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Unit 2 : Elements and Compounds : Atoms, Molecules & Ions

CHM 1045 : General Chemistry and Qualitative Analysis. Unit 2 : Elements and Compounds : Atoms, Molecules & Ions. Dr. Jorge L. Alonso Miami-Dade College – Kendall Campus Miami, FL. Textbook References : Modules #2 & 5. The Early Development of the Atomic Theory. Ancient Atomic Theory.

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Unit 2 : Elements and Compounds : Atoms, Molecules & Ions

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  1. CHM 1045: General Chemistry and Qualitative Analysis Unit 2:Elements and Compounds: Atoms, Molecules & Ions Dr. Jorge L. Alonso Miami-Dade College – Kendall Campus Miami, FL • Textbook References: • Modules #2 & 5

  2. The Early Development of the Atomic Theory

  3. Ancient Atomic Theory Leucippus of Miletus & Democritus of Abdera (Gk. 5th Cent BC) • More philosophical than experimental in origin. • Matter is made up of very small individual atomos- objects that are indivisible. • Everything is made up of these atoms, which move around in a void (a vacuum). • The different physical properties -- color, taste, and so on -- of materials come about because atoms in them are different shapes and/or arrangements and orientations with respect to each other.

  4. Medieval Alchemy(الخيمياء, al-khimia)Science as an early form of investigation, with occult philosophical and spiritual traditions. Merlin the Magician Jabir ibn Hayyan • Principal aim of Alchemist: • the transmutationofcommon metals intogoldorsilver. • Cinnabar (red powder)  Hg • Zn, Cu, Fe  Au or Ag • the creation of a "panacea," or the elixir of life, a remedy that supposedly would cure all diseases and prolong life indefinitely.

  5. Chemistry in the Age of Enlightment Law of Conservation of Mass: (1743 - 1794) (s) (l) (g) 27 g  25 g + 2 g {HgOMovie}

  6. * Law of Constant Composition (or Definite Proportions) In Water Ratio H :O 1 : 8 Compared masses of different elements within the same compound. In H-Peroxide Ratio H : O 1 :16 (1754–1826)

  7. Dalton’s Law of Multiple Proportion John Dalton (1766–1844). When two elements form two different compounds, the mass ratio of the elements in one compound is related to the mass ratio in the other compound by a small whole number. C + O2 (high oxy.conc.)  CO2 C + O2 (low oxy.conc.) CO

  8. * Dalton’s Atomic Theory (1801) • All matter is made up of small indivisible particles called atoms. • The properties of the atoms of one element differ from those of all other elements. • Atoms can neither be created nor destroyed. • All atoms of the same element are identical in mass, size, and physical properties. • Atoms combine in small whole number ratios to form compounds. • All matter is made up of small indivisible particles called atoms. • The properties of the atoms of one element differ form those of all other elements. • Atoms can neither be created nor destroyed. • All atoms of the same element are identical in mass, size, and physical properties. • Atoms combine in small whole number ratios to form compounds.

  9. The Electron Electrically charged particles can be rubbed-out of many substances such as glass rods, hair, shoes, rubber tires and shoes. Excess of electrons Rubber band rubs metals inside {Electroscope Movie} {Electroscope Movie*} Lack of electrons Rubber band Electric motor Van de Graaff Generator

  10. The Electron Batteries (chemicals) Electron Rays • cathode ray tubes produce negatively charged particles (electrons) from chemicals in batteries. • J. J. Thompson is credited with their discovery (1897). Movie 1 Movie 2 Movie 3

  11. The Atom, circa 1900 • “Plum pudding” model, put forward by J. J. Thompson. • Positive sphere of matter with negative electrons imbedded in it.

  12. 4 2 He * Discovery of the Nucleus In 1909 ErnestRutherford shot  particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.

  13. The Nuclear Atom Since some particles were deflected at large angles, J. J. Thompson’s model could not be correct. {*Rutherfords Experiment}

  14. The Nuclear Atom According to Rutherford • He postulated a very small, dense nucleus with the electrons around the outside of the atom. • Most of the volume of the atom is empty space. Difference of 105 = 100,000

  15. Historical Development of Atomic Theory {Bohr Planetary vs. Quantum Model}

  16. Scanning Tunneling Microscopy (STM)

  17. ‘Seeing’ Atoms: Scanning Tunneling Microscope

  18. ‘Seeing’ Atoms: Scanning Tunneling Microscope

  19. Subatomic Particles: Charge & Mass • Protons were discovered by Rutherford in 1919. • Neutrons were discovered by James Chadwick in 1932. {ProtonDiscoveryMovie} • Protons and neutrons have essentially the same mass. • The mass of an electron is so small we ignore it.

  20. Elements IA VIIIA Pure substances that cannot be broken down into more elemental particles by ordinary chemical means. IIA IIIA IVA VA VIA VIIA A. Atomic number B. Mass number C. Isotopes D. Atomic mass or weigh IIB IVB VB VIB VIIB VIII IB IIB

  21. Elements & Atomic terminology • Atomic number(Z) = #p • Mass number (A) = (#p+) + (#no) A = Z + N • Isotopes • Atomic mass or weigh (a.m.u. or grams) #p = #e- in a neutral atom Identifies the element Atoms of the same element (same at. #), having different number of neutrons. The average mass of the isotopes of an element, considering their natural % abundance.

  22. * 11 6 12 6 13 6 14 6 C C C C Isotopes (atomic number =1) 99.985% 0.015% 0% Natural Abundance • Atoms of the same element with different masses (mass number) • Isotopes have different numbers of neutrons. (atomic number =2) 0.000137% 99.999863% Natural Abundance (atomic number =3) Atomic mass or weigh? 7.59% 92.41% Natural Abundance Isotopes of Carbon: (isotope notation) Mass number Atomic number Natural Abundance: negligible 98.89% 1.11% negligible

  23. Isotopes of Hydrogen 1H 2H 3H Natural Abundance 99.985% 0.015% negligible unstable, radioactive

  24. Isotopes of Carbon 14C 13C 12C Natural Abundance 98.89% 1.11% negligible

  25. Table of Isotopes (partial)

  26. Determination of Atomic Mass Mass spectrometer: can separate isotopes of an element based on their charge and mass, & measure their % abundance. South pole (-) of magnet attracts lighter isotope more easily than heavier isotope. Ionizing chamber

  27. Mass Spectrometer Mass Spectrum of Germanium (Ge)

  28. * Atomic Mass (Weight): the average mass of isotopes of an element, considering their natural abundance Isotope% Abund. f.Abund. X Mass # = 3He 0.01 (0.0001 x 3) = 0.0003 4He 99.99 (0.9999 x 4) = +3.9996 4.000 35Cl 75.77 37Cl 24.23 24Mg 78.99 25Mg 10.00 26Mg 11.01 Atomic Mass (Weigh): AM = (f1 x M#1) + (f2 x M#2) + …

  29. 200? 2007 (B)#2 More exact Atomic Mass (Weigh): AM = (f1 x M#1) + (f2 x M#2) + …

  30. Properties of Metal, Nonmetals,and Metalloids

  31. Organization or the Periodic Table: Groups (Families) The following four groups are known by their names: Representative Elements Representative Elements Transition Metals

  32. Diatomic Molecules of Elements • Alonso’s Rule of 7 + 1: • Start with element #7, Nitrogen, trace a 7 and count 7 elements. • Plus 1 more element, #1 Hydrogen At2 These seven + one elements occur naturally as molecules containing two atoms (diatomic). In compounds they may combine in other ratios. NaCl, BaCl2, AlCl3,CCl4.

  33. Compounds • Pure substances (cannot be separated by physical means). • Compounds can be broken down (decomposed) into more elemental particles (elements) by ordinary chemical means. Molecular Formulas: H2O CO2 H2O2 CO CH4 Structural Formulas: (Space-filling )

  34. Classification of Compounds Covalent (Molecular) Compounds: Non-Metals + Non-Metals. Acids: Hydrogen + Nonmetals (polar covalent) Ionic Compounds Salts: Metal + Non-Metal Bases: Metals + Hydroxide Ion (OH-) Acid Salts: Metal + Acid (Hydrogen + Nonmetal) Organic Compounds: covalent compounds containing carbon (C) atom chains, with mostly H & O atoms attached to the chain.

  35. * Covalent (Molecular) Compounds: H2O(g) • Composed of a Non-Metal combined with another Non-Metal. • Are mostly gases, liquids, and sometimes amorphous solids. • Have low melting points • Bonded atoms share electrons. H2O(s) H2O(l)

  36. Nomenclature:Molecular Compound Prefix*- (name of 1st Element) Prefix- (root of 2nd Element) - IDE * Prefix mono- not used for 1st element Molecular Formulas: H2O CO2 H2O2 CO CH4 Structural Formulas: (Space-filling )

  37. * Ionic Compounds

  38. Ionic Compounds (Salts): • Composed of a Metal ion (cation, M+) combined with an Non-Metal ion (anion, N-); atoms exchange transfer electrons. • Are CrystallineSolids. • Have high melting points • Smallest component particle is called a formula unit, not a molecule. Cations Anions

  39. Ionic Nomenclature: Binary Salts 1+ (A) Ions with Fixed Charges (OxidationNumbers) 1- 2+ 3+ 3- 2- Zn2+ Ag1+ Cd2+ Aluminum Nitride Aluminum Oxide Aluminum Bromide Al3+ N3-  AlN Al3+ O2-  Al2O3 Al3+ Br-  AlBr3 Name of 1st Element Root of 2nd Element -IDE

  40. Ionic Nomenclature: Binary Salts (B) Ions with Variable Oxidation Numbers(Mostly Transition Metals) Iron (II) Nitride Iron (III) Nitride Fe2+ N3-  Fe3N2 Fe3+ N3-  FeN IUPAC Nomenclature: Traditional Names: Name of 1st Element (Roman Numeral)(or–ous, -ic) Root of 2nd Element -IDE

  41. Ionic Nomenclature: Binary Salts (B) Ions with Variable Oxidation Numbers • Iron (Ferrum): Fe 2+ (Iron II or Ferrous) Fe 3+ (Iron III or Ferric) • Copper (Cuprum): Cu 1+ (Copper I or Cuprous) Cu 2+ (Copper II or Cupric) • Mercury (Hydragyrum): Hg 1+ (Mercury I or Mercurous) Hg 2+ (Mercury II or Mercuric) Exceptions

  42. Ionic Nomenclature: Salts with Polyatomic Ions Common Representative “-ate” Oxyanions Na+ Na+ Na+ ClO3- chlorate Na+

  43. Ionic Nomenclature: Salts with Polyatomic Ions Common Representative “-ate” Oxyanions Name of 1st Element (Roman Numeral or –ous, -ic) Name of Polyatomic ion (all end in -ATE)

  44. Polyatomic Ion Mnemonics Common Representative “-ate” Oxyanions Pattern in # oxygens: Per-(oxyanion)-ate (oxyanion)-ate (oxyanion)-ite Hypo-(oxyanion)-ite ClO4- ClO3- ClO2- ClO - Every step down the pattern ion has one less oxygen. Name of 1st Element (Roman Numeral or –ous, -ic) Name of Polyatomic ion (Per- Hypo –ATE –ITE)

  45. Ionic Nomenclature: other Polyatomic Ions

  46. Writing Ionic Formulas Mg2+ OH- Mg 2+ O2- Mg 2+ PO43- Mg 2+ SO32- MgOH2 Mg(OH)2 Mg2O2 MgO • If these subscripts are not in the lowest whole-number ratio, divide them by the greatest common factor. Mg3(PO4)2 MgSO3

  47. Given Chemical Names, determine Chemical Formula • Calcium Nitrate • Iron (III) Hydroxide • Zinc Phosphate • Copper (II) Oxide Ca(NO3)2 Ca2+ NO3- Fe3+ OH- Fe(OH)3 Zn2+ PO43- Zn3(PO4)2 CuO Cu2+ O2-

  48. Given Chemical Formula, determine Chemical Names • K2O • NO2 • KMnO4 • Fe(OH)2 • Cu2O • Zn(NO3)2 • Cr2(SO3)5 Potassium Oxide Nitrogen Dioxide Potassium Permanganate Iron (II) Hydroxide Copper (I) Oxide Zinc Nitrate Chromium (V) Sulfite

  49. Hydrates (Hydrated Salts) • Ionic substances containing water molecules incorporated into their crystalline structure • release water upon heating, absorb water under cool, humid conditions H2O H2O Heat (∆) - 4H2O - 2H2O Anhydrous +4 H2O + 2H2O Name of Ionic Salt PREFIX- Hydrate Cu(NO3)2.5H2O

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