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chemistry > module 4 > Flashcards

module 4 Flashcards

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1
Q

what breaking bonds

A

energy is absorbed

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2
Q

forming bonds

A

energy is released

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3
Q

exothermic reactions

A
  • release energy
  • temperature increases
  • bond forming is greater than bond breaking
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4
Q

endothermic reactions

A
  • absorb energy
  • temperature decreases
  • bond breaking is greater than bond forming
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5
Q

solute dissolves in solvent to form a solution

A

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

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6
Q

exothermic and endothermic dissolution of salts

A
  • 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)
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7
Q

salt that dissolves endothermically

A
  • less energy required to break solute solute and solvent solvent interactions
    eg. KCl
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7
Q

salt that dissolves exothermically

A
  • more energy required to break solvent solvent and solute solute
    eg. NaOH
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8
Q

molar heat of a solution (dissolution)

A

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

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9
Q

enthalpy

A

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)

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10
Q

law of conservation of energy

A

energy cannot be created nor destroyed oonly reaaranged

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11
Q

transition state/activated complex

A

point in the reaction pathway between breaking and forming new covalent bonds

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12
Q

bond breaking

A
  • absorbs energy
  • delta H > 0
  • endo
  • positive
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13
Q

bond forming

A
  • releases energy
  • delta H < 0
  • exo
  • negative
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14
Q

energy profile diagrams

A
  • 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)
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15
Q

heat released (q)

A

q = mcAt
- q = J
- m = g
- c = J
- T = K

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16
Q

calorimetry

A

study of enthalpy changes in a reaction

17
Q

percentage error

A

%error = 100% X |theoretical value - experimental value|/|theoretical value|

18
Q

do chemical bonds store energy

A

no

19
Q

bond energy

A

amount of energy required to break a bond
- bonded has less energy than those that are unbonded

20
Q

hess’s law

A

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

21
Q

hess’s cycles/energy cycle diagram

A

represents the enthalpy changes in a system of reactions

22
Q

standard enthalpy of formation

A

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

23
Q

HAVE NO FEAR OF ICE COLD BEER

A

diatomic atoms

24
Q

enthalpy change

A

delta H = sum of products - sum of reactants

25
Q

enthalpy in combustion

A

exothermic –> negative

26
Q

enthalpy in photosynthesis

A

endothermic –> positive (reversed = cellular respiration)

27
Q

two types of catalysts

A

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

28
Q

intermediate

A

temporary species in a chemical reaction formed from a sub-reaction
- not considered a reactant or a product

29
Q

bond breaking

A

h reaction = h bond breaking + h bond forming
- reverse the sign of bond forming

30
Q

haber process

A

production of ammonia from its constituent gaseous elements

31
Q

entropy

A

amount of randomness or lack of ordered structure
- J

delta s = sum of s products - sum of s reactants

32
Q

differences between entropy and enthalpy

A
  • enthalpy = change
  • entropy = absolute
33
Q

spontaneous

A

a process that will naturally occur without any ongoing external energy

34
Q

exothermic and endothermic entropy

A

exothermic:
- heat released
- increases kinetic energy
- particles move faster
- increases entropy
endothermic:
- heat absorbed
- decreases kinetic energy
- decreases entropy

35
Q

what are the drivers of reactions

A

enthalpy drive and entropy drive

36
Q

forward enthalpy/entropy drive

A
  • enthalpy: delta H < 0
  • entropy: delta S > 0
  • spontaneous
37
Q

reverse entropy/enthalpy drive

A
  • enthalpy: delta H > 0
  • entropy: delta S < 0
  • non-spontaneous
38
Q

gibbs free energy

A
  • 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

39
Q

H> 0 S<0

A

are never spontaneous as it will always be positive

40
Q

H< 0 S> 0

A

always be spontaneous because delta G will always be negative

chemistry (9 decks)

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