Chemistry: Energetics
This flashcard set introduces the concept of enthalpy, symbolized as H, and explains how it is measured through temperature changes during reactions. It covers the meaning of ΔH, its units (kJ/mol), and describes heat flow in exothermic reactions where energy is released to the surroundings.
what is enthalpy
the heat energy that is stored in a chemical system
Key Terms
what is enthalpy
the heat energy that is stored in a chemical system
how is enthalpy shown
H
how is enthalpy measured
from temperature changes when a chemical reaction takes place
delta H
heat energy change at constant pressure
what are the units of heat energy change (delta H)
kJ per mole
what happens in terms of heat in an exothermic reaction
heat energy is given out to the surroundings
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| Term | Definition |
|---|---|
what is enthalpy | the heat energy that is stored in a chemical system |
how is enthalpy shown | H |
how is enthalpy measured | from temperature changes when a chemical reaction takes place |
delta H | heat energy change at constant pressure |
what are the units of heat energy change (delta H) | kJ per mole |
what happens in terms of heat in an exothermic reaction | heat energy is given out to the surroundings |
what happens to the temperature of the surroundings when an exothermic reaction taxes place | temperature of surroundings increases |
why is delta H negative in exothermic reactions | the chemicals lose heat energy |
important exothermic reactions | combustion of fuels and respiration |
what happens in terms of heat in endothermic reactions | heat energy is taken in from the surroundings |
what happens to the temperature of the surroundings in endothermic reactions | the temperature of the surroundings decreases |
why is delta H positive in endothermic reactions | the chemicals gain heat energy |
what do endothermic reactions require | the input of heat energy |
important endothermic reactions | thermal decomposition of calcium carbonate photosynthesis |
what can enthalpy profile diagrams be used to illustrate | the enthalpy change for a reaction |
what do enthalpy profile diagrams show | the products, reactants and enthalpy change |
activation energy | the minimum energy required to start a reaction by the breaking of bonds |
do the products or reactants have a higher enthalpy in exothermic reactions | reactants |
do the products or reactants have a higher enthalpy in endothermic reactions | products |
what is required to start the exothermic reaction | energy is required to break the bonds, even though the products have a lower energy than the reactants |
what will provide the energy needed to continue to overcome the activation energy | the net energy, once the barrier has been overcome |
| the break the bonds and start the reaction |
what do most endothermic reactions need to provide the necessary energy | to be heated continuously |
what is bond enthalpy | the energy required to break one mole of a given covalent bond in the molecules in the gaseous state |
why are bond enthalpy values always positive | bond breaking is endothermic |
What do bond enthalpies give an indication of | The relative strength of a covalent bond |
How do bond enthalpies give an indication of the relative strength of a covalent bond | The stronger the bond, the more endothermic the bond enthalpy |
Why do chemists use mean bond enthalpy values | The same covalent bond may appear in different compounds and the value of the bond will be slightly different in each compound |
Mean bond enthalpy | The energy required to break a covalent bond, averaged for that type of bond in a range of different compounds |
What happens in terms of the bonds in a chemical reaction | Bonds in the reactants are broken and new bonds are formed to make the products |
What is required to break bonds | Energy |
What is released when new bonds form | Energy |
What is the enthalpy change of a reaction | Sum of bonds broken - some of bonds formed |
Limitations of bond energy calculations | Using mean bond enthalpies to calculate the enthalpy change of a reaction often leads to a value that is less accurate than a value obtained from Hess's Law |
Why is using mean bond enthalpies to calculate enthalpy change less accurate than a value obtained from Hess's law | Bond enthalpies used are only an average value- not specific to compound in question We are assuming all species are in the gaseous state- lots of compounds won't be eg. Ethanol |
What are the standard conditions for standard enthalpy of formation | 298k and 100kPa |
Standard state | The physical state of a substance under standard conditions |
Standard enthalpy of formation | The enthalpy change that occurs when one mole of a compound is formed from its constituent elements with all reactants and products in their standard states |
What is the standard enthalpy of formation of an element in its standard state | 0 |
What do all products of standard enthalpy of formation need to be | 1 mole |
What is the relationship between negativity of standard enthalpy of formation of a compound and the stability of it | More negative = more stable |
Standard enthalpy of combustions | The enthalpy change that occurs when one mole of a compound reacts completely in oxygen with all reactants and products in their standard states |
Calorimetry | The method used to determine enthalpy changes by experiment |
What does calorimetry involve | Measuring the temperature change of a given amount of water as the reaction occurs- and converting this to a quantity of heat energy |
Specific heat capacity of water | 4.18 J required to hear 1g of water |
What does heat change = | Mass of water x specific heat capacity x temperature change Q=mc DELTA T |
What is q | Heat energy released or absorbed in joules |
What is m | Mass of water in grams |
What is the mass of water in grams equal to | Volume in cm3 |
What is c | The specific heat capacity of water |
What is delta t | Temperature change in K | Same as change in celcius |
What must you do to q when you've initially calculated it using mcDELTA T | Divide it by 1000 to convert J to kJ |
Heat given out or taken in by one mole= | Q(in kJ)/ no of moles |
What are the units of delta H for experimental determination of enthalpy changes | kJmol-1 |
When is delta H negative in calorimetry | If temp increases during reaction (exothermic reaction) |
When is delta H positive during calorimetry | If temperature decreases during a reaction (endothermic) |
Why may the experimental value for enthalpy of combustion be different to the data book value |
|
What can cooling curves be used for | As a method of accounting for heat loss with reactions in solutions |
How to use data to find maximum temperature rise | Plot a graph of temperature against time using results and determine max temperature change accompanying reaction Extrapolate back to where lines meet (ish) to establish max temperature rise |
First law of thermodynamics | Energy cannot be created or destroyed but it can be changed from one form to another |
Where does hess's law apply the first law of thermodynamics | To chemical reactions |
Hess's law | The enthalpy change for a chemical reaction is independent of the route taken |
Why is it not always possible to measure the enthalpy change of a reaction directly |
| - practically impossible to measure |
What is Hess's law used for | To calculate enthalpy changes which cannot be measured directly using an enthalpy cycle |
What can we do using Hess's law if we know two of the enthalpy changes | Calculate the third |
What does delta H1 equal | Delta H 2 + delta H 3 |
How to remember enthalpy cycles of combustion | C=R-P |
Where should you always do the vectors for enthalpy of combustion | Products to element |
Where should you do vectors in calculations for reactions from enthalpy of formation values | Elements to products |
How to remember enthalpy cycles of formation | F=P-R |