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Chapter 2: Atoms, Molecules, and Ions

Chapter 2: Atoms, Molecules, and Ions

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Chapter 2: Atoms, Molecules, and Ions

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  1. Chapter 2: Atoms, Molecules, and Ions Renee Y. Becker Valencia Community College CHM 1045

  2. Dalton’s Atomic Theory • Elements are made of tiny particles called atoms. • Each element is characterized by the mass of its atoms. Atoms of the same element have the same mass, but atoms of different elements have different masses. • Chemical reactions only rearrange the way atoms are combined; the atoms themselves are unchanged.

  3. The Structure of Atoms • Cathode-Ray Tube (Thomson, 1856–1940): Cathode rays consist of tiny negatively charged particles, now called electrons.

  4. The Structure of Atoms • Deflection of electron depends on three factors: • Strength of electric or magnetic field • Size of negative charge on electron • Mass of the electron • Thomson calculated the electron’s charge to mass ratio as 1.758820 x 108 Coulombs per gram.

  5. The Structure of Atoms • Oil Drop Experiment (Millikan, 1868–1953): Applied a voltage to oppose the downward fall of charged drops and suspend them. • Voltage on plates place 1.602176 x 10-19 C of charge on each oil drop. • Millikan calculated the electron’s mass as 9.109382 x 10-28 grams.

  6. The Structure of Atoms • Discovery of Nucleus:Rutherford irradiated gold foil with a beam of alpha particles to search for positive charged particles. Most of the particles passed through but some were deflected at large angles, why?

  7. The Structure of Atoms

  8. Periods: Seven horizontal rows Groups: 18 vertical columns, based on similar chemical properties

  9. Periodic Table

  10. The Structure of Atoms Standard Format

  11. The Structure of Atoms • Isotopes: Atoms with identical atomic numbers, but different mass numbers. • Average Isotopic Mass: A weighted average of the isotopic masses of an element’s naturally occurring isotopes. • Atomic Mass: A weighted average of the isotopic masses of an element’s naturally occurring isotopes.

  12. Example 1: Periodic Table What are the atomic numbers for the following elements? • Copper • Sodium • Sulfur • Oxygen

  13. Example 2: Periodic Table What are the atomic masses for the following elements? • Iron • Magnesium • Bromine • Xenon

  14. Example 3: Periodic Table What are the mass numbers for the following elements? • Chlorine • Nitrogen • Carbon • Zinc

  15. 75 75 Se Se 34 34 Example 4: The Structure of Atoms • The isotope is used medically for diagnosis of pancreatic disorders. How many protons, neutrons, and electrons does an atom of have?

  16. Example 5: The Structure of Atoms • An atom of element X contains 47 protons and 62 neutrons. Identify the element, and write the symbol for the isotope in the standard format.

  17. 35 37 Cl Cl 17 17 Example 6: The Structure of Atoms • Chlorine has two naturally occurring isotopes: with an abundance of 75.77% and an isotopic mass of 34.969 amu, and with an abundance of 24.23% and an isotopic mass of 36.966 amu. What is the atomic mass of chlorine?

  18. Example 7: Periodic Table

  19. 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)

  20. Atoms, Molecules, and Ions Naming Binary Molecular Compounds: • The more cationlike element uses its elemental name. • The more anionlike element substitutes the second half of its elemental name with –ide. • Use the Greek prefixes to express the number of each element present.

  21. Atoms, Molecules, and Ions

  22. Greek Prefixes Nona- 9 Deca- 10

  23. Example 8: Atoms, Molecules, and Ions Examples: CO carbon monoxide CO2 carbon dioxide SF4 sulfur tetrafluoride Name: NCl3 P4O6 S2F2 Write formulas: Disulfur dichloride Iodine monochloride Nitrogen trioxide

  24. Atoms, Molecules, and Ions • Ionic Bonding (Ionic Solids): These are formed by a transfer of one or more electrons from one atom to another. (metal -- non-metal)

  25. Example 9 Which of the following drawings represents an ionic compound? Molecular compound?

  26. Atoms, Molecules, and Ions Naming Binary Ionic Compounds: • Identify the positive ion and then the negative ion. • The positive ion uses its elemental name. • The negative ion substitutes the second half of its elemental name with –ide. • Do not use Greek prefixes such as mono–, di–, or tri–. • Use roman numerals for transition metals

  27. Atoms, Molecules, and Ions

  28. Atoms, Molecules, and Ions

  29. Example 10: Name • NaCl • MgS • Ba3N2 • CaO • K2S • FeCl2 • FeCl3 • CrO2 • ZnCl2 • V2O3

  30. Example 11: Draw • Calcium chloride • Copper (II) sulfide • Sodium nitride • Silver bromide • Nickel (II) phosphide • Cesium oxide • Strontium iodide • Cobalt (I) sulfide

  31. Atoms, Molecules, and Ions

  32. Atoms, Molecules, and Ions • Naming Ionic Compounds Containing Polyatomic Ions : • Same as binary ionic compounds • But use the name provided for the polyatomic ion

  33. Example 12: Atoms, Molecules, and Ions • Examples • CaCO3 Calcium carbonate • FeCrO4 Iron(II) chromate • KOH Potassium hydroxide Name: Write Formulas: Ba3(PO4)2 Iron(II) permanganate Na2SO4 Cesium nitrate Sn(ClO4)4 Zinc acetate

  34. Example 13: Which of the following is a covalent compound? • NaCl • NaOH • H2O • AlCl3

  35. Atoms, Molecules, and Ions Acid: A substance that provides H+ ions in H2O Base: A substance that provides OH- in H2O Oxoacid: Contain oxygen and hydrogen and another element

  36. Atoms, Molecules, and Ions Naming acids: When acid is dissolved in water gives one or more H+ and a polyatomic oxoanion, (has to have (aq)) Name of acid is based on the oxoanion Get pink sheet out!!

  37. Atoms, Molecules, and Ions

  38. Example 14: Atoms, Molecules, and Ions Name the following acids: (a) HBrO3(aq) (b) HCN(aq)(c) HIO3(aq)(d) HMnO4(aq)(e) H2CrO4(aq)

  39. Balancing Chemical Equations • A balanced chemical equation represents the conversion of the reactants to products such that the number of atoms of each element is conserved. Calcium carbonate  calcium oxide + carbon dioxide CaCO3(s)  CaO(s) + CO2(g)

  40. Balancing Chemical Equations • Balancing Equations: write unbalanced equation A2 + B2 A2B • Use coefficients to indicate how many formula units are required to balance the equation: 2 A2 + B22 A2B

  41. Balancing Chemical Equations • Method 1 (suggested) • Balance those atoms which occur in only one compound on each side • Balance remaining atoms • Reduce coefficients to smallest whole integers • Check your answer

  42. Balancing Chemical Equations • Method 2 • Identify most complex compound • Balance this compound by placing 1 before it • Balance remaining compounds using fractions • Multiply fractions to obtain integers

  43. Example 15: Balancing Chemical Equations • Balance the following equations C6H12O6 C2H6O + CO2 Fe + O2 Fe2O3 NH3 + Cl2 N2H4 + NH4Cl KClO3 + C12H22O11 KCl + CO2 + H2O