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Definition
| a property of a system that depends on the initial state and the final state of the system |
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| Examples of State Functions |
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Definition
| Enthalpy (H), Entropy (S), Free Energy (G), Temperature, Pressure, Volume |
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Definition
| the amount of energy needed to raise 1 gram of water 1°C |
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| Calorie --> Joule Conversion Factor |
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Definition
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Definition
| the heat content of a substance or a system |
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Definition
| the heat energy absorbed or released during a chemical or physical process |
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Definition
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Definition
| ΔH = [sum of enthalpies of products]-[sum of enthalpies of reactants] |
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Definition
| chemical potential energy is converted to heat energy and released to the surroundings |
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Term
| Examples of Exothermic Reactions |
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Definition
| combustion; alkali or alkaline earth metal + H2O; Gas --> Liquid --> Solid |
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| Graph of Exothermic Reaction |
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Definition
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Definition
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Definition
| the system absorbs heat from surroundings and it's converted to chemical potential energy |
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| Examples of Endothermic Reactions |
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Definition
| photosynthesis, any exothermic reaction performed in the reverse direction |
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| Graph of Endothermic Reactions |
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Definition
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Term
| Standard Reference Point for Enthalpy Values |
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Definition
| Any free element in its most stable state at 25°C and 1atm has 0 enthalpy. |
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Term
| Three Characteristics of Enthalpy |
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Definition
1. It is an extensive porperty (mass dependent)
2. ΔH for a reaction depends on the physical state of the reactants and products
3. If a reaction is endothermic in the forward direction, it is always exothermic in the reverse direction. |
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| The Five Ways to Calculate ΔH |
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Definition
1. From tables of enthalpy values
2. Stoichiometry
3. Hess's Law
4. Calorimetry
5. Bond Energies |
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Term
| Enthalpies (heats) of Formation (Hf) |
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Definition
| The enthalpy change associated with the formation of 1 mol of a compound from its constituent elements. |
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Definition
| Fe2O3 + Al --> Al2O3 + Fe |
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Definition
| If a reaction is carried out in a series of steps, ΔH for the reaction will be equal to the sum of the enthalpy changes for the individual steps. |
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Definition
| Hess's Law provides a useful means of calculating energy changes that are difficult to measure directly. |
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Term
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Definition
1. Look at the overall reaction for which you want the enthalpy change. Decide how to rearrange the given equations.
2. If an equation is multiplies by a coefficient, the ΔH value is multiplied too.
3. If a reaction is reversed, the sign on ΔH is also reversed.
4. Check to be sure that everything cancels except what you want. |
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Definition
| measures the heat flow accompaning chemical and physical processes. |
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Term
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Definition
| device used to measure heat flow (page 182) |
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Term
| Equation for Calculating Heat Released |
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Definition
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| Secondary Equation About Heat Released (Involving ΔH) |
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Definition
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Term
| Specific Heat Capacity (cp) |
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Definition
| the amount of heat needed to raise the temperature of 1 gram of a substance by 1°C. It is intensive, so it is the same for any mass of an element. For water, it is 4.18 J/g.c or 1.00 cal/g.c |
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| Heat Capacity (Unofficial HC) |
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Definition
| the energy needed to raise the temperature of AN OBJECT by 1°C. Unit is J/C. |
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Term
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Definition
| the energy required to change the temperature of 1 mole of a substance by 1°C. The unit is J/mol.C |
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Definition
| the amount of energy required to break a given bond in 1 mol of a gasseous substance. Unit: kJ/mol |
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| Equation Used for Bond Energy Enthalpy Change |
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Definition
| ΔHrxn = [the sum of the enthalpies of bonds broken]-[the sum of the enthalpies of bonds formed]. |
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Term
| What is the first step to solving an enthalpy problem using bond energies? |
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Definition
| Draw the molecules involved and count the number of bonds. |
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Definition
| A process that occurs without outside intervention once the activation energy is provided. |
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Definition
| the measure of the disorder in a substance or system. |
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| Second Law of Thermodynamics |
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Definition
| For any spontaneous process, the disorder of the universe must increase. |
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Definition
| ΔS = [sum of the entropies of the products] - [sum of the entropies of the reactants]. |
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Definition
| Increased entropy, more disorder. |
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Definition
| Decreased entropy, less disorder. |
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Definition
| A system that cannot exchange matter or energy with the surroundings (+ΔS) |
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Term
| Why does the universe naturally go to a state of increased disorder? |
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Definition
| There is only one way to be ordered, billions of ways to be disordered. Probability. |
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Term
| What are the general trends in ΔS? |
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Definition
1. Solid --> Liquid --> Gas: +ΔS
2. Gas --> Liquid --> Solid: -ΔS
3. If temperature increases, +ΔS
4. When a gas expands, +ΔS
5. When a gas dissolves in liquid: -ΔS
6. When a solid dissolves in liquid: +ΔS
7. If the moles of gas in the reactants is less than the moles of gas in the product: +ΔS |
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| Third Law of Thermodynamics |
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Definition
| Any pure crystalline substance at absolute zero (0K = -273C) has an entropy of 0. |
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| Formula for Calculating ΔSsurr |
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Definition
| ΔSsurr = -(ΔHsys/Tsurr(K)) |
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Definition
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| Chart Comparing ΔHsys and ΔSsurr |
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Definition
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| Chart Comparing ΔSsys, ΔSsurr, ΔSuniv, and spontenaeity |
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Definition
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Term
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Definition
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Term
| What does a -ΔG tell you? |
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Definition
| it means the reaction is spontaneous. The process can do work. The magnitude of ΔG indicates the maximum amount of work that can be obtained from a process. (Think: a gallon of gas can only do so much work before it runs out) |
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Term
| What does a +ΔG tell you? |
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Definition
| It means the process is nonspontaneous. The surroundings must do work on the system to make it go. The magnitude of ΔG indicates the theoretical minimum of work needed to drive the process. |
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Term
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Definition
1. It is a state function.
2. The unit is kJ/mol
3. Any free element in its most stable state at 25°C and 1am has a free energy value of 0.
4. ΔG = [sum of the free energy of the products] - [sum of the free energy of the reactants]. |
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| Chart about ΔG and Temperature (*******) |
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Definition
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