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Atoms, Molecules, and Ions

Atoms, Molecules, and Ions

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Atoms, Molecules, and Ions

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

  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).

  3. 2.1 The Atomic Theory of Matter Of what is matter comprised? • Democritus (400 BC) – tiny, indivisible particles, atomos • Plato, Aristotle – NOT! • Newton (17th 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)

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

  5. 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)

  6. 2.2 The Discovery of Atomic Structure Subatomic particles • 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), g (lightest, high E, 0 charge) • James Chadwick (1932) - neutrons

  7. Separation of Radioactive Particles

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

  9. 2.3 The Modern View of Atomic Structure • Subatomic particles

  10. 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

  11. 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

  12. 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

  13. 2.4 Atomic Weights • Based on 12C (assigned a mass of exactly 12 amu) • 1 amu = 1.66054 x 10-24 g (1/12 mass of a 12C atom) • Weighted average atomic mass = Σ(% abundance)(mass of isotope) • Atomic mass determined using a mass spectrometer (p. 49)

  14. Mass Spectrometer

  15. Mass spectrum of Cl

  16. Calculate the (weighted) average mass of magnesium (in amu).

  17. 2.5 The Periodic Table • 1st table developed by Mendeleev and Meyer in 1869 • Group, period, regions, group names • Physical properties of metals and nonmetals • Seaborg (p. 52)

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

  19. 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. H2O, CO2, NH3, C2H6 Diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2(p. 53) Polyatomic elements: O3, S8, P4 (allotropes – different forms of the same element in the same state, p. 273)

  20. Molecules, cont. • Representation of molecules, CH4

  21. Empirical & Molecular Formulas • Molecular formula – actual number of atoms in a compound • Empirical formula – smallest whole number ratio of atoms

  22. 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)

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

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

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

  26. 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.

  27. Ionic compounds - cations • Cations (Table 2.4, p. 60) • Single metal, single charge • Na+, sodium ion • Al3+, aluminum ion • Single metal, multiple charges • Cr2+, chromium(II) ion • Cr4+, chromium(IV) ion • Polyatomic ions

  28. Ionic compounds - anions • Anions (Table 2.5, p. 63) • Monoatomic, -ide ending • Cl-, chloride ion • O2-, oxide ion • Oxyanions • NO3-, nitrate ion • NO2-, nitrite ion • H+ + oxyanion

  29. NO2– Nitrite ion PO33- Phosphite ion P3– Phosphide ion

  30. Polyatomic Ions to Memorize(formula & charge)

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

  32. 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

  33. Molecular Compounds • Name as written in formula. • Prefixes denote number of each atom. • Exceptions: H2O water NH3 ammonia CH4 methane

  34. 2.9 Some Simple Organic Compounds • Hydrocarbon – contain only C and H • Alkanes – saturated hydrocarbons with only C−C single bonds • Alkane derivatives: −OH alcohol −COOH carboxylic acid −COOC− ester −COC− ketone

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