2. CHEMICAL BONDS Combining elements

3. COMBINING ELEMENTS How is possible to have so many different materials from a limited number of elements?

4. COMBINING ELEMENTS How do you get so many different cheeses? What is the key ingredient in all cheeses? Milk is combined with a small number of ingredients under different conditions to make a huge number of different cheeses. In a similar way, a small number of elements can be combined in many different ways to make a huge number of different compounds.

5. WHAT IS A COMPOUND? A compound is the substance produced when two or more elements combine in a chemical reaction. A compound is always made up of two or more different types of atom. Two elements, hydrogen (H) and oxygen (O), combine to make the compound, water. Which two elements combine to make the compound carbon dioxide?

6. MAKING A COMPOUND – CARBON DIOXIDE A compound has very different properties to the elements from which it is made. carbon dioxide A colourless gas which is used to put out fires. to make carbon A black solid which can be used as a fuel. combines with compound elements oxygen A colourless gas which is essential for life.

7. MAKING A COMPOUND – WATER What are the elements which make up water? In what ways are the elements different to their compound? to make combines with compound elements water A liquid which is essential to our lives and has many different uses. hydrogen A colourless gas which is used in hot air balloons. A colourless gas which is essential for life. oxygen

9. ELEMENT OR COMPOUND?

10. BONDING BASICS

11. ELECTRON SHELLS a)Atomic number = number of Electrons b)Electrons vary in the amount of energy they possess, and they occur at certain energy levels or electron shells. c)Electron shells determine how an atom behaves when it encounters other atoms

12. ELECTRONS ARE PLACED IN SHELLS ACCORDING TO RULES: The 1st shell can hold up to two electrons, and each shell thereafter can hold up to 8 electrons.

13. 12 The Octet Rule Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has eight electrons in its highest occupied energy level. The same number of electrons as in the nearest noble gas The first exception to this is hydrogen, which follows the duet rule. The second exception is helium which does not form bonds because it is already “full” with its two electrons

14. Octet Rule = atoms tend to gain, lose or share electrons so as to have 8 electrons C would like to N would like to O would like to Gain 4 electrons Gain 3 electrons Gain 2 electrons

15. WHY ARE ELECTRONS IMPORTANT? Elements have different electron configurations  different electron configurations mean different levels of bonding

17. IONIC BONDS Losing and gaining electrons

18. IONIC BOND BOND FORMED BETWEEN TWO IONS BY THE TRANSFER OF ELECTRONS

19. IONIC BOND FORMATION Neutral atoms come near each other. Electron(s) are transferred from the Metal atom to the Non-metal atom. They stick together because of electrostatic forces, like magnets.

21. IONIC BOND Between atoms of metals and nonmetals with very different electronegativity Bond formed by transfer of electrons Produce charged ions. Examples: NaCl, CaCl 2, K 2 O

22. Ionic bond – electron from Na is transferred to Cl, this causes a charge imbalance in each atom. The Na becomes (Na + ) and the Cl becomes (Cl - ), charged particles or ions.

23. FORMATION OF IONS FROM METALS Ionic compounds result when metals react with nonmetals Metals lose electrons to match the number of valence electrons of their nearest noble gas Positive ions form when the number of electrons are less than the number of protons Group 1 metals  ion 1+ Group 2 metals  ion 2+ Group 13 metals  ion 3+

24. IONS FROM NONMETAL IONS In ionic compounds, nonmetals in 15, 16, and 17 gain electrons from metals Nonmetal add electrons to achieve the octet arrangement Nonmetal ionic charge: 3-, 2-, or 1-

25. IONIC BONDING Metals will tend to lose electrons and become POSITIVE CATIONS Normal sodium atom loses one electron to become sodium ion

26. IONIC BONDING Nonmetals will tend to gain electrons and become NEGATIVE ANIONS Normal chlorine atom gains an electron to become a chloride ion

27. Ionic Bonds: One Big Greedy Thief Dog!

28. POLYATOMIC IONS-- a group of atoms that act like one ion NH 4 +1 --ammonium ion CO 3 -2 --carbonate ion PO 4 -3 --phosphate ion IONIC BONDING

29. PROPERTIES OF IONIC COMPOUNDS Crystalline structure. A regular repeating arrangement of ions in the solid. Ions are strongly bonded. Structure is rigid. High melting points- because of strong forces between ions.

30. CRYSTALLINE STRUCTURE + + ++ + + + + + - - - - - - - - - The POSITIVE CATIONS stick to the NEGATIVE ANIONS, like a magnet.

31. DO THEY CONDUCT? Conducting electricity is allowing charges to move. In a solid, the ions are locked in place. Ionic solids are insulators. When melted, the ions can move around. Melted ionic compounds conduct. First get them to 800ºC. Dissolved in water they conduct.

32. IONIC SOLIDS ARE BRITTLE +-+- + - +- +-+- + - +-

33. FORMATION OF SODIUM ION Sodium atom Sodium ion Na  – e   Na + 2-8-1 2-8 ( = Ne) 11 p + 11 p + 11 e - 10 e - 0 1 +

34. FORMATION OF MAGNESIUM ION Magnesium atom Magnesium ion  Mg  – 2e   Mg 2+ 2-8-2 2-8 (=Ne) 12 p + 12 p + 12 e- 10 e - 0 2 +

35. SOME TYPICAL IONS WITH POSITIVE CHARGES (CATIONS) Group 1Group 2Group 13 H + Mg 2+ Al 3+ Li + Ca 2+ Na + Sr 2+ K + Ba 2+

36. LEARNING CHECK A. Number of valence electrons in aluminum 1) 1 e - 2) 2 e - 3) 3 e - B. Change in electrons for octet 1) lose 3e - 2) gain 3 e - 3) gain 5 e - C.Ionic charge of aluminum 1) 3- 2) 5- 3) 3 +

37. SOLUTION A. Number of valence electrons in aluminum 3) 3 e - B. Change in electrons for octet 1) lose 3e - C.Ionic charge of aluminum 3) 3 +

38. LEARNING CHECK Give the ionic charge for each of the following: A. 12 p + and 10 e - 1) 02) 2+3) 2- B. 50p + and 46 e- 1) 2+2) 4+3) 4- C. 15 p + and 18e- 2) 3+ 2) 3-3) 5-

39. COVALENT BONDS Sharing electrons

40. COVALENT BOND BOND FORMED BY THE SHARING OF ELECTRONS

41. COVALENT BOND FORMATION When one nonmetal shares one or more electrons with an atom of another nonmetal so both atoms end up with eight valence electrons

43. COVALENT BOND Between nonmetallic elements of similar electronegativity. Formed by sharing electron pairs Stable non-ionizing particles, they are not conductors at any state Examples; O 2, CO 2, C 2 H 6, H 2 O, SiC

44. Oxygen Atom Oxygen Molecule (O 2 ) Oxygen Molecule (O 2 )

45. SINGLE COVALENT BOND A sharing of two valence electrons. Only nonmetals and Hydrogen. Different from an ionic bond because they actually form molecules. Two specific atoms are joined. In an ionic solid you can’t tell which atom the electrons moved from or to.

46. DIATOMIC MOLECULES Diatomic elements are elements that do not exist singularly in nature because they are highly reactive. “Which elements are diatomic?” “HON, it’s the halogens!” H 2, O 2, N 2, F 2, Cl 2, Br 2, I 2

47. WHEN ELECTRONS ARE SHARED EQUALLY NONPOLAR COVALENT BONDS H 2 or Cl 2

48. WHEN ELECTRONS ARE SHARED BUT SHARED UNEQUALLY POLAR COVALENT BONDS H2OH2O

49. Polar Covalent Bonds: Unevenly matched, but willing to share.

50. - water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

51. FANCY BONDING Sometimes, atoms share more than one electron. Occasionally, they can share 2 or even 3 electrons. These are called double and triple bonds.

52. “M” IS FOR MOLECULE A group of atoms held together by covalent bonds is called a molecule. Water is a molecule, and so is sugar. Other examples of molecules are  methane (CH 4 )  ammonia (NH 3 )  oxygen (O 2 )  nitrogen (N 2 ).

53. PROPERTIES OF COVALENT MOLECULES  Low melting point and boiling point temperatures  Relatively soft solids as compared to ionic compounds  Nonconductors of electricity in any phase

54. CHEMICAL FORMULAS Molecules are represented by a chemical formula. The chemical formula tells you the exact number of each kind of atom in the molecule. For example, the chemical formula for water is H 2 O. The subscript 2 indicates there are two hydrogen atoms in the molecule. The chemical formula also tells you that water always contains twice as many hydrogen atoms as oxygen atoms.

55. CHEMICAL FORMULA Water is a simple molecule, so the formula is pretty easy. Let’s look at a more complex molecule. Baking soda, or sodium bicarbonate, is NaHCO 3. That means it has:  1 Sodium (Na)  1 Hydrogen (H)  1 Carbon (C)  3 Oxygen (O)

56. YOU TRY IT! Let’s see how you do it. Next to each formula, write the name and number of each element. Chemical FormulaElements - # C6H6C6H6 NH 3 Al(OH) 3 CO(NH 2 ) 2

57. LEWIS DOT STRUCTURES Representing Bonds

58. LEWIS DOT MOLECULES We’ve already seen how you draw a Lewis dot structure. The dots represent the valence electrons of an atom. We can draw Lewis dot structures for molecules too. Each element forms bonds to reach one of the magic numbers of valence electrons: 2 or 8. In dot diagrams of a happy molecule, each element symbol has either 2 or 8 dots around it.

59. 58 1.Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center. 2.Count total number of valence e -. Add 1 for each negative charge. Subtract 1 for each positive charge. 3.Use one pair of electrons to form a bond (a single line) between each pair of atoms. 4.Arrange the remaining electrons to satisfy an octet for all atoms (duet for H), starting from outer atoms. 5.If a central atom does not have an octet, move in lone pairs to form double or triple bonds on the central atom as needed. Rules for Writing Lewis Structures

60. LEWIS STRUCTURES OF MOLECULES Single Bond: Two atoms sharing one electron pair. Example: H 2 Double Bond: Two atoms sharing two pairs of electrons. Example: O 2 Triple Bond: Two atoms sharing three pairs of electrons. Example: N 2 Resonance Structures: More than one Lewis Structure can be drawn for a molecule. Example: O 3

61. COVALENT BONDING Fluorine has seven valence electrons F

62. COVALENT BONDING Fluorine has seven valence electrons A second atom also has seven FF

63. Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons FF

64. Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons FF

65. Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons FF

66. Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons FF

67. Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons FF

68. Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons l Both end with full orbitals FF

69. Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons l Both end with full orbitals FF 8 Valence electrons

70. Covalent bonding l Fluorine has seven valence electrons l A second atom also has seven l By sharing electrons l Both end with full orbitals FF 8 Valence electrons

71. 70 A chemical bond in which two or more electrons are shared by two atoms. How should two atoms share electrons? FF + 7e - FF 8e - F F F F Lewis structure of F 2 lone pairs single covalent bond Covalent Bond

72. 71 8e - H H O ++ O HH O HHor 2e - Lewis structure of water Double bond – two atoms share two pairs of electrons single covalent bonds O C O or O C O 8e - double bonds Triple bond – two atoms share three pairs of electrons N N 8e - N N triple bond or

73. 72 Write the Lewis structure of nitrogen trifluoride (NF 3 ). Step 1 – N is less electronegative than F, put N in center FNF F Step 2 – Count valence electrons N - 5 (2s 2 2p 3 ) and F - 7 (2s 2 2p 5 ) 5 + (3 x 7) = 26 valence electrons Step 3 – Draw single bonds between N and F atoms. Step 4 – Arrange remaining 20 electrons to complete octets

74. 73 Write the Lewis structure of the carbonate ion (CO 3 2- ). Step 1 – C is less electronegative than O, put C in center OCO O Step 2 – Count valence electrons C - 4 (2s 2 2p 2 ) and O - 6 (2s 2 2p 4 ) -2 charge – 2e - 4 + (3 x 6) + 2 = 24 valence electrons Step 3 – Draw single bonds between C and O atoms Step 4 - Arrange remaining 18 electrons to complete octets Step 5 – The central C has only 6 electrons. Form a double bond. 22

75. 18, 20 Oct 97BONDING AND STRUCTURE74 BOND AND LONE PAIRS Electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. H Cl This is a LEWIS ELECTRON DOT structure. shared or bond pair Unshared or lone pair (LP)

76. LEWIS DOT MOLECULE - EXAMPLE Draw the dot diagram for carbon tetrachloride, CCl 4. 1.List the elements in the molecule Carbon Chlorine

77. LEWIS DOT MOLECULE - EXAMPLE Draw the dot diagram for carbon tetrachloride, CCl 4. 1.List the elements in the molecule 2.Determine how many valence electrons each element has. Carbon - 4 Chlorine - 7

78. LEWIS DOT MOLECULE - EXAMPLE Draw the dot diagram for carbon tetrachloride, CCl 4. 1.List the elements in the molecule 2.Determine how many valence electrons each element has. 3.Match the elements so that each atom has 8 (or 2 for H & He) electrons. Carbon - 4 Chlorine - 7

79. LEWIS DOT MOLECULE - EXAMPLE Notice that with this molecule, each atom has 8 electrons. The shells are all full!!! Each chlorine atom shares an electron with carbon. In return, carbon shares its electrons with chlorine. We can change the drawing to look like this…

80. LEWIS DOT MOLECULE - EXAMPLE Eventually, this drawing changes into… This one… And finally, into this one….

81. WATER H O Each hydrogen has 1 valence electron Each hydrogen wants 1 more The oxygen has 6 valence electrons The oxygen wants 2 more They share to make each other happy

82. WATER Put the pieces together The first hydrogen is happy The oxygen still wants one more H O

83. WATER The second hydrogen attaches Every atom has full energy levels H O H

84. MULTIPLE BONDS Sometimes atoms share more than one pair of valence electrons. A double bond is when atoms share two pair (4) of electrons. A triple bond is when atoms share three pair (6) of electrons.

85. CARBON DIOXIDE CO 2 - Carbon is central atom ( I have to tell you) Carbon has 4 valence electrons Wants 4 more Oxygen has 6 valence electrons Wants 2 more O C

86. CARBON DIOXIDE Attaching 1 oxygen leaves the oxygen 1 short and the carbon 3 short O C

87. Carbon dioxide l Attaching the second oxygen leaves both oxygen 1 short and the carbon 2 short O C O

88. Carbon dioxide l The only solution is to share more O C O

89. Carbon dioxide l The only solution is to share more O C O

90. Carbon dioxide l The only solution is to share more O CO

91. Carbon dioxide l The only solution is to share more O CO

92. Carbon dioxide l The only solution is to share more O CO

93. Carbon dioxide l The only solution is to share more O CO

94. Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom gets to count all the atoms in the bond O CO

95. Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom gets to count all the atoms in the bond O CO 8 valence electrons

96. Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom gets to count all the atoms in the bond O CO 8 valence electrons

97. Carbon dioxide l The only solution is to share more l Requires two double bonds l Each atom gets to count all the atoms in the bond O CO 8 valence electrons

98. HOW TO DRAW THEM Add up all the valence electrons. Count up the total number of electrons to make all atoms happy. Subtract. Divide by 2 Tells you how many bonds - draw them. Fill in the rest of the valence electrons to fill atoms up.

99. EXAMPLES NH 3 N - has 5 valence electrons wants 8 H - has 1 valence electrons wants 2 NH 3 has 5+3(1) = 8 NH 3 wants 8+3(2) = 14 (14-8)/2= 3 bonds 4 atoms with 3 bonds N H

100. NHH H EXAMPLES Draw in the bonds All 8 electrons are accounted for Everything is full

101. EXAMPLES HCN C is central atom N - has 5 valence electrons wants 8 C - has 4 valence electrons wants 8 H - has 1 valence electrons wants 2 HCN has 5+4+1 = 10 HCN wants 8+8+2 = 18 (18-10)/2= 4 bonds 3 atoms with 4 bonds -will require multiple bonds - not to H

102. HCN Put in single bonds Need 2 more bonds Must go between C and N NHC

103. HCN l Put in single bonds l Need 2 more bonds l Must go between C and N l Uses 8 electrons - 2 more to add NHC

104. HCN l Put in single bonds l Need 2 more bonds l Must go between C and N l Uses 8 electrons - 2 more to add l Must go on N to fill octet NHC

105. POLAR MOLECULES Molecules with ends

106. 105 H F F H A covalent bond with greater electron density around one of the two atoms electron rich region electron poor region e - riche - poor ++ -- Polar Covalent Bond

107. POLAR MOLECULES Molecules with a positive and a negative end This can be determined from differences in electronegativity.

108. ELECTRONEGATIVITY A measure of how strongly the atoms attract electrons in a bond. The bigger the electronegativity difference the more polar the bond. 0.0 - 0.3 Covalent nonpolar 0.3 - 1.67 Covalent polar >1.67 Ionic

109. 108 Electronegativities (EN) The ability of an atom in a molecule to attract shared electrons to itself Linus Pauling 1901 - 1994

110. HOW TO SHOW A BOND IS POLAR Isn’t a whole charge just a partial charge  means a partially positive  means a partially negative The Cl pulls harder on the electrons The electrons spend more time near the Cl H Cl  

111. 110 Covalent share e - Polar Covalent partial transfer of e - Ionic transfer e - Increasing difference in electronegativity Classification of Bonds Difference in ENBond Type 0Covalent  2 Ionic 0 < and <2 Polar Covalent

112. 111 Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H 2 S; and the NN bond in H 2 NNH 2. Cs – 0.7Cl – 3.03.0 – 0.7 = 2.3Ionic H – 2.1S – 2.52.5 – 2.1 = 0.4Polar Covalent N – 3.0 3.0 – 3.0 = 0Covalent Classification of Bonds

113. IS IT POLAR? HF H2O H2O NH 3 CCl 4 CO 2

114. INTERMOLECULAR FORCES What holds molecules to each other

115. INTERMOLECULAR FORCES They are what make solid and liquid molecular compounds possible. The weakest are called van der Waal’s forces - there are two kinds Dispersion forces Dipole Interactions  depend on the number of electrons  more electrons stronger forces  Bigger molecules

116. DIPOLE INTERACTIONS  Depend on the number of electrons More electrons stronger forces Bigger molecules more electrons  Fluorine is a gas  Bromine is a liquid  Iodine is a solid

117. DIPOLE INTERACTIONS Occur when polar molecules are attracted to each other. Slightly stronger than dispersion forces. Opposites attract but not completely hooked like in ionic solids. HFHF  HFHF 

118. DIPOLE INTERACTIONS    

119. HYDROGEN BONDS

120. HYDROGEN BONDING Are the attractive force caused by hydrogen bonded to F, O, or N. F, O, and N are very electronegative so it is a very strong dipole. The hydrogen partially share with the lone pair in the molecule next to it. The strongest of the intermolecular forces.

121. HYDROGEN BONDING H H O ++ -- ++ H H O ++ -- ++

122. H H O H H O H H O H H O H H O H H O H H O

123. METALLIC BONDS

124. METALLIC BOND BOND FOUND IN METALS; HOLDS METAL ATOMS TOGETHER VERY STRONGLY

125. METALLIC BONDS How atoms are held together in the solid. Metals hold onto their valence electrons very weakly. Think of them as positive ions floating in a sea of electrons.

126. IONIC BOND, A SEA OF ELECTRONS

127. SEA OF ELECTRONS ++++ ++++ ++++ Electrons are free to move through the solid. Metals conduct electricity.

128. PROPERTIES Good conductors at all states, lustrous, very high melting points Malleable- Hammered into shape (bend). Ductile - drawn into wires.

129. MALLEABLE ++++ ++++ ++++

130. ++++ ++++ ++++ Electrons allow atoms to slide by.

131. Metallic Bonds: Mellow dogs with plenty of bones to go around.

132. MOLECULAR SHAPES OF COVALENT COMPOUNDS

133. LINEAR Number of Bonds = 2 Number of Shared Pairs of Electrons = 2 Bond Angle = 180° EXAMPLE: BeF 2

134. TRIGONAL PLANAR Number of Bonds = 3 Number of Shared Pairs of Electrons = 3 Number of Unshared Pairs of Electrons = 0 Bond Angle = 120° EXAMPLE: GaF 3

135. BENT #1 Number of Bonds = 2 Number of Shared Pairs of Electrons = 2 Number of Unshared Pairs of Electrons = 2 Bond Angle = < 120° EXAMPLE: H 2 O

136. BENT #2 Number of Bonds = 2 Number of Shared Pairs of Electrons = 2 Number of Unshared Pairs of Electrons = 1 Bond Angle = >120° EXAMPLE: O 3

137. TETRAHEDRAL Number of Bonds = 4 Number of Shared Pairs of Electrons = 4 Number of Unshared Pairs of Electrons = 0 Bond Angle = 109.5° EXAMPLE: CH 4

138. TRIGONAL PYRAMIDAL Number of Bonds = 3 Number of Shared Pairs of Electrons = 4 Number of Unshared Pairs of Electrons = 1 Bond Angle = <109.5° EXAMPLE: NH 3

139. TRIGONAL BIPYRAMIDAL Number of Bonds = 5 Number of Shared Pairs of Electrons = 5 Number of Unshared Pairs of Electrons = 0 Bond Angle = <120° EXAMPLE: NbF 5

140. OCTAHEDRAL Number of Bonds = 6 Number of Shared Pairs of Electrons = 6 Number of Unshared Pairs of Electrons = 1 Bond Angle = 90° EXAMPLE: SF 6

141. REVIEW

142. TYPES OF CHEMICAL BONDS Ionic Compound: A compound resulting from a positive ion (usually a metal) combining with a negative ion (usually a non-metal). Example: M + + X -  MX Covalent Bond: Electrons are shared by nuclei. Example: H-H Polar Covalent Bond: Unequal sharing of electrons by nuclei. Example: H-F Hydrogen fluoride is an example of a molecule that has bond polarity.

143. TYPES OF CHEMICAL BONDING 1. Metal with nonmetal:  electron transfer and ionic bonding 2. Nonmetal with nonmetal:  electron sharing and covalent bonding 3. Metal with metal:  electron pooling and metallic bonding 142

144. 143 Three models of chemical bonding Electron transferElectron sharingElectron pooling Ionic CovalentMetallic