century) – Mass is neither created nor destroyed Law of Definite Proportion (Joseph Proust, 19 th Century) – A given compound always contains exactly the same proportion of elements by mass. This principle of constant composition of compounds is a law of definite proportion. Carbon tetra chloride is always 1 atom carbon per 4 atoms chlorine

4. Fundamental Chemical Laws Law of Multiple Proportions (John Dalton, 19 th Century) – When two elements form a series of compounds, the ratio of the masses of the second element that combine with 1g of the first element can always be reduced to small whole numbers. The ratio of the masses of oxygen in H 2 O and H 2 O 2 will be a small whole number (“2”).

5. 2 2.1

6. Dalton’s Atomic Theory (1808) 1. Each element is made up of tiny particles called atoms 2. The atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways 3. Chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative numbers and types of atoms 4. Chemical reactions involve reorganization of the atoms – changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction.

7. 8 X 2 Y 16 X8 Y + 2.1

8. Avogadro’s Hypothesis (1811) At the same temperature and pressure, equal volumes of different gases contain the same number of particles 5 liters of oxygen and 5 liters of nitrogen contain the same number of particle

9. Thomson Atomic Model (1903) An atom consists of a diffuse cloud of positive charge with the negative electrons embedded randomly in it. This model is often called plum (or raisin) pudding model. Observed cathode ray (produced at the negative electrode and repelled by the negative pole of an applied electric field. Cathode ray was a stream of negatively charged particles now called electrons

10. Deflection of Cathode Rays by an Applied Electric Field

11. 2.2

12. J.J. Thomson, measured mass/charge of e - (1906 Nobel Prize in Physics) 2.2

13. The Plum Pudding Model of the Atom

14. Rutherford Atomic Model (1911) An atom with a dense center of positive charge (the nucleus) with electrons moving around the nucleus at a distance that is large relative to the nuclear radius. Nucleus is very small compared with the overall size of the atom. Nucleus is extremely dense, accounts for almost all of the atom’s mass.

15. Radioactivity Spontaneous emission of radiation Gamma (  ) rays: high energy light Beta (  ) particles: high speed electron Alpha (  ) particles (He 2+ ): 2 + charge, charge twice that of electron and with opposite sign. The mass of an  -particle is 7300 times that of the electron

16. 1.atoms positive charge is concentrated in the nucleus 2.proton (p) has opposite (+) charge of electron (-) 3.mass of p is 1836 x mass of e - (1.67 x 10 -24 g)  particle velocity ~ 1.4 x 10 7 m/s (~5% speed of light) (1908 Nobel Prize in Chemistry) 2.2

17. Expected and Actual Results of Rutherford’s Experiment

18. (Uranium compound) 2.2

19. atomic radius ~ 100 pm = 1 x 10 -10 m nuclear radius ~ 5 x 10 -3 pm = 5 x 10 -15 m Rutherford’s Model of the Atom 2.2

20. e - charge = -1.60 x 10 -19 C Thomson’s charge/mass of e - = -1.76 x 10 8 C/g e - mass = 9.10 x 10 -28 g Measured mass of e - (1923 Nobel Prize in Physics) 2.2

21. Chadwick’s Experiment (1932) H atoms - 1 p; He atoms - 2 p mass He/mass H should = 2 measured mass He/mass H = 4  + 9 Be 1 n + 12 C + energy neutron (n) is neutral (charge = 0) n mass ~ p mass = 1.67 x 10 -24 g 2.2

22. Subatomic Particles (Table 2.1) mass p = mass n = 1840 x mass e - 2.2

23. The Chemists’ Shorthand Atomic Symbols Atomic number (Z): number of protons, gives the symbol of the element (X) Mass number (A): Total number of protons and neutrons Elemental form = Zero net charge Therefore, # electrons = # of protons K  Element Symbol 39 19 Mass number  Atomic number 

24. 2.3

25. How many protons, neutrons, and electrons are in C 14 6 ? How many protons, neutrons, and electrons are in C 11 6 ? 6 protons, 8 (14 - 6) neutrons, 6 electrons 6 protons, 5 (11 - 6) neutrons, 6 electrons Understanding Isotopes 2.3

26. Isotopes Atoms with the same number of protons but different number of neutrons. In nature most elements contain mixtures of isotopes 23 11 Na : 11 protons, 11 electrons, and 12 neutrons 24 11 Na : 11 protons, 11 electrons, and 13 neutrons

27. Molecules and Ions Chemical Bonds: The forces that hold atoms together in compounds. H 2 O, NO, CO 2 Covalent bonds: Covalent bonds result from atoms sharing electrons. Cl 2 Ionic bonds: Force of attraction between oppositely charged ions. Molecule: A collection of covalently-bonded atoms. H 2, O 2

28. Ions Ions: An ion is an atom or group of atoms that has a net positive charge or negative charge particle (an unequal number of protons and electrons) is obtained by removing or adding electrons. Na +, Cl - Cation: A positive ion (Na +, Mg 2+, NH 4 + ) Anion: A negative ion (Cl -, SO 4 2- )

29. Period Group Alkali Metal Noble Gas Halogen Alkali Earth Metal 2.4

30. A molecule is an aggregate of two or more atoms in a definite arrangement held together by chemical bonds H2H2 H2OH2ONH 3 CH 4 A diatomic molecule contains only two atoms H 2, N 2, O 2, Br 2, HCl, CO A polyatomic molecule contains more than two atoms O 3, H 2 O, NH 3, CH 4 2.5

31. An ion is an atom, or group of atoms, that has a net positive or negative charge. cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. Na 11 protons 11 electrons Na + 11 protons 10 electrons Cl 17 protons 17 electrons Cl - 17 protons 18 electrons 2.5

32. A monatomic ion contains only one atom A polyatomic ion contains more than one atom 2.5 Na +, Cl -, Ca 2+, O 2-, Al 3+, N 3- OH -, CN -, NH 4 +, NO 3 -

33. 13 protons, 10 (13 – 3) electrons 34 protons, 36 (34 + 2) electrons Do You Understand Ions? 2.5 How many protons and electrons are in Al 27 13 ? 3+3+ How many protons and electrons are in Se 78 34 2- ?

34. Atoms, Molecules, and Ions Covalent Bonding (Molecules): The most common type of chemical bond is formed when two atoms share some of their electrons. (non-metal -- non-metal) 34

35. 2.5

36. Formulas Chemical Formula: In which the symbols for the elements are used to indicate the types of atoms present and subscripts are used to indicate the relative numbers of atoms. CO 2 indicates each molecule contains 1 atom of carbon and 2 atoms of oxygen. Structural Formula: In which the individual bonds are shown by lines. It may or may not indicates the actual shape of the molecules. O=C=O

37. 2.6

38. A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance An empirical formula shows the simplest whole-number ratio of the atoms in a substance H2OH2O H2OH2O molecularempirical C 6 H 12 O 6 CH 2 O O3O3 O N2H4N2H4 NH 2 2.6

39. ionic compounds consist of a combination of cations and an anions the formula is always the same as the empirical formula the sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero The ionic compound NaCl 2.6

40. Formula of Ionic Compounds Al 2 O 3 2.6 2 x +3 = +63 x -2 = -6 Al 3+ O 2- CaBr 2 1 x +2 = +22 x -1 = -2 Ca 2+ Br - Na 2 CO 3 2 x +1 = +21 x -2 = -2 Na + CO 3 2-

41. The Modern View of Atomic Structure The atom contains: Electrons: move around the nucleus (mass: 9.11 X 10 -31 kg, Charge 1 - ) Protons: found in the nucleus, they have a positive charge equal in magnitude to the electron’s negative charge (mass: 1.67 X 10 - 27 kg, charge 1 + ) Neutrons: found in the nucleus, virtually same mass as a proton but no charge. (mass: 1.67 X 10 -27 kg, charge: 0)

42. Nuclear Atom Viewed in Cross Section

43. Two Isotopes of Sodium

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