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Ions and Ionic Compound

When atoms lose or gain electrons, they become ions, charged particles. Ions and Ionic Compound. Cation Formation. e -. Note: Nucleus is unchanged!!!!!!!. Anion Formation. e -. Note: Nucleus is unchanged!!!!!!!. Ions. Learn common ions listed in syllabus p 12!!.

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Ions and Ionic Compound

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  1. When atoms lose or gain electrons, they become ions, charged particles. Ions and Ionic Compound

  2. Cation Formation e- Note: Nucleus is unchanged!!!!!!!

  3. Anion Formation e- Note: Nucleus is unchanged!!!!!!!

  4. Ions Learn common ions listed in syllabus p 12!! • When atoms lose or gain electrons, they become ions. • Cations are positive and are formed by elements on the left side of the periodic chart: metals (H can form 1+ ion) • Anions are negative and are formed by elements on the right side of the periodic chart:nonmetals (H can form 1- ion)

  5. Ions of Transition Metals

  6. More about ions later now:closer look at atoms and electron arrangements in atoms

  7. Isoelectronic Series • When atoms ionize, they form ions with the same number of electrons as the nearest (in atomic number) noble gas. Na = 1s22s22p63s1 = [Ne]3s1 Na+ = 1s22s22p6 = [Ne] Cl = 1s22s22p63s23p5 = [Ne]3s23p5 Cl- = 1s22s22p63s23p6= [Ar]

  8. Isoelectronic Series • N (7 e-): 1s22s22p3 • O (8 e-): 1s22s22p4 • F (9 e-): 1s22s22p5 • N3- (10 e-): 1s22s22p6 = [Ne] • O2- (10 e-): 1s22s22p6 = [Ne] • F- (10 e-): 1s22s22p6 = [Ne]

  9. Isoelectronic Series • Na (11 e-): 1s22s22p63s1 • Mg (12 e-): 1s22s22p63s2 • Al (13 e-): 1s22s22p63s23p1 • Na+ (10 e-): 1s22s22p6 = [Ne] • Mg2+ (10 e-): 1s22s22p6 = [Ne] • Al3+ (10 e-): 1s22s22p6 = [Ne]

  10. 1A Ionsof the highlighted elements are isoelectronicwith Ne. 8A H 2A He 3A 4A 5A 6A 7A Li Be B C N O F Ne Na Mg Al Si P S Cl Ar 7B 8B 8B 8B 1B 2B 3B 4B 5B 6B K Ca Sc Ti V Cr Fe Co Ni Cu Zn Ga Ge As Se Br Kr Mn Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Rf Db Sg Bh Hs Mt Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np Pu Am Cm Bk Cf Es Md No Lr Fm

  11. Isoelectronic Series • Isoelectronic:having the same number of electrons • N3-, O2-, F-, Ne, Na+, Mg2+, and Al3+ form an isoelectronic series. • A group of atoms or ions that all contain the same number of electrons

  12. Isoelectronic Series • Examples of isoelectronic series: • N3-, O2-, F-, Ne, Na+, Mg2+, Al3+ • Se2-, Br-, Kr, Rb+, Sr2+, Y3+ • Also: Cr, Fe2+, and Co3+

  13. Sizes of Ions - Trends • In an isoelectronic series, ions have the same number of electrons. • Ionic size decreases with an increasing nuclear charge.

  14. Chemical Symbols of Ions Mass Number Charge (=0 for atoms) Atomic Number X Charge = # p - # e-

  15. 16 8 O Chemical Symbol of Neutral O • Using nuclear symbols to determine the number of p, n, e, and total charge Mass Number = Atomic Number = 16 8 # protons = atomic number = 8 # neutrons = Mass # - Atomic # = 16 - 8 = 8 # electrons = # protons = 8

  16. 16 8 2- O Chemical Symbol of O-ion Mass Number = Atomic Number = 16 8 # protons = atomic number = 8 # neutrons = Mass # - Atomic # = 16 - 8 = 8 # electrons = # protons - charge = 8 - (-2) = 10

  17. 137 56 2+ Ba Chemical Symbol of Ions Mass Number = Atomic Number = 137 56 # protons = atomic number = 56 # neutrons = Mass # - Atomic # = 137 - 56 = 81 # electrons = # protons - charge = 56 - (+2) = 54

  18. Chemical Symbols - Ions Practice writing nuclear symbols from information given: • 53 p, 74 n, 54 e- 53 proton (= atomic number)  I (Iodine) 74 neutrons + 53 proton  mass number = 127 54 electrons (one more than protons)  1- 127I1- 53

  19. More practice: Given information: 23 e-, 30 n, net charge = +3 # protons? 23 electrons, but charge of 3+ ie 3 more protons than electrons  p= 26  Atomic number = 26  element = Fe 56 3+ Fe 26

  20. What atom/ion does the following show? Write the chemical symbol.

  21. Practice writing chemical symbols from given information:

  22. Periodic Properties: Cation and anion sizes • Trends to know: • Cations (+) are smaller than their parent atoms. • Electrons are removed from the outer shell. • Anions (-) are larger than their parent atoms. • Electron-electron repulsion causes the electrons to spread out more in space.

  23. Sizes of Ions – Periodic Trends • Ions increase in size as you go down a column/group in periodic table.

  24. Ionization Energy = amount of energy required to remove an electron of a gaseous atom or ion to form a cation or more positively charged cation. • The first ionization energy is the energy required to remove first electron. • The second ionization energy is the energy required to remove second electron, etc.

  25. Ionization Energy • The higher the ionization energy, the harder it is to remove an electron. • It requires more energy to remove each successive electron. • When all valence electrons have been removed, the ionization energy takes a quantum leap. Na (g)  Na+ (g) + e- 2nd electron 1st electron

  26. Trends in First Ionization Energies • As one goes down a column, less energy is required to remove the first electron. • valence electrons are farther from the nucleus. Within each row, the ionization energy increases from left to right

  27. Ionization Energy Which element has the higher ionization energy, Br or Ca? Which one will lose an electron easier? • Br has the higher ionization energy • further to the right • Ca will lose an electron easier because its ionization energy is lower.

  28. Electron Affinity • The energy change that occurs when an electron is added to a gaseous atom is called the electron affinity. Cl (g) + e- Cl- (g) • The electron affinity becomes increasingly negative as the attraction between an atom and an electron increases • more negative electron affinity = more likely to gain an electron and form an anion

  29. Electron Affinity • Trends: • Halogens have the most negative electron affinities. • Electron affinities become increasing negative moving from the left toward the halogens. • Electron affinities do not change significantly within a group. • Noble gases will not accept another electron. • To do so would require adding an electron to a new electron shell (significantly higher in energy)

  30. So: Why do ions form to begin with?

  31. Ionic Compounds Bonds occur between atoms as a result of interactions among the electrons. When the interaction is to strip electrons, the resulting bond is said to be ionic and the entity formed is an ionic compound. “Atoms” – now ions – are held together by electrostatic interactions, ionic bonds. This is a crystal of NaCl. Na+ Cl- NaCl formula unit AND empirical formula

  32. Ionic Bonds Ionic compounds (such as NaCl) are generally formed between metals and nonmetals.

  33. Ionic Compounds • Ionic compounds are made of cations and anions, held together by electrostatic attraction: • opposite electrical charges attract each other • like electrical charges repel each other. • Ionic compounds do not exist as discrete molecules, but as structured aggregates (crystals). • In NaCl, an ionic compound, Na exists as Na+ and Cl exists as Cl-. • Important: The overall charge of the ionic compound is ZERO! Equal negative and positive charges.

  34. Properties of Molecular Compounds Properties of Ionic Compounds More about molecular compounds in Unit 3 • held together by covalent bonds • form discrete molecules • soft • low melting point • generally nonconductive • includes all organic compounds • held together by ionic bonds • do NOT form discrete molecules • hard, rigid, brittle • high melting points • conductive when melted or when dissolved in water

  35. Identify Which of the following compounds would you expect to be ionic: N2O, Na2O, CaCl2, SF4? Which of the following compounds are molecular: CBr4, FeS, P4O6, PbF2?

  36. Ion Charges • Metal ions typically have a positive charge. • Group 1A metals always have a +1 charge: • Li+, Na+, K+, etc. • Group 2A metals always have a +2 charge: • Mg2+, Ca2+, Ba2+, etc. • Some metal ions can form differently charged ions (Fe2+ and Fe3+)

  37. Ion Charges • Nonmetal ions typically have a negative charge. • Group 7A nonmetals typically have a -1 charge: • F-, Cl-, Br -, etc. • Group 6A nonmetals typically have a -2 charge: • O2-, S2-, Se2-, etc. !!! Knowing what the Groups mean and knowing where the metal/nonmetal boundary is on the periodic table is a BIG help when dealing with ions and ionic compounds !!!

  38. Practice • Give the chemical symbol, including mass number, for the ion with 22 protons, 26 neutrons, and 19 electrons: Metal or nonmetal? Anion or cation?

  39. Polyatomic Ions • A group of atoms that is covalently bonded yet still has an overall charge is a polyatomic ion. • NO3- • SO42- • PO43- • ClO2- 3- O O P O O phosphate ion !!! You are responsible for knowing the names, symbols, and correct charges for the ions listed in Unit 2 of the syllabus !!!

  40. Naming Binary Ionic Compounds • Name the cation (metal). • Name the anion (nonmetal). • replace the end of the nonmetal with –ide • oxygen becomes oxide • fluorine becomes fluoride • sulfur becomes sulfide • more: NaCl BaI2 Ba3P2 K2S sodium chloride barium iodide barium phosphide potassium sulfide

  41. Naming Metal Ions When More Than One Ion is Possible • Two methods • Stock system (Roman numeral is the charge of the cation) • Fe2+ is iron(II) • Fe3+ is iron(III) • Sn2+ is tin(II) • Sn4+ is tin(IV) • Classic (-ic, -ous) system • -ic is for the ion with the higher charge • -ous is for the ion with the lower charge • Fe2+ is ferrous • Fe3+ is ferric • Sn2+ is stannous • Sn4+ is stannic

  42. Naming Binary Ionic Compounds • MgCl2 • CuS • Cu2S • Fe2O3 • Na2O • magnesium chloride • copper(II) sulfide or cupric sulfide • copper(I) sulfide or cuprous sulfide • iron(III) oxide or ferric oxide • sodium oxide

  43. Writing Formulas for Binary Ionic Compounds • The overall ionic compound MUST BE electrically neutral (have a net charge of 0). • If you do not know the charges of the ions in the compound, you will not be able to write the correct formula for the compound! Write the formula for potassium fluoride. 1. Write the two elements K F 2. Write their charges K+ F- 3. If the charges are equal and opposite, then just put the two elements together: KF Note: there are NO charges in the formula!

  44. Writing Formulas for Binary Ionic Compounds Write the formula for silver oxide. 1. Write the two elements Ag O 2. Write their charges Ag+ O2- 3. When the charges are different, perform a swap: Ag+ O2- Ag2O Why do we not write Ca2O2 for calcium oxide?

  45. Writing Formulas for Binary Ionic Compounds calcium iodide titanium(II) nitride lead(IV) chloride iron(III) oxide CaI2 Ti3N2 PbCl4 Fe2O3

  46. Naming Polyatomic Ions and Polyatomic Oxyanions • Polyatomic ions – memorize list in your syllabus. The names and formulas for other polyatomic ions will be provided to you. • Polyatomic oxyanions • sulfate: SO42- (more O’s, -ate) • sulfite: SO32- (less O’s, -ite) • perchlorate: ClO4- (one more O, per-) • chlorate: ClO3- • chlorite: ClO2- • hypochlorite: ClO- (one less O, hypo-)

  47. Writing Formulas for Ionic Compounds with Polyatomic Ions Write the formula for magnesium sulfate. 1. Write the two ions with their charges. Mg2+ SO42- 2. If the charges are equal and opposite, put the two ions together, DO NOT include the charges in the formula. MgSO4

  48. Writing Formulas for Ionic Compounds with Polyatomic Ions Write the formula for ammonium sulfate. 1. Write the two ions with their charges. NH4+ SO42- 2. If the charges are not equal and opposite, do the “swap.” NH4+ SO42- (NH4)2SO4 Note the parentheses!

  49. Writing Formulas for Ionic Compounds with Polyatomic Ions sodium hydroxide magnesium hydroxide aluminum hydroxide NaOH Mg(OH)2 Al(OH)3 aluminum phosphate sodium phosphate ammonium phosphate calcium phosphate AlPO4 Na3PO4 (NH4)3PO4 Ca3(PO4)2

  50. Practice • Name the following compounds: • (a) K2SO4 • (b) Ba(OH)2 • (c) FeCl3

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