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I. Ion Formation Ionic Formulas Ionic Nomenclature

Ch. 9 – Chemical Names and Formulas. I. Ion Formation Ionic Formulas Ionic Nomenclature. Ionic Bonds. When oppositely charged ions attract, electrostatic force that holds them together = ionic bond Compounds containing ionic bonds = ionic compounds

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I. Ion Formation Ionic Formulas Ionic Nomenclature

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  1. Ch. 9 – Chemical Names and Formulas I. Ion FormationIonic FormulasIonic Nomenclature

  2. Ionic Bonds • When oppositely charged ions attract, electrostatic force that holds them together = ionic bond • Compounds containing ionic bonds = ionic compounds • Electrons are transferred from cations to anions • Typically bonds between metals and nonmetals, however polyatomic ions also involved.

  3. A. Vocabulary • Chemical Bond • attractive force between atoms or ions that binds them together as a unit • bonds form in order to… • fulfill octet rule • increase stability

  4. Octet Rule • Atoms will gain or lose electrons so that they have 8 electrons in their highest energy level • The noble gases already have a full octet, 8 valence electrons, so are chemically stable.

  5. A. Vocabulary • Cation • Positively charged ion formed when an atom loses one or more valence e- • Number of protons stays the same, but less electrons gives + charge Loses an e- Sodium ion

  6. A. Vocabulary • Anions • Nonmetals easily gain e- to form negative ions to get to 8 valence e- • Name is changed to root + -ide Gains an e- Chloride ion

  7. A. Vocabulary ION 2 or more atoms 1 atom Monatomic Ion Polyatomic Ion Na+ NO3-

  8. Monatomic Ions • Ions formed from a single atom of an element • Most main group metals form one type of monatomic ion • Group 1 metals: +1 ions (Na+, Li+…) • Group 2 Metals: +2 ions (Mg2+, Ca2+…) • Group 13 Metals: +3 ions (Al3+, Ga3+…)

  9. Monatomic Ions • Main group nonmetals form one type of anion • Group 15: -3 ions (N3-, P3-) • Group 16: -2 ions (O2-, S2-…) • Group 17: -1 ions (F-, Cl-…)

  10. Monatomic Ions • Some metals, mainly transition metals, can form more than one type of ion • Iron: Fe2+ or Fe3+ • Copper: Cu+ or Cu2+ • Lead: Pb2+ or Pb4+

  11. D. Common Ions

  12. B. Formula Unit

  13. Chemical Formulas • Chemical formulas • identify the elements present in a compound using element symbols • the number of atoms of each element is indicated with numbers in subscript Fe2O3 3 Oxygen atoms 2 Iron atoms

  14. Binary Ionic Compounds • Binary ionic compounds are ionic compounds formed between two elements. • Always formed between a metal (positive ion) and nonmetal (negative ion) • Metal is written first in formula and then nonmetal

  15. D. Ionic Formulas • Writing Ionic Formulas: • Calcium chloride • Ca2+ Cl1- • charges do not cancel, must criss-cross charges • Rewrite as complete formula without charges • CaCl2 2 1

  16. D. Ionic Formulas  KF • K+ F- • Mg2+ N3-  Mg3N2  BaCl2 • Ba2+ Cl-

  17. D. Ionic Formulas • Ca2+ O2- CaO Al2S3 • Al3+ S2-  • Mg2+ Br- MgBr2

  18. C. Lewis Structures • Ionic – show transfer of electrons

  19. G. Ionic Nomenclature Naming Binary Ionic Compounds • Write names of both ions, cation (metal) first • Change ending of monatomic anions (nonmetal) to -ide • Use Romannumerals to show the ion’s charge if more than one is possible

  20. G. Ionic Names with Type II Cations • You must write the charge in parentheses using Roman numerals. To determine charge know that overall charge of compound = 0 • Cr2O3 • CrO Chromium (III) oxide Formula: Element: # atoms x charge = total charge O: 3 x -2 = - 6 Cr: 2 x ___ = +6 +3 Chromium (II) oxide O: 1 x -2 = - 2 Cr: 1 x ___ = +2 +2

  21. CaCl2 Al2O3 Na3N KI Calcium Chloride Aluminum Oxide sodium nitride Potassium Iodide Name the following Compounds

  22. FeN Cu2S ZnCl2 V3P5 MnO3 Iron (III) nitride Copper (I) sulfide Zinc chloride Vanadium (V) phosphide Manganese (VI) oxide Name the following cont.

  23. Strontium nitride Lithium sulfide Gallium bromide Barium oxide Sr3N2 Li2S GaBr3 BaO Write the formula for

  24. Tin (IV) oxide Gold (III) Chloride Chromium (II) Nitride Mercury (I) sulfide SnO2 AuCl3 Cr3N2 (Hg2)2S Write the formulas

  25. F. Ionic Formulas with Type II Cations • Copper (II) bromide • Tin (IV) oxide • Manganese (II) chloride • Cu2+ + Br  CuBr2  SnO2 • Sn4+ + O2 Sn2O4 • Mn2+ + Cl MnCl2

  26. E. Polyatomic Ions • Polyatomic Ions • Ions made of more than one atom • Acts as an individual ion and its charge applies to the entire group of atoms • NEVER change the subscripts – add parentheses and subscripts outside, if necessary • Listed on the back of your periodic table

  27. E. Ionic Formulas with PA Ions • potassium chlorate • magnesium nitrate • ammonium phosphate  KClO3 • K+ ClO3- • Mg2+ NO3-  Mg(NO3)2  (NH4)3PO4 • NH4+ PO43-

  28. E. Ionic Formulas with PA Ions • calcium oxalate • aluminum perchlorate • strontium phosphate • Ca2+ C2O42- CaC2O4 • Al3+ ClO4-  Al(ClO4)3 • Sr2+ PO43- Sr3(PO4)2

  29. Nomenclature with PA • When naming compounds with polyatomic ions use the polyatomic ions name and follow all other rules for ionic compounds • Ca3(PO4)2 calcium phosphate • Fe(NO3)3  iron (III) nitrate

  30. G. Ionic Nomenclature • CaBr • Na2CO3 • NH4OH • calcium bromide • sodium carbonate • ammonium hydroxide

  31. G. Ionic Names with Transition Metals • Cr2(SO4)3 • Cu(NO3)2 • FeCl3 • Chromium (III) sulfate • Copper (II) nitrate • iron(III) chloride

  32. Ch. 9 – Chemical Names and Formulas II. Covalent Bond FormationCovalent Compound Names & Formulas

  33. H2O A. What is a covalent bond? • A chemical bond that results from the sharing of electrons • Molecule = two or more atoms that are held together by covalent bonds • Majority of covalent bonds form between nonmetals (CLOSE together on periodic table)

  34. B. Examples: • Which of the following are covalent compounds? • NaBr • SiO2 • CO2 • AlCl3 • CH4

  35. Cl2 C. Covalent Bonding Formation • Diatomic molecule • molecule containing only two atoms • Some elements always exist this way because they are more stable than the individual atoms

  36. D. Diatomic Elements • The Seven Diatomic Elements Br2I2N2Cl2H2O2F2 H N O F Cl Br I

  37. E. Molecular Nomenclature • Prefix System (binary molecules) • Add prefixes to indicate # of atoms. Omit mono- prefix on first element. • Change the ending of the second element to -ide. • Second element ALWAYS gets a prefix.

  38. PREFIX mono- di- tri- tetra- penta- hexa- hepta- octa- nona- deca- NUMBER 1 2 3 4 5 6 7 8 9 10 E. Molecular Nomenclature

  39. E. Molecular Nomenclature • CCl4 • N2O • SF6 • carbon tetrachloride • dinitrogen monoxide • sulfur hexafluoride

  40. E. Molecular Nomenclature • arsenic trichloride • dinitrogen pentoxide • tetraphosphorus decoxide • AsCl3 • N2O5 • P4O10

  41. Lewis Structures • The Lewis Structures of Covalent compound represents elements in a compound, their valence electrons and how they are shared, and how the elements orient themselves around each other.

  42. A. Drawing Lewis Structures • Count ALL Valence electrons on all atoms in the molecule. • For an anion ion, add one electron for each negative charge. • For a cation, subtract one electron for each positive charge.

  43. A. Drawing Lewis Structures 2. The atom with the least amount is central atom and place the other atoms around the central atom. • Hydrogen is never central atom • Draw a line connecting the peripheral atoms to the central atom • Each line represent 2 electrons • Check for octet

  44. A. Drawing Lewis Structures • Place pairs of valence electrons around each peripheral atom, except hydrogen, until octet is reached. • If any electrons remain place around the central atom until octet is reached. • If central atom still does not have an octet, use a lone pair of electrons on a neighboring atom to form a multiple bond to the central atom.

  45. Examples • CF4 • CO2 • HCN • ClO- • NH4+

  46. B. Drawing Lewis Diagrams • CF4 1 C × 4e- = 4e- 4 F × 7e- = 28e- 32e- + F F C F F 2 16 pairs of e- - 4 pairs of e- 12 pairs of e-

  47. A. Drawing Lewis Diagrams • CO2 1 C × 4e- = 4e- 2 O × 6e- = 12e- 16e- + OCO 2 8 pairs of e- -2 pairs of e- 6 pairs of e-

  48. B. Polyatomic Ions • To find total # of valence e-: • Add 1e- for each negative charge • Subtract 1e- for each positive charge • Place brackets around the ion and label the charge

  49. B. Polyatomic Ions O O Cl O O • ClO4- 1 Cl × 7e- = 7e- 4 O × 6e- = 24e- 31e- + + 1e- 32e- 16 e- pairs - 4 e- pairs 12 e- pairs 2 =

  50. B. Polyatomic Ions H H N H H • NH4+ 1 N × 5e- = 5e- 4 H × 1e- = 4e- 9e- + - 1e- 8e- 2 = 4 pairs of e- -4 pairs of e- 0 pairs of e-

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