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Atomic Structure, Isotopes, and Ionic Compounds

Explore atomic theory, subatomic particles, isotopes, periodic table, molecules, and ionic compounds. Learn about naming species, molecular formulas, and the organization of elements. Understand how ions form and the naming conventions for inorganic compounds.

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Atomic Structure, Isotopes, and Ionic Compounds

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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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