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Understanding Ionic Bonds and Naming Ions

Learn about ionic bonds, valence electrons, formation of ions, and naming conventions for cations and anions in chemistry. Explore oxidation numbers, polyatomic ions, and properties of ionic compounds. Practice assigning oxidation numbers and writing formulas.

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Understanding Ionic Bonds and Naming Ions

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  1. Chapter 7 and 9: Ionic Bonds

  2. 7.1/9.1 Ions/Naming Ions • Valence Electrons: • Electrons in the highest occupied energy level • Determines the chemical properties • = to the group number (groups 1,2, 13-18 minus 10)

  3. Electron dot structures (Lewis structures): • Shows valence electrons as dots Ex. Ca

  4. Octet Rule: • In compounds – atoms want noble gas (octet) configuration • Metallic atoms – lose electron(s) = noble gas (octet) configuration from next lowest energy level • ex. Na+ = Ne configuration • Non-metallic atoms – gain or share electron(s) = noble gas (octet) configuration for that energy level • ex. F- = Ne configuration

  5. Exceptions – transition metals achieve pseudo noble gas configurations (18 electrons in outer energy level) • ex. Ag= 1s22s22p63s23p6 4s23d104p65s1 4d10 Ag+ =1s22s22p63s23p6 4s23d104p64d10

  6. Formation of Ions: • Neutral atom loses or gains electrons • Produces cations (+) charge or anions (-) charge  • Monotomic Ions- composed of 1 atom with a +/- charge • Polyatomic ions- composed of 2 or atoms with a +/- charge

  7. Formation of Cations: • + charge • Lose electron(s) • Mostly metals

  8. Examples: Ionization of Na: Na• loss of valence electron Na+ + e- Ionization Na 1s22s22p63s1 Na+ 1s22s22p6 Ionization of Mg: Mg loss of 2 valence electron Mg+2+ 2e- Ionization Mg 1s22s22p63s2 Mg+2 1s22s22p6

  9. Formation of Anions: • - charge • Gain electron(s) • Mostly non-metals

  10. Examples: Cl + e-gain one valence electron Cl-1 chlorine atom chloride ion Cl 1s22s22p63s23p5Cl-1 1s22s22p63s23p6 O t 2e-gain 2 valence electron O -2 oxygen atom oxide ion O 1s22s22p4 O -2 1s22s22p6

  11. Naming of Cations: • Name = element’s name + ion (ex. Na+ = sodium ion) • Transition Metals-Stock system – use Roman numerals to ID the charge (ex. Fe+2 = iron(II) ion, Fe+3 = iron(III) ion)

  12. Naming of Anions: • Name = ending of the element’s name dropped and ide added (ex. fluorine (F-) = fluoride)

  13. Ch 20.2 Oxidation Numbers • Oxidation number: • + or – # that indicates how many electrons an atom has gained, lost or shared to become stable • Some elements have more than one oxidation #

  14. Rules for Assigning Oxidation #s: • Monatomic ions = + or – what it takes to gain the octet (1-2 group #, 13-17 group # -10) = ion’s charge • Hydrogen = +1. Exception = Metal hydrides (ex. LiH), = -1. • Oxygen = -2 (exception- when combined with F then = +2) • Neutral compounds = sum oxidation #s = 0 • Polyatomic ions-the sum of the oxidation # = the charge of the ion.

  15. Examples: +2-2=0 NO O = -2, so N = +2 (+2-2=0 rules: 1,3,4) +4 –2(2) = 0 NO2O = -2(2), so N = +4 (+4 – 2(2) = 0 rules: 1,3,4) +6-2(4)=-2 SO4 -2 O = -2(4) S -2(4) = -2  S - 8 = -2 S = +6 (+6-2(4)=-2 rules: 1,3,5)

  16. Practice • Assign oxidation numbers to each element in the following: H2S MgF2 PO4-3

  17. Ch 9.1 • Polyatomic Ions: • Tightly bound groups of atoms behaving as a unit and carry a charge • Most – anions (except ammonium)

  18. Naming Polyatomic Ions: • Those containing oxygen (oxyanions) in different ratios have different endings (iteor ate) depending on the # of oxygens presents:

  19. –ite = one less oxygen then ate • ex, chlorite- ClO2-1; chlorate- ClO3-1 • Hypo = one less oxygen then ite • Ex. Hypochlorite- ClO-1 • Per = one more oxygen then ate • Ex. Perchlorate- ClO4-1

  20. 7.2 Ionic Bonds and Ionic Compounds • Formation of Ionic Compounds: • Composed of cations and anions • Electrically neutral

  21. Ionic Compounds Properties: • Crystalline solids • Metals/Nonmetals • High melting points • Good conductors of electricity when in water or melted • Soluble in water

  22. Ionic Bonds: • Electrostatic force that holds ions together Ex. Na gives away its lone electron to chlorine (7 valence electrons) = both have noble (octet) configs in a 1-1 ratio

  23. Chemical formula: • Shows kinds and # s of atoms in the smallest representative unit • Subscript – indicates the # of atoms of each element in the compound = Oxidation #

  24. Formula Units: • Lowest whole number ratio of ions in an ionic compound Ex. NaCl lowest whole number ratio = 1:1 Ex. MgCl2 lowest whole number ratio = 1:2 (1 Mg+2 ion to 2 Cl-1)

  25. 9.2 Naming and Writing Formulas for Ionic Compounds: • Naming Ionic Compounds: • Cation name first followed by anion name Ex. Cs2O = cesium oxide • Cations that form more than 1 oxidation # - include roman numeral in name Ex. CuO = copper(II) oxide

  26. Writing Formulas for Ionic Compounds: • Cation symbol followed by anion symbol • Add subscripts • Assign Oxidation #s • Crisscross method- oxidation # (charge) crosses over to become the subscript

  27. ex. Fe+3 O-2 Fe2 O3

  28. Naming Polyatomic Compounds: • Cation name first followed by the polyatomic anion name (exception – ammonium ion = + 1 charge) Ex. NaClO = sodium hypochlorite • Cations that form more than 1 ion – include roman numeral Ex. Cr(NO2)3 = Chromium(III) nitrite

  29. Writing Polyatomic Compounds: • Cation symbol followed by the polyatomic ion symbol • Add subscripts • ( )needed if more than 1 unit of the same polyatomic ion is needed Ex. Ca+ 2 (NO3)2-1 Ca (NO3)2 Ca(NO3)2

  30. 9.4 Naming and Writing Formulas for Bases and Acids • Bases: • Ionic compound that produces hydroxide ions (OH-1) when dissolved in water

  31. Naming Bases: • Named same way ionic compounds • Anion =hydroxide ion (OH-1) Ex. NaOH = sodium hydroxide

  32. Writing Formulas for Bases: • Same way as ionic compounds Ex. Iron(III) Hydroxide = Fe(OH)3

  33. Acids: • Ionic compound that contains 1 or more hydrogen atoms • Produces hydrogen ions (H+) when dissolved in water

  34. Naming Acids: • Consists of a cation (hydrogen ions) and an anion • HnX formula – H = hydrogen; n= # of Hydrogens; X= monatomic or polyatomic anion

  35. Writing Formulas for Acids: • Use the naming rules in reverse: Ex. 1 Hydrochloric acid = HCl Ex. 2 Sulfurous acid = H2SO3 Ex. 3 Nitric acid = HNO3

  36. 7.3 Bonding in Metals • Metallic bonds: • Metals made up of closely packed cations rather than neutral atoms • Valence electrons – “sea of electrons” • Metallic bonds: • Consist of the attraction of the free-floating valence electrons • Forces of attraction that hold metals together

  37. Metallic Properties: • Good conductors b/c electrons can flow freely in them – as electrons enter one end = # leave the other end • Ductile and malleable b/c of mobility of valance electrons- metal cations easily slide past one another like ball bearings immersed in oil when subjected to pressure.

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