module 4 Flashcards
what breaking bonds
energy is absorbed
forming bonds
energy is released
exothermic reactions
- release energy
- temperature increases
- bond forming is greater than bond breaking
endothermic reactions
- absorb energy
- temperature decreases
- bond breaking is greater than bond forming
solute dissolves in solvent to form a solution
solute -solute bonds: broken (requires energy –> endothermic):
- ionic lattice held together by ionic bonds –> forces
of electrostatic attraction between positively and
negatively charged ions
solvent-solvent bonds: broken (requires energy –> endothermic):
- liquid state (water molecules are held together by
intermolecular forces eg. hydrogen bonding, dipole
- dipole interactions, dispersion forces)
solute-solvent interactions: formed (between ions and water molecules –> energy is released –> exothermic)
- solvation/hydration(case of water) –> solvent
molecules form hydration spheres around
individual ions
- the ion - dipole forces are formed between the
ions and the water molecules are. formed –>
stronger than hydrogen bonding in water
molecules
exothermic and endothermic dissolution of salts
- more energy required to break solute - solute and solvent - solvent (salit dissolves exothermically)
- less energy required to break solute - solute and solvent - solvent (salt dissolves endothermically)
salt that dissolves endothermically
- less energy required to break solute solute and solvent solvent interactions
eg. KCl
salt that dissolves exothermically
- more energy required to break solvent solvent and solute solute
eg. NaOH
molar heat of a solution (dissolution)
molar enthalpy change (delta H) (quantity of heat released upon the dissolution of one mole of a particular substance)
- J or kJ
- delta H = -q/n
- q = J
enthalpy
amount. of stored heat energy within a substance
H reaction = h products - h reactants
- reactants > products (exo) (more energy for bond forming)
- reactants < products (endo) ( more energy for bond breaking)
law of conservation of energy
energy cannot be created nor destroyed oonly reaaranged
transition state/activated complex
point in the reaction pathway between breaking and forming new covalent bonds
bond breaking
- absorbs energy
- delta H > 0
- endo
- positive
bond forming
- releases energy
- delta H < 0
- exo
- negative
energy profile diagrams
- change in enthalpy as the reaction proceeds
- larger the activation complex th emore enery that is required to start the reaction
x - axis: reaction pathway
y - axis: potential energy (kj)
heat released (q)
q = mcAt
- q = J
- m = g
- c = J
- T = K
calorimetry
study of enthalpy changes in a reaction
percentage error
%error = 100% X |theoretical value - experimental value|/|theoretical value|
do chemical bonds store energy
no
bond energy
amount of energy required to break a bond
- bonded has less energy than those that are unbonded
hess’s law
states that the enthalpy change accompanying a chemical change is independent of the route by which the chemical change occurs
- used to measure the enthalpy changes of a reaction
hess’s cycles/energy cycle diagram
represents the enthalpy changes in a system of reactions
standard enthalpy of formation
change in enthalpy associated with the formation of 1 mol of a substance in its standard state from its constituent elements in their standard state
HAVE NO FEAR OF ICE COLD BEER
diatomic atoms
enthalpy change
delta H = sum of products - sum of reactants
enthalpy in combustion
exothermic –> negative
enthalpy in photosynthesis
endothermic –> positive (reversed = cellular respiration)
two types of catalysts
homogeneous: same state as reactants
- usually all in aqueous/liquid/gaseous state
- reacts with one or more of the reactants to form a
temporary immediate
- immediate reacts with the other reactants and
decomposes to form products of reaction
- process regenerates the catalyst –> not consumed
therefore
- eg. decomposition of ozone
heterogeneous: different state as reactants
- usually solid catalyst in a liquid/gaseous/aqueous reaction
- reactants adsorbs to the catalyst at active sites
- chemical reaction takes place on the surface of the catalyst
- product the desorbs from the surface
intermediate
temporary species in a chemical reaction formed from a sub-reaction
- not considered a reactant or a product
bond breaking
h reaction = h bond breaking + h bond forming
- reverse the sign of bond forming
haber process
production of ammonia from its constituent gaseous elements
entropy
amount of randomness or lack of ordered structure
- J
delta s = sum of s products - sum of s reactants
differences between entropy and enthalpy
- enthalpy = change
- entropy = absolute
spontaneous
a process that will naturally occur without any ongoing external energy
exothermic and endothermic entropy
exothermic:
- heat released
- increases kinetic energy
- particles move faster
- increases entropy
endothermic:
- heat absorbed
- decreases kinetic energy
- decreases entropy
what are the drivers of reactions
enthalpy drive and entropy drive
forward enthalpy/entropy drive
- enthalpy: delta H < 0
- entropy: delta S > 0
- spontaneous
reverse entropy/enthalpy drive
- enthalpy: delta H > 0
- entropy: delta S < 0
- non-spontaneous
gibbs free energy
- refers to the maximum amount of work that can be performed by a thermodynamic system during a chemical process
delta G = delta H -T delta S - usually J
delta G < 0 = spontaneous
- reaction will go to completion without an external energy input (usually a flame at the start) (exergonic)
delta G > 0 = non-spontaneous
- continuous energy input is required to force the reaction to occur (endergonic)
delta G = 0
- neither spontaneous or non-spontaneous
H> 0 S<0
are never spontaneous as it will always be positive
H< 0 S> 0
always be spontaneous because delta G will always be negative
