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Chapter #10

Chapter #10. Chemical Bonding. CHAPTER 12 Forces Between Particles Noble Gas Configurations Ionic Bonding Covalent Bonding VSEPR Theory and Molecular Geometry Electronegativity Polar Bonds and Molecules. Atomic Stability.

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Chapter #10

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  1. Chapter #10 Chemical Bonding

  2. CHAPTER 12 • Forces Between Particles • Noble Gas Configurations • Ionic Bonding • Covalent Bonding • VSEPR Theory and Molecular Geometry • Electronegativity • Polar Bonds and Molecules

  3. Atomic Stability 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. He 1s2 Ne 1s22s22p6 Ar 1s22s22p63s23p6 Kr 1s22s22p63s23p64s23d104p6 Xe 1s22s22p63s23p64s23d104p65s24d105p6

  4. E Lewis Diagrams 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

  5. E E E   E   Lewis Diagrams • 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.

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

  7. S Lewis Diagrams 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

  8. Lewis Diagrams S 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

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

  14. 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

  15. Lewis Diagrams 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.

  16. Lewis Diagrams 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.

  17. Na+ + 1e- Cl- + 1e- Cl [Ne]3s23p5 Ion Formation 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-

  18. Ion Formation • 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

  19. Ion Formation • 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

  20. 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.

  21. 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

  22. 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

  23. 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.

  24. 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

  25. + - + - Molecular Compounds 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.

  26. - + + - Molecular Compounds The atomic orbitals overlap to form a newmolecular orbital. This is a stable configuration as each H atom can have a full 1s subshell (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.

  27. Molecular Compounds - + + - 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.

  28. Molecular Compounds 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.

  29. Lewis Structures 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

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

  31. Lewis Structure Construction Step 1 Connect each element with a single line Step 2 Use the “P” formula to determine 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 SO3 NO2+ CO2

  32. Lewis Structure of Carbon Dioxide First, connect atoms with lines O C O

  33. 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

  34. Lewis Structure of Carbon Dioxide Third, add extra lines (Three possible locations) C O O O O O O C C

  35. Lewis Structure of Carbon Dioxide Third, add extra lines (three possible locations) C O O O O O O C C Fourth, give each atom an octet of electrons

  36. Lewis Structure of Carbon Dioxide Third, add extra lines (three possible locations) 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

  37. Lewis Structure of Carbon Dioxide Third, add extra lines (three possible locations) 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

  38. 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 has more than its valence, then it is negative If the element owns less than its valence, then it is positive

  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 O C O O O C C O O Fifth, give each an atom a formal charge If the element has more than its valence, then it is negative If the element owns less than its valence, then it is positive O C O C O O O C O

  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 Fifth, give each an atom a formal charge If the element has more than its valence, then it is negative If the element owns less than its valence, then it is positive - O C O C O O O C 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 has more than its valence, then it is negative If the element owns less than its valence, then it is positive - + 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 has more than its valence, then it is negative If the element owns less than its valence, then it is positive - + - 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 has more than its valence, then it is negative If the element owns less than its valence, then it is positive - + + - O C O C O O O C O

  44. O O O O O O S S S O O O There are actuallythreepossible 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”. Resonance demonstrates how the loosely held pi electrons are free to move about. Or another way to look at it is that the pi electrons are spread over the entire molecule, making it more stable

  45. Molecular Geometry So far we have been considering how electrons are distributed between atoms in molecules and polyatomic ions. An important question is: How can we predict the shape of molecules and polyatomic ions?

  46. The only one way to join two atoms is with a line. All diatomic molecules and ions have a lineargeometry. Molecular shape is the geometry defined by the atoms making up the molecule. Molecular Shape The simplest polyatomic ion or moleculeis made of twoatoms: What is the shape of this type of molecule or ion?

  47. Predicting Molecular Shape Valence shell electron pair repulsion theory(VSEPR theory) allows us to predict the 3 dimensional shapeof molecules and polyatomic ions with >2 atoms. VSEPR theorystates that electrons in lone pairs andbondsmove as faraway from one another as possible to minimize repulsive interactions.

  48. Repulsion Angles

  49. Predicting Molecular Shape For a central atom with three electron regionsthere are two possibilities.

  50. Predicting Molecular Shape For a central atom with four electron regionsthere are three possibilities.

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