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Chap 6: Covalent Compound

Chap 6: Covalent Compound. HRW 6.1 Chemical Bonding. Define chemical bond. Explain why most atoms form chemical bonds. Describe ionic & covalent bonding. Explain why most bonding is not purely covalent or ionic. Classify bonding type according to electronegativity differences.

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Chap 6: Covalent Compound

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  1. Chap 6: Covalent Compound

  2. HRW 6.1 Chemical Bonding • Define chemical bond. • Explain why most atoms form chemical bonds. • Describe ionic & covalent bonding. • Explain why most bonding is not purely covalent or ionic. • Classify bonding type according to electronegativity differences. • Define molecule & molecular formula. • Explain relationships between potential energy, distance between approaching atoms, bond length and bond energy. • State the octet rule. • List steps to write Lewis structures. • Determine Lewis structures with single or multiple bonds or both. • Describe how resonance structures are used.

  3. 6-1: Types of Chemical Bonding • Ionic - results from the _______________ between large numbers of _________ and __________. • Covalent - results from the __________of electron pairs between two ________. • Metallic - results from chemical attraction between metal atoms and the sea of electrons surrounding them.

  4. What about covalent compounds? • The electrons in each atom are attracted to the ___________of the other. • The electrons ___________ each other, • The nuclei ____________each other. • They reach a distance with the __________________ energy. • The distance between is the ______________.

  5. + + How does H2 form? • The nuclei repel

  6. + + How does H2 form? • The nuclei repel • But they are attracted to electrons • They share the electrons

  7. Figure 6-2 • _____________ = space that shared electrons move within

  8. Energy and Stability • Energy is __________ when atoms from a covalent bond • Potential Energy = stored energy • Potential energy determines ___________ • Bond length is the distance between two atoms at their __________ potential energy • Figure 6-4

  9. Potential Energy 0 Internuclear Distance

  10. Tr32 p 165 Bond Length & Stability What’s H2 molecule’s bond length? What does the dashed line represent? How much energy must be added to the bond to break it? Why doesn’t He form a diatomic molecule?

  11. Bond Energy • Bond forming = __________________ • _____________ , Negative • Bond breaking = ______________ • Endothermic, __________ • As bond length____________, more energy is _______. Therefore, more energy is needed to break the bond. • As bond length _________, _________ the bond energy

  12. Visual Concepts Chapter 6 Bond Energy

  13. Tr 32G p. 173 What is the trend between number of bonds and bond energy (bond strength) ? What is the trend between bond number and bond length?

  14. Check Your Understanding! • Why is the H2 molecule more stable than the individual atoms that bond to form it? • What is created when two atomic orbitals overlap?

  15. 3. What happens to the potential energy of two atoms as they approach each other to form a covalent bond? 4. What name is given to the distance between two atoms in a covalent bond at which the potential energy is minimum? 5. List an example of substances that have covalent bonds

  16. Covalent Bonding • Covalent bond: __________________ • Nonpolar covalent bond- electrons are ______________________ • Polar covalent bond- shared ____________

  17. E. Chart of electronegativities: • Element with highest electronegativity: ___ • Lowest electronegativity: _______

  18. Visual Concepts Chapter 6 Electronegativity

  19. Dipole • A molecule that has one partial positive end and one partial negative end • The symbol is used to mean ____________. • _____is used to show a partial positive charge • _____ is used to show a partial negative charge charge • example: __________ • Because the F atom has a partial negative charge, the electron pair is more likely to be found nearer to the fluorine atom

  20. H H H Cl VI. Bond Polarity A. Non-Polar Covalent (_______ Covalent) Bond 1. Electrons are shared _______. 2. The bond has 100% covalent character and 0% _____ character. 3. The electronegativity difference between atoms in the bond is __. • Example: H2 B. Polar Covalent Bond 1. Electrons are shared _________. 2. The electronegativity difference between atoms in the bond is < ___. 3. Example: HCl

  21. (Rule of 2.0) C. Ionic Bond 1. Electrons are transferred from one atom to the other. 2. The electronegativity difference between atoms in the bond is ≥ ___. • Example: LiF D. Examples: What type of bonds are each of the following: (a) KF (b) HBr (c) Br2 (d) H2S (e) CsCl The electronegativity difference is so great that the electron is transferred to the non-metal.

  22. Tr31 Predicting Bond Type from Electronegativity Differences EN for K = 0.8 & Cl = 3.0. What type bond for KCl? H = 2.1; S = 2.5. For H2S? P = 2.1; H = 2.1. For PH3?

  23. Electronegativity Problems • Determine the electronegativity differencebond typeand the more-negative atom for the following: • C and H • C and S • O and H • Na and Cl • Cs and S

  24. Tr30 Fig. 6.1 p. 162 Ionic vs. Covalent Bonding Why does the ionic bonding occur between atoms of different sizes? Why are similar-sized atoms covalently bonded?

  25. Comparison of electron density of nonpolar H-H bond with polar H-Cl bond Why is the electron density greater around chlorine?

  26. Chapter 6 Lecture Notes Part 2

  27. Bonding in Metals • Metallic Bonds: • Metallic bonds- consists of the attraction of the __________________________ for the ___________charged _______________ • Properties of metals- • __________ • ___________ • _____________ • : Conducts heat & electricity • : Can be drawn into wires • : Can be hammered or forced into shapes

  28. 6.2 Drawing and Naming MoleculesElectron Dot diagrams • A way of keeping track of _____________ electrons. • Write the symbol. • Put one dot for each valence electron. • Don’t pair up until they have to. X

  29. Tr26A Fig 6.10 p 170 Electron Dot Notation Why are we interested in the outermost e-s? How does electron-dot notation help us determine bond formation?

  30. The Electron Dot diagram for Nitrogen • Nitrogen has 5 valence electrons. • First we write the symbol. • Then add 1 electron at a time to each side. • Until they are forced to pair up.

  31. Electron Dot Structures Symbols of atoms with dots to represent the valence-shell electrons 1 2 13 14 15 16 17 18 H He:      LiBe B  C  N  O : F :Ne :      

  32. Learning Check A. X would be the electron dot formula for 1) Na 2) K 3) Al   B.  X  would be the electron dot formula  1) B 2) N 3) P

  33. The Octet Rule: The Diatomic Fluorine Molecule F 1s 2s 2p Each has seven valence electrons F 1s 2s 2p

  34. The Octet Rule: The Diatomic Oxygen Molecule O 1s 2s 2p O 1s 2s 2p

  35. The Octet Rule: The Diatomic Nitrogen Molecule N 1s 2s 2p N 1s 2s 2p

  36. Lewis structures show how __________________ are arranged among atoms in a molecule. Lewis structures reflect the central idea that stability of a compound relates to _______________ electron configuration. Shared electrons pairs are covalent bonds and can be represented by two dots (:) or by a single line ( - ) Lewis Structures

  37. The HONC Rule Hydrogen (and Halogens) form __________________ bond Oxygen (and sulfur) form __________________ bonds One double bond, or two single bonds Nitrogen (and phosphorus) form _______________ bonds One triple bond, or three single bonds, or one double bond and a single bond Carbon (and silicon) form ________________ bonds. Two double bonds, or four single bonds, or a triple and a single, or a double and two singles

  38. Covalent Bonding • Single Covalent Bonds: a bond in which ____atoms share a pair of ______________ • Unshared pairs- _____________ or _________________ pairs; • pairs of e- not shared between atoms • Water molecule

  39. H H •• •• •• •• O O X. Lewis Structures of Covalent Compounds: A representation of the covalent bonding in a molecule. A. Covalent bonds are shown as ______. Example: H2 B. Lone pairs of electrons are shown as _____. Example: O2 C. ONLY ________ electrons are shown. D. General steps for drawing Lewis structures: 1. Sum the valence electrons in the compound. 2. Add __ for each negative charge. Subtract 1 for each __ charge. 3. Generally place the element that makes the _________ number of bonds in the center. 4. Draw ______ bonds to the other atoms off of the central atom. 5. Place electrons around the ______ atoms until an _____ is reached. 6. If you run out of electrons, start forming ______ or _____ bonds. 7. If you have EXTRA electrons after all have octets, place them on the ______ atom. 8. In the end, all atoms should have an octet that need an octet (____ is an exception), and the total number of electrons should be placed on the molecule. (least electronegativity)

  40. 3. CH3NH2 Try These! 1. HBr 2. NH3

  41. Deficient Electrons Deficient 2 e-: Make a ____________ into a ___________. Deficient 4 e-: Make one single bond into a ____________ or make 2 single bonds into 2 _____________

  42. 1. Count the total number of valence electrons 2. Place the atom that makes most bonds in the middle. (least electronegative other than hydrogen) 3. Draw single bonds to the other atoms off of the central atom. 4. Place electrons around peripheral atoms to fill octet, then to central atom 5. if electrons run out, start making double, triple bonds 5. CO2 4. CH4 5. O3

  43. 1. Count the total number of valence electrons Add 1 electron for each – charge , subtract 1 from each + charge 2. Place the atom that makes most bonds in the middle. 3. Draw single bonds to the other atoms off of the central atom. 4. Place electrons around peripheral atoms to fill octet, then to central atom 5. if electrons run out, start making double, triple bonds 1. SO42- 2. NH4+1 3. IO-1

  44. Resonance Occurs when more than one valid Lewis structure can be written for a particular molecule. O O O • These are resonance structures. • the double side arrow is used to show that the actual molecule is the AVERAGE of the two.

  45. Resonance in Benzene, C6H6

  46. Binary Covalent Compounds Compounds between two _____________ First element in the formula is named first. Keeps its element name Gets a prefix if there is a subscript on it Second element is named second Use the root of the element name plus the _________ suffix Always use a prefix on the second element

  47. List of Prefixes • 1 = ______ • 2 = _______ • 3 = _______ • 4 = _______ • 5 = ______ • 6 = hexa • 7 = hepta • 8 = octa • 9 = nona • 10 = deka

  48. Naming Covalent Compounds P2O5= CO2 = CO = N2O =

  49. Practice – Write the Formula

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