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Chapter 8

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  1. Chapter 8 Chemical Bonding

  2. Bonds • Forces that hold groups of atoms together and make them function as a unit. • We will consider three major categories of chemical bonds: Ionic Bonds, Covalent Bonds, and Metallic Bonds.

  3. Examples of Bonding Types

  4. Lewis Structures • Since bonds hold atoms close together, then the valence electrons are responsible for bonding since they are on the outside of an atom. • It has been recognized for a long time that the noble gases have great chemical stability. With few exceptions they are unreactiveorinert. • The noble gases have 8 valence electrons with the exception of He which has 2.

  5. Lewis Structures He 1s2 Ne 1s22s22p6 Ar 1s22s22p63s23p6 Kr 1s22s22p63s23p64s23d104p6 Xe 1s22s22p63s23p64s23d104p65s24d105p6

  6. E Lewis Structures The electronic configuration of the noble gases is described as being energetically stable. We can draw a Lewis diagram to illustrate the number of valence electrons an atom has. In a Lewis diagram valence electrons are represented by dots placed above, below and to the left and right of the atoms symbol. e.g. element with 4 valence electrons

  7. E E E   E   Lewis Structures • There are two simple rules to keep in mind when drawing Lewis diagrams: • Place one dot in each of the four locations before doubling up. • There can be only a maximum of 2 dots in any one location.

  8. First write the electron configuration: • 1s1 • Identify the number of valence electrons. • 1 valence electron. H For a representative element it is easy to identify the number of valence electrons as this is equal to the group number. Lewis Structures What is the Lewis diagram for H?

  9. S Lewis Structures What is the Lewis diagram for S? • First write the electron configuration: • [Ne]3s23p4 • Identify the number of valence electrons. • 6 valence electrons Alternatively you can recognize that S is in group VIA so has six valence electrons

  10. S Lewis Structures What is the Lewis diagram for S? • First write the electron configuration: • [Ne]3s23p4 • Identify the number of valence electrons. • 6 valence electrons Alternatively you can recognize that S is in group VIA so has six valence electrons

  11. S Lewis Structures What is the Lewis diagram for S? • First write the electron configuration: • [Ne]3s23p4 • Identify the number of valence electrons. • 6 valence electrons Alternatively you can recognize that S is in group VIA so has six valence electrons

  12. S Lewis Structures What is the Lewis diagram for S? • First write the electron configuration: • [Ne]3s23p4 • Identify the number of valence electrons. • 6 valence electrons Alternatively you can recognize that S is in group VIA so has six valence electrons

  13. S Lewis Structures What is the Lewis diagram for S? • First write the electron configuration: • [Ne]3s23p4 • Identify the number of valence electrons. • 6 valence electrons Alternatively you can recognize that S is in group VIA so has six valence electrons

  14. S Lewis Structures What is the Lewis diagram for S? • First write the electron configuration: • [Ne]3s23p4 • Identify the number of valence electrons. • 6 valence electrons Alternatively you can recognize that S is in group VIA so has six valence electrons

  15. LEWIS STRUCTURES OF THE ELEMENTS He H C B F Li N O Ne Be S Cl P Al Mg Si Na Ar

  16. LEWIS STRUCTURES OF IONS (AFTER REMOVAL OR ADDITION OF ELECTRONS) 1+ He H B C F Li N O Ne Be 3+ S Cl P Mg Si Na Ar Al

  17. Lewis Structures The octet rule states that: “Atoms interact in order to obtain a stable octet of eight valence electrons” The octet rule works extremely well at describing the interactions of the representative elements.

  18. Lewis Structures One way in which atoms can interact to satisfy the octet rule is by transferring electrons between each other. Transferring of electrons results in the atoms acquiring net positive and negative charges. When an atom loses or gains electrons a simple ion is formed. Cations have more protons than electrons and are positive. Anions have more electrons than protons and are negative.

  19. Ionic Bonds • Formed from electrostatic attractions of closely packed, oppositely charged ions. • Formed when an atom that easily loseselectronsreacts with one that has a highelectron affinity.

  20. Na+ + 1e- Cl- + 1e- Cl [Ne]3s23p5 Ionization A Review Consider a Na atom what happens if it loses one electron? I.E. Na [Ne] [Ne]3s1 11 P and 10 e- 11 P and 11 e- Consider a Cl atom would you expect it to lose or gainelectrons? E.A. [Ne]3s23p6 17 P and 18 e- 17 P and 17 e-

  21. Metals tend to lose electrons forming positively charged ions called cations. • A representative metal will lose its group number of electrons to obtain a stable octet. • Na → Na+ + 1e-( Isoelectronic with Ne) • Mg → Mg2+ + 2e- (isoelectronic with Ne) • What would the charge be of the ion formed by a Li atom? • And which Noble gas is it isoelectronic with? +1 The ion formed would be Li+ Isoelectronic with He

  22. Noble Stability • Non-metals tend to gain electrons forming negatively charged ions called anions. • A representative non-metal will gain (8 - group number) electrons to obtain a stable octet. • O + 2e-→ O2- (isoelectronic with Ne) • S + 2e- → S2- (isoelectronic with Ar) • What would the charge be of the ion formed by a I atom? • Which Noble gas is it isoelectronic with? -1 The ion formed would be I- Isoelectronic with Xe

  23. Lewis Structure of NaCl Na+Cl-Na+Cl-Na+Cl- Cl-Na+Cl-Na+Cl-Na+ Forces between oppositely charged ions are called Ionic bonds. Each ion is surrounded by an octet of Electrons, thus making the ions stable.

  24. Crystal Lattice of NaCl Ionic compounds do not exist as discrete molecules. Instead they exist as crystals where ions of opposite charges occupy positions known as lattice sites. Ions combine in the ratio that results in zero charge to form ionic compounds. Which ions are the smaller ones? Crystal Lattice of NaCl

  25. Crystal Lattice of NaCl Ionic compounds do not exist as discrete molecules. Instead they exist as crystals where ions of opposite charges occupy positions known as lattice sites. Ions combine in the ratio that results in zero charge to form ionic compounds. Which ions are the smaller ones? Sodium Crystal Lattice of NaCl

  26. Sodium Chloride Lattice

  27. Molecular Compounds In our early lectures we defineda molecule as “as a compound Made of nonmetals.” Molecules exist as particles containing the number of atoms specified by their formula. e.g.a water molecule is a particle containing 2 hydrogen atoms and one oxygen atom and has the formula H2O.

  28. Molecular Compounds Non-metals may also complete their octetsby sharing electrons. This may occur between non-metal atoms of the same type: e.g. H2, O2, N2, Cl2, F2, I2, etc Or between different types of non-metal atoms: e.g. CO2, H2O, CH4, etc

  29. + - + - Covalent Bond Formation Consider two hydrogen atoms separated by a large distance. Each has 1 electron in a 1s atomic orbital. Why does the electron stay around the nucleus? Now lets bring the two atoms together so there orbitals overlap.

  30. - + + - The atomic orbitals overlap to form a newmolecular orbital. This is a stable configuration as each H atom can have a full 1s susbshell (like He) where the electrons spend most of their time shared between the atoms. In this arrangement each nucleus feels an inwards attraction to the two electrons. This is called covalent bonding.

  31. - + + - This new arrangement of protons and electrons is more stable than separate hydrogen atoms since the attraction of a proton to two electrons is a stronger attraction compared to one proton to one electron of a hydrogen atom.

  32. Lewis Structures • A single bond results when two atoms share one pair of electrons. • A lone pair, or unshared pair, of electrons is a pair of electrons that is not shared. • A bonding pair of electrons is a pair of electrons shared between two atoms.

  33. Multiple Bonds • A double bond results when two atoms share two pairs of electrons. • A triple bond results when two atoms share three pairs of electrons • Bond length is the distance between the nuclear centers of the two atoms jointed together in a bond.

  34. Covalent Bond Formation We can draw Lewis diagrams showing the arrangement of valence electrons in covalent compounds. In these diagrams we represent each pair of electrons between atoms as a line. So for the H2 molecule discussed previously the Lewis diagram would be: H – H All other electrons are represented by dots as described previously.

  35. Molecular Compounds Draw Lewis Structures of the following molecular compounds H H Nonbonding electons O a. H2O H H O Note each element has a Noble gas structure by electron sharing b. NH3 H N N H H H H H Covalent bonding e’s

  36. Simplified Lewis Structures Straight lines are used to indicate a shared pair, or a covalent bond. O H H Nonbonding electrons

  37. Lewis Structure Construction Step 1 Connect each element with a single line Step 2 Use the “P” formula to determine the number of extra bonds Step 3 Insert the extra bonds, to make double or triple bonds. Step 4 Give each atom an octet of electrons, except hydrogen Step 5 Determine the formal charge of each element N = number of atoms in molecule Q = number of hydrogen atoms V = total number of valence electrons P = 8(n-q) +2q - 2(n-1) - v Examples: Give Lewis Structures for the following H2CO3 CO2 SO3 NO2+

  38. Lewis Structure of Carbon Dioxide First, connect atoms with lines O C O Second, use “p” formula to determine the number of extra bonds. P = 8(n-q) + 2q – 2(n-1) - v P = 8(3-0) + 2(0) – 2(3-1) - 16 P = 24 + 0 – 4 - 16 4 extra bonding electrons 2 extra bonds 2 extra lines P = 4

  39. Lewis Structure of Carbon Dioxide Third, add extra lines C O O O O O O C C Fourth, give each atom an octet of electrons

  40. Lewis Structure of Carbon Dioxide Third, add extra lines C O O O O O O C C Fourth, give each atom an octet of electrons O C O O O C C O O

  41. Lewis Structure of Carbon Dioxide Third, add extra lines C O O O O O O C C Fourth, give each atom an octet of electrons O C O O O C C O O Fifth, give each an atom a formal charge If the element owns less than its valence, then it is positive If the elementhas more than its valence, then it is negative O C O C O O O C O

  42. Lewis Structure of Carbon Dioxide Third, add extra lines C O O O O O O C C Fourth, give each atom an octet of electrons O C O O O C C O O Fifth, give each an atom a formal charge If the element owns less than its valence, then it is positive If the elementhas more than its valence, then it is negative - O C O C O O O C O

  43. Lewis Structure of Carbon Dioxide Third, add extra lines C O O O O O O C C Fourth, give each atom an octet of electrons O C O O O C C O O Fifth, give each an atom a formal charge If the element owns less than its valence, then it is positive If the elementhas more than its valence, then it is negative - + O C O C O O O C O

  44. Lewis Structure of Carbon Dioxide Third, add extra lines C O O O O O O C C Fourth, give each atom an octet of electrons O C O O O C C O O Fifth, give each an atom a formal charge If the element owns less than its valence, then it is positive If the elementhas more than its valence, then it is negative - + - O C O C O O O C O

  45. Lewis Structure of Carbon Dioxide Third, add extra lines C O O O O O O C C Fourth, give each atom an octet of electrons O C O O O C C O O Fifth, give each an atom a formal charge If the element owns less than its valence, then it is positive If the elementhas more than its valence, then it is negative - + + - O C O C O O O C O

  46. Practice Draw the most stable Lewis structure for CO2.

  47. Practice Draw the most stable Lewis structure for CO2. O C O

  48. O O O O O O S S S O O O Sulfur Trioxide Lewis Structure There are actually three possible Lewis structures for SO3. - - 2+ 2+ - 2+ - - - Each of these three structures is equivalent. We say they are in “resonance” or that they are “resonance structures”.

  49. Resonance Form Rules

  50. Resonance Form Rules