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Bonding

Bonding. Chapter 7. Names of Elements (don’t write). The most common source for element names is a property of the elements. EX: nitrogen: Greek for nitron (niter) and genes (to be born). Niter was the name for any naturally occurring substance that contained nitrogen.

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Bonding

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  1. Bonding • Chapter 7

  2. Names of Elements (don’t write) • The most common source for element names is a property of the elements. EX: nitrogen: Greek for nitron (niter) and genes (to be born). Niter was the name for any naturally occurring substance that contained nitrogen. • Some elements get their name from their place of discovery, person of honor, or from the mineral from which they are obtained.

  3. Chemical Symbols (don’t write) • J. J. Berzelius, a Swedish chemist, is generally given credit for creating the modern symbols for the elements. • In some cases, the Latin form of the element was used. EX: Iron is Fe for ferric

  4. Chemical Formulas p56 • Warm up: why don’t some of the symbols match the names of the elements? • Formula:a combination of symbols that represents the composition of a compound. • Symbols C8H18 atoms of each

  5. Types of Bonds p56 • 1. Ionic bond Transfer of e- from a metal to a nonmetal forms an ionic compound. EX: Na+ + Cl- NaCl • Ionic compounds are crystal solids with a high melting point.

  6. Types of Bonds p56 • 2. Covalent bond sharing of e- between nonmetals resulting in a molecule. EX: H2 + O  H2O • Molecules have lower melting points and are mostly gases or liquids at room temp.

  7. Types of Bonds p56 • 3. Metallic bond share electrons among many atoms

  8. Bond Character p56 • Electrons are transferred when the electronegativity difference is high. • They are shared when the difference is low • A difference of 2.0 and higher will produce ionic bonds. • 0.4 and below= covalent (sharing) • 0.41 and 1.99= polar covalent (in between sharing and transferring)

  9. Example p56 EX: Magnesium + Oxygen |1.2 – 3.5| = 2.3 =Ionic (on periodic table)

  10. Bond Character p55 • Predict the bond character for the following pairs of elements using a CREW statement: (electronegativities on periodic table) B – P: |2.0 – 2.1|=0.1 covalent Br-Cl : |2.8-3.0| = 0.2 covalent Na-F: |0.9 – 4.0| =3.1 ionic

  11. Valence Electrons: Review 1 8 2 3 4 5 6 7 2 2

  12. Lewis Dot Diagrams p58 • Warm up: What kind of bond will form between Cs and O? • Step 1: Write the symbol. • Step 2: find number of valence electrons (table) • Step 3: Start at the bottom and place dots around symbol clockwise. One at a time, then paired.

  13. Lewis Dot Diagrams p58 • EX: carbon step 1: C step 2: group 4, 4 valence e- step 3: C

  14. Lewis Dot Diagrams p58 • EX: bromine step 1: Br step 2: 7 ve- step 3: Br

  15. Drawing Ionic Bonds p58 • Draw Lewis dots for each element. Na + Cl • Draw an arrow to show the e- exchange between atoms.

  16. Drawing Simple Ionic Bonds p58 • Now draw them together. Na Cl

  17. Exchange of Electrons

  18. Covalent Bonding (ch. 8) p60 • Warm up: what is the difference between an ionic and covalent bond? • We will represent covalet bonding by drawing Lewis Structures.

  19. Lewis Structures of Molecules p60 1. Count the total number of valence e- 2. Determine the central atom. • Often the unique atom (only one of it), and is usually written first. • Or the least electronegative element 3. Arrange the other atoms around the central atom.

  20. 4. Connect all atoms with one bond (structural formula). 5. Subtract the number of electrons used; two for each bond. 6. Distribute the remaining electrons in pairs around the atoms, satisfy the octet rule. the most electronegative atom gets them first. (Hydrogen only gets 2 electrons!!!)

  21. 7. If you run out of electrons, you need to form double or triple bonds. 8. If you have extra electrons, put them on the central atom in pairs (unshared pair). Q. How do you think the lengths and strengths of single, double, and triple bonds compare?

  22. Lewis Structures • EX: AsI3 1. Count valence e- [5 + (3 x 7)]= 26 2. Place As (least in number) in center. 3. Place the three Iodines around As. 4. Draw lines (bonds connecting them) I –As –I I

  23. 5. Subtract # used for bonds (26 – 6) = 20 e-. 6. Place e- around the three I atoms first because they are the most electronegative. I –As –I I 7. We did not run out of e- so we do not have any double or triple bonds.

  24. 8. We have 2 extra e- (20 starting – (3x6) = 2. Place them around the central atom. I –As –I I

  25. Lewis Structure of CH2O 1. Count total valence e- (C=4, H=1x2, O=6)= 12 valence e- 2. Place C in the center. 3. Place the 2 hydrogens and one oxygen around C. 4. Draw lines connecting H and O to C. 5. Subtract number of bonded e- from total. (12-6) = 6

  26. 6. Place e- around the oxygen atom first because it is the most electronegative. 7. We ran out of e- before carbon satisfied its octet. This means that we will have a double bond between the carbon and oxygen. (Hydrogen cannot form double bonds). 8. There will be no unshared e-.

  27. Lewis Structure of CH2O H C O H

  28. Exceptions to the Octet Rule P62 • Warm up: how do you think the length and strength of a single, double and triple bond compare? Hydrogen only needs 2 ve- • Al and B only need 6 ve- • Ex) BCl3 • Elements above atomic number 15 can have more than an octet. • Ex) SF4

  29. Lewis Structures of Polyatomic Ions p62 • Lewis Diagrams for the polyatomic ions is the same, except the difference in charge (+ or -) must be accounted for. • Put final structure in brackets

  30. Lewis Structure of ClO31- 1. Count valence e-, including charge 7 + (3x6) + 1 = 26 total 2. Place Cl in the center. 3. Arrange the three O around it. 4. Draw bonds from O to Cl. 5. Subtract e- used in bonds (26 – 6) = 20 e-. 6. Place remaining e- around the three oxygens to satisfy octet.

  31. 7. There are two e- left over so we will not have multiple bonds. 8. Place the last two e- on the central atom. *For polyatomic ions: place structure in brackets with the charge indicated on outside as demonstrated

  32. 1- O Cl O O

  33. Resonance Structures p62 • Some molecules and ions cannot be represented by a single Lewis Structure. EX: Ozone O3 O=O—O O—O=O • Draw a 2-way arrow to show resonance

  34. Resonance of NO2

  35. Shapes of Molecules p64 • Warm up: what’s the difference between 2-dimensional and 3-dimensional? • video

  36. VSEPR Theory p64 • Valence shell, electron-pair repulsion (VSEPR) Theory states that repulsion between valence-level electrons causes them to be as far apart as possible.

  37. VSEPR Theory to Predict Molecular Geometries • Step 1: Write the Lewis Structure for the molecule. • Step 2: Represent the central atom in molecule by the letter A. • Step 3:Represent the atoms bonded to it by the letter X • Step 4: Represent the # of unshared pairs with the letter E • Step 5: Now refer to Molecular Structures handout

  38. Ex. AlCl3 • What are the bond angles?

  39. Ex. PCl3 • Bond angles? • Why do they differ?

  40. Ex. H2O • Notice that there are several bond angles… why? video

  41. Polar bonds • Occur if the electronegativity difference between 2 atoms is high. • Electrons are pulled to one end of the bond, making a negative end and positive end (dipole). Ex) H-F • An arrow is drawn to show direction electrons pulled. • Ex) H-F • Nonpolar bonds have electrons evenly dispersed. Ex) F-F

  42. Polar Molecules • When one end of a molecule has a high concentration of electrons and the other has lost them. • Ex. H2O, CHCl3 video • Nonpolar molecules can have polar bonds but are not polar overall. • Ex) CF4 video

  43. Lewis Structure of AlCl3 • EX: AlCl3 1. Lewis: Al has 3 valence and Cl has 7 x 3 for a total of 24 e-. 2. Place Al in center. 3. Place Cl around Al. 4. Draw bonds connecting Al and Cl. 5. Subtract bonds from total e- (24 – 6) = 18 e-

  44. 6. Now put electrons around the most electronegative atoms (Cl) first, try to satisfy octet rule. 7. There are no remaining e-. Note: this molecule is an exception to the octet rule because in this case Al only forms three bonds.

  45. Cl Cl Al Cl

  46. Predicting Molecular Geometry of AlCl3 1. Refer to structure. 2. Al will be represented by letter A. 3. Cl will be represented by letter X. 4. We have AX3. Refer to Table 1 and look for AX3. 5. According to the Table, the molecular geometry for AlCl3 is Trigonal-planar (aka- equilateral triangle).

  47. Next Example: PCl3 1. Lewis Structure: Cl Cl P Cl 2. P is represented by A. The unshared pair is represented by E. 3. Cl is represented by X. 4. We have AX3E. Refer to Table 3.

  48. Molecular Geometry of H2O • Remember, H2O has a Lewis Structure of H—O—H • The O is represented by A and the two H are represented by X. There are two unshared pairs represented by E2.

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