Module 3 - [ch9] (enthalpy) Module 5 - [ch22a] (lattice energy) Flashcards
what is a system in a chemical reaction
atoms and bonds involved in a chemical reaction
explain law of conservation
energy in an isolated system remains the same
energy cannot be created nor destroyed. only transferred from one form to another
endothermic reaction
energy taken in to break bonds
ΔH is positive
What is meant by the term exothermic? [1]
energy released to make bonds
ΔH is negative
activation energy
minimum energy required for a reaction to take place
standard conditions
1atm - 101kPa
298K - 25 degrees
what is standard state
the state an element/compound exists in at standard conditions
Define the term enthalpy change of formation. [2]
energy change when1 mole of substance is formed from its constituent elements in their standard state under standard conditions
define the term enthalpy change of combustion [2]
change in energy when 1 mole of substance is completely combusted
enthalpy change of neutralisation
energy change when 1 mole of water is produced from a neutralisation reaction
enthalpy change of reaction
energy change associated with a given reaction
enthalpy change {}
Q = mcΔT
energy change = mass * specific heat capacity * change in temperature
shc = 4.18g^-1K^-1
what are the advantages to using a bomb calorimeter
minimises heat loss to surroundings
pure oxygen used to ensure complete combustion
Suggest three reasons why standard enthalpy changes of combustion determined
experimentally are less exothermic than the calculated theoretical values [3]
- heat is released to surroundings
- non-standard conditions
- incomplete combustion (reaction may not go to completion)
average bond enthalpy [2]
average energy required to break 1 mole of bonds in gaseous molecules
Use ideas about the enthalpy changes that take place during bond breaking and bond
making to explain why some reactions are exothermic. [2]
bond breaking absorbs energy whereas bond making releases energy
in some reactions more energy is released than absorbed, these reactions are exothermic
State le Chatelier’s principle [1]
The position of equilibrium will shift so as to minimise the effect of any change in conditions
It is very difficult to determine the standard enthalpy change of formation of
hexane directly. Suggest a reason why [1]
many different hydrocarbons would form
OR
activation energy very high
Why do bond enthalpies have positive values? [1]
bond breaking is endothermic so energy has to be put in to break a bond
Write the equation, including state symbols, that represents the standard
enthalpy change of formation for carbon monoxide, CO [2]
C(s) + 1/2O2(g) -> CO(g)
what is Hess’ law
enthalpy change is independent of route taken
what is lattice enthalpy
measure of the strength of the ionic bonding in a giant ionic lattice
the enthalpy change when 1mole of an ionic compound is formed from its gaseous ions under standard conditions
born haber cycle
gaseous ions at top
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gaseous atoms |
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elements in |
standard states |
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ionic lattice
define standard enthalpy of atomisation
the enthalpy change when 1 mole of gaseous atoms are formed from the elements in its standard state under standard conditions
why is the value of enthalpy of atomisation always positive
endothermic process
bonds are broken to form gaseous atoms
Atomisation of :
sodium
chlorine
Na(s) -> Na(g)
1/2 Cl2(g) -> Cl(g)
why are ionisation energies always positive
energy is required to overcome the attraction between a negative electron and the positive nucleus
endothermic
what is electron affinity
measure of energy to gain electrons
( opposite of ionisation energy )
first electron affinity
the energy change that takes place when one electron is added to each atom in one mole of gaseous atoms to form one mole of gaseous 1- ions
why are first electron affinities exothermic
electron being added is attracted in towards the nucleus
why are second electron affinities endothermic
a second electron is being gained by a negative ion which repels the electron
energy must be put in to overcome this repulsion and force another electron in
define standard enthalpy change of solution
the enthalpy change that takes place when one mole of a solute dissolves in a solvent
in an enthalpy change of solution, when would the products be in their aqueous state and why
when the solvent is water
ions from ionic lattice surrounded by water molecules
delta positive H+ attracted to negative ions
delta negative O~ attracted to positive ions
in enthalpy of solution what mass is used
mass of the solution
mass of the water + mass of the solid
define enthalpy change of hydration
the energy change that accompanies the dissolving of gaseous ions in water to form one mole of aqueous ions
describe the dissolving process of a solid ionic compound
the ionic lattice breaks up
water molecules are attracted to and surround the ions
lattice energy born haber cycle , fewer steps
Gaseous ions
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| aqueous ions
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ionic lattice
2 factors affecting lattice enthalpy
ionic size
ionic charge
what is the effect of ionic size in ionic compounds
as ionic radius increases
attraction between ions decreases
lattice energy increases (less negative)
melting point decreases
what is the effect of ionic charge in ionic compounds
as ionic charge increases
attraction between ions increases
lattice energy becomes more negative ( decreases )
melting point increases
what are the factors affecting hydration
ionic size
ionic charge
how does ionic charge affect hydration
as ionic charge increases
attraction with water molecules increases
hydration energy becomes more negative
how does ionic size affect hydration
increasing ionic radius
decreased attraction between ion and water molecules
hydration energy less negative
how much energy is required to dissolve an ionic compound in water
attraction between ions in the ionic lattice must be overcome
energy = to lattice enthalpy needed
when would a compound dissolve
(in terms of enthalpies)
if the sum of the hydration enthalpies is larger than the magnitude of the lattice enthalpy
compound should dissolve
enthalpy change of solution will be exothermic
