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Chapter Two: Atoms, Molecules & Ions

Chapter Two: Atoms, Molecules & Ions. Atomic Theory & Structure Isotopes, Numbers & Masses Periodic Table Molecules, Ions, Compounds & Formulas Naming Species. Atomic Theory and Structure. What is the smallest piece of matter possible?

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Chapter Two: Atoms, Molecules & Ions

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  1. Chapter Two: Atoms, Molecules & Ions • Atomic Theory & Structure • Isotopes, Numbers & Masses • Periodic Table • Molecules, Ions, Compounds & Formulas • Naming Species

  2. Atomic Theory and Structure • What is the smallest piece of matter possible? • Democritus called the smallest particles “atomos” • Dalton’s atomic theory of matter: • elements are composed of small particles -- atoms • all atoms of an element are identical • atoms are not created or destroyed chemically • compounds formed by chemical combination of two or more elements • a given compound has same relative number & type of atoms (law of constant composition) • atoms retain character during chemical rxns. only undergo rearrangement (conservation of matter)

  3. C O Law of Multiple ProportionsIf two elements, A & B, form more than one compound, the masses of B that can combine with a given mass of A are in a ratio of small whole numbers carbon dioxide 12 g of Carbon & 32 g of Oxygen O C O 2 x carbon monoxide 12 g of Carbon & 16 g of Oxygen

  4. alpha particle source detector Subatomic Particles -J.J. Thompson determined charge:mass ratio of e-, 1897-Robert Millikan measured charge of e-, 1909-Thompson developed “plum pudding” model of atom-Rutherford developed “nuclear” model of atom

  5. Modern Atomic Structure * unit charge = 1.602 x 10-19 C (coulomb)  amu (u) -- atomic mass unit = 1.66054 x 10-24 g

  6. Atomic Number • number of protons in an atom • defines an element • shown as the symbolsubscript6C • Mass Number • total number of protons plus neutrons • will vary between isotopes • shown as the symbol superscript12C • Isotopes • elements which have the same atomic number but different mass numbers 12C613C614C6are isotopes

  7. Periodic Table • Allows for organization of elements • Allows for grouping of elements in terms of physical and chemical characteristics • Metals, Non-metals & Metalloids • Group 1A Alkali Metals • Group 2A Alkaline Earth Metals • Group 6A Chalcogens • Group 7A Halogens • Group 8A Nobel Gases • B Groups Transition Metals Know these !!

  8. Molecules and Molecular Compounds • Molecule • the smallest particle of a compound that can be identified as that compound • chemical combination of two or more atoms • a pure substance • Chemical Formula • a symbol representation of a molecule/compound • shows the type and ratio of atoms in a molecule • type is given by symbol • ratio is given by a subscript to right of symbol

  9. Examples: Molecule Ratio 2 : 1 2 : 2 1 : 2 1 : 1 - heteroatomic heteroatomic heteroatomic heteroatomic homoatomic H2O H2O2 CO2 CO O2

  10. Formulas • Molecular FormulasGive the type and exact number of each type of atom • Empirical FormulasGive only the type and simplistratio of atoms Molecular Formula Empirical Formula H2O H2O2 C6H6 C2H6 H2O HO CH CH3

  11. H H H H H C O C H H C C O H H H H H • Structural FormulasShow which atoms are attached to which atoms C2H6O dimethylether ethanol

  12. Ions & Ionic Compounds • Some elements will either lose or gain one or more electrons to become charged species • Metals • typically lose electrons, become +, cations • Non-Metals • typically gain electrons, become -, anions

  13. Na Na+ Monatomic Ions • made from a single element • Na  Na+ + 1e- • Cl + 1e- Cl- 11 p+ 11 e- 1e- + 11 p+ 10 e- 17 p+ 17 e- 17 p+ 18 e- + 1e- Cl- Cl

  14. Hints to Determine Ion Charges • Hydrogen +1 • Oxygen - 2 • Group IA +1 • Group IIA +2 • Group VIA - 2 • Group VIIA - 1

  15. Polyatomic Ions -- “molecules” which have a net positive or negative charge • CO32-carbonate ion • NH4+ammonium ion • OH-hydroxide ion • Prediction of Charges -- all species tend toward the most stable state • Nobel gases are very stable • Elements add or lose electrons to “mimic” nobel gases

  16. Ionic Compounds • Oppositely charged ions form ionic compounds • held together by ionic bonds due to the electrostatic attraction between the opposite charges • Ionic compounds are alwaysneutral species • Mg2+ and Cl- form MgCl2 not MgCl or Mg2Cl

  17. Naming Inorganic Compounds • Names of Monatomic Ions • cations are named for the elementsNa+ is sodium ion Al+3 is aluminum ion Fe+2 is iron(II) ion Fe+3 is iron(III) ion (ferrous ion) (ferric ion) Cu+ is copper(I) ion Cu2+ is copper(II) ion (cuprous ion) (cupric ion) • anions are named for the root name of the element with the ending -ideO-2 is oxide ion Cl- is chloride ion H- is hydride ion N-3 nitride ion

  18. Naming Polyatomic Ions • Know the names, charges and formulas of the important polyatomic ions • NH4+ ammonium ion • CO3-2 carbonate ion • SO4-2 sulfate ion • OH- hydroxide ion • NO3- nitrate ion • Polyatomic ions are treated as separate entities or units • Naming and formula rules are the same as for compounds with monatomic ions

  19. Naming Binary Ionic Compounds • Cations always named first • Anions always named last • NaCl sodium chloride • BaCl2 barium chloride • for cations which have more than one possible charge, the charge of the ion must be given in the name • Fe2O3 iron(III) oxide • FeO iron(II) oxide • Combinations must be neutral!

  20. Examples: • 2 Na+ and 1 CO3-2 is sodium carbonateNa2CO3 • 2 NH4+ and 1 S-2 is ammonium sulfide(NH4)2S • 1 Ba+2 and 2 OH- is barium hydroxideBa(OH)2 • 3 Mg+2 and 2 PO4-3 is magnesium phosphateMg3(PO4)2 • 1 Na+ , 1 H+ and 1 CO3-2 is sodium hydrogen carbonate or sodium bicarbonate, NaHCO3

  21. Acids • A compound that produces hydrogen ions (H+) when dissolved in water • tastes sour • turns litmus red • has a pH less than 7 • typically the formula begins with one or more H’s • HCl(aq) hydrochloric acid • H2SO4(aq) sulfuric acid • HC2H3O2(aq) acetic acid

  22. Binary Acids • Acids which contain H and another non-metallic element • Naming -- to the root name of the non-metallic element: • add the prefix hydro- • add suffix -ic acid • HF(aq) hydrofluoric acid • HBr(aq) hydrobromic acid • HCl(aq) hydrochloric acid Note!

  23. Oxyacids • Acids which contain H and O and another element (or H and a polyatomic anion containing O) • Naming -- to the polyatomic ion name • if the suffix is -ate, change it to -ic • if the suffix is -ite, change it to -ous • add acid to the end of the name • HNO3nitric acidHNO2nitrous acid • H2SO4sufuric acid H2SO3 sulfurous acid • You must know polyatomic ion names/charges

  24. Binary Molecular Compounds • Chemical combinations of non-metals and non-metals (no ions involved) • The more metallic element is named first • The second element (less metallic) is named with the ending -ide • Because there are no ions to use to determine relative ratio of atoms we must indicate the number of each atom by a prefix • N2O3dinitrogen trioxide • SO3 sulfur trioxide

  25. Name the Following: calcium iodide • CaI2 • Cu2O • CuO • Cl2O7 • HClO3 copper(I) oxide copper(II) oxide dichlorine heptaoxide note chloric acid

  26. Write Formulas for the Following: Ca(ClO)2 • calcium hypochlorite • Mg+2 and ClO2- • carbon tetrachloride • NH4+ and SO4-2 Mg(ClO2)2 CCl4 (NH4)2SO4

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