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Ch04 lecpptchem1012011f


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  • 1. 06:33 PM
  • 2. Chapter 4: Chemical Bonds06:33 PM
  • 3. Overview of Chapter 4 Electron configuration- valence electrons and stability. Electron-dot structures of atoms and ions; use to describe reactions. Bonding in compounds: Ionic and covalent. Nomenclature of ionic and covalent compounds. Electronegativity. Polar versus nonpolar compounds. Writing electron-dot structures of molecules. Simple geometries of molecules. Intermolecular forces in states of matter and in mixtures: Dipole forces, hydrogen bonding, dispersion forces, forces in solutions.06:33 PM
  • 4. Bonding and Valence Electrons Forces that hold atoms together within a molecule, or ions together in crystals. Involve valence electrons = Outermost electrons. Inner electrons (core electrons) are generally not involved in bonding. Electron configuration- arrangement of electrons. Valence electrons + Core06:33 PM electrons
  • 5. Noble Gas Configurations Example: Sodium (Na) losing an electron- 11p 11p +  Na: 11 e-1 Na+1: 10 e-1 lost e-1 The sodium ion and neon are 10pisoelectronic- They have the same electron configuration. Ne: 10 e-1
  • 6. Noble Gas ConfigurationsExample: Chlorine (Cl) gaining an electron- 17p +  17p Cl: 17 e-1 gained e-1 Cl-1: 18 e-1 18p Are the chlorine ionand argon isoelectronic?. Ar: 18 e-1 Ar: 18 e-1
  • 7. Electron-Dot Structures Electron-dot structures (EDSs) = Lewis-dot structures (LDSs)? Represent the no. of valence e-1 around an atom (dots around symbol). Example 1: Give the electron-dot structure for sodium. What is the symbol for sodium? Na How many valence electrons does it contain? 1 Write the symbol, then place electrons around the symbol (in pairs if possible). Na
  • 8. Electron-Dot Structures of Atoms Give the electron-dot structure for the following: 2.) K  K 3.) Mg  Mg 4.) Al  Al 5.) O  O 6.) Br  Br
  • 9. More on Electron Configurations Noble gases are most stable group of elements (least reactive) Why? 8 valence electrons = stable octet  Stable electron configuration. Octet rule- Atoms attempt to obtain 8 valence electrons. Exceptions: Group 1A, 2A, 3A elements and helium.
  • 10. Electron-Dot Structures of IonsExample: Sodium forming sodium ion. 11p 11p +  Na: 11 e-1 Na+1: 10 e-1 lost e-1 Na Na +1
  • 11. Electron-Dot Structures of IonsExample: Chlorine forming chloride ion. 17p +  17p Cl: 17 e-1 gained e-1 Cl-1: 18 e-1 Cl Cl -1
  • 12. Ionic Bonds A chemical bond between two ions; Transfer of electrons between a metal and a nonmetal. Ionic compounds contain ionic bonds; usually crystals. Very strong bond indeed!
  • 13. Sodium and Chlorine: Boom! Example 1: Sodium will react with chlorine to form sodium chloride (NaCl). Na + Cl ---- -----> NaClNa + Cl Na +1 Cl -1
  • 14. Ionic CompoundsExample 2: A potassium atom reacts with a bromine atom to form potassium bromide (KBr). K + Br  KBr K + Br K +1 + Br -1
  • 15. Ionic CompoundsExample 3: A potassium atom reacts with an oxygen atom to form potassium oxide (K2O). 2 K + O  K 2O K K +1 O + O -2 K + K+1
  • 16. Nomenclature of Binary Ionic CompoundsRoman numeral system. See Table 5.2 (CFCT).Rules: Names of the elements in order according to chemical formula. Change the ending of the last element to “-ide”. Consider ionic state of metal: If the metal has only one ionic state, you are finished naming. If the metal has more than one ionic state, state in parentheses after the metal the size of the charge on the metal ion.
  • 17. Nomenclature of Binary Ionic Compounds: Formula to NameExamples: 1.) LiCl Li+1 and Cl-1 lithium chlorine  lithium chloride 2.) CuO Copper oxygen Cu+2 and O-2 copper (II) oxygen  copper (II) oxide
  • 18. Nomenclature of Binary Ionic Compounds: Name to FormulaExamples: 1.) What is the chemical formula for sodium iodide. Na and I  Na+1 + I-1  NaI 2.) What is the chemical formula for potassium sulfide? K and S  K+1 + S-2  Cross-over method: K2S
  • 19. Name to FormulaExamples:3.) Copper (II) oxide Cu+2 O-2  One Cu for one O Cu2O2  lowest-whole number ratios! CuO4.) Iron (III) chloride FeCl3
  • 20. Covalent BondingElectrons strongly shared, not transferred.Covalent bonds are weaker bond thanionic bonds; nonmetal bonded tononmetal.Examples: SO2 (sulfur dioxide) H2S (hydrogen sulfide) NH3 (ammonia)
  • 21. Types of Covalent BondsThree major types of covalent bonds:- Single: one pair of electrons Ex: C:H or C–H- Double: two pairs of electrons Ex: C::O or C=O- Triple: three pairs of electrons Ex: N:::N or N≡NExample: How many pairs of electrons are in thehydrogen-oxygen bond in water? H-O-H Answer: 1 pair
  • 22. Rules for NamingBinary Covalent Compounds: Ex 1. Start with the chemical formula: CO2 Elements: carbon oxygen First element- Use a prefix for elements in quantity greater than one: carbon oxygen Second element- Use prefix for elements of any quantity: carbon dioxygen Add –ide carbon dioxide
  • 23. Nomenclature for Covalent Compounds: Name from FormulaExample 2: What is the chemical name for CO? carbon monoxide.Example 3: What is the chemical name for PCl3? phosphorus trichloride.Example 4: What is the chemical name for N2O? dinitrogen monoxide.
  • 24. Nomenclature for Covalent Compounds: Formula from NameExample 1: What is the chemical formula for dihydrogen monoxide? H2OExample 2: What is the chemical formula for sulfur trioxide. SO3Example 3: What is the chemical formula for tetraphosphorus trisulfide? P4S3
  • 25. Electronegativity The ability for a nucleus to attract electrons. Atoms of different elements have different abilities of attracting electrons. See table in text, and lecture guide.
  • 26. Applying Electronegativities: Overview Take differences between the two elements in a bond to determine predominant character of bond.  Nonpolar covalent bond: < 0.5.  Polar covalent bond: between 0.5 and 2.0.  Ionic: > 2.0.
  • 27. Applying ElectronegativitiesExamples:Hydrogen (H2): H—H Difference = 0; nonpolar covalentHydrogen chloride (HCl): H—Cl Difference = 0.96; polar covalentSodium chloride (NaCl): Difference = 2.23; ionic  Na+1 Cl-1
  • 28. Applying ElectronegativitiesExamples: How is the electron density around the molecule distributed for a hydrogen molecule? Hydrogen atoms  hydrogen molecule H + H  H–H
  • 29. Applying ElectronegativitiesExample: Give the partial charges on the atoms in the hydrogen-chlorine bond in HCl. Show the dipole moment in the bond. How is the electron δ+ δ- density distributed? H – Cl 2.2 3.16 H – Cl
  • 30. Applying Electronegativities Example: Show ions given their electronegativities. Na Cl 0.93 3.16 Na+1 Cl-1
  • 31. Polyatomic molecules See LG (p. 169) for rules on preparing molecules. Go through examples in the Lewis-Dot Structures Worksheet for Covalent Molecules in the LG (p. 171-175) Consult rules for writing Lewis-dot structures as reference.
  • 32. Polyatomic Ions Polyatomic ions: ions that contain more than one atom. CFCT: Table 5.4 on p. 235. Be familiar with these ions. Examples include:  Carbonate (CO3-2)  Bicarbonate (HCO3-1)  Phosphate (PO4-3)  Sulfate (SO4-2)  Hydroxide (OH-1)  Nitrate (NO3-1)  Ammonium (NH +1)
  • 33. Writing Lewis-dot Structures SeeRules for “Writing Lewis-Dot Structures” (p. 203 in Lecture Guide).
  • 34. Determining the Central Atom inLewis-Dot Structures of Molecules Consider how many free pairs of electrons an atom has before bonding. The higher this number, the more potential for bonding. This means it is more likely to be a central atom. Abridged Periodic Table of Elements – Lewis Dot Structures 1A 2A 3A 4A 5A 6A 7A 8A H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar
  • 35. Free Radicals Molecules/atoms with an unpaired electron. Cl N O Importance in terms of health: Can cause damage to tissue in the body. Antioxidants are used to counteract free radicals: Sources include blueberries and green tea.
  • 36. Types of Molecular Geometries
  • 37. Polar vs. Nonpolar MoleculesDipole: a molecule that hasunequally distributed charges. Polar- having unequally distributed charge. Nonpolar- having equally distributed charge. Polar Example: Hydrogen Chloride δ+ δ- δ+ δ- H Cl H Cl
  • 38. Water: Notice how the electron pairs spread out and cause a bent geometry. This also causes the formation of a dipole, making water polar.
  • 39. Water as a Polar Molecule O H H Molecular Geometry: Bent
  • 40. Water as a Polar Molecule δ- Bond Dipole moment O H H δ+ δ+ Molecular Geometry: Bent
  • 41. Water as a Polar Molecule δ- Overall Dipole O Moment = Polar H H δ+ δ+ Molecular Geometry: Bent
  • 42. Ammonia (NH3): Notice how the electron pairs spread out and cause a pyramidal geometry. This also causes the formation of a dipole, making ammonia polar.
  • 43. Ammonia as a Polar Molecule N H H H Molecular Geometry: Pyramidal
  • 44. Ammonia as a Polar Molecule δ- N H H δ+ H δ+ δ+ Molecular Geometry: Pyramidal
  • 45. Ammonia as a Polar Molecule δ- Overall Dipole Moment = Polar N H H H δ+ δ+ δ+ Molecular Geometry: Pyramidal
  • 46. Methane (CH4): Notice how the electron pairs spread out and cause a tetrahedral geometry. This causes the molecule to be nonpolar since the overall dipole moments (see +) to counteract each other.
  • 47. Methane as a Nonpolar Molecule δ+ H δ- C H H δ+ δ+ H δ+Molecular Geometry: Tetrahedral
  • 48. Methane as a Nonpolar Molecule H C H H H No Dipole Moment = Nonpolar! No positive or negative side overallMolecular Geometry: Tetrahedral