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Chapter 9-Covalent Bonds

This chapter covers the basics of covalent bonds, including the formation of molecules, Lewis dot structures, multiple bonds, bond strength and energy, naming covalent compounds, and molecular structures.

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Chapter 9-Covalent Bonds

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  1. Chapter 9-Covalent Bonds Agenda- Lab - Review - Quiz – Review –Chapter 8 / 9 Test – Chapter 8/9

  2. Covalent Bonds Section 1 • Why do atoms bond? • To become noble or stable • To achieve an octet (are exceptions)

  3. Covalent Bonds What is a covalent Bond? • Elements share electrons • Majority form between nonmetallic elements Result? • A Molecule is formed

  4. Lewis Dot • S • Cl • Ar Review: In your notes draw the following dot structures • H • N • O • C

  5. Lewis structures Groups and Bonds • Group 17 = 1 Bond HCl • Group 16 = 2 Bonds H2S • Group 15 = 3 Bonds PH3 • Group 14 = 4 Bonds • CCl4

  6. Sigma Bond • The single covalent bond is calls the… “Sigma Bond” • Shared electrons between two atoms

  7. Multiple Covalent Bonds Why Multiple Bonds? Hint: Think Noble. • To achieve an Octet! Example: C2H4 • Draw the central atoms • Attach the surrounding atoms • Make sure each atom has an octet

  8. Sigma and pi Bonds Sigma Bonds • Two atoms share electrons pi Bonds • Parallel orbitals over lap • Forms double bonds Example: C2H4

  9. Let’s take a closer look

  10. Lets take a closer look

  11. Strength and Energy Bond Strength • The shorter the bond length, the stronger the bond, the greater the bond-dissociation energy Bond Energy • Endothermic – more energy is needed to break the bond than is released • Exothermic – more energy is released during bond formation than is required to break it.

  12. Naming Covalent Section 2:Naming Covalent Molecules • Different than Ionic • First element = entire name • Second element = root + ide • Prefixes used to indicate the # of each type present in compound

  13. Prefixes Prefixes for Covalent Molecules

  14. Example • P205 • Follow your rules! di Phosphorus pent oxide

  15. Naming Acids 2 types of acids • Binary Acids • HCl, H2S, HBr, HCN • Oxyacids • H2SO4, HClO3, HClO2

  16. Binary Acids Example: HCl • Hydro + root of second element • Hydrochlor… • Add –ic then acid • Hydrochloric acid Name the following: HBr, HI, HF, HCN

  17. Oxyacids Example: H2SO4 and H2SO3 • Root of oxyanion present • Sulfur… • If oxyanion ends in …ate add -ic to the end • H2SO4 = sulfuric acid • If oxyanion ends in …ite add –ous to the end • H2SO3 = sulfurous acid

  18. Section 3 • Molecular Structures

  19. Section 3: Molecular Structures • Structural Formula • Uses letter symbols and bonds to show relative positions of atoms.

  20. Lewis Structures Step 1… • Predict the location of atoms • H is always terminal (end) • Central atom has the least attraction for shared electrons (closer to the left of the periodic table) • Determining Lewis Structures

  21. Step 2… • Find total # of valence electrons Step 3… • Determine # of bonding pairs. Divide # of valence electrons by 2 Step 4… • Place 1 pair (single bond) between the central atom and terminal atoms Lewis Structures

  22. Step 5… • Subtract pairs used from total possible pairs (step 3) • Place remaining pairs around terminal and central atom (octet) Step 6… • If central atom does not have octet, use lone pairs as double bonds. Lewis Structures

  23. Lewis Structures C O O 16 8 : Example: Carbon dioxide • Step 1- predict location • Step 2 – Total Valence Electrons = • Step 3 – Divide by 2 = pairs • Step 4 – Central Atom bonds • Step 5 – Place remaining pairs • Step 6 – Check Octet rule : : : : : : : : : : : : : : :

  24. Lewis Structures • WHAT’S WRONG WITH THIS PICTURE? • Carbon ~ octet? • Move electron pairs on each O to achieve octet around C : : O C O : : : :

  25. Positive Charges… • You must remove electrons from the total electrons available for bonding according to the charge. • Example NH4+ Total Valence Electrons = Subtract the charge Divide by 2 (8/2 = 4) ~ bonding pairs Charge Molecules 9 (9-1 = 8) + H N H H H

  26. Negative Charges… • You must add electrons to the total electrons available for bonding according to the charge. • Example PO43- = Total Valence Electrons = ADD the charge Divide by 2 (32/2 = 16) ~ bonding pairs Charged Molecules 29 (29+ 3 = 32) 3- : : O : : : O : O : P : : : : O :

  27. VSEPR Model Valence Shell Electron Pair Repulsion • Electrons are located as far apart as they can be • Shared electron pairs repel one another • Lone pairs also repel (even more) Hybrid Orbitals • S and p orbitals change to form new equal orbits • Each bond between atoms represents an s, p or d orbit

  28. Visualizing the Models Example #1: BeCl2 • 1st Draw the Lewis dot. • Determine the # of shared pairs and lone pairs around the central atom. • Shared pairs = 2 • Lone pairs = 0 • 2 Total hybrid bonds • S and p (sp)

  29. Visualizing the Models Example #1: AlCl3 (Exception to the Octet Rule) • 1st Draw the Lewis dot. • Determine the # of shared pairs and lone pairs • Shared pairs = 3 • Lone pairs = 0 • 3 Total hybrid bonds • s, p and another p (sp2)

  30. Visualizing the Models Example #1: CH4 • 1st Draw the Lewis dot. • Determine the # of shared pairs and lone pairs around the central atom! • Shared pairs = 4 • Lone pairs = 0 • 4 Total hybrid bonds • s, p, p and another p (sp3)

  31. Visualizing the Models Example #1: PH3 • 1st Draw the Lewis dot. : • Determine the # of shared pairs and lone pairs • Shared pairs = 2 • Lone pairs = 1 • 2 Total hybrid bonds • S and p (sp3) • Shape • Trigonal Pyramidal P H H H

  32. Electronegativity and Polarity Even or uneven sharing of electrons. Determined by the electronegativity Identical atoms share evenly Bonds between to different atoms. • one atom pulls the electrons closer • creates a relative negative and positive side of the molecule RULES: difference between electronegative #s 0.0 - 0.4 = Nonpolar covalent 0.4 - 1.7 = polar covalent > 1.7 = ionic bond

  33. Properties of Covalent Compounds Solubility • Polar Molecules soluble in polar substances • Non-polar in non-polar Intermolecular force between molecules is called the “van der Waals” force 3 types of intermolecular foces 1. Nonpolar (weak) = dispersion forces 2. Polar (weak) = dipole-dipole force 3. Hydrogen Bonds • Very Strong • between H and (F, N, O)

  34. Molecular Shapes sp 180o sp2 120o 109.5o sp3 107.3o 104.5o

  35. Molecular Shapes sp3d 90o / 120o sp3d2 90o

  36. Does the compound contain a metal? Ionic or Covalent NO YES The compound is covalent; use prefixes Is the metal a transition metal? YES NO Use I, II, III, IV, V – to indicate the charge of the metal Don’t use roman numerals: Don’t use prefixes Example: N2O – dinitrogen monoxide P2O5 – diphosphoruspentoxide Example: FeO – Iron (II) oxide Cu2S – Copper (I) Sulfide Example: NaCl – sodium chloride CaCl2 – calcium chloride

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