1. Chemical Bonding

2. Do Now Define a compound. What is a compound made of? What are some examples of compounds?

3. Compounds A compound is a pure substance made of more than one kind of atom, and it can be broken down into its elemental components. Ex: CO 2 Ex: C 9 H 8 O 4

4. Chemical Bonds Almost everything is made up of a combination of atoms that are held together by chemical bonds. Chemical Bond = an attraction between atoms that allows the formation of chemical substances that contain two or more atoms The bond results from simultaneous attraction of electrons to 2 nuclei

5. Why do atoms bond? As individual atoms, many atoms are considered unstable (if they do not have 8 valence electrons). To achieve 8 valence electrons, atoms bond with each other in a variety of ways. Bonded atoms acquire 8 valence electrons and therefore become more stable. As a result of bond formation, both atoms achieve a valence shell that is identical to one of the noble gases. 8 valence electrons is associated with maximum stability. It is known as a stable octet. Remember: Most atoms also want to have 8 valence electrons!

6. Energy changes Transferring electrons involves energy change. Energy based on position is called potential energy. Chemical Energy = the particular form of potential energy involved in the making and breaking of chemical bonds Remember: Forming a chemical bond causes a chemical change, because a new substance is produced with new properties.

7. How do bonds change energy? As individual atoms, many atoms have a very high potential energy and are considered unstable (if they do not have 8 valence electrons). Bonded atoms acquire 8 valence electrons and therefore become more stable, their energy decreases. 8 valence electrons is associated with maximum stability and minimum potential energy content. Potential energy decrease occurs when atoms form chemical bonds with one another.

8. Energy changes Remember: Ionization energy = the energy required to remove an electron from an atom Ex: Na + energy  Na + + e- Remember: Electron affinity = the energy needed to add an electron onto a neutral atom Ex: Cl + e-  Cl - + energy The energy released is less than the energy required to remove an electron, but there is also energy released by forming the salt, so overall, energy is released!

9. Review Chemical reaction: – Reactants → Products Exothermic Reaction = energy release – Reactants  Products + energy Endothermic Reaction = energy absorbed – Reactants + energy  Products

10. Energy Changes and Stability in Bonding an Exothermic Process When chemical bonds are formed, energy is released (an exothermic process). The products are more stable than the reactants. Whenever large amounts of energy are released in the formation of a bond, the bond is said to be strong and very stable. Weak bonds and unstable systems are associated with the release of small amounts of energy. The potential energy of the products is lower than the potential energy of the reactants. BARF!

11. Energy Changes and Stability in Bonding in an Endothermic Process When chemical bonds are broken, energy is absorbed (an endothermic process). The reactants are more stable than the products. The potential energy of the products is higher than the potential energy of the reactants. BARF!

12. Practice Ex: Which statement best describes the production of a chlorine molecule according to the reaction Cl + Cl  Cl 2 + 58 Kcal? a. A bond is broken, and the reaction is exothermic. b. A bond is broken, and the reaction is endothermic. c. A bond is formed, and the reaction is exothermic. d. A bond is formed, and the reaction is endothermic.

13. Practice Explain in terms of electron configuration, why is an oxygen molecule more stable than oxygen atoms. The oxygen molecule has achieved a noble gas electron configuration or has a stable octet of valence electrons

14. Which of the bonds is most stable? The greater the energy released during bond formation the more stable the bond. BondEnergy released in Formation (kcal/mol) H-F135 H-Cl103 H-Br87 H-I71

15. Do Now Is energy released or absorbed when a bond is formed? Is energy released or absorbed when a bond is broken? Which is the endothermic potential energy diagram? Which is the exothermic potential energy diagram? A B

16. INTRAMOLECULAR BONDS Bonds within the compound, hold atoms together. TYPES OF INTRAMOLECULAR BONDS – 1. IONIC BONDS – 2. COVALENT BONDS When atoms bond, valence electrons are redistributed. The way they are redistributed determines the type of bond.

17. IONIC BONDING

18. COVALENT BONDING

19. Reactivity Review Remember: most atoms are satisfied when there outer energy level is complete with 8 valence electrons = octet rule. Noble gases are the least reactive elements, because they have completely filled outer energy levels. Alkali metals (group 1) and halogens (group 17) are the most reactive elements, because they will react to lose or gain electrons so the outer orbital is full. – Alkali metals (group 1) have 1 electron in their outer energy level, so they can easily give up that 1 electron. – Halogens (group 17) have 7 electrons in their outer energy level, so they can easily gain 1 electron.

20. CATION (ca+ion) ANION (aNion)

21. Some stable ions

22. Some stable ions without noble gas configuration

24. Atoms vs. Ions Ions and their parent atoms have very different properties. – Ex: Sodium and chlorine atoms are very reactive, producing a violent reaction when brought together. However, after the reaction is complete, a stable, less reactive solid (sodium chloride) remains (made of sodium ions and chlorine ions).

26. Metals vs Nonmetals Most metals form cations, because it requires less energy to lose a few electrons compared to gaining a lot of electrons. Most nonmetals form anions, because it requires less energy to gain a few electrons compared to losing a lot of electrons. Remember: Ionization energy = the energy required to remove an electron from an atom – PERIOD TREND: Tends to increase across a period – GROUP TREND: Tends to decrease down a group

27. Which element has the greatest tendency to gain electrons? FBrClI In the compound XCl, X represents which element? RbSIBr

28. What did you learn today?

29. When a bond is broken, energy is absorbed. When a bond is formed, energy is released. Atoms attain a stable valence electron configuration by bonding with other atoms. Noble gases have stable valence configurations and tend not to bond.

30. Do Now What is the difference between an atom and an ion? How does an atom become an ion?

31. 1. IONIC BONDS Usually exist between a metal (cation) and a nonmetal (anion). Chemical bond that results from the electrical attraction between cations and anions (opposite charges attract). One atom loses electrons and becomes a positively charged ion and its radius decreases. The atom that receiving additional electrons becomes negatively charged ion and its radius increases. The force of attraction that holds the oppositely charged ions together is called an ionic bond = strong bonds All ionic compounds are called salts.

32. IONIC BONDING

33. Properties of Ionic Solids The ions produced as a result of electron transfer usually have the electron configuration identical to those of noble gases. Have high melting and boiling points due to strong bonds In the solid phase ( s ) Held together in very rigid, fixed crystalline structures. Poor conductors of electricity Melted ( l ) or aqueous ( aq ) state Ions are free to move and are capable of conducting electricity (mobile ions) Ex: Light bulb demo

34. Ionic Salts Most ionic compounds are salts. The salt is electrically neutral composed of cations and anions held together by ionic bodes in a simple, whole number ratio. – NaCl = 1:1 ratio However, there is more than 1 sodium ion and 1 chlorine ion. There are billions of cations and anions attracted and being pulled together into a tightly packed structure, producing a distinctive crystal structure.

35. IONIC SOLIDS ARE CRYSTALLINE SOILDS

36. Using ELECTRONEGATIVITY to determine whether a bond is IONIC For an IONIC bond to form the atoms involved in the bond must have an ELECTRONEGATIVITY difference between 1.7 to 4.0. Ex. NaCl Na Electronegativity _______ Cl Electronegativity _______ Difference of:__________

37. Neutral (Formula) Units An ionic compound is composed of positive and negative ions that are combined so that the number of positive and negative charges are equal. The chemical formula of an ionic compound is the simplest collection of atoms (ratio of ions). IT IS CALLED A FORMULA UNIT! To achieve neutral units: 1. Na + must bond with ______ to form ______. 2. Ca +2 must bond with ______ to form _____. 3. Al +3 must bond with ______ to form ______.

38. Neutral Ionic Compounds Ionic compounds never have an excess of positive or negative charge. The positive and negative charges need to balance out, so the compound has no overall charge!

39. Naming Ionic Compounds Naming salts is very easy, because they are binary ionic compounds (made up of two elements). – The cation is named by borrowing the name of the element. – The anion named by combining the name of the element with an –ide ending. The name of compound is made up of both the cation and anion name – Ex: NaCl = sodium chloride – Ex: ZnS = zinc sulfide – Ex: K 2 O = potassium oxide – Ex: Mg 3 N 2 = magnesium nitride – Ex: Al 2 S 3 = aluminum sulfide

40. Another Naming System For metallic elements that can form two positive ions, the suffixes –ous and –ic may be attached to the Latin name of the element: – Ex: FeCl 2 = ferrous chloride – Ex: FeCl 3 = ferric chloride – Ex: Cu 2 O = cuprous oxide – Ex: CuO = cupric oxide – Ex: Hg 2 Br 2 = mercurous bromide – Ex: HgBr 2 = mercuric bromide We will focus on naming in the next unit…

41. Do Now What do Lewis dot diagrams represent? Draw a Lewis dot diagram for Na and Cl.

42. Electron Dot Diagram of Ions Na + Ca +2 Al +3 Cl -

43. ELECTRON DOT NOTATION WITH IONIC BONDS 1. NaCl

44. 2. CaCl 2

45. 3. AlCl 3

46. 4. CaO

47. 5. Na 2 O

48. What did you learn today?

49. When an atom gains one or more electrons, it becomes a negative ion and its radius increases. When an atom loses one or more electrons, it becomes a positive ion and its radius decreases. Electron-dot diagrams (Lewis structures) can represent the valence electron arrangement in elements, compounds, and ions. Electronegativity indicates how strongly an atom of an element attracts electrons in a chemical bond. Electronegativity values are assigned according to arbitrary scales. Metals tend to react with nonmetals to form ionic compounds.

50. Do Now What are the two main different bonds? How are they different?

51. COVALENT BONDING

52. 2. COVALENT BONDS Valence electrons are shared between atoms to make all atoms satisfied! The shared electrons move in the space surrounding the nuclei = molecular orbital Results when a nonmetal bonds with a nonmetal. Results when a metalloid and a nonmetal bond. » Remember: Hydrogen = nonmetal

53. COVALENT BONDING

54. Elements with COVALENT BONDS are MOLECULAR COMPOUNDS: Molecule = a neutral compound held together by covalent bonds Molecules can exist as solids, liquids or gases depending on the degrees of attraction that exists among molecules. Molecules may consist of identical atoms bonded together (O 2 ) or different atoms bonded together (H 2 O)

55. Molecular Compounds Have covalent Bonds Electrons are shared between atoms Nonmetal + Nonmetal or Nonmetal + Metalloids

56. Ionic Compounds Have Ionic Bonds Electrons are transferred between atoms Metal + Nonmetal

57. Properties of Molecular Solids (have covalent bonds) Relatively soft and low melting and boiling points due to weak intermolecular forces, easy to overcome. Do not conduct electricity due to no free moving ions (no mobile ions) Good insulators (ex. lipids) Nonpolar molecules contain weak bonds due to van der waals forces (weak intermolecular forces)

58. COMPARISION of IONIC and MOLECULAR COMPOUNDS IONIC COMPOUNDS 1. Forces between ions are very strong. 2. High melting point 3. High boiling point 4. Hard (solids) 5. Soluble in water 6. Good conductors of electricity only when dissolved in water ( aq ) or melted ( l ) (not in solid form). MOLECULAR COMPOUNDS 1. Forces between atoms are not as strong as ionic bonds. 2. Low melting point 3. Low boiling point (many are gases at STP) 4. Soft 5. Poor conductors of electricity.

59. Covalent Bond Energy When atoms bond, they become more stable. The bond causes a decrease in potential energy, and this energy is released to the surroundings (exothermic).

60. Covalent Bond Strength Covalently bonded atoms are not rigidly bonded at a fixed distance! Covalent bonds are more flexible and the two nuclei vibrate back and forth. Bond length = average distance between the two nuclei Bond energy = energy required to break the bond (dissociation) Usually higher bond energy results in shorter bond length and stronger bonds. ElementsBond energy (KJ/mol) Bond length (pm) H-F57092 C-F552138 O-O498121 H-H43675 H-Cl432127 C-Cl397177 H-Br366141

61. Using ELECTRONEGATIVITY to determine if a bond is COVALENT For a COVALENT bond to form the atoms involved in the bond must have an ELECTRONEGATIVITY difference less than 1.7. Ex. CO 2 C Electronegativity _______ O Electronegativity _______ Difference of:__________

62. MOLECULAR FORMULA: Indicates the number and kinds of atoms in a molecule – Ex: H 2 O – Ex: CO 2 – Ex: CH 4

63. Naming Covalent Compounds Covalent compounds are named in a similar way to ionic compounds – The first element in the formula is usually written first in the name – The second element has an –ide ending Ex: SO 2 = sulfur oxide – However, this is not completely correct…

64. Naming Covalent Compounds Since multiple covalent compounds can be made from the same elements, the name must distinguish them as different. – Prefixes are used to indicate the number of atoms of each element in the molecule. Ex: SO 2 = sulfur dioxide Ex: SO 3 = sulfur trioxide

65. Naming Covalent Compounds Some prefixes include: – Mono- – Di- – Tri- – Tetra- – Penta- We will revisit naming of compounds in our next unit…

66. Do Now What is the Lewis dot diagram for Hydrogen? What is the Lewis dot diagram for Carbon? What type of bond would these elements form?

67. OCTET RULE Compounds form so that each atom, by gaining, losing or sharing electrons, has 8 valence electrons. Electron Dot Diagrams with F 2 Molecules: The pair of dots between the symbols (in the middle) represents the covalent bond. You can also use lines to represent the bonds.

68. Covalent Bonds Each atom contributes one electron to a shared electron pair. The electrons that are not part of the bond are called unshared pairs or lone pairs.

69. Drawing Lewis Structures Step 1. Gather information – Determine the total number of valence electrons in the compound. (Sum total the valence electrons of all atoms in the compound.) – Draw Lewis structures for each atom in the compound.

70. Drawing Lewis Structures Step 2. Arrange the atoms – Arrange the Lewis structures to show how the atoms bond in the molecule. (Chemical formulas are often written in the order in which the atoms are connected. If a central atom has a group of other atoms bonded to it, the central atom is usually written first as in CH 4 ) Halogens and Hydrogen often bind to one other atom and usually at the end of the molecule. Carbon is usually in the center of the molecule.

71. Drawing Lewis Structures Step 3. Distribute the dots – Distribute the electrons dots so that each atom (except hydrogen) satisfies the octet rule. – Place any leftover electrons on the central atom.

72. Drawing Lewis Structures Step 4. Draw the bonds. – Change each pair of dots that represents a shared pair of electrons to a long dash.

73. Drawing Lewis Structures Step 5. Verify the structure – Count the number of electrons surrounding each atom (except hydrogen). – If there are not enough electrons to give the central atom an octet, try multiple bonds. (For example, use one or more of the unshared pairs of electrons on the atoms bonded to the central atom to form double or triple bonds). YOU CANNOT USE MORE OR LESS THAN THE TOTAL NUMBER OF VALENCE ELCTRONS OF THE ATOMS IN THE MOLECULE!!!

74. HCl H 2 O

75. NH 3 CH 4

76. O 2 N 2 CO 2

77. Covalent Bonds Single bond = one shared pair of electrons (two electrons) Double bonds = two shared pairs of electrons (four electrons) Triple bonds = three shared pairs of electrons (six electrons)

81. Sigma and Pi Bonds Most single bonds involve sigma bonds because the bond lies along a line that joins the nuclei of the two bonding atoms

82. Sigma and Pi Bonds Most double and triple bonds involve pi bonds due to the dumbbell shape of the p orbitals overlapping Pi bonds are more reactive than sigma bonds

83. Sigma and Pi Bonds A double bond consists of one sigma bond and one pi bond. A triple bond consists of one sigma bond and two pi bonds.

84. Kind of Bond Covalent bond DescriptionDrawing Single bond Double bond Triple bond

85. Resonance Structures When a molecule has two or more possible Lewis structures, the different structures are called contributing resonance structures.

86. Resonance Structures The actual structure of the molecule is a composite of the contributing resonance structures called a resonance hybrid.

87. Electron Dot Diagrams Molecule with Covalent BondCompound with Ionic Bond

88. Polyatomic Ions Monoatomic = “one-atom” Diatomic = “two atoms” Polyatomic = “many atoms” A polyatomic ion is a charged group of two or more bonded atoms that can be considered a single ion Polyatomic ions have an overall positive or negative charge (just like any other ion) – Ex: ammonium = NH 4 +

89. POLYATOMIC IONS (see Table E) Contain BOTH IONIC and COVALENT BONDS. They have more or less electrons than normal. Mg 3 (PO 4 ) 2 Mg 2+ (PO 4 ) 3- Ionic bondCovalent bonds

90. Review Ionic Bonds Remember: Ionic bonds form between cations and anions. Do the following compounds contain ionic bonds? NH 4 Cl NaNO 3 CH 3 OH yes no

91. Exceptions to the Octet Certain atoms may have more or less than 8 electrons in their valence shell when they bond. Ex: BeCl 2 Ex:BF 3 Ex: PCl 5 Ex: SF 6

92. What did you learn today?

93. In a multiple covalent bond, more than one pair of electrons are shared between two atoms. Electron-dot diagrams (Lewis structures) can represent the valence electron arrangement in elements, compounds, and ions. Nonmetals tend to react with other nonmetals to form molecular (covalent) compounds.

94. Quiz Quiz on Ionic and Covalent Bonding

95. Lab Bonding lab

96. Do Now Think of the north and south poles of our planet. What do these poles mean? What do you think “polar bonding” means?

97. Ionic vs. Covalent In general: – Bonds between metals and nonmetals are ionic – Bonds between nonmetals and nonmetals are covalent – Bonds between metalloids and nonmetals are covalent

98. Using Electronegativity Bonding is RARELY purely ionic or covalent. It usually forms somewhere between the 2 extremes. Where it falls depends on how strongly one atoms attracts the electrons of the other atom. Atoms covalently bonded can share electrons equally or unequally. To determine the type of bond between 2 atoms, calculate the electronegativity differences!

99. Polar vs. Nonpolar Nonpolar covalent bond = covalent bond in which the shared electron pair in the molecular orbital is shared equally. Polar covalent bond = covalent bond in which the shared electron pair is shared, but not equally. – Result when one atom has a higher electronegativity, thus, attracts the electrons more, so the electrons are more likely to be found closer to the atom with higher electronegativity.

100. Polar Bond Nonpolar Bond

101. Fig. 9.15a, p.413 δ means “slightly”

102. Covalent Bonds can be Polar or Nonpolar NONPOLAR BOND- POLAR BOND-

103. Covalent vs. Ionic If the electronegativity difference is great enough, the atom with the higher value may remove an electron from another atom. This forms ions that will bond in ionic bonds.

104. Bond Strength The greater the electronegativity difference, the greater the polarity of the bond and the stronger the bond.

105. Electronegativity Difference determines type of bond 0 0.3 1.74.0 NonpolarPolarIonicCovalent

106. Calculating Electronegatvity Differences (IONIC) Covalent

107. Electronegativity difference Type of Bond 0-0.3NONPOLAR COVALENT 0.3-1.6POLAR COVALENT 1.7-4.0IONIC BOND

108. Nonpolar Polar Ionic Bond Polarity

109. Review Questions 1. What type of bond exists between that atoms of LiF? 2. What type of bond exists between the atoms of H 2 O? 3. What type of bond exists between the atoms of Cl 2 ?

110. What type of bond exists between the following elements: Sulfur and:Electronegativity Difference Bond Type Hydrogen Cesium Oxygen Which is most polar?

111. What type of bond exists between the following elements: Chlorine and: Electronegativity Difference Bond Type Cesium Nitrogen Bromine Which is least polar?

112. Do Now Predict which elemental group will bond with Cl to from XCl 2 compounds.

113. Periodic Table and Bonding We can use our understanding of the periodic table to predict what bonds will form between elements and in what proportions. Remember: – Ionic bonds form between cations (metals) and anions (nonmetals) – Covalent bonds form between nonmetals (and metalloids and nonmetals)

114. Group 1 Group 1 metals will form +1 cations.

115. Group 2 Group 2 metals will form +2 cations.

116. Group 16 Group 16 nonmetals will form -2 anions.

117. Group 17 Group 17 nonmetals will form -1 anions.

118. Which element will bond with Cl to from XCl 2 compounds?

119. Element X is in Group 1, and element Y is in group 16. Which represents a compound formed from these two elements? X 2 Y XY 2 X 2 Y 7 X 7 Y 2

120. Which element reacts with fluorine to produce XF 2 ?

121. Which element bonds with oxygen to form X 2 O 3. Al Ba Cd Na

122. Which element group is represented by X in the formula X 2 O 5 ?

123. Which group of elements represent the metal X in the formula X 2 (SO 4 ) 3 ?

124. Element M has an electronegativity of less than 1.2 and reacts with bromine to form MBr 2. What could Element M be? Al Na Ca K

125. What element covalently bonds with hydrogen to form XH 4 ?

126. Do Now Draw the electron dot diagram for H 2 O. Draw the electron dot diagram for CO 2.

127. The Shapes of Molecules These notes only refer to MOLECULES (COVALENT COMPOUNDS), not to ionic compounds. All ionic compounds form CRYSTALS!!!

128. The Shapes of Molecules The physical and chemical properties of molecules depend not only on the bonding of atoms, but also on the MOLECULAR GEOMETRY or 3D arrangement of the atoms. Remember: a molecular formula indicates the number and kind of atoms in a particular molecule.

129. Formula vs. Shape Sucrose and Sucralose have similar shapes, so both stimulate nerve receptors of the tongue and your brain interprets them as sweet. However, sucrose and sucralose have different chemical formulas, so they are processed by the body differently. (sugar substitute)

130. SHAPES OF MOLECULES A structural formula indicates two- dimensional arrangement of the bonds and lone pairs of electrons in a molecule. In order to determine the shape of a molecule we must think on the three dimensional level. – Lewis dot diagrams can help predict molecular shapes. – Shared electron pairs can determine molecular shapes.

131. VSEPR THEORY Valence Shell Electron Pair Repulsion Theory The valence electrons in a molecule are arranged as far apart from each other as possible. Unshared electron pairs (lone pairs) affect the orientation of the shared pairs, therefore you MUST be able to correctly sketch Lewis Dot diagrams of compounds in order to determine molecular shape.

132. Predicting Molecular Shapes Step 1: Gather information – Draw the Lewis dot diagram of the molecule. Step 2: Count the shared and unshared pairs. Step 3: Apply VSEPR theory and verify it. – Find the shape that allows the shared and unshared pairs of electrons to be as far apart as possible. Linear Trigonal Planar Tetrahedral Trigonal Pyramidal Bent

133. What are the possible shapes of molecules? LINEAR Molecules with only 2 atoms are linear. Sometimes molecules with three atoms are also LINEAR. In these, cases, the central atom is double bonded to the atoms on either side and there are NO LONE PAIRS of electrons on the central atom. Ex: H 2 & CO 2 Bond angle = 180˚

134. What are the possible shapes of molecules? Trigonal Planar In general, a trigonal planar molecule has a central atom bonded to three other atoms, and the central atom has no lone pairs of electrons. Ex: SO 3 Bond angle = 120˚

135. What are the possible shapes of molecules? TETRAHEDRAL Tetra- = four In general, a tetrahedral molecule has a central atom (usually a carbon) bonded to four other atoms. Ex: CH 4 & CH 2 Cl 2 Bond angle = 109.5˚

136. What are the possible shapes of molecules? TRIGONAL PYRAMIDAL Three triangle sides with a triangle base “pyramidal”. In general, a trigonal pyramidal molecule has a central atom bonded to three other atoms and a LONE PAIR of electrons on the central atom. (The lone pair of electrons REPELS other electrons so it creates the tetrahedral shape without the top). Ex: NH 3 Bond angle = 107˚

137. What are the possible shapes of molecules? BENT In general, a bent molecule has a central atom bonded to two other atoms and TWO LONE PAIRS of electrons on the central atom. Ex: H 2 O & O 3 Bond angle = 105˚

138. What are the possible shapes of molecules? TRIGONAL BIPYRAMIDAL 5 pairs of electrons are shared Ex: PCl 5

139. What are the possible shapes of molecules? OCTAHEDRAL 6 pairs of electrons are shared Ex: SF 6

142. Classify the following shapes of molecules:

146. Regions of High Electron Density

147. Practice

148. Do Now What is the difference between a polar and nonpolar bond? What do you think molecular polarity is? How do you think molecular shape affects molecular polarity? Which types of molecule do you think has polar ends?

149. Molecular Properties Molecular shape affects both physical and chemical properties. One property that shape determines is the polarity of a molecule.

150. Molecules can be Polar or Nonpolar Molecules Whether a molecule is polar or nonpolar depends on the SYMMETRY of the molecule. – Determine the polarity of each bond. – Then determine the arrangement of the atoms in space (molecular shape)

151. Polar vs. Nonpolar Molecules Nonpolar molecules Have symmetry Can have polar or nonpolar bonds Can have only one kind of element or more than one kind of element in the compound Do not have poles Polar Molecules Do not have symmetry Have only polar bonds Have more than one kind of element in the compound Have poles They are dipoles

152. NONPOLAR MOLECULES Have symmetry Can have polar or nonpolar bonds Can have only one kind of element (ex. diatomic molecules) or more than one kind of element in the compound Do not have poles (no positive or negative charges in molecule) If a molecule contains ONLY nonpolar bonds, then it is a nonpolar molecule. Any molecule composed of only one kind of atom is NONPOLAR.

153. NONPOLAR MOLECULES H 2 O 2 Oxygen Molecule

154. POLAR MOLECULES Do not have symmetry Have polar bonds Have more than one kind of element in the compound Have poles They are dipoles A molecule that has one end with a slightly positive charge and one end with a slightly negative charge.

155. What is a Dipole? DIPOLE- a molecule that has both a positive and a negative pole There is an uneven distribution of electrical charges (electrons). Polar molecules are known as DIPOLES.

156. Fig. 9.15a, p.413 δ means “slightly”

157. Practice: HCl The chlorine is more electronegative; therefore the Cl is holding the electrons tighter than the hydrogen is. This creates a slightly negative charge on the Cl atom and a slight positive charge on the hydrogen atom.

158. Practice: HF

159. Practice: NH 3

160. Practice: A molecule that contains polar bonds is not necessarily a polar molecule. Examples: CO 2

161. Practice CCl 4

162. Practice: H 2 O

163. Practice: CH 4

164. Practice: Br 2

165. Practice CH 2 Cl 2

166. Review S = symmetric N = nonpolar A = asymmetric P = polar

167. Structure of Molecule ShapeBond (Polar or Nonpolar) Molecule (Polar or nonpolar) H2H2 O2O2 N2N2 HCl

168. Structure of Molecule ShapeBond (Polar or Nonpolar) Molecule (Polar or nonpolar) HF NH 3 CO 2 CCl 4

169. Structure of Molecule ShapeBond (Polar or Nonpolar) Molecule (Polar or nonpolar) H2OH2O CH 4 Br 2 CO

170. Review Which molecular shapes tend to be nonpolar molecules? Which molecular shapes tend to be polar molecules?

171. Covalent Bonds Polar Bonds Electrons shared unequally Uneven distribution of charge Electronegativity difference 0.3-1.6 UNEQUAL PULL Nonpolar Bonds Electrons shared equally Balanced distribution of charge Electronegativity difference 0-0.3 EQUAL PULL Molecules Polar Molecules DO NOT HAVE SYMMETRY Only have polar bonds Nonpolar Molecules HAVE SYMMETRY Can have polar bonds Or nonpolar bonds

172. Polyatomic Ions Monoatomic = “one-atom” Diatomic = “two atoms” Polyatomic = “many atoms” A polyatomic ion is a charged group of two or more bonded atoms that can be considered a single ion Polyatomic ions have an overall positive or negative charge (just like any other ion) – Ex: ammonium = NH 4 +

173. POLYATOMIC IONS (see Table E) Contain BOTH IONIC and COVALENT BONDS. They have more or less electrons than normal. Mg 3 (PO 4 ) 2 Mg 2+ (PO 4 ) 3- Ionic bondCovalent bonds

174. NH 4 + CO 3 2-

175. Which of the following contains both Ionic and Covalent Bonds? Al 2 S 3 HCl Mg(OH) 2 H 2 O yes no

176. Review Ionic Bonds Remember: Ionic bonds form between cations and anions. Do the following compounds contain ionic bonds? NH 4 Cl NaNO 3 CH 3 OH yes no

177. What have you learned this week?

178. Molecular polarity can be determined by the shape of the molecule and the distribution of charge. Symmetrical (nonpolar) molecules include CO 2, CH 4, and diatomic elements. Asymmetrical (polar) molecules include HCl, NH 3, and H 2 O. The electronegativity difference between two bonded atoms is used to assess the degree of polarity in the bond. Metals tend to react with nonmetals to form ionic compounds. Nonmetals tend to react with other nonmetals to form molecular (covalent) compounds. Ionic compounds containing polyatomic ions have both ionic and covalent bonding.

179. Lab Molecular Models

180. Quiz Quiz on Molecular shapes and polarity

181. Do Now What does metallic mean? What are some properties of a metallic? What is the difference between ionic and covalent bonds?

182. Additional Types of Bonding 1. Metallic Bonding 2. Covalent Network Bonds 3. Coordinate Covalent Bond

183. 1. METALLIC BONDING

184. 1. Metallic Bonding Chemical bonding in metals is different than it is in ionic or molecular (covalent) compounds. Properties of Metals: – Conduct electricity and heat – Malleable – Ductile

185. Ex: Li or Na or K Ex: Be or Mg or Ca

186. Why are metals such good conductors of electricity? Metals have fairly empty valence shells. – Ex: Mg Electrons can bounce around a lot because of this! High mobility of electrons gives metals their properties!

187. In metals, vacant orbitals overlap and the electrons that occupy these levels can roam from the orbitals surrounding one nucleus to the orbitals surrounding other nuclei. Electrons are DELOCALIZED, which means they do not belong to any one atom but they move about freely. BUZZ PHRASE: “Sea of mobile electrons” surrounding positive nuclei.

188. Properties of Metallic Bonds Excellent Conductors of Electricity Metals constitute more than ¾ of the elements on the Periodic Table. Valence electrons moving around the positive metal ion “electron sea” Sea of mobile electrons e- are moving around/ uniform distribution/ free migration Force of attraction between electron and positive charged ion

190. 2. Covalent Network Bonds Certain atoms are able to form covalent bonds with other atoms in a 3D arrangement called a network solid. Ex: carbon, SiC, SiO 2

191. Covalent Network Compounds Diamond, a network of covalently bonded carbon atoms

192. Covalent Network Compounds Graphite, a network of covalently bonded carbon atoms

193. Allotropes of Carbon Allotrope = 2 or more different forms (structural/molecular arrangement) that an element can exist in and results in different properties

194. Graphite to diamond animation This animation shows how graphite becomes diamond under extreme heat and pressure

195. Other Allotropes Oxygen gas vs. Ozone gas (O 2 vs. O 3 ) – They have different molecular structures, resulting in different properties Phosphorus (P 4 vs P 8 ) – They have different molecular structures, resulting in different colors: white, black or red

196. Covalent Network Compounds While there are strong covalent bonds between carbon atoms in each layer, there are only weak forces between layers. This allows layers of carbon to slide over each other in graphite (and layers are transferred to paper as you write). Rotate the structure of diamond - Notice that each diamond atom is the same distance to each of its neighboring carbon atoms. There is a rigid network of bonds within the diamond crystal. http://www.edinformatics.com/interactive_molecule s/graphite.htm http://www.edinformatics.com/interactive_molecule s/graphite.htm

197. Properties of Covalent Network Solids In some solids, atoms are covalently bonded but DO NOT form molecules. Instead they form networks extending throughout the entire crystal (share with neighboring atoms) Very hard High melting points Complex – very strong Poor conductors of electricity. Share electrons, but compressed so much that there is a network. Examples: Diamond, Graphite

198. 3.COORDINATE COVALENT BOND Special type of bond formed by an attraction of an atom, usually hydrogen, to a lone pair of electrons Remember: lone pair of electrons = unshared pairs of electrons Ex: NH 4 +

199. Types of bonds we have discussed: 1. Ionic 2. Covalent 3. Metallic 4. Network Covalent 5. Coordinate Covalent Bonds These are all INTRAMOLECULAR FORCES: Forces that exist within/inside the molecule Exist within the molecule holding the atoms together.

200. Do Now What do you think the prefixes intra- and inter- mean?

201. Intra vs. Inter INTRA- WITHIN INTER- BETWEEN

202. Intramolecular vs. Intermolecular INTRAMOLECULAR FORCES- within molecule INTERMOLECULAR FORCES- between molecules Intermolecular forces are weaker than bonds and so they require less energy to break.

203. INTERMOLECULAR FORCES Intermolecular forces = forces of attraction between neighboring molecules

204. Intermolecular Forces Physical properties of substances (conductivity, malleability, solubility, hardness, melting point, boiling point) can be explained in terms of chemical bonds and intermolecular forces. Whether a substance is a solid, liquid or gas does not depend on the strength of the intramolecular forces but rather on how strongly the molecules are held to each other.

205. Intermolecular Forces What effect do Intermolecular forces have on molecules? The strength of these forces is responsible for the boiling and melting points of substances. – Strong IMF = high boiling and melting points – Weak IMF = low boiling and melting points The presence of these forces is responsible for whether one substance dissolves in another.

206. Generally, “like dissolved like”. So, a polar substance will dissolve in a polar substance as shown below: Starch (polar) in H 2 O (polar) vs. Styrofoam (nonpolar) in H 2 O (polar)

207. Nonpolar substances will dissolve in nonpolar substances. Acetone (nonpolar) and Styrofoam (nonpolar) vs. Acetone (nonpolar) and Starch (polar)

208. Dry ice weak IMF http://www.stevespanglerscience.com/experi ment/awesome-dry-ice-experiments http://www.stevespanglerscience.com/experi ment/awesome-dry-ice-experiments

209. INTRAMOLECULAR FORCE- within molecule INTERMOLECULAR- between molecules 1. Ionic 2. Covalent 3. Metallic 4. Network Covalent 5. Coordinate Covalent Bonds 1. Dipole-Dipole Attraction 2. Hydrogen Bonding 3. Weak Intermolecular Forces/Van Der Waal’s Forces/London Dispersion Forces 4. Molecule-Ion Attraction (Ion-Dipole)

210. 1. DIPOLE-DIPOLE ATTRACTIONS INTERMOLECULAR FORCE Exist between polar molecules. Exist between dipoles. The positive end (pole) of one molecule attracts the negative end (pole) of a neighboring molecule. This creates a dipole- dipole force. Use dotted lines to represent intermolecular forces. http://intro.chem.okstate.edu/1515SP01/Lect ure/Chapter12/HCldipole.html http://intro.chem.okstate.edu/1515SP01/Lect ure/Chapter12/HCldipole.html

211. 2. HYDROGEN BONDING INTERMOLECULAR FORCE Hydrogen Bonds form between molecules that contain covalent bonds between hydrogen and another atom that has a small radius and high electronegativity. The positive pole of one dipole is attracted to the negative pole of another dipole. http://www.kentchemistry.com/links/bonding /bondingflashes/bond_types.swf http://www.kentchemistry.com/links/bonding /bondingflashes/bond_types.swf

212. Hydrogen Bonding

213. Examples The positive pole of one dipole is attracted to the negative pole of another dipole! H 2 O and NH 3 H 2 O and HF

215. Hydrogen Bonds ice and liquid water In liquid water each molecule is hydrogen bonded to approximately 3.4 other water molecules. In ice each each molecule is hydrogen bonded to 4 other molecules. Compare the two structures below. Notice the empty spaces within the ice structure http://www.edinformatics.com/interactive_molecules/ic e.htm http://www.edinformatics.com/interactive_molecules/ic e.htm Ice Liquid water

216. Water: the Universal Solvent One side of water is negatively charged because the oxygen atom keeps the shared electrons longer than the hydrogen atoms. As a result the oxygen side is negatively charged and the hydrogen side of water is positively charged. O

217. Water: the Universal Solvent Like a magnet that pulls on things that are magnetic, water pulls on things that are electrically charged. Magnets have north & south poles, water has positive and negative poles and thus called a polar solvent. Since unlike charges attract, the negative end of water will be attracted to the positive sodium ion. The positive end of water will be attracted to the negative chloride ion. Since water is always in motion, it will pull on the ionic compound and move the ions away from each other. This dissolves the ionic compound.

218. Water: the Universal Solvent O O Na + Cl - O O O

219. Wax does not repel water We’ve heard that wax or oils repel water. But that isn’t true. Water is so attracted to other water molecules that anything between them is squeezed out of the way. O O Oil droplet O O O

220. Hydrogen Bonding Hydrogen Bonding intermolecular forces are SO strong that they account for the unusually high boiling points of substances.

221. Properties of Water Hydrogen bonds also account for the unique properties of water: – Formation of a meniscus – Surface tension – Adhesion – Cohesion If hydrogen bonds did not exist, water would be a gas at STP and life as we know it would not exist!!!

223. Water is always trying to pull itself into a tight ball as long as there is nothing nearby that has a charge on it. Therefore, this surface is not repelling water; it’s simply not attracting it and keeping water from doing what it does naturally.

224. We see the same effect on waxy leaves. Water pulls on itself so much that it forms a “skin.” It’s called surface tension. http://www.youtube.com/watch?v=45yabrnryXk

225. We are lucky that water has this strong attraction force otherwise we’d never see raindrops. The water would just breakup into a mist as it fell. Very few liquids would remain as drops if they fell from a large height.

226. Do Now Why does water have surface tension?

227. 3. VAN DER WAAL’S FORCES (aka: London Dispersion Forces or Weak Intermolecular Forces) Even though the average sharing of electrons in nonpolar molecules is equal, because the electrons are in constant motion, at any moment, the distribution of electrons may be slightly uneven. http://intro.chem.okstate.edu/1515SP01/Lect ure/Chapter12/LondonDisp.html http://intro.chem.okstate.edu/1515SP01/Lect ure/Chapter12/LondonDisp.html

229. VAN DER WAAL’S FORCES MOMENTARY DIPOLE = A momentary dipole in one molecule can INDUCE a momentary dipole in another nonpolar molecule, and this creates a weak force of attraction called Van Der Waal’s Forces (also known as: London Dispersion Forces or Weak Intermolecular Forces) Induced dipole: http://waterontheweb.org/curricula/bs/student/oxy gen/shock2.html http://waterontheweb.org/curricula/bs/student/oxy gen/shock2.html http://antoine.frostburg.edu/chem/senese/101/liqui ds/faq/h-bonding-vs-london-forces.shtml http://antoine.frostburg.edu/chem/senese/101/liqui ds/faq/h-bonding-vs-london-forces.shtml

230. VAN DER WAAL’S FORCES The strength of these forces INCREASES with: 1. INCREASED MOLECULAR SIZE H-H…H-H = weighs less, weaker force of attraction O-O…O-O = weighs more, stronger force of attraction 2. and with DECREASED DISTANCE BETWEEN THE MOLECULES. H-H…H-H = closer, stronger force of attraction H-H…………H-H = farther, weaker force of attraction

231. Use by animals GeckoGecko climbing glass using its natural setæsetæ The van der Waals forces is the force to which the gecko's climbing ability is attributed. A gecko can hang on a glass surface using only one toe. Efforts continue to create a synthetic "gecko tape" that exploits this knowledge. So far, research has produced some promising results - early research yielded an adhesive tape product, which only obtains a fraction of the forces measured from the natural material, and new research has yielded a discovery that purports 200 times the adhesive forces of the natural material.geckoadhesive tape Researchers at Stanford University recently developed a gecko-like robot which uses synthetic setae to climb wallsStanford Universityrobot synthetic setae The van der Waals force is the force to which the gecko's unique ability to cling to smooth surfaces is attributed. A gecko can hang on a glass surface using only one toe. In 2003 a kind of synthetic adhesive tape was created using this principle. geckoadhesive tape

232. 4. MOLECULE ION ATTRACTION (Ion-Dipole) Intermolecular Bond An attraction between polar molecules and ions. These attractions exist in solution when an ionic compound dissolves in a polar solvent. The ions are attracted to the different ends of the polar molecule. For example: NaCl in H 2 O Ions are freed from their crystal lattice and are then surrounded by water molecules.

233. Molecule Ion Attraction occurs in AQUEOUS SOLUTIONS!!!!!!

234. What have you learned this week?

235. Chemical bonds are formed when valence electrons are: transferred from one atom to another (ionic) shared between atoms (covalent) mobile within a metal (metallic) Physical properties of substances can be explained in terms of chemical bonds and intermolecular forces. These properties include conductivity, malleability, solubility, hardness, melting point, and boiling point. Intermolecular forces created by the unequal distribution of charge result in varying degrees of attraction between molecules. Hydrogen bonding is an example of a strong intermolecular force.

236. Lab “Water’s Weird” lab

237. Lab Bonding Type worksheet/lab

238. Test Test on Chemical Bonding