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Chemical Bonding

Chemical Bonding. Ch. 8. IONIC. Atoms in ionic compounds transfer electrons when they bond and become charged. This occurs when nonmetals bond with metals. COVALENT. Covalent bonds – atoms held together by sharing electrons.

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Chemical Bonding

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  1. Chemical Bonding Ch. 8

  2. IONIC • Atoms in ionic compounds transfer electrons when they bond and become charged. • This occurs when nonmetals bond with metals.

  3. COVALENT • Covalent bonds – atoms held together by sharing electrons. • Bonding allows atoms to fill their energy level and become more stable. • Sharing electrons form covalent bonds and occurs when nonmetals bond with nonmetals. Each bond shares 2 electrons.

  4. Atoms share electrons in covalent bonds to attain a noble-gas electron configuration. • Valence electrons are electrons located on the outer energy level. We can represent the valence electrons with dots. • Atoms usually acquire a total of 8 electrons in the outer level. This is called the octet rule. • H and He, need 2 electrons to fill their valence. • All other elements need 8 electrons to fill their valence.

  5. For a noble gas, the outermost energy level for the electrons is filled. • In forming the molecule HF, the F atom attains the electron configuration of Ne.

  6. Molecule – atoms joined together by covalent bonds • Atoms in molecules share electrons when they bond.

  7. NH3 is a covalent molecule. • Diatomic molecule – consists of 2 atoms • Br.I.N.Cl.H.O.F. • Chlorine (Cl) exists as a diatomic molecule • The diatomic molecule F2 contains a single covalent bond.

  8. Electron Dots and Bonding • Lewis structures use electron-dot diagrams to show how electrons are arranged in atoms. • Electrons are represented by dots. • A pair of electrons may be represented by 2 dots or a dash. • The electron dot structure for H2 is H:H or H―H

  9. The pairs of valence electrons that do not participate in bonding in diatomic oxygen molecules are called unshared pairs.

  10. Use NASB to draw dot diagrams. • N – electrons needed to fill valence (8 or 2) • A- electrons available in atom • S – electrons shared = N – A • B – formed bonds (S 2)

  11. In Cl2, the total number of unshared pairs of electrons is 6. • Cl + Cl  Cl―Cl or Cl:Cl • The diatomic molecule N2 contains a triple covalent bond. • N + N  N≡N or N⋮⋮N • In the N2 molecule, there is only one unshared pair of electrons in each nitrogen atom.

  12. The HI molecule contains only one single covalent bond. • H• + I  H―I or H:I • There are 4 single covalent bonds in a molecule of CH4. H H • C + 4H•  H:C: H or H C H H H • There are 2 double covalent bonds in a molecule of CO2. • C 2 O  O═C═O or O::C::O • Carbon monoxide has a triple covalent bond. C O  C≡O OR C⋮⋮O

  13. Bonding Theories • According to VSEPR theory, molecules adjust their shapes to keep pairs of valence electrons as far apart as possible. • VSEPR – • Valence Shell Electron Pair Repulsion • A stereoactive set is a shared pair or an unshared pair of electrons around the central atom.

  14. The shape of a molecule of CO2 is linear.

  15. The shape of a molecule of HCN is linear.

  16. The shape of a molecule of CH4 is tetrahedral.

  17. The shape of a molecule of NH3 is trigonal pyramidal.

  18. The shape of a molecule of H2O is bent.

  19. According to VSEPR theory repulsive forces between unshared pairs of electrons causes water molecules to have their shape.

  20. Bond angle = 109.5 degrees Example: CH4 Bond angle = 106.5 degrees Example: NH3

  21. Bond angle = 104.5 degrees Example: H2O Bond angle = 120 degrees Example: CO32-

  22. Bond angle = 118.6 degrees Example: O3 Bond angle = 180 degrees Example: CO2

  23. Polar Bonds • In a polar covalent bond, electrons are shared unequally. Some atoms (elements) are more electronegative and attract electrons more. They will gain a slightly negative charge. • The electronegativity difference determines the type of bond.

  24. HF contains a polar covalent bond. • F forms a more polar bond with H than Cl, Br, or I. • An ionic bond would be expected in a molecule of LiF.

  25. Intermolecular forces – forces between 2 molecules • Van der Waals forces - 2 weakest attractions between molecules • Dipoleinteractions – polar molecules attracted to one another • Dispersion forces – caused by the motion of electrons

  26. Hydrogen bonds – strongest intermolecular forces • Hydrogen covalently bonded to a very electronegative atom is also bonded to an unshared electron pair of another electronegative atom. • H2O is a polar molecule. • 2 H• + O  :O:H • H

  27. http://enstrophy.colorado.edu/~mohseni/ASEN5519Spring2004/TypesOfChemicalBonds1.htmlhttp://enstrophy.colorado.edu/~mohseni/ASEN5519Spring2004/TypesOfChemicalBonds1.html

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