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REVIEW

REVIEW. We can tell how many electrons and atom will gain or lose by looking at its valence. Metals like to lose electrons. (Cations) Ex. Na + Nonmetals like to gain electrons. (Anions) Ex: O 2- All elements try to have a full valence of 8 electrons(OCTET RULE). REVIEW.

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REVIEW

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  1. REVIEW • We can tell how many electrons and atom will gain or lose by looking at its valence. • Metals like to lose electrons. (Cations) • Ex. Na + • Nonmetals like to gain electrons. (Anions) • Ex: O 2- • All elements try to have a full valence of 8 electrons(OCTET RULE).

  2. REVIEW • Cation- is a positively charged ion. • How do cations form? • When atoms LOSE electrons they become positive. • Anion- is a negatively charged ion. • How do anions form? • When atoms GAIN electrons they become negative.

  3. Na Cl Chemical Bonding Notes A chemical bond is the force of attraction that holds two atoms together. Attractive Force

  4. http://www.wisc-online.com/objects/ViewObject.aspx?ID=GCH2204http://www.wisc-online.com/objects/ViewObject.aspx?ID=GCH2204

  5. Why do elements form chemical bonds? Atoms form chemical bonds in order to fill their outermost energy level with electrons. A full valence shell causes an atom to be more stable. A full valence shell consists of 8 valence electrons.

  6. Ionic Bonding Ionic bonds: Metal atoms transfer electrons to nonmetal atoms. Producing oppositely charged ions (cation & anion) which attract each other. Na + Cl Na+Cl-

  7. Remember: Non-metal atoms take electrons from metal atoms to form an octet.

  8. How to write a formula. • Write cation first, followed by anion • Example: Anion : P3- Cation : Al3+ Formula : AlP

  9. How to write a formula. Compound must be neutral, so all charges must cancel

  10. Write an ionic formula for Na+ bonding with F− Balance the charges. Na+ F− (1+) + (1-) = 0 1 Na+ and 1 F− = NaF

  11. Write an ionic formula for Mg2+ bonding with Cl− Balance the charges. Mg2+ Cl− Cl− (2+) + 2(1-) = 0 1 Mg2+ and 2 Cl− = Mg Cl2

  12. Write an ionic formula for K+ bonding with S2− Balance the charges. K+ S2− K+ 2(1+) + (2-) = 0 2 K+ and 1 S2− = K2S

  13. Write the formula for… an ionic compound composed of: Al3+ and S2- Al2S3

  14. Write an ionic formula for Fe3+ bonding with OH− Balance the charges. Fe3+ OH− OH− OH− (3+) + 3(1-) = 0 1 Fe3+ and 3 OH− = Fe(OH)3

  15. Let’s play the Ionic Bonding Dating Game!

  16. Step 4: AlCl 3 Example: Aluminum Chloride Criss-Cross Rule Aluminum Chloride Step 1: write out name with space Al Cl 3+ 1- Step 2: write symbols & charge of elements Al Cl Step 3: 1 3 criss-cross charges as subsrcipts combine as formula unit (“1” is never shown)

  17. Example: Aluminum Oxide Criss-Cross Rule Step 1: Aluminum Oxide Step 2: Al3+ O2- Step 3: Al O 2 3 Step 4: Al2O3

  18. Example: Magnesium Oxide Criss-Cross Rule Step 1: Magnesium Oxide Step 2: Mg2+ O2- Step 3: Mg O 2 2 Step 4: Mg2O2 Step 5: MgO

  19. Criss-Cross Rule criss-cross rule: charge on cation / anion “becomes” subscript of anion / cation ** Warning: Reduce to lowest terms. In3+ and Br1– Al3+ and O2– Ba2+ and S2– In1 Br3 Al2 O3 Ba2 S2 Al2O3 BaS InBr3 aluminum oxide barium sulfide indium bromide

  20. Lesson Three--Transition Metal Compounds Transition metals have electrons in d orbitals and can donate different numbers of electrons, thus giving them several different positive charges. These can be determined from the Roman numeral which is written next to the metal's name. Example: Cu1+is Copper I Pb2+is Lead II Fe3+is Iron III Sn4+s Tin IV

  21. Metals with more than 1 charge Examples: • Cu + Copper (I) • Cu+2 Copper (II) • Fe+2 Iron (II) • Fe+3 Iron (III)

  22. Practice • K and Cl • K and S • Ca and S • Cu (II) and S

  23. Polyatomic Ions!!!!!!!! • A polyatomic ion is a charged species (ion) composed of two or more atoms covalently bonded. • PO4-3 NH4+1

  24. Lewis Dot Structures for Polyatomic ions H + NH4+ H N H H

  25. Al + PO4-3 K + SO4-2 • Al +3 + PO4-3 K +1 + SO4-2 • Al(PO4) K2 (SO4)

  26. Ca + PO4-3 • Ca +2 + PO4-3 • Ca3(PO4)2

  27. Covalent bonding Notes • Covalent bond: The sharing of a pair of electrons between 2 nonmetal atoms in order to fill its valence shell. • Each atom gains 1 electron from each covalent bond it forms with another atom.

  28. When electron sharing usually occurs so that atoms attain a stable electron configuration and have 8 valence electrons.

  29. Single Covalent Bonds Diatomic Molecules Each chlorine needs to gain one electron by sharing electrons each atom achieves stability . Cl + Cl  Cl Cl The pair of shared electrons is often represented as a dash. Cl-Cl

  30. Single Covalent Bonds Diatomic Molecules The chlorine atoms only share one pair of valence electrons. The electrons pairs not shared are called unshared electron pairs or lone pairs. Cl + Cl  Cl Cl

  31. Single Covalent Bonds in compounds • H20 is a molecule containing three atoms with two single covalent bonds. • Count up the electrons you have!!! • 2 H + O H O H • The hydrogen and oxygen attain stable configurations by sharing electrons.

  32. Your Turn • Example OF2

  33. Double Covalent Bonds Two pair of electrons are being shared. S + S  S S

  34. Triple Covalent Bonds Three pair of electrons are being shared. P + P  P P

  35. Charged Compounds • Some compounds do not satisfy their stable configuration and therefore have a charge on the compound. • Example- NH4+1

  36. Exceptions to the Octet Rule The octet rule cannot be satisfied in molecules whose total number of valence electrons is an odd number. However, these molecules do exist in nature. Examples: Nitrogen dioxide (NO2) Boron trifluoride (BF3) Phosphorus pentachloride (PCl5) = 10 v.e- Expanded octet Sulfur hexafluroride (SF6) = 12 v.e- Expanded octet

  37. Nonpolar Covalent Bond • When atoms bond equally it is considered a nonpolar covalent bond. • Cl2 • O2 • N2 • H2

  38. Polar Covalent Bond • When electrons are shared unequally it is a polar covalent bond. • An atom that strongly attracts electrons is more electronegative and therefore gains a slightly negative charge. • The less electronegative atom has a slightly positive charge. • This results in a polar bond!

  39. An arrow is used to show which element is donating the unshared pair of electrons. The crossed end of the arrow indicates a pos. end and the arrow points in the direction of the neg. end Example: H-Br

  40. polar molecules are also called dipoles. A dipole is a molecule with two partially charged ends or poles.

  41. Examples • H-Br • H2S • SCl2 • CO2

  42. C. Bond Polarity • Nonpolar Covalent Bond • e- are shared equally • symmetrical e- density • usually identical atoms C. Johannesson

  43. - + C. Bond Polarity • Polar Covalent Bond • e- are shared unequally • asymmetrical e- density • results in partial charges (dipole) C. Johannesson

  44. - + + B. Lewis Structures • Nonpolar Covalent - no charges • Polar Covalent - partial charges C. Johannesson

  45. + - H Cl A. Dipole Moment • Direction of the polar bond in a molecule. • Arrow points toward the more e-neg atom. C. Johannesson

  46. F BF3 B F F B. Determining Molecular Polarity • Nonpolar Molecules • Dipole moments are symmetrical and cancel out. C. Johannesson

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