Honors Chemistry Dr. Kevin D. Moore

Transcription

1 Honors Chemistry Dr. Kevin D. Moore

2 Key Properties: Solid is less dense than liquid Water reaches maximum density at 4 C Very high specific heat Dissolves many substances Normal Boiling Point: 100 C Normal Melting Point: 0 C Uses Medium for tons of chemical reactions Used to cool many different processes Sustains life Density of Water

3 Gases Sparse, Random Motion No intermolecular forces Liquid Dense, Restricted Motion Intermolecular forces Solid Dense, Highly-Restricted Motion Intermolecular forces

4 Intramolecular Forces: ionic & covalent bonds Polar Covalent Bond Partial Separation of Charge Electronegativity: H 2.1 Cl 3.0 H Cl δ + δ - Dipole Moment measure of the net polarity in a molecule (µ)

5 Consider H 2 O O H H Dipole Moment is the net of all bond dipoles N H H H

6 Molecules which have no dipole moment are Non-Polar Even if individual bonds have dipoles Cl Cl C Cl Cl No Net Dipole! Highly symmetric molecules tend to be non-polar!

7 Non-polar molecules have µ=0 Compound Dipole Moment (μ) NaCl 9.00 CH3Cl 1.87 H2O 1.85 NH CO2 0 CCl4 0

8 Draw Dipole Moments by estimating positive and negative centers of charge Show 3D structure Draw the Lewis Structure and Dipole Moment for Methanol (CH 3 OH) H H C O H H

9 Hydrocarbons are non-polar C 3 H 8 Symmetric molecules with no lone pairs on the central atom are non-polar Tetrahedral (CCl 4 ) Trigonal Planar (SO 3 ) Linear (CO 2 ) Different atoms connected to central atom is always polar CH 2 F 2 CH 3 COCH 3

10 Solid and Liquid Phase implies a binding force Often called van der Waals forces Electrical in nature Consider an ion interacting with a dipole O H H O H H

11 Weak Force of interaction between dipoles Only relevant when molecules are very close together O S O O S O O S O

12 Strength of interaction is a function of dipole moment Substance Molar Mass Dipole Moment (D) Boiling Point (K) CH3CH2CH CH3OCH CH3Cl CH3CN

13 Benzene (µ=0) Boiling Point = 80.1 C Melting Point = 5.5 C H C H C C H H C C C H H

14 Instantaneous asymmetric distribution of charge Causes a distortion in a molecule closeby Tiny attractive force Increases with the polarizability of electron How easily electron field can be distorted Increases with size (surface area) of molecule ALL substances experience LDF Substance Melting Point (K) Boiling Point (K) F Cl Br I

15 Dominant force of all non-polar compounds Holds molecules together in liquid and solid Large, heavy molecules tend to have higher LDF

16 Interaction between a hydrogen bonded to a high EN atom (N, O & F) and unshared electrons (lone pairs) on another high EN atom. N O H H H H H O H H N H H H

17 Stronger interaction than dipole-dipole Through space interaction Proteins Folding DNA Helix

18 Summary Force Strength Effect Example LDF Very Weak Depends on MW CCl4, C6H6, All Substances Dipole-Dipole Weak Solubility, BP, MP SO2, CH3COCH3 H-bonding Moderate BP, MP, 3D Structure H2O, NH3, Organic Acids, Organic Bases Ion-Dipole Strong Solubility

19 Draw Lewis Structure No Dipole? Yes LDF No Hydrogen Bonding? Yes Dipole-Dipole LDF H-Bonding LDF

20 CCl 4 Tetrahedral, very symmetric Non-Polar (No Dipole) LDF Only CH 3 COCH 3 (Acetone) Trigonal Planar (on central Carbon) Symmetric Dipole across C=O Dipole-Dipole, LDF CH 3 CH 2 OH (Ethanol) Tetrahedral (Bent at C-O-H) Asymmetric Hydrogen Bonding, LDF

21 Resistance of a liquid to spreading out Least Surface Area Sphere

22 Pressure of a liquid on a closed container Depends on Temperature All Liquids and many solids Dynamic Equilibrium between vapor and liquid # of molecules ΔH vap Kinetic Energy

23 Vapor Pressure increases with Temperature Boiling Point Temperature at which the vapor pressure is equal to atmospheric Pressure H 2 O: at 1 atm = 100 C

24 P(mmHg) vs. T(K) H 2 O Pvap (mmhg) T (K) 1/T ln P vap

25 Plot ln P vap vs. 1/T 7 ln P vs. 1/T

26 Passing of a liquid to a vapor without being at the boiling point Liquid cools as evaporation occurs

27 ΔH Enthalpy Heat gained/lost for a reaction or phase change Symbol Name Beginning Phase Ending Phase ΔH fus Fusion (Melting) Solid Liquid -ΔH fus Freezing Liquid Solid ΔH vap Vaporization Liquid Gas -ΔH vap Condensation Gas Liquid ΔH sub Sublimation Solid Gas -ΔH sub Deposition Gas Solid

28 Change of a solid directly to a vapor without forming a liquid Occurs when solid has a high vapor pressure CO 2 sublimes Ice sublimes at low pressure

29 Exothermic Gives off heat Negative value Endothermic Takes in heat Positive value

30 Calorie amount of heat (energy) needed to raise the temperature of 1 gram of water by 1ºC Joule SI Unit for Energy J = 1 calorie How many Joules are in 60.0 calories? cal J 3sig figs J cal J

31 Amount of Heat required to raise1 g of a substance by 1 ºC Substance Specific Heat (J/gºC) Water (liquid) Water (solid) 2.03 Aluminum Iron Copper Silver Gold 0.129

32 s q m T s: Specific Heat q: Heat Added or Removed m: mass of object ΔT: Temperature Change (T f T i )

33 How much heat (in Joules) is required to heat a 45.0 g Aluminum rod from 25.0 ºC to 52.0 ºC? (s Al = J/gºC) T... q s m T 0897 J g C 450 g 27 0 C 1090 J If 2.11 kj is added to a 100. g piece of Copper, what would be the final Temperature if it is initially 22.0ºC? (s Cu = J/gºC) q s m 2110 J JgC 100. g add to C 768. C

34 100. g of water is heated from 11.2 ºC to 29.8 ºC. How much heat was added? (s H2O =4.184 J/gºC) q s m T J g C( 100. g)( C) 7780 J kg block of ice is initially at 0.0 C. If 4.52 kj is removed from it, what is its final Temperature? (s ice = 2.03 J/g C) T T q s m 4520 J J gc 850 g 262. C C

35 To break apart substances Requires Energy Increases Disorder To organize substances Releases Energy Decreases Disorder Disorder Solid < Liquid < Gas Solids are highly ordered gases have no order

36 For a process to be spontaneous, it wants to: Give off energy Become more disordered An endothermic process can be spontaneous Disorder effect is greater than heat (energy) absorption Ice Heat More Energy More Disorder

37 How much heat does it take to melt 50.0 g of ice which has been warmed to 0 C? ΔH fus (H 2 O) = 6.02 kj/mol n ice 50. 0g 1 mol g 277. mol H2O q H n ice fus ice q ( )( mol) kj ice kj mol

38 How much heat is needed to vaporize 50.0 g of H 2 O at exactly 100. C (ΔH vap =40.6 kj/mol)? 1 mol nho g g q H n HO vap HO 2 2 mol q ( )( mol) 112 kj HO 2 kj mol 277. Vaporizing takes more heat than melting 112 kj vs 16.7 kj

39 Changing the direction of the phase change will change the sign of q + is heat absorbed - is heat released How much heat is released when 50.0 g of steam condenses to water at 100. C? -112 kj

40 Heating within a phase Molar Heat Capacity amount of heat required to raise 1 mol of a substance by 1 C Related to specific Heat Capacity Moles rather than grams The specific heat of water is J/g C, what is its molar heat capacity? C C molar mass C mol HO gram J g C g 1 mol J mol C

41 Heat from -10 C to 110 C Temperature ( C) 100 C ice = 36.5 Jmol -1 C -1 C H2O = 75.3 Jmol -1 C -1 C steam = 33.6 Jmol -1 C -1 ΔH vap = kjmol -1 ΔH fus = 6.01 kjmol Heat Added (kj)

42 #1 Warming the ice q 365. J ( 1 mol)( C) 365 J mol C #2 melting the ice kj q ( 1 mol) 6010 J #3 warming the water fus mol q J ( 1 mol)( 100. C) 7530 J mol C #4 vaporizing the water kj q ( 1 mol) J vap mol

43 #5 Heating of steam Total Heat J q 336. ( 1 mol)( C) 336 J steam mol C q q q q q q tot ice fus water vap steam q 365 J 6010 J 7530 J J 336 J tot q J kj tot

44 Atoms/Molecules in fixed positions which significantly restricts motion Only capable of vibrating Melting conversion of structured solid phase to unstructured liquid phase Occurs when vibrations become strong enough to dislodge atoms from the solid array

45 Crystalline Solids Ionic NaCl, KNO 3 Metallic Ag, Zn, Au Network Covalent (Atoms) Graphite, Diamond, Quartz Molecular Sugar Amorphous Rubber, Glass

46 Highly structured form of a solid 3D ordered atoms Many different crystal packing methods Ionic crystals have very high normal melting points NaCl 800 C Ions interact very strongly Molecular crystals have low normal melting points Molecules (no ions) interact very weakly

47 Atomic Ionic Molecular

48 Different forms of the same substance in the same physical state Carbon over 40 known allotropes Graphite Flat sheets of Carbon atoms Diamond 3D web-like connected atoms Buckminster Fullerene C 60 soccer ball structure Silicon Silicon Dioxide Sand, Quartz, Glass

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52 Special form of atomic crystal Malleable Ductile Strong bonding, but non-directional Electron Sea Model Metal atoms are in a sea of valence electrons Electrons easily move between atoms Conducts electricity & heat Atoms easily move around other atoms Ductile, Malleable, Easily scratched

53 Substance containing a mixture of elements and has metallic properties Stronger, harder, less dense (lighter), resist corrosion Substitutional Alloy Host metal atoms are replaced by a different metal Sterling Silver Ag Ag Ag Ag Cu Ag Ag Ag Ag Ag Ag Ag Ag

54 Interstitial Alloy Holes in the metal are packed with very small atoms Prevents atoms from moving across each other Hardens the metal Steel Carbon is in the holes between Iron atoms Carbon creates directional bonds (tetrahedral) Prevents Iron from moving Fe Fe Fe Fe Fe Fe Fe Fe C Fe Fe Fe Fe

55 Pewter Sn/Pb or Sn/Cu Always contains a small amount of Sb Brass Cu/Zn Sterling Silver Ag/Cu Bronze Cu/Sn Stainless Steel Fe/Cr/Ni

56 Solids which have no true structure Randomly arranged atoms or molecules Rubber Glass Plastic Easily shaped

57 Plot of phases for Pressure vs. Temperature Melting SuperCritical Fluid Solid Liquid Boiling Sublimation Gas

58 Triple Point Temp & Pressure at which all phases coexist Triple Point must be below 1 atm for liquid to exist under normal conditions CO 2 Critical Temperature Temp above which a gas cannot be liquified Critical Pressure Pressure above which a liquid cannot be vaporized Super Critical Fluid Slope of Melting Line

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61 Intermolecular Forces cause substances to become liquids and solids Phases change if necessary heat is put into or taken out of substance Vapor Pressure and Evaporation occur because a small portion of the compound has the vaporization energy Heating Curve shows the relationship between heat and Temperature Atoms, ions and molecules form crystals Phase Diagrams show the relationship between phases in a substance