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Covalent Bonds Part 2
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Covalent Bonds Part 2

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  • 1. Covalent Bonds Part 2
  • 2.
    • When two atoms have a covalent bond, the valence orbital of one atom overlaps or merges with the valence orbital of another atom
    • Bonding orbital
      • localized region where bonding electrons can most likely be found
  • 3.
      • Single covalent bonds (one bonding pair)
      • Occurs when electron pair is shared in area centered between two atoms
      • Results if atomic orbitals overlap end to end
      • Sigma bonds form with:
        • Overlap of s-orbital and s-orbital
        • Overlap of s-orbital and p-orbital
        • Overlap of p-orbital and p-orbital
    Sigma bonds ( σ )
  • 4.  
  • 5.  
  • 6.  
  • 7. Multiple Bonds
    • Double Bond
      • Two bonding pairs of electrons
      • Consist of a sigma bond and one pi bond
    • Triple Bond
      • Three bonding pairs of electrons
      • Consist of a sigma bond and 2 wo pi bonds
  • 8. Pi bonds ( ∏ )
    • Formed when parallel orbitals overlap to share electrons
    • Electrons shared in a pi bond occupy the space above and below the line that represents where the two atoms are joined together
    • Double and triple bonds consist of a sigma bond and one or two pi bonds
  • 9.  
  • 10.  
  • 11.  
  • 12. Bond Strength
    • Recap: How are covalent bonds formed
    • Attractive and repulsive forces
    • Covalent bond broken when balance upset
    • Bond strength determined by different factors
  • 13. Bond length
    • The distance from the center of one nucleus to the center of the other nucleus of two bonded atoms during the point of maximum attraction
    • Determined by:
      • Size of atom
      • Number of electrons shared
  • 14. Shared Pairs…
    • INCREASE the number of shared pairs of electrons  DECREASE bond length
    • Triple bond has shorter bond length than single bond length
    • Single bonds are weaker than double bonds, which are weaker than triple bonds
  • 15. Energy in Bonds
    • When bond is formed  E released
    • When bond is broken  E is required
    • Bond Dissociation Energy
      • Amount of energy required to break a specific covalent bond
      • Breaking bonds requires adding energy
      • Positive value kJ/mol
      • Sum of bond dissociation energies for all bonds in a compound determines the chemical potential energy available in a molecule of that compound
  • 16. Relationship Between Bond Energy and Bond Length
    • The closer atoms are bonded together, the more energy is required to break the bond
    • DECREASE bond length = INCREASE bond dissociation energy
    • Which of the following has the greatest bond energy? Which has the least?
      • F 2 , O 2 , N 2
  • 17.
    • Total Energy change of chemical reaction determined by the energy of bonds broken and formed
    • Endothermic Reactions
      • Greater amount of energy required to break the existing bonds in the reactants than is produced in the new bonds formed in the products
    • Exothermic Reactions
      • More energy is released forming new bonds in the products than is required to break the bonds in the initial reactants
  • 18. Naming Molecules
    • Binary molecular compound
      • Covalently bonded compound containing only two different elements
      • Composed of 2 different nonmetals
        • No ions or metals
  • 19. Naming Binary Molecular Compounds
    • The first element in the formula is always named first, using the entire element name
    • The second element in the formula is named using the root of the element and adding the suffix – ide
    • Prefixes are used to indicate the number of atoms of each type that are present in the compound
      • Exception: first element in formula never uses mono -
    • Drop final letter in the prefix when the element name begins with a vowel
    • Hydrogen bonded to 7A halogens (drop mono)
  • 20.
    • 1 - mon(o)
    • 2 - di
    • 3 - tri
    • 4 - tetr(a)
    • 5 - pent(a)
    • 6 - hex(a)
    • 7 - hept(a)
    • 8 - oct(a)
    • 9 - non(a)
    • 10 - dec(a)
    Prefixes
    • H - hyd
    • C - carb
    • N - nitr
    • P - phosph
    • As - arsen
    • O - ox
    • S - sulf
    • Se - selen
    • F - fluor
    • Cl - chlor
    • Br - brom
    • I - iod
    Nonmetal roots
  • 21. Practice problems:
    • CO
    • P 2 O 5
    • CCl 4
    • As 2 O 3
    • NF 3
    • SO 2
    • Carbon monoxide
    • Diphosphorus pentoxide
    • Carbon tetrachloride
    • Diarsenic trioxide
    • Nitrogen trifluoride
    • Sulfur dioxide
  • 22. More practice…
    • H 2 O
    • NH 3
    • N 2 H 4
    • N 2 O
    • NO
    • Di hydrogen monoxide
    • Nitrogen trihydride
    • Dinitrogen tetrahydride
    • Dinitrogen monoxide
    • Nitrogen monoxide
    • Water
    • Ammonia
    • Hydrazine
    • Nitrous oxide
    • Nitric oxide
  • 23. Naming Acids
    • Molecules can be put in solution (water) and they make acids
    • If compound releases H+ ions when put in water solution, it is an ACID
    • Only name acids if molecule is put in water!!!
    • Two types
      • Binary Acids
      • Oxyacids
  • 24. Binary Acids
    • Hydrogen + one other element
    • Sometimes there are more than 2 elements
    • To name hydrogen, use prefix Hydro-
    • Root (or form of root) of the second element followed by suffix –ic
      • If there are more than 2 elements involved, the root of the second part of the name is the root of the polyatomic ion that acid contains
    • Add the word acid to the end
    • Example: HCN
      • Hydrocyanic acid
    • Example: HCl
      • Hydrochloric acid
  • 25. Oxyacids
    • Acids thant contain OXYANION
    • What is an oxyanion?
      • Polyatomic ion that contains oxygen
    • First: Determine anion present
    • Use a form of the root of the anion
    • Add suffix
      • Anion suffix –ate….oxyacid suffix= - ic
      • Anion suffix –ite….oxyacid suffix= -ous
    • Add the word acid
    • Example:HNO 3
      • Oxyanion: nitrate NO 3 -
      • Oxyacid name: nitric acid
    • Example:HNO 2
      • Oxyanion: nitrite NO 2 -
      • Oxyacid: nitrous acid
  • 26. Practice
    • HI
    • HClO 3
    • H 2 SO 4
    • H 2 S
    • HClO 2
    • Hydroiodic acid
    • Chloric acid
    • Sulfuric acid
    • Hydrosulfuric acid
    • Chlorous acid
  • 27. Writing formulas
    • Write the symbols for the elements in the order mentioned in the name.
    • Write subscripts indicated by the prefixes. If the first part of the name has no prefix, assume it is mono-.
    • Prefixes  tell you SUBscripts fro each element
  • 28. Writing Formulas for Binary Covalent Compounds: Examples nitrogen dioxide NO 2 diphosphorus pentoxide P 2 O 5 xenon tetrafluoride XeF 4 sulfur hexafluoride SF 6 * Second element in ‘ide’ from mono 1 * Drop –a & -o before ‘oxide’ deca 10 nona 9 octa 8 heptaa 7 hexa 6 penta 5 tetra 4 tri 3 di 2
  • 29. Air Pollution Class Work
    • Many common air pollutants for acids when dissolved in a water solution
    • Complete the following table
    HNO 3 NO 2 Carbonic acid SO 2 Name of Acid Formula of Acid Name of Molecule Formula of Pollutant

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