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Dalton’s Atomic Theory

Dalton’s Atomic Theory. Elements - made up of atoms Same elements, same atoms. Different elements, different atoms. Chemical reactions involve bonding of atoms. Law of Definite Composition. A compound always contains the same proportion of elements by mass. Law of Multiple Proportions.

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Dalton’s Atomic Theory

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  1. Dalton’s Atomic Theory • Elements - made up of atoms • Same elements, same atoms. • Different elements, different atoms. • Chemical reactions involve bonding of atoms

  2. Law of Definite Composition • A compound always contains the same proportion of elements by mass

  3. Law of Multiple Proportions • Compounds form from specific combinations of atoms • H2O vs H2O2

  4. Chemical Bonds • Holds compounds together • Need to be broken for chemical and physical changes to occur

  5. The Atom • Made up of: • Protons – (+) charged • Electrons – (-) charged • neutrons

  6. Periodic Table • Alkaline Metals – Grps. I & II • Transition Metals • Non-metals • Halogens – Group VII • Noble Gases –Group VIII - little chemical activity

  7. Periodic Table • Atomic Mass - # at bottom • how much element weighs • Atomic Number - # on top • gives # protons = # electrons

  8. Periodic Table • Atomic Mass • number below the element • not whole numbers because the masses are averages of the masses of the different isotopes of the elements

  9. Ions • Are charged species • Result when elements gain electrons or lose electrons

  10. 2 Types of Ions • Anions – (-) charged • Example: F- • Cations – (+) charged • Example: Na+

  11. Highly Important! • Gain of electrons makes element (-) = anion • Loss of electrons makes element (+) = cation

  12. Charges • When elements combine, they have to be in the form of IONS. • Cations and anions combine to form compounds. • For a neutral compound, the sum of the charges must be ZERO. • For a polyatomic ion, the sum of the charges must equal the charge of the ION.

  13. Examples • In CO2, the charge of C is + 4 • In CO, the charge of C is +2. • In KMnO4, since the charge of K is +1, O is -2 so -2 x 4 = -8, Mn must be +7. • In (PO4)3-, the charge of O is -2, so -2 x 4= -8, then P must have a charge of +5, so the sum when the charges are added will be -3.

  14. Isotopes • Are atoms of a given element that differ in the number of neutrons and consequently in atomic mass.

  15. Example Isotopes % Abundance 12C 98.89 % 13C 1.11 % 14C 11C

  16. For example, the mass of C = 12.01 a.m.u is the average of the masses of 12C, 13C and 14C.

  17. Determination of Aver. Mass • Ave. Mass = [(% Abund./100) (atomic mass)] + [(% Abund./100) (atomic mass)]

  18. Take Note: • If there are more than 2 isotopes, then formula has to be re-adjusted

  19. Sample Problem 1 • Assume that element Uus is synthesized and that it has the following stable isotopes: • 284Uus (283.4 a.m.u.) 34.6 % • 285Uus (284.7 a.m.u.) 21.2 % • 288Uus (287.8 a.m.u.) 44.20 %

  20. Solution • Ave. Mass of Uus = • [284Uus] (283.4 a.m.u.)(0.346) • [285Uus] +(284.7 a.m.u.)(0.212) • [288Uus] +(287.8 a.m.u.)(0.4420) • = 97.92 + 60.36 + 127.21 • = 285.49 a.m.u (FINAL ANS.)

  21. Periodic Table • Mendeleev – arranged elements in the (.) table

  22. Periodic Table • Atomic Mass • number below the element • not whole numbers because the masses are averages of the masses of the different isotopes of the elements

  23. For example, the mass of C = 12.01 a.m.u is the average of the masses of 12C, 13C and 14C.

  24. Oxidation Numbers • Is the charge of the ions (elements in their ion form) • Is a form of electron accounting • Compounds have total charge of zero (positive charge equals negative charge)

  25. Oxidation States • Are the partial charges of the ions. Some ions have more than one oxidation states.

  26. Oxidation States • - generally depend upon the how the element follows the octet rule • Octet Rule – rule allowing elements to follow the noble gas configuration

  27. Nomenclature • - naming of compounds

  28. Periodic Table • Rows (Left to Right) - periods • Columns (top to bottom) - groups

  29. Rule 1 – IONIC COMPOUNDS • Metals w/ Fixed Oxidation States • Name metal or first element asis - Anion always ends in “–ide”

  30. Terminal element or anion • O - oxide P - phosphide • N - nitride Se - selenide • S - sulfide Cl - chloride • F - fluoride I - iodide • Br - bromide C - carbide

  31. Note • Only elements that come directly from the periodic table WILL end in –IDE. • POLYATOMIC IONS will be named AS IS.

  32. Name the following: • CaO- • NaCl - • MgO - • CaS - • Na3N -

  33. Answers: • CaO- calcium oxide • NaCl - sodium chloride • MgO - magnesium oxide • CaS - calcium sulfide • Na3N - sodium nitride

  34. Where do the subscripts come from? • Answer: From the oxidation states of the ions. • Remember: Ions are the species that combine. • Target: Compounds! (No charges!)

  35. Second Rule • II. Ionic Compounds - Metals with no fixed oxidation states (Transition Metals) except for Ag, Zn and Al • Metal(Roman #) + 1st syllable + ide • Use Roman numerals after the metal to indicate oxidation state

  36. Name the following: • Copper (I) sulfide • Iron (II) oxide • Tin (II) iodide • Iron (III) nitride

  37. Answers: • Copper (I) sulfide Cu2S • Iron (II) oxide FeO • Tin (II) iodide SnI2 • Iron (III) nitride FeN

  38. What about…….? • Cesium hydroxide • Iron (III) acetate • Lithium phosphate • Aluminum Sulfite • Lead (II) sulfate • Silver nitrate

  39. POLYATOMIC IONS • Consist of more than 1 element. • Have charges. • Ex. SO42-, SO32-, PO43-,PO33-

  40. Rule 3 – Covalent Compounds • III. For Non-metals (grps IV, V, VI VII), use prefixes. Mono – 1 Hepta - 7 Di - 2 Octa - 8 Tri – 3 Nona - 9 Tetra – 4 Deca - 10 Penta – 5 Hexa - 6

  41. Rule 3 – Covalent Compounds (only have Non- Metals) • Name 1st element as is. Use prefix, if necessary. • Prefix + 1st element + prefix + 1st syllable of anion + ide

  42. Name the following compounds • CO2 - carbon dioxide • N2O – dinitrogen oxide • SO3 – sulfur trioxide • N2O5 – dinitrogen pentoxide • P2S5 – diphosphorus pentasulfide • CO – carbon monoxide

  43. Naming Acids • I. Acids without Oxygen • Use hydro + 1st syllable + “- ic acid” • Example: HCl = hydrochloric acid HCN = hydrocyanic acid HBr = hydrobromic acid

  44. II. Acids with oxygen • Polyatomic “ate” converts to “ic” + acid • Polyatomic “ite” converts to “ous” + acid - H2SO3 sulfurous acid • H2SO4 sulfuric acid • HNO3 nitric acid • HNO2 nitrous acid • H3PO4 phosphoric acid

  45. Trick! • If anion ends in “ – ate”, acid ends in “ – ic” • Example: • HClO4 perchlorate perchloric acid • HClO3 chlorate chloric acid

  46. Trick! • If anion ends in “ – ite”, acid ends in “ – ous” • Example: • HClO2 chlorite chlorous acid • HClO hypochlorite hypochlorous acid

  47. Name the following: • HBrO4 (perbromate) • HBrO3 (bromate) • HBrO2 (bromite) • HBrO (hypobromite)

  48. Fundamental laws • Law of Conservation of Mass • Mass is neither created or destroyed • Conversion from one form to another

  49. Determination of Aver. Mass • Ave. Mass = [(% Abund./100) (atomic mass)] + [(% Abund./100) (atomic mass)]

  50. Sample Problem 1 • Assume that element Uus is synthesized and that it has the following stable isotopes: • 284Uus (283.4 a.m.u.) 34.6 % • 285Uus (284.7 a.m.u.) 21.2 % • 288Uus (287.8 a.m.u.) 44.20 %

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