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Chapter Seven

Chapter Seven. Ionic Compounds & Metals. Section One. Ion Formation. Valence Electrons & Chemical Bonds. Chemical Bond Force that holds two atoms together What is a valence electron? Electrons in highest energy level How can we depict valence electrons? Lewis Dot Structures.

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Chapter Seven

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  1. Chapter Seven Ionic Compounds & Metals

  2. Section One Ion Formation

  3. Valence Electrons& Chemical Bonds • Chemical Bond • Force that holds two atoms together • What is a valence electron? • Electrons in highest energy level • How can we depict valence electrons? • Lewis Dot Structures

  4. Valence Electrons • Ionization energy • How easily an atom loses an electron • Electron Affinity • How much attraction an atom has for electrons • Reactivity • Related to valence electrons

  5. Positive Ion Formation • Cation • Positively charged ion • Atom loses an electron • Metal Ions • Lose valence electrons • Groups 1 and 2 most reactive • Group 13 or 3A can form ions

  6. Positive Ion Formation • Transition Metal Ions • In general have outer energy level ns2 • Tend to lose s electrons resulting in 2+ ions • Also able to lose d electrons resulting in 3+ • Pseudo-noble gas configurations • Octet most stable configuration • Elements in 11-14 lose e- to form an outer energy level containing full s, p, d sublevels • Zinc

  7. Negative Ion Formation • Anion • Negatively charged ion • Gains electrons • Suffix -ide added to the root name of the element • Example: Chloride • Nonmetal Ions • Gain electrons up to 8

  8. Section Two Ionic Bonds&Ionic Compounds

  9. Formation of an Ionic Bond • Ionic Bond • Electrostatic force that holds oppositely charged particles together in an ionic compound • Ionic Compound • Compounds that contain an ionic bond • Between metals and oxygen oxides form • Most others are salts

  10. Binary Ionic Compounds • Contain only two different elements • Contain metallic cation and nonmetallic anion

  11. Compound Formation & Charge • Let’s look at • Calcium Fluoride • Sodium Chloride • Aluminum Oxide

  12. Properties of Ionic Compounds • Physical Structure • Ionic compound • Large # of ions • Ratio determined by # of e- transferred • Packed into repeating pattern that balances forces of attraction and repulsion • Crystal Lattice • 3D geometric arrangement of particles

  13. Properties of Ionic Compounds • Physical Properties • MP, BP, hardness depend on how strongly particles are attracted • Electricity depends on ability of freely moving particles • Ionic solids do not conduct electricity • Melts into liquid or dissolves in solution conduct electricity • Electrolyte • Ionic compound whose aqueous solution conducts an electric current

  14. Energy and the Ionic Bond • Lattice energy (L.E.) • Energy required to separate 1 mol of the ions in an ionic compound • Ionic compounds are arranged in crystal lattices • Directly related to the size of the ions bonded • Electrostatic force of attraction increases as the distance between the charges decreases • Smaller ions produce stronger interionic attractions and greater lattice energies • Example: Li compound greater than K, Li+ is smaller than K+

  15. Energy and the Ionic Bond • Value of L.E. affected by charge of ion • Larger charges = greater L.E.

  16. Section Three Names and Formulas for Ionic Compounds

  17. Formulas for Ionic Compounds • Formula Unit • Represents the simplest ratio of the ions involved • Monatomic Ions • One atom ion • Oxidation Numbers • Charge of a monatomic ion • Is equal to the number of electrons transferred to from the atom to form the ion • Most transition metals have more than one (See Table 7.8) • Fe2+ Iron(II) • Fe3+ Iron(III)

  18. Formulas for Polyatomic Ionic Compounds • Polyatomic ions • Ions made up of more than one atom • Acts as an individual ion in a compound; charge applies to whole group • Follow same rules for monatomic balancing • Ions have specific names used in naming compounds

  19. Transition Metals • How can we tell the charge of the transition metal from the name? • In the name, the oxidation number will tell you • Example: Iron(II) Sulfate • Formula: FeSO4 • How can we tell the charge from the formula? • Math • Let’s look at Cr(NO3)3

  20. Ammonium Nitrite Nitrate Hydroxide Cyanide Permanganate Hydrogen carbonate Hypochlorite Chlorite Chlorate Perchlorate Bromate Iodate Acetate Periodate Dihydrogen phosphate Carbonate Sulfite Sulfate Thiosulfate Peroxide Chromate Dichromate Hydrogen phosphate Phosphate Arsenate Common Polyatomic Ions (their names)

  21. Practice • Which of the following is incorrect? OxygenAluminde KBr CuCl Zinc Bromide

  22. Very Good! Oxygen Aluminide • First aluminum is a cation, so it should not have an –ide ending • Oxygen is an anion, so it should have the –ide ending • Aluminum should go first Let’s do another. . .

  23. Let’s look at those again

  24. Practice Makes Perfect • Which of the following is correct? AlCl2 Mg2S Cs2N3 Be3P2

  25. Very Good! Be3P2 • Be has a +2 charge • P has a -3 charge

  26. Let’s look at those again

  27. Polyatomic Time • Work on naming these in your copybook: • 1. NH4Cl • 2. HClO2 • 3. Ca(BrO3)2 • 4. BeSO4 • 5. (NH4)3N

  28. Polyatomic Time • Work on writing the formula for these in your copybook: • sodium chromate • barium nitrate • ammonium sulfate • aluminum hydroxide • calcium phosphate

  29. cesium cyanide sodium nitrite calcium acetate beryllium chlorite rubidium sulfite NH4NO3 Sr3(PO4)2 Zn(ClO3)2 AgIO3 K2Cr2O7 Homework

  30. Multiple Charges • Name these. . . • CuF • CuF2 • Cr2O3 • PbI2 • PbCl4

  31. Multiple Charges • Write the formula for these. . . • manganese (VII) oxide • niobium (V) chloride • titanium (III) phosphide • palladium (IV) sulfide • platinum (II) fluoride

  32. Oxyanion • A polyatomic ion composed of an element, usually a nonmetal, bonded to one or more oxygen atoms • Identify the ion with the greatest number of oxygens • -ate ending • The ion with less oxygen • -ite ending

  33. Section Four Metallic Bonds and the Properties of Metals

  34. Metallic Bonds • Metals are NOT ionic • Share properties with ionic compounds • Based on the attraction of particles with unlike charges • Sea of electrons • Electron sea model • Proposes that all the metal atoms in a metallic solid contribute their electrons to form a ‘sea’ of electrons

  35. Metallic Bonds • Electron sea model • Delocalized electrons • Free to move, not held by any specific atom • Metallic bond • Attraction of a metallic cation for delocalized electrons

  36. Properties of Metals • Melting and Boiling Points (MP and BP) • Vary • Generally have moderately high MPs and BPs • MP are lower than BP • Electrons and cations are mobile in metal • When boiling atoms must be separated

  37. Properties of Metals • Malleability, ductility, and durability • Metals are generally durable • Metallic cations are mobile and strongly attracted to electrons  not easily removed from metal • Thermal conductivity and electrical conductivity • Mobile electrons allows for good conductivity • Electrons interact with light absorbing and releasing photons (creates luster)

  38. Properties of Metals • Hardness and strength • Mobile electrons consist of s and d electrons • As # of mobile electrons increases stronger metallic bonds • Alkali metals are ‘soft’

  39. Metal Alloys • Alloy • Mixture of elements that has metallic properties • Examples: stainless steel, brass, cast iron • Properties • Steel is iron with another element • Has increased strength • Vary; can vary based on heating and cooling

  40. Metal Alloys • Substitutional alloys • Some of the atoms of the original metallic solid are replaced by other metals of similar atomic size • Example: Sterling Silver • Copper replaces some of the silver resulting in properties of both silver and copper • Interstitial alloys • Small holes (interstices) are filled with smaller atoms • Example: Steel • Holes in the iron are filled with carbon

  41. Everyday Chemistry • Read page 229 in your textbook

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