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Covalent Bonding

Covalent Bonding. Mr. Solsman Chapter 8. Apply the octet rule to atoms that form covalent bonds. Describe the formation of single, double, and triple covalent bonds. Contrast sigma and pi bonds. Relate the strength of a covalent bond to its bond length and bond dissociation energy.

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Covalent Bonding

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  1. Covalent Bonding Mr. Solsman Chapter 8

  2. Applythe octet rule to atoms that form covalent bonds. • Describethe formation of single, double, and triple covalent bonds. • Contrastsigma and pi bonds. • Relatethe strength of a covalent bond to its bond length and bond dissociation energy.

  3. Atoms gain stability when they share electrons and form covalent bonds. • Lower energy states make an atom more stable. • Gaining or losing electrons makes atoms more stable by forming ions with noble-gas electron configurations. • Sharing valence electrons with other atoms also results in noble-gas electron configurations.

  4. Atoms in non-ionic compounds share electrons. • The chemical bond that results from sharing electrons is a covalent bond. • A moleculeis formed when two or more atoms bond.

  5. Diatomic molecules (H2, F2 for example) exist because two-atom molecules are more stable than single atoms.

  6. The most stable arrangement of atoms exists at the point of maximum net attraction, where the atoms bond covalently and form a molecule.

  7. When only one pair of electrons is shared, the result is a single covalent bond. • The figure shows two hydrogen atoms forming a hydrogen molecule with a single covalent bond, resulting in an electron configuration like helium.

  8. In a Lewis structure dots or a line are used to symbolize a single covalent bond. • The halogens—the group 17 elements—have 7 valence electrons and form single covalent bonds with atoms of other non-metals.

  9. Atoms in group 16 can share two electrons and form two covalent bonds. • Water is formed from one oxygen with two hydrogen atoms covalently bonded to it .

  10. Atoms in group 15 form three single covalent bonds, such as in ammonia.

  11. Atoms of group 14 elements form four single covalent bonds, such as in methane.

  12. Sigma bondsare single covalent bonds. • Sigma bonds occur when the pair of shared electrons is in an area centered between the two atoms.

  13. Double bonds form when two pairs of electrons are shared between two atoms.

  14. Triple bonds form when three pairs of electrons are shared between two atoms.

  15. A multiple covalent bond consists of one sigma bond. The pi bondis formed when parallel orbitals overlap and share electrons.

  16. The strength depends on the distance between the two nuclei, or bond length. • As length increases, strength decreases.

  17. The amount of energy required to break a bond is called the bond disassociation energy. The shorter the bond length, the greater the energy required to break it.

  18. An endothermic reactionis one where a greater amount of energy is required to break a bond in reactants than is released when the new bonds form in the products. • An exothermic reactionis one where more energy is released than is required to break the bonds in the initial reactants.

  19. Naming Binary Molecular Compounds • The first element is always named first using the entire element name. • The second element is named using its root and adding the suffix –ide.

  20. Prefixes are used to indicate the number of atoms of each element in a compound.

  21. Name each of the covalent compounds below: • CO2 • SO2 • NF3 • CCl4 • N2O4 • Write the formula for diarsenic trioxide

  22. Naming Acids • The first word has the prefix hydro- followed by the root of the element plus the suffix –ic. • The second word is always acid (hydrochloric acid is HCl in water). • HBr • HI • HF • HCN

  23. An oxyacid is an acid that contains both a hydrogen atom and an oxyanion. • Identify the oxyanion present. • The first word is the root of the oxyanion and the prefix per- or hypo- if it is part of the name, plus the suffix -ic if the anion ends in -ate or -ous if the oxyanion ends in -ite.

  24. Name the following oxyacids: • HClO • HClO2 • HClO3 • HClO4

  25. Structural Formulas • A structural formulauses letter symbols and bonds to show relative positions of atoms.

  26. Lewis Structures • Predict the location of certain atoms. • Determine the number of electrons available for bonding. • Determine the number of bonding pairs. • Place the bonding pairs. • Determine the number of bonding pairs remaining. • Determine whether the central atom satisfies the octet rule.

  27. Determine and draw the Lewis structure: • LiH

  28. BeH2

  29. BH3

  30. CH4

  31. NH3

  32. OH2

  33. FH

  34. CO2

  35. N2

  36. NO3

  37. Resonance Structures

  38. Two or more correct Lewis structures that represent a single ion or molecule are resonance structures. • The molecule behaves as though it has only one structure. • The bond lengths are identical to each other and intermediate between single and double covalent bonds.

  39. Try drawing resonance structures for: • SO2 • NO2- • O3 • SO32-

  40. Exceptions to the Octet Rule • Some molecules do not obey the octet rule. • A small group of molecules might have an odd number of valence electrons. • NO2 has five valence electrons from nitrogen and 12 from oxygen and cannot form an exact number of electron pairs.

  41. A few compounds form stable configurations with less than 8 electrons around the atom—a suboctet. • A coordinate covalent bond forms when one atom donates both of the electrons to be shared with an atom or ion that needs two electrons.

  42. A third group of compounds has central atoms with more than eight valence electrons, called an expanded octet. • Elements in period 3 or higher have a d-orbital and can form more than four covalent bonds.

  43. Draw the Lewis structure for: • ClF3 • XeF4 • SF6

  44. Shape: VSEPR Model • The shape of a molecule determines many of its physical and chemical properties. • Molecular geometry (shape) can be determined with the Valence Shell Electron Pair Repulsion model, or VSEPR model which minimizes the repulsion of shared and unshared atoms around the central atom.

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