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Chemistry Chemical Bonding. The Development of Atomic Models 1. Dalton – solid, indivisible mass.

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Presentation on theme: "Chemistry Chemical Bonding. The Development of Atomic Models 1. Dalton – solid, indivisible mass."— Presentation transcript:

1 Chemistry Chemical Bonding

2 The Development of Atomic Models 1. Dalton – solid, indivisible mass

3 2. Thomson – “plum-pudding” model -Negatively charged e- (raisins) stuck in positively charged proton dough -No neutrons

4 3. Rutherford – electrons surrounding a dense nucleus

5 4. Bohr model – elctrons arranged in symmetrical orbits around the nuclues -“planetary model” -Electrons in a given path have a fixed energy level

6 5. Quantum mechanical model – modern mathematical description of the atom Sodium atom:

7 Energy level – region around the nucleus where the electron is likely to be moving. An electron can jump from one level to another by absorbing energy.

8 Quantum – the amount of energy required to move an electron from its present energy level to the next higher one “quantum leap”

9 Quantum mechanical model – uses mathematical equations to describe the location and energy of electrons in an atom Developed by Erwin Schrodinger Electrons are not in definite paths Their location is described in terms of probability of being in a certain region Electron cloud (ceiling fan) Conventionally, the border is drawn at 90% probability

10 Atomic orbital – region in space that an electron is likely to be in Electrons can be described by a series of 4 quantum numbers.

11 1. Principle quantum number (n) Describes the energy level Values of 1, 2, 3, 4, etc.

12 2. Azimuthal quantum number ( l ) Describes the shape of atomic orbitals Sublevels Values of 0 to n-1 0 = s, 1 = p, 2 = d, 3 = f

13 s = spherical, p = peanut shape, d&f = more complex shapes d = “daisy” f = “fancy” So if n = 1, then l can be 0 (s) = 1 sublevel n = 4, then l can be 0 (s), 1 (p), 2(d), 3(f), = 4 sublevels

14 s orbitals – spherical s orbitals – spherical

15 p orbitals – dumbbell-shaped

16 d orbitals – daisy-shaped

17 f orbitals – fancy shapes

18 3. Magnetic quantum number (m l ) Orientation of the orbital in space Values of – l to + l So s has 1 orbital p has 3 d has 5 f has 7

19 4. Spin quantum number (m s ) Values of +½ and -½ Each orbital can hold 2 electrons with opposite spins Since spinning charged objects create a magnetic field, the electrons must spin opposite directions to minimize repulsion

20

21 Ex. How many orbitals are in the following? A. 3pD. 4p B. 2sE. 3d C. 4fF. 3 rd energy level How many electrons can be in each of the above?

22 Are these possible? nlmlml msms 100+1/ /2 21+1/2

23 Electron Configuration – ways in which electrons are arranged around nuclei of atoms.

24 Rules that govern filling of atomic orbitals: 1. Aufbau principle – electrons enter orbitals of lowest energy first.

25

26 2. Pauli exclusion principle – An atomic orbital can describe at most two electrons. They must have opposite spins.

27 3. Hund’s rule – When electrons occupy orbital of equal energy, one electron enters each orbital until all orbitals contain one electron with parallel spins.

28

29 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 4f

30 Na = 11 1s 2 2s 2 2p 6 3s 1 Cd =48 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10

31 Practice: Write the electron configuration for the following elements: Li O Sc

32 More practice: Identify each of the following atoms on the basis of its electron configurations. a) 1s 2 2s 2 2p 6 neon b) 1s 2 2s 2 2p 6 3s 1 sodium c) [Kr] 5s 2 4d 2 zirconium d) [Xe] 6s 2 4f 6 samarium

33 Ground state – lowest energy level for an electron. (Normal, nonexcited state)

34 Exceptional Electron Configurations: Cr:expected: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 actual: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Cu:expected: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 actual: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 Half-filled energy levels are more stable than other partially filled energy levels. There are other exceptions.

35 Light and Atomic Spectra Electromagnetic radiation – a series of energy waves that includes radio waves, microwaves, visible light, infrared, and ultraviolet light, X-rays, and gamma rays.

36

37 Wavelength, crest trough Parts of a wave:

38 Amplitude – height of the wave from the origin to the crest. Wavelength - - distance between the crests Frequency – f  or  – the number of wave cycles to pass a given point per unit of time. The units of frequency are 1/s, s -1, or Hertz (Hz)

39 c= f where c = speed of light = 3.00x10 8 m/s or 3.00x10 10 cm/s As increases, f decreases.

40 As wavelength increases, frequency decreases. As wavelength decreases, frequency increases.

41

42 Ex. A certain wavelength of yellow light has a frequency of 2.73x10 16 s -1. Calculate its wavelength. Convert to nm. C = f  c  f  3.00x10 8 m/s  2.73x10 16 s -1   x   m

43 Spectrum – series of colors produced when sunlight is separated by being passed through a prism. ROY G. BIV Red: longest wavelength, lowest frequency Violet: shortest wavelength, highest frequency

44

45 Atomic emission spectrum – series of lines of colored light produced by passing light emitted by an “excited atom through a prism.” This can be used to identify the element. The atomic emission spectrum of hydrogen shows three series of lines. The lines in the UV region (Lyman series) represent electrons falling to n=1, lines in the visible region (Balmer series) represent electrons falling to n=2 and lines in the IR region (Paschen series) represent electrons falling to n=3.

46

47

48 Max Planck found that the energy emitted or absorbed by a body changes only in small discrete units he called quanta. He determined that the amount of radiant energy, E, absorbed or emitted by a body is proportional to the frequency of the radiation. E=h f E = energy (J) f = frequency h = Planck’s constant, 6.626x Js

49 Einstein studied the photoelectric effect whereby light of sufficient frequency shining on a metal causes current to flow. The amplitude of the radiation was not important, the frequency was. This told him that light must be in particles, each having a given energy. Einstein proposed that electromagnetic radiation can be viewed as a stream of particles called photons: E=h f

50 Photoelectric Effect Electron (photons) light metal

51 Energy of a photon: E = h f Energy of a photon: E = h f

52 Example: Calculate the energy of an individual photon of yellow light having a frequency of 2.73x10 16 s -1. E=h f E = (6.626x Js)(2.73x10 16 s -1 ) E = 1.81x J

53 Einstein’s special theory of relativity: E=mc 2 Matter and energy are different forms of the same entity.

54 Going Further:

55 Louis deBroglie suggested that very small particles like electrons might also display wave particles and he came up with: deBroglie’s equation: = h mv m = mass in kg v = velocity in m/s h = Planck’s constant, 6.626x Js

56 DeBroglie’s equation is used to find the wavelength of a particle. It was determined that matter behaves as through it were moving in a wave. This is important in small object such as electrons but is negligible in larger objects such as baseballs. Heavy objects have very short wavelengths.

57 Example: Calculate the wavelength of an electron traveling at 1.24x10 7 m/s. The mass of an electron is 9.11x g. = h mv 9.11x g 1 kg = 9.11x kg 1000g = (6.626x Js) = (9.11x kg)(1.24x10 7 m/s) = 5.87x m

58 End Going Further.

59 In the photoelectric effect, electrons (called photoelectrons) are ejected by metals (esp. alkali) when light of sufficient frequency shines on them. Red light won’t work. Photoelectric cells convert light energy into electrical energy. They are used in automatically opening doors and security systems.

60 Heisenberg’s Uncertainty Principle – it is impossible to determine accurately both the momentum and the position of an electron simultaneously. We detect motion by electromagnetic radiation. This interaction disturbs electrons.

61 Einstein and Heisenberg!

62 Mendeleev- arranged elements in order of increasing atomic mass Moseley- arranged elements in order of increasing atomic number

63 Periodic Law - When the elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties.

64 Be able to locate noble gases, representative elements, transition metals, inner transition metals.

65 Noble Gases have completely filled shells of electrons similar electronic structures ◦ He1s 2 ◦ Ne1s 2 2s 2 2p 6 ◦ Ar1s 2 2s 2 2p 6 3s 2 3p 6 ◦ Kr1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 4p 6 ◦ etc.

66 Representative Elements elements in A groups on periodic chart representative - because they best represent what we know about elemental structure & periodicity

67 d - Transition Elements  elements in B groups on periodic chart  metals  have d electrons  transition from metals to nonmetals

68 f - Transition Elements  inner transition metals

69 Electron configuration using the periodic table:

70

71 Columns are called groups or families. Rows are called periods or series.

72 Shorthand electron configuration:

73

74

75 Periodic Trends in atomic size: covalent atomic radius - half the distance between the nuclei of atoms in a homonuclear diatomic molecule (like Cl 2 ) Trends: 

76

77

78 Atomic size increases going down a group because electrons are added to higher energy levels that are farther from the nucleus. It decreases going across a period because as each proton is added to the nucleus an electron is being added to the same energy level. This “shell” of electrons is pulled closer in towards the nucleus.

79 Size changes little in the transition metals because the electrons being added are core electrons.

80 Z eff = effective nuclear charge -actual pull of the nucleus on the valence electrons. Z eff = Z actual - effect of e - repulsions

81 Trends:  Increases from H to He Decreases from He to Li because 1s electrons shield 2s electrons Increases from Li to Be Decreases from Be to B because 2s shield 2p

82 Increases from B to C to N Decreases from N to O because of repulsion due to doubly occupied orbitals Increases from O to F to Ne Decreases from Ne to Na because 1s,2s & 2p shield 3s

83 Know exceptions to Z eff trends and reasons for these.

84 Ionization energy (IE) - energy required to remove the highest energy electron from a gaseous atom Li(g) + energy  Li + (g) + e - - depends on Z eff and size

85

86 Trends:  Ionization energy decreases down a a group because the valence electrons are farther from the nucleus and are thus held less tightly. Ionization energy increases across a period because atomic size decreases and valence electrons are held more tightly. Z eff increases.

87

88 1 st ionization energy = energy required to remove the first electron Al + energy  Al + + e - 2 nd ionization energy = energy required to remove the second electron Al + + energy  Al 2+ + e - 3 rd ionization energy = energy required to remove the third electron Al 2+ + energy  Al 3+ + e -

89 For Na, the 1 st ionization energy is fairly low but the 2 nd would be high.

90 Know exceptions to ionization energy trends and reasons for them.

91 Ionic Size Anions are larger than the atoms from which they were formed. Know why!!! Cations are smaller than the atoms from which they were formed. Know why!!!!

92 Sizes of Ions Related to Positions of the Elements in the Periodic Table

93 Isoelectronic ions- a group of ions with the same number of electrons The one with the highest atomic number is the smallest in size (More protons pulling on the same # of electrons).

94 Na +, Mg 2+, Ne, F -, O 2-, N 3- are isoelectronic. They all have 10 electrons. Mg 2+ is the smallest because it has 12 protons pulling on 10 electrons. (The protons win the “tug of war”) N 3- is the largest because it has 7 protons and 10 electrons. (The electrons win the “tug of war”)

95 Electron Affinity -energy change that occurs when an electron is added to a gaseous atom -usually exothermic Cl(g) + e -  Cl - (g)+ energy Trends:  (but many exceptions)

96

97 Electronegativity -relative tendency of an atom to attract shared electrons to itself Trends:  FONCl (Phone Call)

98

99 Elements with high electronegativity (nonmetals ) tend to gain electrons to form anions. Elements with low electronegativities (metals) often lose electrons to form cations.

100 Valence electrons- electrons in the highest occupied energy level of an atom. Valence electrons are the only electrons involved in the formation of chemical bonds.

101 Electron dot structures for atoms: -each dot represents a valence electron p p X s p Electron dot structures for atoms: -each dot represents a valence electron p p X s p

102 Examples: N

103 N

104 O O

105 O

106 Xe

107

108 Al Al

109

110 Na Na

111

112 I

113 I

114 Si Si

115

116 One of the major “driving forces” in nature is the tendency to go to lower energy. Atoms lose, gain or share electrons to become lower in energy and thus more stable.

117 Metals lose electrons easily to become positively charged cations. They will usually lose their valence electrons to achieve a noble gas electron configuration. Na  Na + + e - [Ne]3s 1 [Ne] Al  Al e - [Ne]3s 2 3p 1 [Ne]

118 Some transition metals lose their highest energy level s and p electrons but still have d electrons remaining. Their electron configuration is not quite that of a noble gas but is still stable. It is called a pseudo-noble gas electron configuration. For example, zinc loses its two electrons in 4s but keeps the ten electrons in 3d.

119 Transition metals always lose their highest numerical energy level electrons first. Transition metals in the 4th period lose their 4s and 4p electrons before losing any from 3d. Metals in groups 3, 4, & 5 do this also. Transition metals always lose their highest numerical energy level electrons first. Transition metals in the 4 th period lose their 4s and 4p electrons before losing any from 3d. Metals in groups 3, 4, & 5 do this also.

120 Example: Fe [Ar]4s 2 3d 6 Fe 2+ [Ar]3d 6 Fe 3+ [Ar]3d 5

121 Nonmetals tend to gain electrons to become stable and form negatively charged anions. They achieve a noble gas electron configuration. Example: Cl + e -  [ Cl ] - [Ne]3s 2 3p 5 [Ne]3s 2 3p 6 N + 3e -  [ N ] 3- 1s 2 2s 2 2p 3 1s 2 2s 2 2p 6

122 Ionic Bonding- the attraction of oppositely charged ions (cations and anions) When the electronegativity difference between two elements is large, the elements are likely to form a compound by ionic bonding (transfer of electrons). The farther apart across the periodic table two Group A elements are, the more ionic their bonding will be.

123 We can use Lewis dot formulas to represent the formation of ionic compounds. Na + Cl  Na + [ Cl ] - or NaCl Mg: N Mg 2+ [ N ] 3- Mg:  Mg 2+ or Mg 3 N 2 Mg: N Mg 2+ [ N ] 3-

124 Properties of Ionic Compounds: They are usually crystalline solids with high melting points (>400 o C) Their molten compounds and aqueous solutions conduct electricity well because they contain mobile charged particles.

125 Metals Metals form metallic solids that consist of positively charged metal cations in a “sea” of loosely held valence electrons. This arrangement allows metals to have their unique properties.

126

127 Metals are ductile (can be pulled into a wire) and malleable (can be hammered into a thin sheet) because the valence electrons act as “grease”, allowing the cations to slide past each other without colliding with each other and shattering. When ionic compounds such as NaCl are hammered, like-charged ions collide causing repulsion and the crystal shatters.

128 Metals can conduct electricity easily. Electricity is a flow of electrons. As electricity (electrons) enters one end of a piece of metal, an equal number of electrons exit the other end.

129 Alloys- solutions of solids in solids

130 Substitutional alloy- atoms of one metal are substituted for atoms of a similar-sized metal in a metallic crystal. Ex. brass, sterling silver, pewter Interstitial alloy- smaller metal atoms fit into holes in the crystal structure of a metal with larger atoms Ex. steel (carbon in iron)

131

132 Amalgam- alloy which contains mercury

133 Covalent Bonding

134 Hydrogen and nonmetals of Groups 4,5,6 & 7 often become stable and gain noble gas electron configurations by sharing electrons to form covalent bonds. Atoms will usually share electrons to follow the octet rule (eight electrons, like most noble gases) or the duet rule (2 electrons, like helium).

135 When atoms share one pair of electrons to form a covalent bond, it is called a single covalent bond. The electrons shared between the atoms are a “shared pair”. A dash can be used instead of two dots to represent the shared pair. Any other electrons on the atoms are “unshared pairs” or “lone pairs”.

136 Ex. H 2 H-H Cl 2 Cl-Cl HCl H-Cl Ex. H 2 H-H Cl 2 Cl-Cl HCl H-Cl

137 Atoms must sometimes share more than one pair of electrons to become stable. When two pair of electrons are shared between two atoms, it is called a double bond. If three pair are shared, it is a triple bond. Ex. O 2 N 2 O = O N  N

138 Rules for Writing Lewis Structures (electron dot structures): (Use pencil!) 1.Add up the valence electrons from all the atoms. Don’t worry about keeping track of which electrons come from which atoms. If you are working with an ion, you must add or subtract electrons to equal the charge. 2.Use a pair of electrons to form a bond between each pair of bound atoms.

139 3.Arrange the remaining electrons to satisfy the duet rule for hydrogen and the octet rule for everything else. 4.If necessary, change bonds to double or triple. 5.Remember, we cannot create or destroy electrons!

140 H 2 O H 2 O 8 electrons H-O-H

141 NH 3 NH 3 8 electrons H-N-H H

142 NH = 8 electrons H + H-N-H H

143 CO 2 CO 2 16 electrons O-C-O This used 20 electrons! BAD!!!

144 CO 2 O=C=O

145 CCl 4 CCl 4 32 electrons

146 CCl 4 CCl 4 32 electrons Cl Cl-C-Cl Cl

147 CN = 10 electrons C-N BAD! BAD!

148 CN = 10 electrons C  N

149 SO 4 2- SO electrons O 2 - O S O O

150 CO electrons CO electrons O O C O O 2- 2-

151 Coordinate Covalent Bond- Bond in which both electrons came from the same atom. This bond is not really any different than any other single bond.

152 Exceptions to the Octet Rule: A few compounds are stable with less than an octet. They include beryllium or boron. These electron deficient compounds are very reactive.

153 Ex. BF 3 BeCl 2 Ex. BF 3 BeCl 2 24 electrons 16 electrons F Cl-Be-Cl F-B-F

154 Elements in the third period and below can exceed the octet rule. They can place extra electrons in empty d orbitals. Elements in the second period can not exceed the octet rule because there is no 2d orbital for the extra electrons to go into. If it is necessary to exceed the octet rule, place the extra electrons on the central atom.

155 Ex. PCl 5 SF 6 Cl Cl F F F P S Cl Cl Cl F F F

156 I3-I3-I3-I3- 22 electrons [ I-I-I ] -

157 More practice: NF 3

158 F F - N - F F F - N - F

159 OF 2 20 electrons OF 2 20 electrons

160 F - O - F

161 KrF 4 36 electrons

162 F F- Kr - F F F F- Kr - F F

163 BeH 2 4 electrons BeH 2 4 electrons

164 H - Be - H H - Be - H

165 SO electrons

166 O O - S - O O O - S - O

167 NO electrons NO electrons

168 O O - N - O O O - N - O

169 H 2 O 2 14 electrons H 2 O 2 14 electrons

170 H - O - O - H H - O - O - H

171 Resonance occurs when more than one valid Lewis structure can be written for a molecule. The actual structure is an average of all of the resonance structures.

172 Ex. NO 3 - O O O O - N - O O - N - O O - N - O

173 In nitrate, the experimental bond length is in-between that of a single bond and a double bond. It acts like a 1 1/3 bond.

174 Ex. Benzene, C 6 H 6

175

176 VSEPR

177 Lewis structures can be used to determine the shapes of molecules. Their shapes will tell us a lot about their chemical behavior.

178 The valence shell electron pair repulsion (VSEPR) theory tells us that valence electrons on the central atom repel each other. They are arranged as far apart as possible around the central atom so that repulsions among them are as small as possible. When we are using VSEPR to determine molecular shape, we are really looking for regions of electron density. Double and triple bonds count the same as single bonds in determining molecular shape.

179 In CO 2, there are only two regions of electron density (effective electron pairs) around the central atom. These regions arrange themselves as far apart as possible, making the bond angle 180 o and the molecular shape linear. O = C = O

180 In CO 3 2-, there are three effective electron pairs around the central atom. The bond angle will be 120 o and the shape will be trigonal planar. 2 - O C O O In CO 3 2-, there are three effective electron pairs around the central atom. The bond angle will be 120 o and the shape will be trigonal planar. 2 - O C O O

181 In CH 4, there are four effective electron pairs. We might expect the bond angle to be 90 o. Actually, since molecules are three-dimensional, the electron pairs are o apart (further than 90 o ) and take a tetrahedral arrangement. H H C H H

182

183 In NH 3, there are four effective electron pairs. Three are shared but one is unshared. Unshared pairs of electrons take up more space than shared pair because they are pulled closer to the nucleus. The presence of the unshared pair distorts the other bond angles, making them less than o and the shape is called trigonal pyramidal. (The bond angles in ammonia are about 107 o.) H - N - H H

184

185 In H 2 O, there are four effective electron pairs, also. Two are shared and two are unshared. Since the unshared pairs repel more than shared pair, the bond angle is less than o (actually o for H 2 O) and the shape is bent. H – O – H

186

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188 In PF 5, there are five effective pairs, all shared. The bond angles are 90 o and 120 o and the shape is called trigonal bipyramidal. It is like two trigonal pyramids with their bases touching. F F F - P - F F

189

190 In SF 6, there are six effective pairs, all shared. The bond angles are 90o and the shape is called octahedral. It is like two square pyramids with their bases touching. In SF 6, there are six effective pairs, all shared. The bond angles are 90 o and the shape is called octahedral. It is like two square pyramids with their bases touching.

191

192 Molecules that exceed the octet rule and have unshared electrons can have more complex shapes such as T-shaped, see-saw, and square pyramidal.

193 Practice Determining Molecular Shape:

194 H 2 S H - S - H bent, <109.5 o

195 CCl 4 Cl Cl-C-Cl Cl tetrahedral, o

196 NH 4 + H H - N - H H NH 4 + H H - N - H H tetrahedral, o

197 BF 3 F B F F trigonal planar, 120 o

198 NO 2 - O - N = O bent, <120 o

199 PF 6 - F F F P F F F octahedral, 90 o

200 SbCl 5 Cl Cl Cl Sb Cl Cl SbCl 5 Cl Cl Cl Sb Cl Cl trigonal bipyramidal, 90 o and 120 o

201 SO 2 O - S = O bent, <120 o

202 POLAR AND NONPOLAR MOLECULES Covalent bonds within a molecule can be polar (electrons are shared unequally) or nonpolar (electrons are shared equally). We can predict polar bonds by looking up the electronegativity values of each element in the bond and subtracting the smaller value from the larger value to determine the electronegativity difference.

203 If the electronegativity difference (  EN) is zero, the bond is nonpolar covalent. If the  EN is between 0 and 2.0, we can predict that the bond is polar covalent. If the  EN is 2.0 or greater, the bond is usually considered to be ionic. If the electronegativity difference (  EN) is zero, the bond is nonpolar covalent. If the  EN is between 0 and 2.0, we can predict that the bond is polar covalent. If the  EN is 2.0 or greater, the bond is usually considered to be ionic.

204 There are no strict dividing lines between covalent and ionic bonds. Physical properties are used to help determine whether something is covalent or ionic. Some covalent substances are so polar that they ionize partially or completely in water.

205

206 Polarity of a bond or molecule can be represented by arrows or lowercase delta (  ) symbols to show a partial charge. Ex. H F  +  - This shows that the hydrogen end of the molecule is more positive (less electronegative) and the fluorine end is more negative (more electronegative). A charge difference such as this is called a dipole.

207 An entire molecule is polar if it has polar bonds that do not cancel. This happens in nonsymmetrical molecules. Molecules with all nonpolar bonds such as H 2, O 2, and Cl 2, are always nonpolar. This means that there are no positive and negative ends on the molecule.

208 Heteronuclear diatomic molecules such as HCl, BrCl, or HF are always polar because the polar bonds can not cancel each other out.

209 Linear triatomic molecules are polar only if the two outside legs are different. CO 2 is nonpolar because the two C=O bonds cancel each other out. Bent molecules are polar.

210 Trigonal planar molecules are nonpolar if the three legs are the same. Trigonal pyramidal molecules are polar. Tetrahedral molecules are nonpolar if the four legs are all the same. Trigonal bipyramidal and octahedral molecules are nonpolar if all legs are the same element.

211 Practice: Polar or Nonpolar? NH 3 H H - N - H Polar 1 st : Draw dot structure 2 nd : Determine shape 3 rd : Symmetrical = Nonpolar Nonsymmetrical = Polar Nonsymmetrical = Polar

212 SO 2 SO 2 O = S - O Polar

213 H 2 O H 2 O H - O - H Polar

214 BF 3 BF 3 F F - B - F Nonpolar

215 CH 4 CH 4 H H - C - H H Nonpolar

216 INTERMOLECULAR FORCES

217 Intramolecular bonding- sharing electrons Intermolecular bonding- interactions between particles (atoms, molecules or ions)

218 Changes in state are due to changes in intermolecular bonding, not intramolecular bonding.

219 Dipole-dipole attraction- attraction of polar molecules for each other. (negative-positive) *approx. 1% as strong as covalent or ionic bonds. *molecules orient themselves to minimize repulsion and maximize attractions

220

221 Hydrogen bonding- unusually strong dipole- dipole attractions involving hydrogen atoms which are covalently bonded to a very electronegative element and a very electronegative atom (F,O,N only) with unshared electrons.

222

223 Two reasons for strength of hydrogen bonds: 1. small size of H atom allows closeness 2. great polarity Substances with much H bonding have high boiling points compared to similar substances. Ex. H 2 O, NH 3, HF

224

225 London dispersion forces ( LDFs)- relatively weak forces (usually) that exist between noble gas atoms and nonpolar molecules. LDFs also exist in compounds that have dipole-dipole and/or hydrogen bonding. LDFs may be the most important force in large molecules of these types. LDFs occur because of momentary electron imbalance (temporary dipole) which can induce the same to occur in adjacent molecules.

226

227 This force is often very weak, thus the low freezing point of noble gases. The freezing point of noble gases increases going down the group because heavier atoms have more electrons and an increased chance of temporary dipoles. They also have a lower velocity and have more opportunity for attractions. This causes London dispersion forces to increase going down a group on the periodic table.

228 Covalent network solid- Group 4 substances such as diamond, Si, SiC, and Ge form extensive covalent bonds and result in giant molecules. They have an atom at each lattice point and are held together by covalent bonds. These substances have the strongest attractions.

229 General trends in strength of attraction: (weakest) LDF dipole-dipole hydrogen-bonding metallic bonding ionic bonding covalent network bonding (strongest)


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