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Section 6.2 Covalent Bonding and Molecular Compounds PowerPoint Presentation
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Section 6.2 Covalent Bonding and Molecular Compounds

Section 6.2 Covalent Bonding and Molecular Compounds

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Section 6.2 Covalent Bonding and Molecular Compounds

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  1. Section 6.2Covalent Bonding and Molecular Compounds Sharing is caring!

  2. Chapter 6 Vocabulary • molecule (molecular compound) • molecular formula • bond energy • electron-dot notation • Lewis structure • structural formula • single bond • resonance Reading Guide Section 2 • write answers on same paper as Section 1 • due at end of hour Laughing Gas

  3. Section 6.2Covalent Bonding • Main Ideas: • Covalent bonds form from shared electrons • Bond lengths and energy vary from molecule to molecule • Atoms tend to form bonds to follow the octet rule

  4. Section 6.2Covalent Bonding • Main Ideas: • Dots placed around an element’s symbol can represent valence electrons • Electron-dot notations can represent compounds. • Some atoms can share multiple pairs of electrons.

  5. Section 6.2Covalent Bonding • Main Ideas: • Resonance structures show hybrid bonds • Some compounds are networks of bonded atoms

  6. Covalent Bonds • The word covalent is a combination of : • co- (latin) “together” • valere - “to be strong” • Two electrons shared together have the strength to hold two atoms together in a bond.

  7. Covalent bonds • Nonmetals hold on to their valence electrons. • They can’t give away electrons to bond. • But still want noble gas configuration. • Get it by sharing valence electrons with each other = covalent bonding • By sharing, both atoms get to count the electrons toward a noble gas configuration.

  8. Molecules • Many elements found in nature are in the form of molecules: • a neutral group of atoms joined together by covalent bonds. • Example: air contains oxygen molecules, consisting of two oxygen atoms joined covalently • Called a “diatomic molecule” (O2) • 7 diatomics: H2, N2, O2, F2, Cl2, Br2, I2

  9. Molecule

  10. Molecular Compounds • Compounds that are bonded covalently (like in water, or carbon dioxide) are called molecular compounds • Molecular compounds tend to have relatively lower melting and boiling points than ionic compounds

  11. Molecular Compounds • Thus, molecular compounds tend to be gases or liquidsat room temperature • Ionic compounds are solids (crystals) • A molecular compound has a molecular formula: • Shows how many atoms of each element a molecule contains

  12. Molecular Formulas • The formula for water is written as H2O • The subscript “2” behind hydrogen means there are 2 atoms of hydrogen; if there is only one atom, the subscript 1 is omitted • Molecular formulas do not tell any information about the structure (the arrangement of the various atoms).

  13. 3. The ball and stickmodel is the BEST, because it shows a 3-dimensional arrangement. These are some of the different ways to represent ammonia: 1. The molecular formulashows how many atoms of each element are present 2. The structural formula ALSO shows the arrangement of these atoms!

  14. + + + + How does H2 form? • The nuclei repel each other, since they both have a positive charge (like charges repel). (diatomic hydrogen molecule)

  15. + + How does H2 form? • But, the nuclei are attracted to the electrons • They share the electrons, and this is called a “covalent bond”, and involves only NONMETALS!

  16. Bond Dissociation Energies... • The total energy required to break the bond between 2 covalently bonded atoms • High bond energy = strong bond • High bond energy usually means the chemical is relatively unreactive, because it takes a lot of energy to break it down.

  17. Section 6.2Covalent Bonding • Main Ideas: • Covalent bonds form from shared electrons • Atoms tend to form bonds to follow the octet rule

  18. Section 6.2Covalent Bonding • Main Ideas: • Dots placed around an element’s symbol can represent valence electrons • Electron-dot notations can represent compounds. • Some atoms can share multiple pairs of electrons.

  19. OBJECTIVES: • Determine the number of valence electronsin an atom of a representative element. • Explain how atoms tend to form bonds to follow the octet rule. • Use electron-dot notation to represent valence electrons and represent compounds

  20. Valence Electrons are…? • The electrons responsible for the chemical properties of atoms, and are those in the outer energy level. • Valence electrons- The s and pelectrons in the outer energy level • the highest occupied energy level • Core electrons – are those in the energy levels below.

  21. Determining Valence Electrons • Atoms in the same column... • Have the same outer electron configuration. • Have the same valence electrons. • The number of valence electrons are easily determined. It is the group number for a representative element • Group 2: Be, Mg, Ca, etc. • have 2 valence electrons

  22. Valence electron – electron in the highest occupied energy level of an atom. *To find the number of valence e- in an atom of a representative element, simply look at its group number

  23. Electron Dot diagrams… (px) X (py) (s) (pz) shows valence e-’s symbol represents nucleus & core e-’s Each side = orbital (s or p) dot = valence e- (8 max) don’t pair up until they have to (Hund’s rule)

  24. The Electron Dot diagram for Nitrogen N • Nitrogen has 5 valence electrons to show. • First we write the symbol. • Then add 1 electron at a time to each side CCW. • Now they are forced to pair up. • We have now written the electron dot diagram for Nitrogen.

  25. Sec 6.2 Practice problems 19. Page 200 • Li e. C • Ca f. P • Cl g. Al • O h. S

  26. The Octet Rule • In Chapter 5, we learned that noble gases are unreactive in chemical reactions • In 1916, Gilbert Lewis used this fact to explain why atoms form certain kinds of ions and molecules • The Octet Rule: in forming compounds, atoms tend to achieve a noble gas configuration; 8 in the outer level is stable • Each noble gas (exceptHe, which has 2) has 8 electrons in the outer level

  27. Electron dot structure– a notation that depicts valence electrons as dots around the atomic symbol of the element. Octet rule – atoms react by gaining or losing electrons so as to acquire stable e- structure of a noble gas, usually 8 valence e-

  28. F Covalent bonding • Fluorine has seven valence electrons (but would like to have 8)

  29. F F Covalent bonding • Fluorine has seven valence electrons • A second atom also has seven

  30. F F Covalent bonding • Fluorine has seven valence electrons • A second atom also has seven • By sharing electrons…

  31. F F Covalent bonding • Fluorine has seven valence electrons • A second atom also has seven • By sharing electrons…

  32. F F Covalent bonding • Fluorine has seven valence electrons • A second atom also has seven • By sharing electrons… • …both end with full orbitals

  33. Covalent bonding • Fluorine has seven valence electrons • A second atom also has seven • By sharing electrons… • …both end with full orbitals F F 8 Valence electrons

  34. Covalent bonding • Fluorine has seven valence electrons • A second atom also has seven • By sharing electrons… • …both end with full orbitals F F 8 Valence electrons

  35. Lewis Structure • Replace the two bonded electrons with a straight line… F F F F

  36. A Single Covalent Bond is... • A sharing of two valence electrons. • Only nonmetals and hydrogen. • Different from an ionic bond because they actually form molecules. • Two specific atoms are joined. • In an ionic solid, you can’t tell which atom the electrons moved from or to

  37. Carbon covalent bonding C ….. special example - can it really share 4 electrons: 1s22s22p2? 2p 1s 2s C Yes, due to e- promotion!

  38. Carbon covalent bonding C ….. special example - can it really share 4 electrons: 1s22s22p2? 2p 1s 2s C Yes, due to e- promotion!

  39. How to show the formation… • It’s like a jigsaw puzzle. • You put the pieces together to end up with the right formula. • Carbon is a special example - can it really share 4 electrons: 1s22s22p2? • Yes, due to electron promotion! • Another example: lets show how water is formed with covalent bonds, by using an electron dot diagram

  40. H O Water • Each hydrogen has 1 valence electron - Each hydrogen wants 1 more • The oxygen has 6 valence electrons - The oxygen wants 2 more • They share to make each other complete

  41. O Water • Put the pieces together • The first hydrogen is happy • The oxygen still needs one more H

  42. O Water • So, a second hydrogen attaches • Every atom has full energy levels Note the two “unshared” pairs of electrons H H

  43. Another way of indicating bonds • Often use a line to indicate a bond • Called a structural formula • Each line is 2 valence electrons H O H H O H =

  44. Electron Dot/Lewis Structure Rules • Central atom is least electronegative (furthest left and down) • Count up the total number of electrons to make all atoms happy=8 except H=2 (Need). • Add up all the valence electrons (Have). • Subtract; then Divide by 2 = number of bonds • Fill in the rest of the valence electrons so each atom, except for hydrogen, beryllium, and boron, satisfies the octet rule. • Check that the number of valence electrons used is the same as added in step 3.

  45. Lewis Structure for Water

  46. Example • NH3, which is ammonia • N – central atom; has 5 valence electrons, wants 8 • H - has 1(x3) = 3 valence electrons, wants 2(x3) = 6 • NH3 has 5+3 = 8 • NH3 wants 8+6 = 14 • (14-8)/2= 3 bonds • 4 atoms with 3 bonds N H

  47. Examples • Draw in the bonds; start with singles • All 8 electrons are accounted for • Everything is full – done with this one. H H N H

  48. Lewis Structural Example H N H H

  49. Sec 6.2 Practice problems Set C p.176 #1-4 1. Draw the Lewis structure for ammonia, NH3 2. Draw the Lewis structure for hydrogen sulfide, H2S

  50. 3. Draw the Lewis structure for silane(silicon tetrahydride), SiH4 4. Draw the Lewis structure for phosphorus trifluoride, PF3