1 / 36

Chapter 9-Covalent Bonds

Chapter 9-Covalent Bonds. Agenda- Lab - Review - Quiz – Review –Chapter 8 / 9 Test – Chapter 8/9. C ovalent Bonds. Section 1. Why do atoms bond? To become noble or stable To achieve an octet (are exceptions). C ovalent Bonds. What is a covalent Bond?. Elements share electrons

avon
Download Presentation

Chapter 9-Covalent Bonds

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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 In your notes draw the following dot structures • H • N • O • C

  5. Lewis structures Draw the Lewis structures • HCl • CCl4 • PH3 • H2S • SiH4

  6. 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

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

  8. Let’s take a closer look

  9. Lets take a closer look

  10. Naming Covalent 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

  11. Prefixes Prefixes for Covalent Molecules

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

  13. 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

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

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

  16. 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

  17. Molecular Structures • Structural Formula • Uses letter symbols and bonds to show relative positions of atoms.

  18. 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

  19. 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

  20. 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

  21. 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 : : : : : : : : : : : : : : :

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

  23. 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

  24. 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 :

  25. 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

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

  27. 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)

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

  29. 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

  30. 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

  31. 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)

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

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

More Related