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Covalent Bonds

Covalent Bonds. Chapter 6 p. 188. 6.1 Covalent Bonding. 6.1 What is a covalent bond? 6.2 Lewis Structures 6.3 Shapes of Molecules Molecular modeling lab. 6.1 Covalent Bonds. Think about it …………. What does it means when we use the word bond to describe the relationship between people?.

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Covalent Bonds

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  1. Covalent Bonds Chapter 6 p. 188

  2. 6.1 Covalent Bonding 6.1 What is a covalent bond? 6.2 Lewis Structures 6.3 Shapes of Molecules • Molecular modeling lab

  3. 6.1 Covalent Bonds Think about it…………. • What does it means when we use the word bond to describe the relationship between people?

  4. 6.1 Standard 2a Students know atoms exchangeelectrons to form ionic bonds or, combine to form molecules by sharing electrons to form covalent bonds…….. ‘ELECTROSTATIC’ Teacher notes: cf ionic bond = transfer

  5. Learning target • How is a covalent bond different from an ionic bond? • What is a covalent bond?

  6. 6.1 Covalent Bonds • Shared pair of electrons holds the atoms together • ‘diatomic’ molecule (H2)

  7. 6.1 Covalent Bonds Think about it….. What is this diagram telling us?

  8. 6.1 Covalent Bonds p. 192(a)(c) • the bonded arrangement of atoms is lower(lower/higher) energy than the individual atoms) • Energy is released when bonds form

  9. 6.1 Covalent Bonds p. 192c  a • ‘low energy (c)  high energy (a); ‘up-hill’ • Energy is required to break bonds

  10. Think about it….. Read para 3 p. 193 Study Table 1 - Bond Energies and Bond Lengths for Single Bonds What is Table 1 telling you about covalent bonds? • Are all covalent bonds equally strong? • Are all covalent bonds equal length? • Trends/patterns? • HONS: how can this information be used to explain why energy is released when methane burns?* CH4 + O2 CO2 + H2O + energy

  11. Frayer: Electronegativity Definition: Ameasureoftheabilityofanatominachemicalcompoundtoattractelectrons. A measure of the ability of an atom in a chemical compound to attract electrons. (p. 137) Facts: Read/Notes ‘Electronegativity’ p. 137 (first 3 paragraphs)

  12. 6.1 Electronegativity • Scale: 0-4

  13. 5.1 Electronegativity p. 194 Frayer: examples: • Fluorine has the strongest demand for electrons • Francium has the lowest demand for electrons

  14. 5.1 Electronegativity p. 195 • Facts: Electronegativity difference between two atoms tells us what kind of bond will form between them.

  15. 5.1 Electronegativity p. 194 • Calculate the EN difference between sodium and chlorine

  16. 5.1 Electronegativity p. 195 • What kind of bonding occurs between sodium and chlorine?

  17. 6.1 Covalent Bonds: Electronegativity difference NOTES: Ionic bonds • Electronegativity difference is large (small/medium/large?) • atoms gain/lose electrons to form ions

  18. 6.1 Word Storm What is the word: POLAR “The difference between the electronegativity values of hydrogen and fluorine show that H and F atoms form a polar covalent bond.”

  19. 6.1 Word Storm: POLAR What does the word polar mean?

  20. 6.1 Word Storm: POLAR

  21. 6.1 Covalent Bonds Learning Target What is the difference between a polar covalent bond and a nonpolar covalent bond?

  22. Covalent bonds: Electronegativity difference Notes: Nonpolar covalent e.g H-H • electrons are shared equally between atoms • Electronegativity difference is small (small/medium/large?)

  23. Covalent bonds: Electronegativity difference Notes: Polar covalent bondse.g. C-O • electrons are unevenly shared between atoms • Electronegativity difference is medium (small/medium/large?) P.194

  24. 6.1 Practice Predict the type of chemical bonds (either ionic, polar covalent, non-polar covalent) the following compounds form: • KF • O2 • PBr P.194 YouTube: Dogs Teaching Chemistry - Chemical Bonds

  25. 6.1 Dipole Notes: Amoleculeorpartofamoleculethatcontainsbothpositivelyandnegativelychargedregions p. 195 Draw it!! PhET sim: molecule polarity

  26. 6.1 Bond Dipole: Draw it!! • δ = delta • δ+ = partially positive

  27. 6.1 Molecular Dipoles - NOTES • Individual bond dipoles add together to form an overall molecule dipole. Vid: covalent bonding (weiner 7000)

  28. 6.1 Molecular Dipoles - NOTES Some molecules are polar and some are non-polar. Examples: • Carbon dioxide = nonpolar molecule (b/c bond dipoles cancel) • Water = polar molecule (b/c of bent shape)

  29. What are the physical properties of metals? Make a list Solids (high melting points) Hardness Shiny (lustre) Electrical conductivity Malleable (can be shaped) Ductile (can be drawn into a wire) 6.1 Metallic Bonds:Think about it…….

  30. Standard c. Analyze and interpret provided data about bulk properties of various substances to support claims about the relative strength of the interactions among particles in the substance.

  31. Metallic Bonding‘Sea of Electrons’ model

  32. What is necessary in order to have an electric current? Electrical charges (electrons or ions) that are free to move (current) Youtube:metallic bonding and metallic properties 6.1 Metallic Bonds: notesElectrical Conductivity

  33. 6.2 Drawing and Naming Molecules Objectives • Draw the Lewis structure of a molecule when given its chemical formula

  34. 6.2 Lewis Structures of atoms Teacher notes: how electrons are arranged on page

  35. 6.2 Lewis Structure Hydrogen molecule, H2

  36. 5.2 Lewis Structures Practice Lets draw the Lewis structures of • Cl2 • HCl What is the type of covalent bond in each of these? (polar or nonpolar covalent) • Nonpolar covalent • Polar covalent

  37. 5.2 Lewis Structures Notes: Unshared Pair OR LONE PAIR: anonbondingpairofelectronsinthevalenceshellofanatom. Draw it p. 200

  38. 6.2 Lewis Structures A Lewis Structure tells us how electrons are shared in a molecule. A Lewis Structure does NOT tell us about the shape of a molecule.

  39. 5.2 Lewis Structures Practice Procedure: Lewis structure rules p. 201 • H2O More Practice p. 202 # 1, 2 Even more Practice p. 207 # 7 a-c p. 217 # 30 a-e

  40. 5.2 Lewis Structures: Multiple Bonds Double Bond = 2 bonding pairs of electrons Examples: • O2 • C2H4 A double bond has four electrons, so it is shorter and stronger than a single bond.

  41. 5.2 Lewis Structures: Multiple Bonds Skills Toolkit see p. 205 Practice Hint: only use multiple bonding if you have tried single bonding first Triple Bond = 3 bonding pairs of electrons Examples: • N2 • A triple bond has six electrons, so it is shorter and stronger than a single or double bond.

  42. 5.2 Lewis Structures: Multiple Bonds Skills Toolkit see p. 205 Practice Hint: only use multiple bonding if you have tried single bonding first Practice: p. 205 # 1,2 p. 217 32 a,b,c

  43. 6.2 Polyatomic Ions Learning Target How do we draw a Lewis structure of a polyatomic ion?

  44. 6.2 Lewis Structures of Polyatomic ions see Practice Hint on p. 203 Example: • For positive charged ions subtract (add/subtract) electrons from the count • For negative charged ions add (add/subtract) electrons from the count • Ammonium ion • Sulfate ion Practice: p. 203 #1,2; p. 207 #7d

  45. 6.2 Lewis Structure of Polyatomic Ions Practice Draw the Lewis structures the all the important polyatomic ions we highlighted on our ions sheet. (see Practice Hint on p. 203) • Ammonium • Sulfate • Hydroxide • Nitrate • Nitrite • Cyanide • Peroxide • Sulfite • Carbonate • Bicarbonate • Phosphate

  46. 6.2 Think about it…… Different non-metal elements have a tendency to form specific numbers of bonds to become isoelectronic with a noble gas. Take a look at all the Lewis structure examples we have done so far. How many covalent bonds do the following elements tend to form? • Hydrogen? • Carbon? • Nitrogen? • Oxygen, Sulfur? • The halogens?

  47. 6.3 Molecules

  48. 6.2 Diatomic molecules Only nobles gases are monatomic Diatomic molecules: H2 = Hydrogen N2 = Nitrogen F2 = Fluorine O2= Oxygen I2 = Iodine Cl2 = Chlorine Br2 = Bromine Memory aid: Have No Fear Of Ice Cold Beer See Ions Sheet

  49. 6.3 Resonance Structures Draw the Lewis structure of ozone, O3. Notes: Anyoneoftwoormorepossibleconfigurationsofthesamecompoundthathaveidenticalgeometrybutdifferentarrangementsofelectrons. Any one of two or more possible configurations of the same compound that have identical geometry but different arrangements of electrons. Read Resonance Structures p. 206

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