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Chapter 7 and 9: Ionic Bonds

Chapter 7 and 9: Ionic Bonds. 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 ). Electron dot structures (Lewis structures): Shows valence electrons as dots

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Chapter 7 and 9: Ionic Bonds

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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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