1 Chapter 2 Atoms, Molecules and Ions Preview: Fundamental Chemical Laws and Atom. Modern View of Atomic Structure, Molecules, and Ions. Periodic Table. Naming Simple compounds, Ionic compounds, Formula from names.

2 The Early History of Chemistry 4 Before 16th Century – Alchemy: Attempts (scientific or otherwise) to change cheap metals into gold 4 17th Century –Robert Boyle: First “chemist” to perform quantitative experiments 4 18th Century –George Stahl: Phlogiston flows out of a burning material. –Joseph Priestley: Discovers oxygen gas, “dephlogisticated air.”

3 Law of Conservation of Mass 4 Discovered by Antoine Lavoisier 4 Mass is neither created nor destroyed 4 Combustion involves oxygen, not phlogiston

4 Other Fundamental Chemical Laws 4 A given compound always contains exactly the same proportion of elements by mass. 4 Carbon tetrachloride is always 1 atom carbon per 4 atoms chlorine: CCl 4 Law of Definite Proportion

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

7 Mass of Nitrogen That Combines With 1 g Oxygen Compound A1.750 g Compound B g Compound C g A/B = 1.750/0.8750= 2/1 B/C = 0.875/ = 2/1 A/C = 1.750/ = 4/1 i.e. amount of nitrogen in A is twice that in B, etc.

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10 Dalton’s Atomic Theory (1808) ÊEach element is made up of tiny particles called atoms. ËThe atoms of a given element are identical; the atoms of different elements are different in some fundamental way or ways.

11 Dalton’s Atomic Theory (continued) ÌChemical compounds are formed when atoms combine with each other. A given compound always has the same relative numbers and types of atoms. ÍChemical reactions involve reorganization of the atoms - changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction.

12 Dalton's theory lead to: 1gm hydrogen + 8gm of oxygen water he assumed that water formula is "OH" and the mass of hydrogen is "1" and of oxygen is "8". Using the same concepts, Dalton's proposed the first table of atomic masses. It has been proved later that Dalton's table contain incorrect. Gay-lussac ( ) found experimentally that:

13 Figure 2.4: A representation of some of Gay-Lussac's experimental results on combining gas volumes. Interpreted in 1811 by Avogadro

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

15 Figure 2.5: A representation of combining gases at the molecular level. The spheres represent atoms in the molecules.

16 Early Experiments to Characterize the Atom H J. J. Thomson - postulated the existence of electrons using cathode ray tubes. Measured charge /mass of e- Received 1906 Nobel Prize in Physics e = x 10 8 C/g m

17 Figure 2.7: A cathode-ray tube. The fast-moving electrons excite the gas in the tube, causing a glow between the electrodes.

18 2.2

19 (Uranium compound)

20 Figure 2.9: Thomson plum pudding model of the atom.

21 H Ernest Rutherford - explained the nuclear atom, containing a dense nucleus with electrons traveling around the nucleus at a large distance.

22 Figure 2.12: Rutherford's experiment on  -particle bombardment of metal foil.

23 Figure 2.13: (a) The expected results of the metal foil experiment if Thomson's model were correct. (b) Actual results.Actual results

24 Millikan oil drop experiment Millikan did another experiment to determine the mass of the –ve particles (electrons). The experiment used mainly to determine the magnitude of the electron charge and using e/m to get m- value.value

25 The Modern View of Atomic Structure l electrons l protons: found in the nucleus, they have a positive charge equal in magnitude to the electron’s negative charge. l neutrons: found in the nucleus, virtually same mass as a proton but no charge. The atom contains:

26 Figure 2.14: A nuclear atom viewed in cross section. Note that this drawing is not to scale.

27 Figure 2.15: Two isotopes of sodium. Both have eleven protons and eleven electrons, but they differ in the number of neutrons in their nuclei.

28 The Mass and Charge of the Electron, Proton, and Neutron

29 Summary J.J. Thompson (1897) “cathode rays are electrons” (e – ) and finds e/m ratio Robert Millikan (1909) measures e and hence m electron known at 9.11  kg E. Rutherford (1906) bounces  (He 2+ ) off Au tissue proving protons (p + ) in nucleus F.A. Aston (1919) “weighs” atomic ions J. Chadwick (1939) observes neutrons (no charge) by decomposition (to p +, e –, and ).

30 The Chemists’ Shorthand: Atomic Symbols K  Element Symbol Mass number  Atomic number 

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32 Q6. In Rutherford’s experiment, most of the alpha particles directed at the thin metal foil, A)passed directly through the foil without being deflected. B)were reflected directly back from the foil. C)were deflected at a 90º angle. D)were absorbed by the foil. E)converted into high-energy gamma particles.

33 Chemical Bonds The forces that hold atoms together in compounds. Covalent bonds result from atoms sharing electrons. Molecule: a collection of covalently-bonded atoms.

34 The Chemists’ Shorthand: Formulas Chemical Formula: Symbols = types of atoms Subscripts = relative numbers of atoms CO 2 Structural Formula: Individual bonds are shown by lines. O=C=OO=C=O

35 Figure 2.16: The structural formula for methane.

36 Figure 2.17: Space-filling model of methane. This type of model shows both the relative sizes of the atoms in the molecule and their spatial relationships.

37 Figure 2.18: Ball-and-stick model of methane.

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39 Ions Cation: A positive ion Mg 2+, NH 4 + Anion: A negative ion Cl , SO 4 2  Ionic Bonding: Force of attraction between oppositely charged ions.

40 Periodic Table Elements classified by:  properties  atomic number Groups (vertical) 1A = alkali metals 2A = alkaline earth metals 7A = halogens 8A = noble gases Periods (horizontal)

41 Earth’s Most Abundant Elements Oxygen 0 46% Silicon Si 27% Aluminum Al 8% Iron Fe 5% Calcium Ca 3% Sodium Na 2.8% Potassium K 2.5% Magnesium Mg 2.0%

42 Figure 2.21: The Periodic Table.

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44 Naming Compounds 1. Cation first, then anion 2. Monatomic cation = name of the element Ca 2+ = calcium ion 3. Monatomic anion = root + -ide Cl  = chloride CaCl 2 = calcium chloride HI = hydrogen iodide Binary Ionic Compounds:

45 Figure 2.19: Sodium metal reacts with chlorine gas to form solid sodium chloride.

46 Naming Compounds (continued)  metal forms more than one cation  use Roman numeral in name PbCl 2 Pb 2+ is cation PbCl 2 = lead (II) chloride Binary Ionic Compounds (Type II):

47 FeCl 2 2 Cl - -2 so Fe is +2 iron(II) chloride FeCl 3 3 Cl - -3 so Fe is +3 iron(III) chloride Cr 2 S 3 3 S so Cr is +3 (6/2)chromium(III) sulfide

48 Figure 2.22: The common cations and anions

49 Naming Compounds (continued)  Compounds between two nonmetals  First element in the formula is named first.  Second element is named as if it were an anion.  Use prefixes  Never use mono- P 2 O 5 = diphosphorus pentoxide Binary compounds (Type III):

50 NF 3 nitrogen trifluoride SO 2 sulfur dioxide N 2 Cl 4 dinitrogen tetrachloride NO 2 nitrogen dioxide N2ON2Odinitrogen monoxide Molecular Compounds 2.7 TOXIC ! Laughing Gas

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55 Figure 2.23: A flowchart for naming binary compounds.

56 Figure 2.24: Overall strategy for naming chemical compounds.

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59 Figure 2.25: A flowchart for naming acids. An acid is best considered as one or more H+ ions attached to an anion.

60 Naming Exercise Al 2 (S 2 O 3 ) 3 P 4 O 10 Cu(NO 2 ) 2 NaMnO 4 CS 2 Fe 2 (CrO 4 ) 3 HCl (gas) PH 4 BrO 2 Aluminum thiosulfate Tetraphosphorous decaoxide Copper(II) nitrite Sodium permanganate Carbon disulfide Iron(III) chromate Hydrogen chloride Phosphonium bromite

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