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Chapter 9

Chapter 9. Covalent Bonding. Section 9.1. Atoms bond together because they want a stable electron arrangement consisting of a full outer energy level. Two ways that atoms can bond together are ionically & covalently.

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Chapter 9

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  1. Chapter 9 Covalent Bonding

  2. Section 9.1 • Atoms bond together because they want a stable electron arrangement consisting of a full outer energy level. • Two ways that atoms can bond together are ionically & covalently. • A covalent bond is a chemical bond that results from the sharing of the valence electrons. Covalent bonds are usually formed between elements close to each other on the periodic table and nonmetallic elements.

  3. Nonmetals bonded with nonmetals

  4. Section 9.1 • A molecule is formed when 2 or more atoms bond covalently. • Example • HI • F2 • Practice problems 1. HBr 2. O2 3. H2O

  5. ` • Diatomic molecules are molecules of like atoms. Diatomic molecules include hydrogen, nitrogen, oxygen, fluorine, chlorine, bromine, & iodine. • Write the formulas & structures for the diatomic molecules.

  6. The number of covalent bonds that these groups form are: • Halogens  single covalent bond • Group 6A  double/two covalent bond • Group 5A  triple/three covalent bond • Group 4A  four covalent bonds

  7. Practice • HF • OCl2 • H2S • N2 • PH3 • SiH4

  8. Section 9.1 • Single covalent bonds are also called sigma bonds, symbolized by the Greek letter s • A sigma bond results when orbitals overlap. • s orbitals can overlap s orbitals • s orbitals can overlap p orbitals • p orbitals can overlap p orbitals

  9. Section 9.1 • Multiple covalent bonds occur when more than one pair of electrons are shared. • A pi bond, p, is formed when parallel orbitals overlap to share electrons. • A double bond is made of 1 sigma bond & 1 pi bond • A triple bond is made of 1 sigma bond & 2 pi bonds

  10. Practice Draw the lewis dot structure for each molecule. • 1. PH3 • 2. H2S • 3. HCl • 4. CCl4 • 5. SiH4 • 6. CO2

  11. If H2O is water, what is H2O4? • Drinking, bathing, washing, swimming…

  12. Section 9.1 • The strength of covalent bonds depend on the length of the bond, the distance between the bonded nuclei. • Bond length is determined by the size of the atoms & the number of shared electron pairs. • As the number of shared pairs increases, the bond length decreases.

  13. The shorter the bond, the stronger the bond. • Energy is released when a bond forms & energy is added to break a bond. • Which has a shorter bond? H2 or S2 • The amount of energy needed to break a bond is the “Bond Dissociation Energy” • Among F2, O2, & N2, which would have the greatest bond dissociation energy? Least?

  14. Section 9.1 • Endothermic reactionsoccur when a greater amount of energy is required to break the existing bonds in the reactants than is released when the new bonds form in the product molecules • Exothermic reactions occur when more energy is released forming new bonds than is required to break bonds in the initial reactants

  15. Exothermic Rxn Endothermic Rxn Measuring heat

  16. Section 9.2 • Naming Binary Molecular Compounds • 1st element—use entire element name • 2nd element—use root word & “ide” • use prefixes to indicate # of atoms for each element • prefixes • mono—1 • di—2 • tri—3 • tetra—4 • penta—5 • hexa—6 • hepta—7 • octa—8 • nona—9 • deca--10

  17. Section 9.2 • **NEVER USE mono WITH 1ST ELEMENT** • Example • N2Cl3 • CO2 • Practice • 1. P2O5 • 2. NO2 • 3. CO • Common names of some molecular compounds—p. 249 ** Worksheet “Naming molecular compounds”

  18. What is the name for H2O? • What is the name for CH2O? • Sea water

  19. Naming acids • Binary acids—hydrogen plus one other element • Hydro + root of 2nd element + ic acid • HClhydrochloric acid • HBrhydrobromic acid • H2S  hydrosulfuric acid • Hydroiodic acid

  20. Oxyacids—contain hydrogen & oxygen and one other element • Root of other element + ic acid • H2CO3 • H2SO4 • HC2H3O2 • Phosphoric acid  • Nitric acid  • ** Worksheet “Naming acids”

  21. Section 9.3 • Structural formulas use letter symbols & bonds to show relative positions of atoms. • Hydrogen is always a terminal atom because it can bond with only 1 other atom. • Resonance is a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion.

  22. Resonance • Example • HCl • SO2 • Practice • 1. SO3 • 2. O3

  23. A coordinate covalent bond forms when one atom donates both electrons to be shared.

  24. 9.4 Molecular Shapes • Molecular shape is determined by the overlap of orbitals that share electrons. • The VSEPR model is used to determine the molecular shape of the molecule. VSEPR stands for • Valence • Shell • Electron • Pair • Repulsion model.

  25. F2 has linear shape • CO2 has linear shape • BH3has trigonal planar shape • CF4has tetrahedral shape • NH3has pyramidal shape • H2O has bent shape

  26. Hydridization is a process where atomic orbitals are mixed to form new, identical hybrid orbitals. • Carbon is the most common element that undergoes hybridization. • The hybrid orbital of carbon is called sp3. • It takes shape to look like a tetrahedral.

  27. 9.5 Electronegativity & Polarity • Remember that electronegativity is defined as the tendency of an atom to attract electrons. • Fluorine has the highest EN at 3.98 and Francium has the lowest EN at 0.7.

  28. Polar/nonpolar • Polar covalent bond is when electrons are unequally shared—large difference in EN. Electrons spend more time around the element that is more EN. The more EN element has a partial negative charge. A polar covalent bond is also called a dipole. • Electronegativity occurs in molecules. For example H2O & HCl • Page 263

  29. Polar/nonpolar • Nonpolar covalent bond is when electrons are shared equally—have small or no EN differences. Symmetrical molecules with balanced charges are nonpolar. (CCl4) These are pure covalent bonds. • Ionic bonds generally form when EN differences are 1.7 or greater.

  30. Polar, Nonpolar, or Ionic? • Example: • HCl • Practice: • SCl2 • H2S • CF4 • CS2

  31. Properties of covalent compounds-p. 266-267 • “like dissolves like” • Many different types of forces between molecules. **Worksheet “Shape/polarity of molecules”

  32. Mainly gases (some liquids), melt/vaporize easily, hardness varies. • Oxygen gas, acids, parafinn, graphite, diamonds

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