1. Atoms, Molecules, and Ions Chapter 2 BLB 12 th

2. Expectations Recognize important steps in the discovery of the atom and its structure. Work with isotopes. Learn about the periodic table. Differentiate between molecular and ionic compounds. Name compounds (molecular and ionic).

4. 2.1 The Atomic Theory of Matter Of what is matter comprised? Democritus (400 BC) – tiny, indivisible particles, atomos Plato, Aristotle – NOT! Newton (17 th century) – favored atoms as invisible particles Boyle (1660) – gas experiments with pressure & volume Priestly (1774) – isolated oxygen Lavoisier (1789) – Law of Conservation of Mass: Mass is neither created or destroyed. (p.78) Proust (1800) – Law of Definite Proportions (or constant composition): A compound always contains the same proportion of elements. (p. 10)

5. Dalton’s Atomic Theory (1808) 1. Elements are composed of small particles called atoms. 2. All atoms of a given element are identical. 3. Atoms of an element are not changed in a chemical reaction. 4. Compounds are formed when different atoms combine. >> Atoms are the building blocks of matter.<<

6. 2.1 The Atomic Theory of Matter Dalton – Law of Multiple Proportions: element mass proportions in a compound are in a ratio of small whole numbers. (p. 40) Avogadro (1811) – equal volumes of gases contain the same number of particles (p. 401)

7. 2.2 The Discovery of Atomic Structure J.J. Thomson (1897) – cathode ray tube experiments; electrons; charge-to-mass ratio of the electron (1) plum-pudding model of atoms (Fig. 2.9, p. 43) Robert Millikan (1909) – oil-drop experiment; charge and mass of electron (9.10939 x 10 -28 g) Henri Becquerel, Marie Curie (1896, 1899) – radioactivity Ernest Rutherford (1911) – gold foil experiment; nucleus & protons (1919); (2) nuclear model of atom 3 types of radioactivity: α (heaviest, 2+ charge), β (high- speed electrons, 1− charge),  (lightest, high E, 0 charge) James Chadwick (1932) - neutrons Subatomic particles

10. Separation of Radioactive Particles

11. Rutherford’s Gold Foil Experiment pp. 42-43

13. 2.3 The Modern View of Atomic Structure Subatomic particles ParticleFunctionPosition (to nucleus) ChargeMass (kg) electronchemistryoutside−19.11 x 10 -31 protonattract electrons inside+11.67 x 10 -27 neutronnuclear glueinside0 (neutral) 1.67 x 10 -27

14. Atoms Atomic masses ~10 -23 g Atomic diameters (e - cloud) ~10 -10 m = 1 Å Atomic nuclei ~10 -4 Å (very small and dense) Atoms are neutral: # protons = # electrons

15. Practice Exercise 2.1 How many carbon atoms can be placed side by side across the width of a pencil line that is 0.20 mm wide? C atom diameter = 1.54 Å

16. Isotopes, atomic and mass numbers Isotopes – atoms with same number of protons but different numbers of neutrons Nuclide – a single atom of a particular isotope

17. 2.4 Atomic Weights Based on 12 C (assigned a mass of exactly 12 amu) 1 amu = 1.66054 x 10 -24 g (1/12 mass of a 12 C atom) Weighted average atomic mass = Σ(% abundance)(mass of isotope) Atomic mass determined using a mass spectrometer (p. 49)

18. Mass Spectrometer

19. Mass spectrum of Cl

20. Calculate the (weighted) average mass of magnesium (in amu). isotope% abund.Mass (amu) 24 Mg78.9923.98504 25 Mg10.0024.98584 26 Mg11.0125.98259

21. 2.5 The Periodic Table 1 st table developed by Mendeleev and Meyer in 1869 Group, period, regions, group names Physical properties of metals and nonmetals Seaborg (p. 52)

22. Physical Properties Metals High electrical conductivity High thermal conductivity Metallic luster Most are solids Malleable, ductile Metallic bonding Nonmetals Poor electrical conductivity Good heat insulator No metallic luster Solids, liquids, and gases Brittle in solid state Covalently bonded molecules; noble gases monoatomic

24. 2.6 Molecules and Molecular Compounds Chemical bonds – forces that hold atoms together in molecules and compounds covalent bonds – sharing of electrons Molecules – discrete units of covalently bonded atoms; typically nonmetals, e.g. H 2 O, CO 2, NH 3, C 2 H 6 Diatomic elements: H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2 (p. 53) Polyatomic elements: O 3, S 8, P 4 (allotropes – different forms of the same element in the same state, p. 273)

25. Molecules, cont. Representation of molecules, CH 4

26. Empirical & Molecular Formulas Molecular formula – actual number of atoms in a compound Empirical formula – smallest whole number ratio of atoms MolecularEmpirical C2H4C2H4 CH 2 P 4 O 10 P2O5P2O5 H2O2H2O2 HO

27. 2.7 Ions and Ionic Compounds Ionic bond – attraction between oppositely charged ions; results from a transfer of electrons cation – positively charged ion (metals) anion – negatively charged ion (nonmetals) Common ions (Fig. 2.20, p. 56)

28. © 2009, Prentice-Hall, Inc. Ionic Bonds Ionic compounds (such as NaCl) are generally formed between metals and nonmetals.

29. Predicting ionic charges Atoms will lose or gain electrons to attain a noble gas configuration. P3–P3–

30. Ionic Compounds Ionic compounds – consist of ions; form crystal lattices + and − charges balance Formula unit – ratio of cation to anion

31. 2.8 Naming Inorganic Compounds 1957 IUPAC (Int’l Union of Pure and Applied Chemistry) – devised systematic rules for naming compounds Binary compounds – consist of two different elements Don’t capitalize compound or element names.

32. Ionic compounds - cations Cations (Table 2.4, p. 60) 1. Single metal, single charge Na +, sodium ion Al 3+, aluminum ion 2. Single metal, multiple charges Cr 2+, chromium(II) ion Cr 4+, chromium(IV) ion 3. Polyatomic ions

34. Ionic compounds - anions Anions (Table 2.5, p. 63) 1. Monoatomic, -ide ending Cl -, chloride ion O 2-, oxide ion 2. Oxyanions NO 3 -, nitrate ion NO 2 -, nitrite ion 3. H + + oxyanion

35. P 3– Phosphide ion PO 3 3- Phosphite ion NO 2 – Nitrite ion

36. Polyatomic Ions to Memorize (formula & charge) NameFormulaNameFormula acetate ionC2H3O2-C2H3O2- hypochlorite ionClO - ammonium ionNH 4 + nitrate ionNO 3 - carbonate ionCO 3 2- nitrite ionNO 2 - chlorate ionClO 3 - perchlorate ionClO 4 - chlorite ionClO 2 - permanganate ionMnO 4 - chromate ionCrO 4 2- phosphate ionPO 4 3- cyanide ionCN - phosphite ionPO 3 3- hydroxide ionOH - sulfate ionSO 4 2- sulfite ionSO 3 2-

37. Ionic compounds Cation first, anion second Charges (+ and -) must balance

38. Acids (p. 64) Acid – substance which produces a H + when dissolved in water If anion ends in ____, acid ends with ____. -ide-ic -ate-ic -ite-ous

39. Molecular Compounds Name as written in formula. Prefixes denote number of each atom. Exceptions: H 2 O water NH 3 ammonia CH 4 methane

40. 2.9 Some Simple Organic Compounds Hydrocarbon – contain only C and H Alkanes – saturated hydrocarbons with only C−C single bonds Alkane derivatives: −OHalcohol −COOHcarboxylic acid −COOC−ester −COC−ketone

41. Organic compounds, cont. Unsaturated hydrocarbons: Alkenes – contain at least one C=C double bond Alkynes – contain at least one C≡C triple bond