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I. Introduction to Bonding

Learn about the fundamentals of chemical bonding, types of bonds, properties of compounds, and how to determine bond types based on electronegativity. Explore ionic compounds, Lewis structures, octet rule, and naming conventions for binary compounds.

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I. Introduction to Bonding

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  1. Unit 4 – Bonding, Nomenclature & Molecular Structure I. Introduction toBonding

  2. Vocabulary • Chemical Bond • attractive force between atoms or ions that binds them together as a unit • bonds form in order to… • decrease potential energy (PE) • increase stability

  3. Vocabulary • Molecule • Smallest electrically neutral unit of a substance that still has the properties of the substance • Made up of 2 or more atoms • Ion • Atom or group of atoms that have a positive or negative charge

  4. Vocabulary ION 2 or more atoms 1 atom Monatomic Ion Polyatomic Ion Na+ NO3-

  5. Vocabulary COMPOUND more than 2 elements 2 elements Binary Compound Ternary Compound NaCl NaNO3

  6. Vocabulary CHEMICAL FORMULA IONIC COVALENT Formula Unit Molecular Formula NaCl CO2

  7. Types of Bonds COVALENT IONIC e- are transferred from metal to nonmetal e- are shared between two nonmetals Bond Formation Type of Structure true molecules crystal lattice Physical State liquid or gas solid Melting Point low high Solubility in Water yes usually not yes (solution or liquid) Electrical Conductivity no Other Properties odorous

  8. Types of Bonds/Compounds Ionic Bonding - Crystal Lattice

  9. Types of Bonds/Compounds Ionic Bonding

  10. Types of Bonds/Compounds Covalent Bonding - True Molecules Diatomic Molecule

  11. Types of Bonds/Compounds Covalent Bonding

  12. Types of Bonds/Compounds METALLIC e- are delocalized among metal atoms Bond Formation Type of Structure “electron sea” Physical State solid Melting Point very high no Solubility in Water yes (any form) Electrical Conductivity malleable, ductile, lustrous Other Properties

  13. Types of Bonds/Compounds Metallic Bonding - “Electron Sea”

  14. Determining Bond Type • Most bonds are a blend of ionic and covalent characteristics. • Difference in electronegativity determines bond type.

  15. Determining Bond Type • Electronegativity • Attraction an atom has for a shared pair of electrons. • higher e-neg atom  - • lower e-neg atom +

  16. Determining Bond Type • Electronegativity Trend • Increases up and to the right.

  17. Determining Bond Type • Nonpolar Covalent Bond • e- are shared equally • symmetrical e- density • usually identical atoms

  18. - + Determining Bond Type • Polar Covalent Bond • e- are shared unequally • asymmetrical e- density • results in partial charges (dipole)

  19. Determining Bond Type • Nonpolar • Polar • Ionic

  20. nonpolar polar ionic 2.5 3.5 0.9 3.0 2.5 2.1 How To Determine the Bond Type Using Electronegativity Differences 0 to 0.4 = ______________ covalent bond 0.5 to 2.0 = _____________ covalent bond Above 2.0 = _______________ bond Practice Problems: Determine the type of bond that forms between the atoms in the following compounds. a) CO2 b) NaCl c) CH4 1.0 = polar covalent 2.1 = ionic 0.4 = nonpolar covalent

  21. Unit 4 – Bonding & Molecular Structure II. Ionic Compounds

  22. Ne Octet Rule • Octet Rule • Most atoms form bonds in order to obtain 8 valence e- • Full energy level stability ~ Noble Gases

  23. Lewis Structures • Electron Dot Diagrams used to show bonding • Ionic Compounds – show transfer of e-

  24. Lewis Structures • cation + anion = ionic compound

  25. Cation or Anion? • How do you determine whether an atom loses or gains electrons? • Look at the group # • What does that tell you? • Ask yourself, “is it easier to LOSE ______ electrons or GAIN ____ electrons to achieve the octet?”

  26. Practice Problem • Aluminum and nitrogen Al has __ valence electrons - will it lose __ or gain __? - Al becomes… N has __ valence electrons - will it lose ___ or gain __? - N becomes…

  27. Barium and Sulfur Ba [Xe] ____ S [Ne] ___ ___ ___ ___ 6s Which noble gases will they look like after they react? 3s 3p

  28. After reacting… Ba2+ [Xe] ____ S2- [Ne] ___ ___ ___ ___ 6s Looks like Argon 3s 3p

  29. REVIEW - Common Ion Charges 1+ 0 2+ 3+ NA 3- 2- 1-

  30. Naming Type I Binary Compounds (single-charge cations) • Write the names of both ions, cation first. • Change ending of monatomic ions to -ide. • Examples: • SrS • AlBr3

  31. Formulas for Type I Binary Compounds (single-charge cations) • Write each ion, cation first. Don’t show charges in the final formula. • Overall charge must equal zero. • If charges cancel, just write symbols. • If not, use subscripts to balance charges. • Some transition metals have only 1 charge: • Ag+ • Zn2+ • Cd2+

  32. Example: Aluminum Chloride Criss-Cross Rule Al3+ Cl1- Step 1: write symbols & charge of elements Al Cl Step 2: 1 3 criss-cross charges as subscripts AlCl Step 3: 3 combine as formula unit (“1” is never shown)

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

  34. PRACTICE • Type I Binary Compounds WS (Single-Charge Cations)

  35. Naming Type II Binary Compounds (multiple-charge cations) • Since the metal ion can have more than one charge, a Roman numeral is used to specify the charge. • Determine the charge on the cation using the charge on the anion • Example: NiBr2 • NiBr2 = nickel (II) bromide

  36. Common Cations and Anions

  37. Naming Type II Binary Compounds (multiple-charge cations) PRACTICE PROBLEMS • CuCl _________________ • HgO _________________ • Fe2O3 _________________ • MnO2 _________________ • PbCl2 _________________ copper (I) chloride mercury (II) oxide iron (III) oxide manganese (IV) oxide lead (II) chloride

  38. Formulas for Type II Binary Compounds (multiple-charge cations) • Roman numerals indicate the ion’s charge. • Examples: • Iron (III) chloride ________ • Tin (IV) oxide ________ • Lead (II) sulfide ________ FeCl3 SnO2 PbS

  39. PRACTICE • Multiple-Charge Cations WS

  40. Formulas for Ternary Compounds (polyatomic ions) • Use parentheses to show more than one polyatomic ion. • Example: Ca3(PO4)2

  41. Naming Ternary Compounds PRACTICE PROBLEMS • NaNO2 sodium nitrite • KClO2 potassium chlorite • Ba3(PO4)2 barium phosphate • Fe(OH)3 iron (III) hydroxide • NaHCO3 sodium bicarbonate ‘sodium hydrogen carbonate’

  42. Criss Cross Rule – Ternary Compounds Example: Magnesium Phosphate Mg2+ PO43- Step 1: Mg (PO4) Step 2: 2 3 Mg3(PO4)2 Step 3:

  43. More Ternary Compound Examples Zn3(PO4)2 • ________________ zinc phosphate • ________________ ammonium carbonate • ________________ aluminum sulfate • Na2SO4 ____________________ • LiCN ____________________ • Ba(ClO3)2 ____________________ • ________________ copper (II) hydroxide (NH4)2CO3 Al2(SO4)3 sodium sulfate lithium cyanide barium chlorate Cu(OH)2

  44. PRACTICE • Ternary Compounds WS

  45. Go Fish For An Ion Acceptable Matches Na+ Na+ SO4 2- Al3+ Cl- Cl- Cl- Unacceptable Matches Li+ Na+ SO4 2- Al3+ O2-

  46. Go Fish Scoring

  47. Unit 4 – Bonding, Nomenclature & Molecular Structure III. Molecular / Covalent Compounds (Type III Binary Compounds)

  48. Energy of Bond Formation • Potential Energy • based on position of an object • low PE = high stability

  49. no interaction increased attraction Energy of Bond Formation • Potential Energy Diagram attraction vs. repulsion

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