In depth and simple to understand class notes taken over the course of 2 years in a highly accredited IB school for Chemistry HL, in regards to the topic of Energetics and thermochemistry
This includes diagrams, annotations, extensive notes and explanations and they are the notes that have solely...
- All definitions and formulas in booklet given by Dr.Raj
- System: Area of interest (eg. beaker and its contents)
- Can be open or closed (whether you trap the reaction or not)
- OPEN: exchange of energy and matter (eg. HCl + Mg→ MgCl2 + H2, releases or
absorbs heat and H2 which is matter)
- CLOSED: Only exchange of energy, no exchange of matter (a reaction with a
capping, heat as energy exchanged, no matter released)
- Surroundings: everything outside the system
- Enthalpy: Amount of energy stored in a chemical (eg. when you burn smt, the amount of
energy released)
- Can only measure changes in energy, enthalpy can not be measured, only the change
- Kj / mol
- Substances with lower enthalpy are more stable than those with higher enthalpy
- In exothermic reactions, enthalpy of product is lower than that of the reactants, temp
increases (Write energy released as negative)
- Delta H = enthalpy change
- In an exothermic reaction, Delta H is negative. Negative means energy released
- Products in exothermic reactions are more stable
- Activation energy (minimum energy need for chem reaction to take place) is small
-
- In endothermic reactions, enthalpy of product is higher than that of reactants, temp decreases
(write energy absorbed as positive)
- Delta H = enthalpy change
- In an endothermic reaction, Delta H is positive, means energy taken in
- Products in endothermic reactions are
less stable
- Activation energy is large
Ea = activation energy
, - In the exothermic reaction, we see on graph, first it absorbs energy, and then it releases
energy. Bonds must first be broken and then new bonds are formed
- To break bond, energy is needed, to make bond energy is released
- Eg. Reactant= 12 Kj then it goes up to 14, then drops to 6
- In endothermic stronger bonds are made and weaker bonds are formed. It absorbs energy, eg
from 6 to 14 and then down to 12. Exact opposite of exothermic reactions
Temperature is the measure of the average kinetic energy of the particles
- Higher the kinetic energy, the higher the temperature
- Not all the particles have the same kinetic energy
- Temp depends on: mass, amount of heat added, specific heat capacity
Specific heat capacity: The amount of energy required to raise the temperature of gram by one kelvin
Unit: JK^-1 g ^-1
Substance Specific Heat
Capacity
J K-1 g-1
Water 4.18
Ethanol 2.44
Air 1.00
Iron 0.450
Copper 0.385
SATP: In data booklet formula
- 298k or 25°C
- Pressure: 100 kPa (1atm)
- Concentration: 1 mol dm^-3
Different enthalpy
- Enthalpy of combustion (delta H small c)
- Enthalpy change when one mole of substance is burned in excess oxygen under
standard conditions (298k, 1 atm)
- Enthalpy of neutralisation (delta H small n)
- Neutralisation means acid and base react
- The enthalpy change when a solution of acid and alkali react together to produce one
mole of water under standard conditions (298k, 1atm)
- NaOH + HCl → NaCl + H2O
- In a neutralisation reaction, it can only produce one mole of water to fit the definition
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