Chemical Thermodynamics States of Matter, Phase Changes, Heat, State Functions, Specific Heats and Latent Heats

Standards 7. Energy is exchanged or transformed in all chemical reactions and physical changes of matter. As a basis for understanding this concept: a. Students know how to describe temperature and heat flow in terms of the motion of molecules (or atoms). 7. b. Students know chemical processes can either release (exothermic) or absorb (endothermic) thermal energy. 7. c. Students know energy is released when a material condenses or freezes and is absorbed when a material evaporates or melts. 7. d. Students know how to solve problems involving heat flow and temperature changes, using known values of specific heat and latent heat of phase change. 7. e.* Students know how to apply Hess’s law to calculate enthalpy change in a reaction. aA + bB = c C + d D ΔHor = [c Hof (C) + dHof (D)] − [a Hof (A) + bHof (B)] 7. f.* Students know how to use the Gibbs free energy equation to determine whether a reaction would be spontaneous. ΔG = Δ H − ΤΔS

States of Matter Plasma aka phases Gas Liquid Solid Deionization Condensation Boiling Liquid Sublimation Deposition Freezing Melting Solid

Condensation (Gas  Liquid)

Boiling (Liquid  Gas)

Sublimation (Solid  Gas)

Deposition (Gas  Solid)

Freezing (Liquid  Solid)

Melting (Solid  Liquid)

Molecular Motion The state of matter depends on how much energy (motion) the molecules, atoms, or ions have. The state of matter also depends on how attracted the atoms, molecules, or ions are to each other.

Molecular Motion of Gases

Molecular Motion State of Matter O Gas O H + – Liquid Cl – Na + Solid

Nonpolar molecules O O

Polar molecules O H + – O H + – O H + –

Cl – Cl – Na + Na + Cl – Cl – Na + Na + Cl – Cl – Na + Na + Ionic compounds Cl – Cl – Na + Na + Cl – Cl – Na + Na + Cl – Cl – Na + Na +

Temperature Scales Water Boils Human Body Water Freezes 212°F 98.7°F 373K 310K 273K

Temperature Scales Surface of Sun Room Temp. Absolute Zero 9,941°F 5,778K 294K 0K

Converting Temperatures Fahrenheit  Celsius °C = (°F – 32)×(5/9) Celsius  Fahrenheit °F = °C ×(9/5) + 32 Celsius  Kelvin K = °C + 273.15

Absolute Zero At Zero Kelvin (0 K or –273.15 °C), atoms and molecules stop moving. There is no temperature lower than absolute zero (0 K).

Heat – it’s what makes things Hot Heat is the energy that raises a substance’s temperature. Heat is transferred from one place to another. Property Symbol Unit Unit Symbol Heat Q calories cal Joules J 1 calorie = 4.186 Joules

Heating Water Temperature (°C) Heat Added, Q, (Joules) 100 steam warming water boiling water warming ice melting ice warming

Specific Heat (aka Heat Capacity) When you heat a substance at a constant rate, its temperature usually increases at a constant rate also. Q = m·Cp·(ΔT) heat = mass × specific heat change in temperature × J = g × °C × J g °C Cool Substance + Warmer Substance heat

ΔT = Tfinal – Tinitial Change in Temperature ΔT = 15˚C – 25˚C Example: A substance starts with a temperature of 25˚C and ends with a temperature of 15˚C. What is the change in temperature? ΔT = 15˚C – 25˚C ΔT = -10˚C

Specific Heat Example #1 What is the amount of heat needed to increase the temperature of 4 grams of ethanol by +6˚C? The specific heat of ethanol is 2.44 J/(g ·˚C). ΔT m Cp Q = m · Cp · ΔT Q = (4) · (2.44) · (+6) Q = 58.6 Joules

Specific Heat Example #2 Q m 45 Joules of heat are removed from a 0.5 g sheet of aluminum foil that was initially at 120˚C. What will the final temperature of the foil be? The specific heat of aluminum is 0.90 J/(g ·˚C). Tinitial Cp Q = m · Cp · ΔT ΔT = Tfinal – Tinitial -45 = (0.5) · (0.90) · ΔT -45 = 0.45 · ΔT -100 = Tfinal – 120 0.45 0.45 +120 +120 ΔT = -100˚C Tfinal = 20˚C

Enthalpy ΔH = enthalpy ΔH = + ΔH = – Enthalpy – A measure of the total internal heat of a substance per mole. ΔH = enthalpy Positive Change in Enthalpy – the process is absorbing heat. ΔH = + Negative Change in Enthalpy – the process is releasing heat. ΔH = –

Specific Types of Enthalpy Changes Heat of Fusion – the heat per mole required to melt a solid. ΔHfus Heat of a Reaction – the heat per mole released or absorbed by a reaction. ΔHr= ΔHproducts – ΔHreactants

Exo/Endothermic Reactions exothermic reaction – the reaction releases heat. (includes freezing, condensation, combustion) ΔHr = – Reactants Products + heat endothermic reaction – the reaction absorbs heat. (includes melting, boiling, charging up a battery) Reactants + Products heat Δ Hr = +

N2 (g) + 3 H2 (g) 2 NH3 (g) + 92kJ nitrogen hydrogen ammonia heat gas gas gas + +

Q = m·ΔHfus Q = m·ΔHvap Latent Heat means hidden Latent Heat – the heat required to cause a phase-change (hidden because this heat does not change the temperature). Latent Heat of Fusion – the heat required to melt a solid. Q = m·ΔHfus Latent Heat of Vaporization – the heat required to boil a liquid. Q = m·ΔHvap

Heat of Solidification Latent Heats Gas / Vapor Heat of Vaporization ΔHvap Heat of Condensation ΔHcond = – ΔHvap Liquid Heat of Fusion ΔHfus Heat of Solidification ΔHsolid = – ΔHfus Solid

Heating Water Temperature (°C) Heat Added, Q, (Joules) 100 Q = m·(ΔT)·Cp steam steam warming Q = m·ΔHvap water boiling Q = m·(ΔT)·Cp water water warming Q = m·ΔHfus ice melting Q = m·Cp ice·(ΔT) ice warming

Using Heats of Formation Heat of Formation (ΔH0f) – A measure of the internal heat of a substance relative to other substances. compared to As a reference point, the heat of formation of an element (ex. Fe, O2, H2) is 0 kJ/mol. ΔHr= ΔH0f products – ΔH0f reactants

ΔHr= - (B.E. products – B.E. reactants) Using Bond Energies Bond Energy (B.E.) – A measure of energy that is locked away in covalent bonds. There are different bond energies depending on the types of atoms and if it is a single bond, double bond, etc. ΔHr= - (B.E. products – B.E. reactants) because we are breaking bonds

Lewis Dot Structures O H C O C O N O O C H

Entropy ΔS = entropy ΔS = + ΔS = – Entropy – A measure of the disorder of a substance. ΔS = entropy Positive Change in Entropy – the substances are becoming more disordered. ΔS = + Negative Change in Entropy – the substances are becoming less disordered. ΔS = – Nature tends to favor a continual increase in entropy unless outside energy is used. (ex. Cleaning a room, photosynthesis, distillation)

Which has the greater entropy? Low Entropy High Entropy

Which has the greater entropy? High Entropy Low Entropy

Which has the greater entropy? High Entropy Low Entropy

Which has the greater entropy? High Entropy Low Entropy

Which has the greater entropy? Low Entropy High Entropy

Gibbs (Free) Energy ΔG = Gibbs energy ΔG = ΔH – ΔS × T ΔG = – ΔG = + Gibbs Energy – A measure of both the enthalpy and entropy of a substance combined. ΔG = Gibbs energy ΔG = ΔH – ΔS × T Negative Change in Gibbs energy – a process can happen. The process is spontaneous, but the process might still be incredibly slow. ΔG = – Positive Change in Gibbs energy – a process will not happen. The process is nonspontaneous. ΔG = +

State Functions Enthalpy (ΔH), Entropy (ΔS), and Gibbs Energy (ΔG) are all state functions. They are called state functions because it doesn’t matter how they got there, it only matters where they started and where they finished.

State Functions X X

X X State Functions Water at 100˚C Ice at -30˚C Water at 20˚C

State Functions Summary Property Symbol Meaning Units Sign change Enthalpy ΔH heat + = endothermic – = exothermic Entropy ΔS disorder + = increase in disorder – = decrease in disorder Gibbs energy ΔG does it happen? + = nonspontaneous – = spontaneous 0 = at equilibrium kJ mol kJ mol·K kJ mol

Exo/Endothermic Reactions exothermic reaction – the reaction releases heat. endothermic reaction – the reaction absorbs heat. H2O (g) + 3 H2 (g) 2 NH3 (g) + 92kJ Solid hydrogen ammonia heat gas gas gas

He H B C N O Ne F Li Be P Al Si S Cl Ar Na Mg Br Kr K Ca I Xe

He H B C N O Ne F Li Be P Al Si S Cl Ar Na Mg Br Kr K Ca I Xe

4 e– in valence shell