Chapters 6


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Chapters 6

  1. 1. Chapter 6: Chemical Bonding1. Use the periodic table to infer the number of valence electrons in an atom and draw its electron dot structure.2. Be able to explain the types of bonds that atoms can form.3. List the characteristics of the different types of chemical bonds.4. Define the vocabulary words.5. Use electronegativity values to classify a bond 1
  2. 2. Valence Electrons  Electrons in the highest occupied energy level of an element’s atoms  For representative elements, the number of valence electrons is the same as the group number of that element (Page 414)  Shown in electron dot structures 3 5 2 Symbol 1 of the 6 element 8 4 7 Right, left, top, bottom (1,2,3,4) Then 12 o’clock and counterclockwise (5,6,7,8) 2
  3. 3. Valence Electrons (cont’d) Electrons in the highest occupied energy level of an element’s atoms Can be figured out using the group numbers in the periodic table.  Ex: The elements of Group 1A (hydrogen, lithium, sodium, etc.) all have a valence number of 1, which means there is 1 electron in the highest occupied energy level. The elements of group 7A (fluorine, chlorine, bromine, etc.) have 7 electrons in the outer energy level.  The valence numbers also tell us the likely oxidation state of that element. More on this later. 3
  4. 4. Oxidation StatesThe oxidation state of an atom is the charge it has when it gains or loses electrons to form it’s most stable electron configuration. Valence Number Oxidation State (charge on the ion) 1 +1 2 +2 3 +3 } CATIONS 5 -3 6 -2 } ANIONS 7 -1 4
  5. 5. The Octet Rule Gilbert Lewis used this to explain why atoms form certain kinds of ions and molecule. In forming compounds, atoms tend to achieve the electron configuration of a noble gas (8 valence e -)  Recall that each noble gas (except He) has 8 electrons in its highest energy level and a general electron configuration of ns2 np6 Exceptions: Molecules with an odd number of electrons, more than an octet (PCl5), and less than an octet (very rare) Example: NO2 has seventeen valence electrons [Nitrogen contributes five and each oxygen 5 contributes 6 (2 x 6 =12)]
  6. 6. The Octet Rule An atom’s loss of an electron produces a cation, or positively charged ion. The most common cations are those produced by the loss of valence electrons from the metals, since most of these atoms have 1-3 valence electrons. Let’s look at sodium (a 1A metal) as an example: -e- Na 1s22s22p63s1 Na+ 1s22s22p6 6
  7. 7. Practice Problems  Write the electron dot structure for each of the following: 1. Na 2. Al 3. N 4. S 5. Kr 6. Chloride ion 7. Oxide ion  Refer to pages 414, 417, and 418 for answers. 7
  8. 8. Practice ProblemsPlease write the oxidation numbers of the following:1) Na 11) Po2) Al 12) Ga3) F 13) Cr4) Cl 14) N5) Mg6) P7) Ca8) Sb9) I10) Sc 8
  9. 9. Common Polyatomic IonsHydroxide: OH- Permanganate: MnO4-Bicarbonate: HCO3- Ammonium: NH4+Carbonate: CO32- Acetate: C2H3O2-Sulfate: SO42- Hydrogen-Sulfite: SO32- Phosphate: HPO42-Phosphate: PO43- Dichromate: Cr2O72-Perchlorate: ClO4-Nitrate: NO3-Nitrite: NO2-Chlorate: ClO3-Cyanide: CN- 9
  10. 10. Chemical Bonding Chemical energy & potential energy stored in chemical bonds Atoms prefer a low energy condition Atoms that are bonded have less energy than free atoms- more stable. To combine atoms: energy is absorbed To break a bond: energy is released (AB) 10
  11. 11. Chemical Bonds Created when two nuclei simultaneously attract electrons When electrons are donated or received, creating an ion (anion, cation) In most elements, only valence electrons enter chemical reactions Atoms of everyday substances are held together by chemical bonds (water, salt anti-freeze) 11
  12. 12. Types of Chemical Bonds1) Ionic Bond: chemical bonding that results fromthe electrical attraction between cations and anionswhere atoms completely give their electron(s) away 12
  13. 13. Types of Chemical Bonds2) Covalent Bond: chemical bonding that resultsfrom the sharing of electron pairs between two atoms.The electrons are “owned” equally by the two atoms. 13
  14. 14. Relative Forces ofAttraction  Ability of a nucleus to hold its valence electrons (Group 7A has a greater ability to hold on to its valence electrons than Group 1A)  Ionization energy: energy required to lose an electron (As atomic number increases down a group, the most loosely bound electrons are more easily removed, so ionization energy decreases. For the most part, it increases along each period.) 14
  15. 15.  Electron affinity – tendency to gain an electron (Energy is released) Electronegativity – measure of the electron attracting power of an atom when it bonds with another atom * Fluorine (4.0) is the highest * Cesium (0.7) is the lowest – least ability to attract bonding electrons and thus the greatest tendency to lose an electron * Noble gases are not assigned electronegativities because these elements do not generally form bonds 15 (inert)
  16. 16.  The periodic trend of the electronegativities is the same as that of the ionization energies. Thus, as the atomic number increases along a period, the electronegativity increases. As the atomic number increases down a group, the electronegativity decreases. In general, metals have a low electronegativity and nonmetals have a high electronegativity 16
  17. 17. Electronegativity and Bond TypesCovalent Bonds: bonding between elements with an electro- negativity difference of 1.7 or less.Nonpolar-Covalent Bonds: covalent bond in which electrons are shared evenly by the bonded atoms with an electronegativity difference of 0 – 0.3.Polar-Covalent Bonds: covalent bond in which the bonded atoms have unequal attraction of the shared electrons, and have an electronegativity difference of 0.4 – 1.7Ionic Bonds: bonding due to difference in electric charge of two elements due to loss/gain of electrons. Must have an electro- negativity of 1.8 – 4.0. 17
  18. 18. Electronegativity and Bond Types Water is a polar molecule, because the electrons are not shared evenly by the hydrogen and oxygen.Ionic Bonds: bonds in which electrons are donated from one atom to another and have an electronegativity difference of 1.8 or higher. 18
  19. 19. Electronegativity and Bond Types Using the electronegativity values found on page 161 of your book, predict the types of bonds the following will form. 1) O2 2) NaCl 3) N24) Knowing that the electronegativity of sulfur is 2.5, what type of bond will sulfur form with: a) hydrogen b) cesium c) chlorine 19
  20. 20. ElectronegativityElectronegativity is a measure of how strongly an element can remove an electron from another element. 20
  21. 21. Ionic (Electrovalent) Bonds chemical bond The strongestComplete transfer of electron(s) from one element to another Generally formed when metals combine with nonmetals (Groups 1-2a w/ 5-7a) Coulombic forces – electrostatic force in which two oppositely charged ions are mutually attracted Usually occurs when the difference in electronegativities is 1.8 or greater 21
  22. 22. NaCl – Ionic Bond  Draw 22
  23. 23. Writing Ionic Compounds Beryllium fluoride Calcium oxide Scandium sulfide Aluminum chloride 23
  24. 24. Ionic Solids  Form crystal lattice (orderly, repeating, three-dimensional pattern)  The charges and relative sizes of the ions determines the crystal structure  The number of ions of opposite charge that surround the ion in a crystal is called the coordination number of the ion. 24
  25. 25.  Poor conductors of electricity (no free electrons) High melting point High boiling point Brittle and break easily under stress Liquid or aqueous: good conductors of electricity but ionic bond is dissolved 25
  26. 26. The NormalArrangement of anIonic Crystal - + - + + - + - - + - + + - + - - + - +Opposite charges attract 26
  27. 27. Arrangement when Stressis Applied - + + - - + - + + - + - - + - + + - - + Adjacent to ions with same charge (repulsion) 27
  28. 28. Crystal Lattice isDestroyed - + + - - + - + + - + - - + - + + - - +Crystal melts, vaporizes, or dissolves in water (ions free to move about)Cleavage – splitting along a definite line 28
  29. 29. Covalent Bonding Electrons are shared One atom does not have enough pull on the electron to take it completely from the other atom Occurs when electronegativity difference is less than 1.8 Covalently Bonded Solids: 1. Softness 2. Poor conductor of electricity and heat 3. Low melting point 29
  30. 30. Lewis Structures Single covalent bond – one shared pair of electrons: H· + ·H H : H or H H Double covalent bond – two shared pairs of electrons : : : : : : : O. + : O. : O: : O: or : O O: . . Triple covalent bond – three shared pairs of electrons : : . N. + . N. : N : : : N : or : N N: . . 30 Note: all of these obey the octet rule
  31. 31.  Coordinate covalent bond – one atom contributes both bonding electrons NH3 + H+ [NH4]+ ammonia hydrogen ion ammonium ion H + : H: N : H + H+ H: N : H : : H H The structural formula shows an arrow that points from the atom donating the electrons to the atom receiving them. Refer to page 444 31
  32. 32. How to Construct Lewis StructuresStep 1: Determine the type and # of atoms in molecule CH3I has 1 Carbon, 3 Hydrogens and 1 IodineStep 2: Write electron dot notation for each type of atom .. . . .. . .. . C H· I .Step 3: Determine the total # of electrons available in the atoms to be combined. C 1 x 4e- = 4e- I 1 x 7e- = 7e- H 3 x 1e- = 3e- 32 14 e-
  33. 33. How to Construct Lewis StructuresStep 4: Arrange the atoms to form a skeleton structure for the molecule. Then connect the atoms by electron-pair bonds. H . .. . . . .. HC I HStep 5: Add unshared pairs of electrons to each non- metal atom so that each is surrounded by 8. H . .. . . . .. . . . ... HC I 33 H
  34. 34. Electron Dot Practice: Compounds1) H2O 5) CCl2H22) H2O2 6) NH33) HCN 7) N24) AlF3 8) CO2 34
  35. 35.  A single water molecule is a good example of covalent bonding between atoms. The hydrogen atoms “share” their electrons with the larger oxygen atom so that oxygen now has a full outer level with 8 electrons and each hydrogen has a full outer level with 2 electrons. Oxygen has a higher electronegativity than hydrogen, so there is actually an uneven sharing of electrons, resulting in a polar molecule. More on this later. e- e- e- e-shared electrons 8p+ shared electrons e - e- 8n0 e- e- e- e- 1p+ 1p+ 35
  36. 36.  Bond dissociation energy: total energy required to break the bond between two covalently bonded atoms (remember that energy is measured in joules or kilojoules) H–H + 435 kJ H . + .H Resonance Structures: refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure. 36
  37. 37. Bond Length vs. Bond EnergyThere is a correlation between bond length and the amount of potential energy stored in that bond. For example: Bond Bond Length (pm) Bond energy (Kj/mol) C C 154 346 C C 134 612 C C 120 835 C N 147 305 C N 132 615 C N 116 887 N N 145 163 N N 125 418 37 N N 110 945
  38. 38.  Molecular orbitals – when two atoms combine and their atomic orbitals overlap Sigma bond - molecular orbital that is symmetrical along the axis connecting two atomic nucleiIn both of these examples, the p orbitals are overlapping and sharing electrons. 38
  39. 39.  pi bond – weaker than sigma bond; usually sausage-shaped regions above and below the bond axis (Page 445) 39
  40. 40. Examples of Sigma and Pi bonds H3C – CH3 H2C = CH2 HC – CH – – 40
  41. 41. VSEPR Theory (page 200)VSEPR Theory (Valence Shell Electron-Pair Repulsion theory):states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far away from each other as possible, thus determining the shape of molecules. 41
  42. 42. VSEPR Theory So then why is H2O bent, but BeF2 is linear? H2O BeF2The answer is the free electron pairs. Oxygen has 2 pairs, beryllium has none. 1 : : O. . Be. 2 . 42
  43. 43. VSEPR Shapes Linear Trigonal-Planer Bent/Angular Tetrahedral Trigonal-Pyramidal Trigonal-Bipyramidal Octahedral(#s 3, 5 and 7 are coordinate covalent bonds!) 43
  44. 44. VSEPR TheoryThese free electron pairs repel each other because they have a negative charge, and so they force those atoms that are . covalently bonded to be pushed as far away as possible. . . . 44
  45. 45.  Hybridization – several atomic orbitals mix to form the same total number of equivalent hybrid orbitals (CH4 – Page 457 ) *Note: An sp3 orbital is an example of a hybrid. 45
  46. 46. Types of CovalentBonding  1. Nonpolar – when atoms have the same or similar electronegativity; when the atoms in the bond pull equally and the bonding electrons are shared equally. (Generally a difference of 0.0- 0.4) * Examples: Diatomic elements (H2 , N2 , O2 , F2 , Cl2 , I2 , Br2) * Nonpolar Covalent: Bonded Hydrogen Atoms 46
  47. 47.  2. Polar – unequal sharing of electrons * Pairing of atoms when one has a stronger attraction for the electrons * Most compounds are polar covalent Examples: H2O , NH3 , HF , HCl *Polar covalent also called dipoles *Creates partial charges Partially + Partially -Example: HCl (0.9 difference of the electro- negativities) H (2.1) Cl (3.0) Electronegativity difference is less than 1.7 47
  48. 48. Example: water.An uneven distribution of the electrons results because the oxygen has a higher electron affinity than the hydrogens. Thus, you have a negative and positive end of the molecule: polarity. Because this molecule has 2 poles, it is called a dipole molecule. δ-δ = delta or e- e- overall Oxygen e - e- e - 8p+ e- 8n0 e- e- Hydrogens e- e- 1p+ 1p+ δ+ 48
  49. 49. Attractions Between Molecules Molecules are often attracted to each other by a variety of forces. The intermolecular attractions are weaker than either an ionic or covalent bond. These attractions are responsiblefor determining whether a molecular compound is a gas, liquid, or solid at a given temperature. Here is a list of these various attractions:1. van der Waals forces: weakest type of intermolecular attractions. Dispersion (London) forces: weakest of all molecular interactions, caused by the motion of electrons. Increases as the # of electrons increases. Ex: Cl & F are gases at STP; Br is liquid at STP; I is solid at STP. 49
  50. 50. 2. Dipole Interactions: attraction of polar molecules to one another. Remember that polar molecules are like magnets; they have a positive and negative end. A glucose molecule in water has many dipole interactions since both water and glucose are polar. The positive poles of the water molecule are attracted to the negative poles on the glucose and vice versa. 50
  51. 51. 3. Hydrogen Bonds: attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshaired pair of electrons on another electronegative. Hydrogen bonding always involves hydrogen. Hence the name. Duh. The hydrogen bonding between water molecules dictates many of the properties of water. It also explains why water is a liquid rather than a gas at room temperature. 51
  52. 52. Network Solids  Macromolecules  Covalent network of atoms bonded  Absence of molecules throughout the solid  Properties 1. Hardness 2. Poor conductor of electricity (electrical insulation) 3. Poor conductor of heat 52
  53. 53.  Examples: diamond graphite (carbon) Does not melt - vaporizes to a gas at 3500 °C Carbon atom Covalent bond Boron nitride (BN), asbestos, silicone carbide (SiC, grindstones), silicone dioxide (SiO2 , quartz) 53
  54. 54. Metallic Bonds  Most metallic elements, except liquid mercury, are solids at room temperature and exhibit a crystal structure (zinc)  Arrangement of stationary positive metal ions surrounded by a “sea of mobile electrons”” - - - - + + + - - - - - + - + - + - - - - - + + + 54
  55. 55.  Properties: 1. Malleability – ability to be hammered into different shapes 2. Ductility – ability to be drawn into wire 3. Conductor of heat 4. Conductor of electricity 5. Luster – shine 6. Tenacity – structural strength (resistance to being pulled apart) 55
  56. 56. Alloys  Mixtures composed of two or more elements, at least one of which is a metal  Properties usually superior to those of the component elements  Sterling silver – silver and copper  Bronze – copper and tin  Steel – iron, carbon, boron, chromium, manganese, molybdenum, nickel, tungsten, vanadium (Interstitial alloy)  Interstitial alloy – smaller atoms fit into spaces between larger atoms  Substitutional alloy – atoms of the components are about the same size (They can replace each other in the structure.) 56
  57. 57. Summary : Types ofBonds1. Ionic – complete transfer of electrons2. Covalent – share electrons A. Nonpolar : same or similar electronegativity B. Polar – unequal sharing Electronegativity Difference: C < 2.0 ≤ I (Know exceptions) *Know table on page 4653. Network solids – covalent network of atoms (absence of molecules)4. Metallic – positive ions around a “sea of mobile electrons” 57
  58. 58. ElectronegativityDifferences and BondTypes  0.0-0.3 Nonpolar covalent  0.4-1.0 Moderately polar covalent  1.0-1.8 Very polar covalent  1.8 or greater Ionic 58
  59. 59. General Trends of theRepresentative Elements  Group 1A - lose one electron  Group 2A - lose two electrons  Group 3A - lose three electrons  Group 4A - share, lose or gain 4 e-  Group 5A - share, gain three electrons  Group 6A - share, gain two electrons  Group 7A - gain one electron  Group 8 - do not react, noble gases 59
  60. 60. Think!  Why is it possible to bend metals but not ionic crystals?  In an ionic compound, ions of like charge do not have mobile electrons as insulation. When forced into contact by physical stress, the ions of like charge repel, causing the crystal to shatter. 60
  61. 61. Attr actionsbetweenMolecules weakName and describe theattractive forces that hold groups ofmolecules together. 61
  62. 62. Van der Waals Forces  Weaker than either an ionic or covalent bond  Responsible for determining whether a molecular compound is a gas, liquid, or solid at a given temperature  Two types: dispersion forces and dipole interactions  Dispersion – caused by motion of electrons; dispersion generally increases as the number of electrons increases 62
  63. 63.  Example of dispersion forces: (Refer to Group 7A) F and Cl are gases at STP; Br is a liquid at STP, and I is a solid at STP Dipole interactions – electrostatic attractions between oppositely charged regions (Example: water) 63
  64. 64. Hydrogen Bonds Strongest of the intermolecular forces Important in determining the properties of water and biological molecules such as proteins Has only about 5% of the strength of an average covalent bond 64
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