Presentation is loading. Please wait.

Presentation is loading. Please wait.

Chapter 10: Structure of Solids and Liquids Chem 1110 Figures: Basic Chemistry 3 rd Ed., Timberlake and Timberlake.

Similar presentations


Presentation on theme: "Chapter 10: Structure of Solids and Liquids Chem 1110 Figures: Basic Chemistry 3 rd Ed., Timberlake and Timberlake."— Presentation transcript:

1 Chapter 10: Structure of Solids and Liquids Chem 1110 Figures: Basic Chemistry 3 rd Ed., Timberlake and Timberlake

2 Electron-Dot Symbols Electron-dot symbols: Show the valence electrons of an atom Electrons are arranged in s and p orbitals around the element symbol Can mix orbitals and create “hybrid” orbitals with equal energy (like in Chapter 6) o sp, sp 2, and sp 3

3 Covalent Bond Revisited Covalent Bond is a chemical bond that results from two nuclei attracting the same shared electron pair Shared electrons in the bond For Example: F 2

4 Bonding and Lone Pair Electrons Bonding Pair are shared to create the covalent bond Non-Bonding Pair or lone pair make up the octet

5 Bonding and Lone Pair Electrons Bonding Pair are shared to create the covalent bond Non-Bonding Pair or lone pair make up the octet H 2 HCl H 2 O

6 Rules for Drawing Electron Dot Formulas 1.Find the total number of valence electrons contributed by all atoms 2.Write chemical symbols in the order that they are bonded 3.Place one covalent bond between each atom Determining the central atom is important First atom is usually the central atom o EXCEPT when it is hydrogen (H)

7 Rules for Drawing Electron Dot Formulas 4.Add lone pairs to the bonded atoms first, then to central atom 5.Check the Octet Rule for ALL atoms May need to make multiple bonds For Example: SF 2

8 Rules for Drawing Electron Dot Formulas 1.Find the total number of valence electrons contributed by all atoms 2.Write chemical symbols in the order that they are bonded 3.Place one covalent bond between each atom 4.Add lone pairs to the bonded atoms first, then to central atom 5.Check the Octet Rule for ALL atoms May need to make multiple bonds

9 Electron Dot Formulas EXAMPLES: CO 2 NH 3 SiCl 4

10 Exceptions to the Octet Rule Boron is stable with a total of 6 electrons Boron will make compounds with only three bonds to B BCl 3 Period 3 (and above) non-metals can form compounds with expanded octets (> 8 valence electrons): P, S, Se

11 Multiple Bonds Multiple bonds involve the sharing of more than one pair of electrons:

12 Bond Order Bond Order denotes the level of bonding in a molecule: Single bond: Share one electron pair Bond Order of 1 Double bond: Share two electron pairs Bond Order of 2 Triple bond: Share three electron pairs Bond Order of 3

13 Bond Length and Bond Energy As we add additional bonds between atoms the atoms are drawn closer together: Bond length decreases with more bonds: Single > Double > Triple As we add additional bonds between atoms the bonds get stronger: Bond energy increases with more bonds: Triple > Double > Single

14 Bond Length and Bond Energy Practicing organic compounds: Single bond (alkane): Ethane, C 2 H 6 Double bond (alkene): Ethene, C 2 H 4 Triple bond (alkyne): Ethyne (Acetylene), C 2 H 2

15 Electron Dot Formulas EXAMPLES: O 2 HCN CS 2 CO

16 Resonance Structures Resonance structures are equivalent structures that show where multiple bonds may occur Electrons are delocalized over entire molecule Resonance changes bond order, bond length, and energy Examples: SO 2 SO 3

17 Draw the three resonance structures of carbonate, CO 3 2- Learning Check

18 Electron Dot Formulas How can we draw electron dot structures for polyatomic ions? NH 4 Br or Na 2 SO 3 These compounds have both an ionic and covalent: [NH 4 ] + [Br] - 2 [Na] + [SO 3 ] 2-

19 Electron Dot Formulas [NH 4 ] + [SO 3 ] 2-

20 Draw the electron dot formula for ClO 3 - Learning Check

21 Write two resonance structures for nitrite. Learning Check

22 Orbital Geometry and Molecular Shape Electrons in bonds and lone pairs want to get as far apart from each other as possible We have seen the tetrahedral shape of carbon Causes the molecule to conform to a set of shapes

23 Determining Orbital Geometry Electron pairs do NOT like to share space with other electron pairs Electrons will move away as far as possible to obtain “personal space” o An electron pair can be in a bond o An electron pair can be a “lone pair” (non-bonding)

24 VSEPR Theory: Valence-Shell Electron Pair Repulsion Theory Areas of electron density around an atom want to maximize their “personal space” A lone pair of electrons is ONE area Two electrons in a single bond are ONE area Four electrons in a double bond are ONE area Six electrons in a triple bond are ONE area

25 VSEPR Theory Determining Orbital Geometry: Count ALL areas of electron density to determine the base geometry: Two is linear (sp) Three is trigonal planar (sp 2 ) Four is tetrahedral (sp 3 )

26 Orbital Geometry Linear (sp)

27 Orbital Geometry Trigonal Planar (sp 2 )

28 Orbital Geometry Tetrahedral (sp 3 )

29 Orbital Geometry

30 Linear and Trigonal Planar Molecules

31 Tetrahedral Molecules

32 Molecular Shape is Altered by Lone Pairs Two different “bent” structures which differ in their bond angle: 120° vs ° Lone pairs take up more room and squeeze bond angles together: o Trigonal planar: 120° o Trigonal planar with lone pair: < 120° o Tetrahedral: 109.5° o Tetrahedral with lone pair: < 109.5°

33 Orbital Geometry and Molecular Shape REMEMBER: 1)Electrons repel each other 2)Electrons in bonds have a fixed location 3)Electrons in lone pair take up more space 4)Two atoms always define a line.

34 Molecular Shape Let’s look at two seemingly similar molecules: CO 2 and H 2 O We first have to draw the molecule to determine their shapes:

35 Molecular Shape We can base molecular shape on orbital geometry: Linear: CO 2 Trigonal Planar: BF 3, CH 2 O Tetrahedral: CH 4, NH 3, H 2 O

36 Learning Check Determine the shape of the N 2 O molecule:

37 Electron Dot Formulas Practice drawing electron dot structures and identify the proper molecular shape for: AlH 3 NF 3 SiCl 4 HCN CO

38 In the preceding structures we had sharing of electrons (covalent bonds) and charged species (transfer of electrons, ionic bonds): How do we know what bond type we have? General Rule of Thumb: o A metal with a non-metal forms an ionic bond o A non-metal binding with a non-metal forms a covalent bond WHY???? Electronegativity

39 The type of bond interaction is determined by differences in electronegativity: Electronegativity is the ability of an atom to attract shared electrons in a bond towards itself Time to revisit Periodic Trends…

40 Electronegativity Increasing Electronegativity Fluorine is the MOST electronegative Cs/Fr are the LEAST electronegative

41 Electronegativity Values We can assign values for electronegativity:

42 Place the following in order of increasing electronegativity: O, K, and C Learning Check

43 Determining Bond Type Difference in Electronegativity (ΔEN) allows us to determine the type of bond formed: 1.Difference > 1.8 → ionic bond 2.Difference = 0.0 → pure covalent bond 3.Difference between 0.0 and 0.4 Non-polar covalent 4.Difference between 0.4 and 1.8 Polar covalent bond

44 Common Pure Covalent Bond Species O 2 Cl 2 Br 2 I 2 H2N2H2N2 S8P4S8P4

45 Bond Polarity Bonds which have an unequal sharing of electrons are said to be Polar: The electron cloud is pulled toward the more electronegative atom This sets up a separation of charges: dipole The more electronegative element attracts the cloud more strongly - partially negative The less electronegative element attracts the cloud less - partially positive

46 Bond Polarity

47 Blue: low electron density Red: high electron density Green: non-polar Bond Polarity

48 A different way of looking at polarity:

49 EXAMPLES Molecule NH 3 O2O2 NaCl SO 2 Δ EN and Bond Type 0.9 → polar covalent 0 → pure covalent 2.3 → ionic 1.0 → polar covalent

50 Electronegativity and Bond Types

51 Predicting Bond Types

52 Use electronegativity differences to classify each of the following bonds as nonpolar covalent (NP), polar covalent (P), or ionic (I): A bond between: 1) K and N 2) N and O 3) Cl and Cl 4) H and Cl Learning Check

53 Drawing Bond Dipoles K – N N – O Cl – Cl

54 Molecular Shape and Polarity Dipoles can add across a molecule Polar molecules – have an overall dipole Non-polar molecules – no overall dipole To determine whether a molecule is polar or non-polar, we must determine the 3-D shape of the molecule

55 Molecular Polarity Additive dipoles across a molecule make it polar: If dipoles cancel, no net dipole and the molecule is non-polar:

56 Molecular Polarity EXAMPLES: CO 2 NH 3 CF 4

57 Learning Check Identify each of the following molecules as (P) polar or (NP) nonpolar: A.PBr 3 B.HBr C.Br 2 D. SiBr 4

58 Learning Check Draw the Lewis Structure for SeCl 2 and determine its shape and molecular polarity:

59 Forces in Matter Intermolecular Forces: are attractive forces between molecules: Weaker than bonds (covalent or ionic) Help to determine the state of matter: solid, liquid or gas Intermolecular forces organize matter and are opposed by motion of molecules which disorganize matter

60 States of Matter Solid Attractive Forces >> Disruptive Forces Liquid Attractive Forces ≈ Disruptive Forces Gas Attractive Forces << Disruptive Forces

61 Intermolecular Forces Three primary types of forces that help to hold covalent molecules together: 1.London Dispersion Forces 2.Dipole – Dipole Interactions 3.Hydrogen Bonding

62 Intermolecular Forces London Dispersion Forces occur in non-polar molecules by forming a temporary dipole: Temporary dipoles induce dipoles in nearby molecules Lasts for milliseconds! ALL matter has London Dispersion Forces Let’s consider H 2 (non-polar molecule, H–H)

63

64 Intermolecular Forces London Dispersion Forces Increase with increasing size Increase with increased number of electrons Examples: CH 4, CF 4, CCl 4, CBr 4 CBr 4 will have the strongest London dispersion forces as it is the largest CH 4 will have the weakest London dispersion forces

65 Intermolecular Forces London Dispersion Forces Increase with increasing size Increase with increased number of electrons o Electrons on larger atoms are further from the nucleus therefore easier to induce temporary dipole Examples: CH 4, CF 4, CCl 4, CBr 4 CBr 4 will have the strongest London dispersion forces as electrons on Br are further from nucleus

66 Intermolecular Forces Dipole-Dipole Interactions occur between polar molecules where δ + end attracts the δ - end of another molecule There is a permanent dipole in polar molecules: attraction of molecules Occur only between polar molecules Consider FCl: Fluorine is more electroneagtive (δ - )

67

68 Intermolecular Forces Dipole-Dipole Interactions Stronger than London dispersion forces Strength of interaction increases with increasing size of the permanent dipole For Example: FCl (ΔEN= 1.0), FBr (ΔEN= 1.2), FI (ΔEN= 1.5) FI will have the strongest dipole–dipole interaction FCl will have the weakest dipole–dipole forces

69 Intermolecular Forces Hydrogen bonds are a special type of dipole- dipole force when we have a VERY large difference in electronegativity To have a Hydrogen Bond: 1.Need to have H 2.Need to have F, O, or N 3.H must be bonded to F, O, or N 4.At least one lone pair on the F, O, or N

70 Intermolecular Forces Hydrogen bond: The hydrogen on one molecule is attracted to the lone pair on another molecule The Hydrogen bond is the STRONGEST intermolecular force! Let’s consider H 2 O!

71 Hydrogen Bonds in Water

72 Hydrogen Bonds

73 Predicting strength of H-bonds: 1)CH 3 OCH 3 vs. C 2 H 5 OH 2)CH 3 NH 2 vs. CH 3 NHCH 3 vs. (CH 3 ) 3 N

74 Learning Check Identify the major type of attractive force in each of the following substances: A. NCl 3 B. H 2 O C. Br 2 D. KCl E. NH 3

75 Equilibrium When the forces acting on matter are in balance, we have Equilibrium Or, Equilibrium occurs when two or more opposing forces balance each other

76 Equilibrium: Balance of Forces Equilibrium occurs when two or more opposing forces balance each other, and is a dynamic process: Represented by the symbol: ⇄ For Example (Changes of State): Solid ⇄ LiquidLiquid ⇄ Gas Melting point Boiling point

77 Changes of State

78 Intermolecular Forces Strength of intermolecular forces increase as previously described: London < Dipole-Dipole < Hydrogen Bond As the strength of intermolecular forces increases, the properties of molecules change: Want to “stick together” more! Polar molecules tend to be liquids or solids at room temperature Requires increased kinetic energy (disruptive forces) to separate them!

79 Intermolecular Forces Strength of intermolecular forces increase as previously described: London < Dipole-Dipole < Hydrogen Bond As the strength of intermolecular forces increases, they change the properties of molecules: o Melting Point Increases o Boiling Point Increases o Vapor Pressure Decreases

80 Learning Check Identify the compound in each pair that has the higher melting point. Explain your choice: A.NCl 3 or NH 3 B.HBr or Br 2 C. KCl or HCl

81 Balance of Forces Let’s consider H 2 O: Hydrogen bonds are the strongest force Requires lots of energy to break apart intermolecular interactions Consequently, water has an increased melting point, decreased vapor pressure, and increased boiling point

82 Liquid Water


Download ppt "Chapter 10: Structure of Solids and Liquids Chem 1110 Figures: Basic Chemistry 3 rd Ed., Timberlake and Timberlake."

Similar presentations


Ads by Google