Unit 7 Intermolecular Forces IMFs EnergeticsThermodynamics Chapter 10

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Unit 7: Intermolecular Forces (IMFs), Energetics/Thermodynamics Chapter 10 & 11

Unit 7: Intermolecular Forces (IMFs), Energetics/Thermodynamics Chapter 10 & 11

Key Concepts • Intermolecular Forces (IMFs) • Hydrogen Bonding, Dipole-Dipole Forces, London Dispersion Forces

Key Concepts • Intermolecular Forces (IMFs) • Hydrogen Bonding, Dipole-Dipole Forces, London Dispersion Forces • Melting Point, Boiling Point, Surface Tension • Energy, Potential Energy/Kinetic Energy • Temperature/Heat • q=mcΔT, Specific Heat Capacity • Heating Curve: Heat of Fusion, Heat of Vaporization • Calorimetry • Endothermic, Exothermic • Enthalpy (ΔH)

Molecular Forces • Intramolecular Forces: strong bonds that hold atoms together in a molecule

Molecular Forces • Intramolecular Forces: strong bonds that hold atoms together in a molecule – Ionic bonds – Covalent bonds – Metallic bonds • Intermolecular Forces: weak forces between molecules – Hydrogen bonds – Dipole-Dipole forces – Dispersion forces (strongest) (medium) (weakest)

Molecular Forces

Molecular Forces

Hydrogen Bonding (Strongest) • H-bonding: only occurs between… O-H N-H O, N, or F

Hydrogen Bonding (Strongest) • H-bonding: only occurs between… O-H N-H O, N, or F F-H }

Hydrogen Bonding (Strongest) – Occurs between liquid molecules like H 2 O – Also

Hydrogen Bonding (Strongest) – Occurs between liquid molecules like H 2 O – Also occurs in DNA between base pairs A-T & C-G

Dipole-Dipole Forces (Medium) • Dipole-Dipole Forces: occurs between polar molecules only

Dipole-Dipole Forces (Medium) • Dipole-Dipole Forces: occurs between polar molecules only

Dispersion Forces (Weakest) • Dispersion Forces: occurs between all molecules – temporary dipoles form

Dispersion Forces (Weakest) • Dispersion Forces: occurs between all molecules – temporary dipoles form when e- are all on the same side of a molecule at the same time

Dispersion Forces (Weakest) • Dispersion Forces are stronger when… 1. molecule has more total

Dispersion Forces (Weakest) • Dispersion Forces are stronger when… 1. molecule has more total e 2. molecule is more spread out vs. Which has stronger dispersion forces, Ne or CH 4?

Intermolecular Forces (IMFs) • Ex. name the IMFs present in the following molecules –

Intermolecular Forces (IMFs) • Ex. name the IMFs present in the following molecules – H 2 – HCl – NH 3 Hint: Draw the structure! Is it polar? Does it have H bonded to O, N, or F?

Intermolecular Forces (IMFs) • IMFs hold molecules together • The stronger the IMFs, the

Intermolecular Forces (IMFs) • IMFs hold molecules together • The stronger the IMFs, the harder it is to separate molecules!

Intermolecular Forces (IMFs) 100% IMFs 80% IMFs

Intermolecular Forces (IMFs) 100% IMFs 80% IMFs

Intermolecular Forces (IMFs) • Stronger IMFs = – higher Melting Point – higher Boiling

Intermolecular Forces (IMFs) • Stronger IMFs = – higher Melting Point – higher Boiling Point – higher Surface Tension Melting Point Boiling Point

Intermolecular Forces (IMFs) • Ex. place the following molecules in order of increasing boiling

Intermolecular Forces (IMFs) • Ex. place the following molecules in order of increasing boiling point – Methanol: CH 3 OH – Methane: CH 4 – Ethanol: C 2 H 5 OH – Hydrogen sulfide: H 2 S Hint: draw the structure, list the IMFs… Increasing IMFs = Increasing BP

Intermolecular Forces (IMFs) Methane Dispersion Hydrogen sulfide Dispersion Dipole-Dipole Methanol Dispersion Dipole-Dipole H-Bond Ethanol

Intermolecular Forces (IMFs) Methane Dispersion Hydrogen sulfide Dispersion Dipole-Dipole Methanol Dispersion Dipole-Dipole H-Bond Ethanol Dispersion Dipole-Dipole H-Bond Same IMFs… Which one has stronger IMFs? Why?

Energy • Energy: the capacity to do work such as move an object •

Energy • Energy: the capacity to do work such as move an object • Energy is involved when there is a change in phase Energy is absorbed and used to break IMFs HIGH ENERGY LOW ENERGY Energy is released when IMFs are formed

Energy • Potential Energy: energy of position/stored energy • Kinetic Energy: energy of motion

Energy • Potential Energy: energy of position/stored energy • Kinetic Energy: energy of motion

Temperature vs. Heat • Temperature: average kinetic energy of particles in motion • Heat:

Temperature vs. Heat • Temperature: average kinetic energy of particles in motion • Heat: energy transferred between objects *Heat flows from HOT systems to COLD systems Temperature: kinetic energy of water molecules Heat: energy transferred from stove to pot

Temperature vs. Heat • You are COLD! You have a cup of hot chocolate

Temperature vs. Heat • You are COLD! You have a cup of hot chocolate and a hot tub full of water at the same temp. Which will warm you up faster? or 60°C

q=mcΔT • Used to calculate the heat that is needed to change the temp

q=mcΔT • Used to calculate the heat that is needed to change the temp of a substance q = m c ΔT mass (g) heat transferred (J) specific heat capacity (J/g°C) change in temp Tfinal-Tinitial (°C)

Specific Heat (c) • Specific Heat: the energy required to raise the temperature of

Specific Heat (c) • Specific Heat: the energy required to raise the temperature of one gram of a substance by one degree Celsius • Every substance has its own specific heat! Ex. the specific heat of water = 4. 18 J/g°C

q=mcΔT • Ex. how much does it take to raise 50 g of water

q=mcΔT • Ex. how much does it take to raise 50 g of water 20°C? • Ex. if 8, 750 J are transferred to 250 g of water, how many degrees will the temperature go up?

ga s Heating Curve l g id liq u Temperature 100°C s l lid.

ga s Heating Curve l g id liq u Temperature 100°C s l lid. ICE o s Time

ga s Heating Curve l g id liq u Temperature 100°C s l 0°C

ga s Heating Curve l g id liq u Temperature 100°C s l 0°C id l so Time

ga s Heating Curve l g id liq u Temperature 100°C q=mcΔT s l

ga s Heating Curve l g id liq u Temperature 100°C q=mcΔT s l 0°C id l so q=mcΔT Time q=mcΔT

Heating Curve • q=mcΔT is used when temperature changes, but it can’t be used

Heating Curve • q=mcΔT is used when temperature changes, but it can’t be used during phase changes because there’s no temperature change • What equation can we use during phase changes?

ga s Heating Curve for Water l g id ΔHvap liq u Temperature 100°C

ga s Heating Curve for Water l g id ΔHvap liq u Temperature 100°C q=mcΔT s l 0°C id l so ΔHfus q=mcΔT Time q=mcΔT

ΔHfus & ΔHvap • Heat of Fusion: the energy required to change between a

ΔHfus & ΔHvap • Heat of Fusion: the energy required to change between a liquid and a solid • Heat of Vaporization: the energy required to change between a liquid and a gas

HOT Heat of Vaporization Heat of Fusion

HOT Heat of Vaporization Heat of Fusion

Heating Curve for Water Heat of Fusion Heat of Vaporization (l - g) (s

Heating Curve for Water Heat of Fusion Heat of Vaporization (l - g) (s - l) Equation q=nΔHf Value ΔHfus= 6. 01 KJ/mol q=m. Lf q=nΔHv q=m. Lv Lf= 334 J/g ΔHvap= 40. 7 KJ/mol Lv= 2, 260 J/g Note: this table will be provided on the test.

ΔHfus & ΔHvap • Ex. how much energy would be required to melt 200

ΔHfus & ΔHvap • Ex. how much energy would be required to melt 200 grams of ice? • Ex. how much heat is required to boil 3. 2 moles of water?

Combined Heat Calculations • Use q=mcΔT during temp change • Use ΔHfus or ΔHvap

Combined Heat Calculations • Use q=mcΔT during temp change • Use ΔHfus or ΔHvap during phase change

Combined Heat Calculations • Ex. If you have 100 g of liquid water at

Combined Heat Calculations • Ex. If you have 100 g of liquid water at 50°C and you bring it up to 100°C and boil it all away, how much energy does this require? • Ex. how much energy is needed to make a 200 g block of ice at 0°C completely boil away?

Calorimetry • Calorimetry: measurement of q (heat/energy transferred) during a reaction

Calorimetry • Calorimetry: measurement of q (heat/energy transferred) during a reaction

System/Surroundings • System: whatever we’re focusing on • Surroundings: everything besides the system These

System/Surroundings • System: whatever we’re focusing on • Surroundings: everything besides the system These terms are just a frame of reference so we know what to focus on!!

System/Surroundings • Heat of the System = Heat of the Surroundings • Ex. Heat

System/Surroundings • Heat of the System = Heat of the Surroundings • Ex. Heat is transferred from the hot metal to the cold water. If the hot metal loses 100 J of heat, then the cold water gains 100 J of heat. -qsystem = qsurroundings HOT metal (-100 J = 100 J) COLD water

Endothermic/Exothermic • Exothermic: heat is released, q is negative • Endothermic: heat is absorbed,

Endothermic/Exothermic • Exothermic: heat is released, q is negative • Endothermic: heat is absorbed, q is positive

Endothermic/Exothermic • Exothermic: heat is released, q is negative A+B C+D+ energy • Endothermic:

Endothermic/Exothermic • Exothermic: heat is released, q is negative A+B C+D+ energy • Endothermic: heat is absorbed, q is positive A+B+ energy C+D

Enthalpy (ΔH) • Enthalpy (ΔH): heat absorbed or released per mol ΔH = q/mol

Enthalpy (ΔH) • Enthalpy (ΔH): heat absorbed or released per mol ΔH = q/mol • ΔH is the energy difference between reactants and products (products - reactants) Exothermic ΔH = -energy Endothermic ΔH = +energy

Lab Based Calorimetry Problem • Jamie wondered how much energy was in her kukui

Lab Based Calorimetry Problem • Jamie wondered how much energy was in her kukui nut. She set up a can of water above the nut and lit it on fire. – What is the system? – What are the surroundings? Your goal is to calculate ΔH = q/mol