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Chemical Bonding

Chemical Bonding. Focus on Covalent Bonds. Molecules. Nonmetals share electrons Covalent bond : chemical bond caused by atoms sharing electrons Molecule : atoms joined by covalent bonds Diatomic : made of 2 atoms

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Chemical Bonding

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  1. Chemical Bonding Focus on Covalent Bonds

  2. Molecules • Nonmetals share electrons • Covalent bond: chemical bond caused by atoms sharing electrons • Molecule: atoms joined by covalent bonds • Diatomic: made of 2 atoms • 7 elements that exist naturally as diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2

  3. Formation of Diatomic Elements

  4. Single Covalent Bonds • When one pair of electrons is shared between 2 atoms

  5. Multiple Covalent Bonds • Double bonds: when 4 electrons are shared between two atoms • Triple bonds: when 6 electrons are shared between two atoms

  6. Strength of Covalent Bonds • Bond dissociation energy: amount of energy needed to break a bond • Stronger bonds need more energy to be broken • Double bonds are stronger than single bonds • Triple bonds are stronger than double bonds

  7. Length of Covalent Bonds • In stronger bonds, the atoms are closer together • Double bonds are shorter than single bonds • Triple bonds are shorter than double bonds

  8. Electronegativity • How much an atom pulls electrons in a covalent bond towards itself • Noble gases are not listed because they do not form compounds

  9. Nonpolar Covalent Bonds • Electrons are pulled equally by the atoms • Nonpolar molecules are not attracted to an electric field • Examples: H2, Cl2, O2

  10. Polar Covalent Bonds • Electrons are pulled unequally to the atoms in a covalent bond • The atom with the higher electronegativity value pulls the electrons more • Polar molecules align in an electric field • Dipole: polar molecule

  11. Electronegativity Difference and Bond Type • We can calculate the difference in Electronegativity to determine whether a bond is a nonpolar, polar, or ionic.

  12. Higher electronegativity = pulls electrons more • δ = partial charge (<1) • Polar molecules: one end of molecule is δ- and the other end is δ+ • Can also be shown w/ arrow pointing towards more electronegative atom

  13. Practice 8-4 Determine the bond type that exists between the following atoms. Show your work! • C-H • Na-F • Br-Cl • H-F • Fe-O

  14. Representing Molecules

  15. Lewis Structures: Another Example of an Electron Dot Structure Lewis structures are representations of molecules showing all electrons, bonding and nonbonding. Represents the valence electron distribution within the compound

  16. Goal of all Bonding • To give every element a full outer shell! • Most elements need 8 electrons but there are some exceptions.

  17. Exceptions to the Octet Rule • There are three types of ions or molecules that do not follow the octet rule: • Ions or molecules with an odd number of electrons. • Ions or molecules with less than an octet. • Ions or molecules with more than eight valence electrons (an expanded octet).

  18. Odd Number of Electrons Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons. Ex: NO: 11 electrons

  19. Fewer Than Eight Electrons • Hydrogen: can only take 2 ve- • Beryllium: can only take 4 ve- • Boron: can only take 6 ve-

  20. Elements that can Expand their Octet • When the central atom is a non-metal in period 3 or greater can expand to hold 8, 10, or 12 electrons.

  21. More Than Eight Electrons • The only way PCl5 can exist is if phosphorus has 10 electrons around it. • It is allowed to expand the octet of atoms on the 3rd row or below. • Presumably d orbitals in these atoms participate in bonding.

  22. Find the sum of valence electrons of all atoms in the polyatomic ion or molecule. If it is an anion, add one electron for each negative charge. If it is a cation, subtract one electron for each positive charge. PCl3 Writing Lewis Structures 5 + 3(7) = 26

  23. Writing Lewis Structures • The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds. Keep track of the electrons: 26  6 = 20

  24. Writing Lewis Structures • Fill the octets of the outer atoms. Keep track of the electrons: 26  6 = 20  18 = 2

  25. Writing Lewis Structures • Fill the octet of the central atom. Keep track of the electrons: 26  6 = 20  18 = 2  2 = 0

  26. Writing Lewis Structures • If you run out of electrons before the central atom has an octet… …form multiple bonds until it does.

  27. Lewis Structure Example Draw Lewis structures for the following molecules: • CH4 • NH3 • CCl4 • CO2

  28. LDS of Ions • When doing the LDS of ions – you need to add or subtract electrons from the total to account for the charge – PRIOR to starting the structure. • Ex: Ammonium = +1 charge • Subtract 1 electron from the total • Ex: Carbonate = -2 charge • Add 2 electrons to the total • Total structure must be surrounded by brackets and charge to show that it is an ion.

  29. Practice 8-5 Draw Lewis structures for the following: • water • BF3 • SO22- • H3O+ • sulfur dichloride • C2H6

  30. Resonance • Single bonds are longer than double bonds, so if we were to draw the LDS for the carbonate ion, we would expect to see one shorter bond and two longer bonds.

  31. Resonance • But in actuality – scientists have measured the length the all three bonds and found them to be identical. All three bonds were about 30% shorter than a single bond – it was a HYBRID of all three bonds.

  32. Resonance of the Carbonate Ion

  33. Practice 8-6 Draw Lewis structures for the following: • Sulfate ion • Carbon dioxide • Sulfur trioxide • Sulfur hexafluoride • Carbon monoxide

  34. Practice 8-7 Draw Lewis structures for the following: • O3 (ozone) • Boron trifluoride • Phosphorus pentafluoride • Nitrate ion

  35. Formal Charge and Lewis Structure Assessing the Best Possible Structure

  36. Definition and Equation • Formal Charge is the difference between the valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis Structure. (Chang 346) • Equation: Formal charge on an atom in a Lewis structure Total number of valence electrons in the free atom - Total number of nonbonding electrons - = ½ (total number of bonding electrons)

  37. Writing Lewis Structures • To assign formal charges. • For each atom, count the electrons in lone pairs and half the electrons it shares with other atoms. • Subtract that from the number of valence electrons for that atom: The difference is its formal charge.

  38. Writing Lewis Structures • The best Lewis structure… • …is the one with the fewest charges. • …puts a negative charge on the most electronegative atom.

  39. EX: Look at BF3 • Consider BF3: • Giving boron a filled octet places a negative charge on the boron and a positive charge on fluorine. • This would not be an accurate picture of the distribution of electrons in BF3.

  40. Using Formal Charge Therefore, structures that put a double bond between boron and fluorine are much less important than the one that leaves boron with only 6 valence electrons.

  41. Look at Phosphate Ion Even though we can draw a Lewis structure for the phosphate ion that has only 8 electrons around the central phosphorus, the better structure puts a double bond between the phosphorus and one of the oxygens.

  42. Phosphate Ion • This eliminates the charge on the phosphorus and the charge on one of the oxygens. • The lesson is: When the central atom is on the 3rd row or below and expanding its octet eliminates some formal charges, do so.

  43. Example • Write formal charges for the carbonate ion

  44. Example Problem • Write formal charges for the carbonate ion Reasoning and Solution: The Lewis Structure for the carbonate ion was developed in: The formal charges on the atoms can be calculated as follows: The C atom: The O atom in C=O : The O atom in C – O : Formal charge = 4 – 0 – 1/2 (8) = 0 Formal charge = 6 – 4 – 1/2 (4) = 0 Formal charge = 6 – 6 – 1/2 (2) = -1

  45. Example Problem Continued • The LDS for (CO32-) with formal charges written out is  • Note that the sum of the formal charges is -2, the same as the charge on the carbonate ion. 0 0 -1 -1

  46. Practice Exercise • Write the formal charges for the nitrite ion (NO2-). See ONE NOTE for the answer and explanation

  47. Important!! • Sometimes there is more than one acceptable Lewis structure for a given species. In such cases, we can often select the most plausible Lewis structure by using formal charges and the following guidelines: • For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present. • Lewis structures with large formal charges (+2, +3, and/or -2, -3…) are less plausible than those with small formal charges. • Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms.

  48. Valence Shell Electron Pair Repulsion Theory (VSEPR) “The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”

  49. What Determines the Shape of a Molecule? • Simply put, electron pairs, whether they be bonding or nonbonding, repel each other. • By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

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