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Chemistry - SCH4U1 > Unit 3 - Energy Changes and Rates of Reaction > Flashcards

Unit 3 - Energy Changes and Rates of Reaction Flashcards

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1
Q

thermochemistry

A

the study of the energy changes that accompany physical or chemical changes in matter

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2
Q

energy (E)

A
  • the ability to do work
  • unit: Joules (J)
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3
Q

potential energy (Ep)

A

the energy of a body or system due to its position or composition

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4
Q

kinetic energy (Ek)

A

the energy of an object due to its motion

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5
Q

thermal energy

A

the total quantity of kinetic and potential energy in a substance

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6
Q

heat

A

the transfer of thermal energy from a warm object to a cooler object

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7
Q

temperature

A

a measure of the average kinetic energy of entities in a substance

e.g. The total quantity of thermal energy of an iceberg is greater than that of a cup of hot water. The average water molecule in the hot water has more kinetic energy (higher temperature) than in the iceberg, but the iceberg contains billions more water molecules than a cup of hot water. (more thermal energy ≠ higher temperature)

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8
Q

Law of Conservation of Energy

A

Energy cannot be created or destroyed.

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9
Q

chemical system

A

a group of reactants and products being studied

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10
Q

surroundings

A

all the matter that is not part of the chemical system

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11
Q

open system

A

a system in which both matter and energy are free to enter and leave the system

e.g. An operating propane barbeque is an open system, since propane gas flows into the burner and thermal energy, and the products of combustion flow out.

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12
Q

closed system

A

a system in which energy can enter and leave the system, but matter cannot

e.g.; A glow stick is a closed system. Light energy is released from a chemical reaction, but the chemicals remain sealed inside.

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13
Q

exothermic reaction

A
  • releasing energy to the surroundings
  • the products of an exothermic reaction have lower potential energy than the reactants
  • the temperature of the surroundings increases
  • e.g. CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) + energy
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14
Q

endothermic reaction

A
  • absorbing energy from the surroundings
  • the products of an exothermic reaction have higher potential energy than the reactants
  • the temperature of the surroundings decreases
  • e.g. N2(g) + O2(g) → energy + 2NO(g)
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15
Q

enthalpy (H)

A

the total amount of thermal energy in a substance

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16
Q

enthalpy change (ΔH)

A
  • the energy released to or absorbed from the surroundings during a chemical or physical change
  • ΔH for physical changes (overcoming intermolecular forces) < ΔH for chemical changes (overcoming chemical bonds) < ΔH for nuclear changes (overcoming the forces holding nuclei together
  • ΔH > 0; endothermic reaction
  • ΔH < 0; exothermic reaction
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17
Q

expressing enthalpy change

A
  1. thermochemical equations with energy terms (e.g. AgNO3(s) + 22.6kJ → Ag+1(aq) + NO3-1(aq))
  2. thermochemical equations with ΔH values (e.g. 2SO2(g) + O2(g) → 2SO3(g); ΔH = -197.8 kJ
  3. molar enthalpy change (e.g. CH3OH(ℓ) + 3/2O2(g) → CO2(g) + 2H2O(ℓ); ΔHcomb = -726 kJ/mol
  4. potential energy diagram
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18
Q

specific heat capacity (c)

A
  • the quantity of thermal energy required to raise the temperature of 1 g of a substance by 1°C
  • unit: J/(g⋅°C)
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19
Q

specific heat capacity of water

A

4.18 J/(g⋅°C)

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20
Q

calorimetry

A

the experimental process of measuring the thermal energy change in a chemical or physical change

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21
Q

calorimeter

A

a device that is used to measure thermal energy changes in a chemical or physical reaction, consisting of:

  • a well-insulated reaction chamber to minimize energy losses
  • a tight-fitting cover with insulated holes for a thermometer
  • a stirring mechanism
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22
Q

types of calorimeter

A
  • simple calorimeter (a.k.a “coffee cup” calorimeter), for reactions in solution
  • flame calorimeter (a.k.a. bomb calorimeter), for combustion reactions
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23
Q

assumptions to be made when analyzing data obtained using a simple calorimeter

A
  • any thermal energy transferred from the calorimeter to the outside environment is negligible
  • any thermal energy absorbed by the calorimeter itself is negligible
  • all dilute, aqueous solutions have the same density (1.00 g/mL) and specific heat capacity (4.18 J/(g⋅°C)) as water
24
Q

quantity of heat transferred

A
  • equation: q = mcΔT
  • unit: kiloJoules (kJ)
  • if q is negative, it is an exotermic change (the temperature of the water increases)
  • if q is positive, it is an endothermic change (the temperature of the water decreases)
25
Q

calorimetry equation

A
  • ΔHsystem = -qsurroundings
  • ΔHsystem = nΔHr
  • qsurroundings = mcΔT
  • nΔHr = -mcΔT
26
Q

Hess’ Law

A

the enthalpy change for the conversion of reactants to products is the same, whether the conversion occurs in one step or several steps

ΔHtarget = ΣHknown

27
Q

rules for enthalpy changes (Hess’ Law)

A
  1. If you reverse a chemical reaction, you must also reverse the sign of ΔH.
  2. If the coefficients in a balanced equation are altered by multiplying or dividing by a factor, the value of ΔH must also be altered by the same factor.
28
Q

standard enthalpy of formation (ΔHf°)

A
  • the change in enthalpy that accompanies the formation of 1 mol of a compound from its elements in their standard states
  • the standard enthalpy of an element in its standard state is zero, unless it exists in more than one form at SATP (e.g. diamond, graphite)
29
Q

standard state

A

the most stable form of a substance under standard conditions (25°C and 100kPA)

30
Q

formation equations

A

formation equations show…
1. One mole of compound being formed from its elements.
2. Reactant elements are in their most stable form at SATP (some elements are diatomic; H2, O2, F2, Br2, I2, N2, Cl2).
3. Coefficients are balanced to yield one mole of product.

e.g. 2C(s) + 3H2(g) + 1/2O2(g) → C2H5OH(ℓ); ΔHf° = -235.2kJ/mol

31
Q

equation for using standard enthalpies of formation to calculate enthalpy change

A

ΔHr° = ΣnproductsΔHf°products - ΣnreactantsΔHf°reactants

32
Q

chemical kinetics

A

the area of chemistry that deals with rates of reactions

33
Q

reaction rate

A

the change in concentration of a reactant or a product of a chemical reaction, per unit time

34
Q

physical properties used to measure reaction rates

A
  1. change in volume and/or pressure of a gas
  2. change in colour
  3. change in transparency
  4. change in mass of a solid reactant or product
  5. change in pH
  6. change in conductivity
  7. change in temperature
35
Q

change in volume and/or pressure of a gas

A

can be measured using a eudiometer, if the number of moles of gas is different on the reactant and product side of the equation

e.g. CaCO3 + 2H+(aq) → 2Ca+2(aq) + H2O(ℓ) + CO2(g)

36
Q

change in colour

A

can be measured using a spectrophotometer, if the amount of coloured substance is changing during a chemical reaction,

examples of coloured species:

  • Cu+2(aq) = blue
  • Cu+(aq) = green
  • MnO4-(aq) = purple
  • IO-(aq) = yellow
  • Cl2(g) = green
  • I2(g) = purple
  • Br2(g) = red-brown
  • NO2(g) = brown

e.g. 2Cu(s) + 2H+(aq) → 2Cu+(aq) + H2(g)

37
Q

change in transparency

A

can be measured using a spectrophotometer, if a fine precipitate forms during the reaction

e.g. Pb(NO3)(aq) + NaI(aq) → PbI2(s) + 2NaNO3(aq)

38
Q

change in mass of a solid reactant or product

A

can be measured when a solid reactant is disappearing or a solid product (of large particle size, not a precipitate) is appearing

e.g. Mg(s) + 2H+(aq) → H2(g) + Mg+2(aq)

39
Q

change in pH

A

can be measured using a pH meter or pH indicator, if the moles of hydrogen ions is different on the reactant and product side of the equation

e.g. SO3(g) + H2O(ℓ) → 2H+(aq) + SO4-2(aq)

40
Q

change in conductivity

A

can be measured using a conductivity meter or voltmeter, if the moles of ions is different on the reactant and product side of the equation

e.g. C4H9Br(g) + H2O(ℓ) → C4H9OH(aq) + H+(aq) + Br-(aq)

41
Q

change in temperature

A

can be measured using a thermometer, as all reactions have some change in enthalpy

42
Q

average reaction rate

A

the change in reactant or product concentration over a given time interval (secant)

rateA = Δ[A] ÷ Δt

43
Q

instantaneous reaction rate

A

the rate of a chemical reaction at a single point in time (tangent)

rateinst = [A] ÷ t

44
Q

collision theory

A

the theory that chemical reactions can occur, only if reactants collide with proper orientation, and with enough kinetic energy to break reactant bonds and form product bonds

45
Q

activation energy (Ea)

A

the minimum energy that reactant molecules must possess for a reaction to be successful; represents the energy difference between reactants and the activated complex

46
Q

purposes of activation energy

A
  1. to overcome the electrostatic repulsive forces between colliding entities
  2. to weaken the bonds of the reactants
47
Q

activated complex (transition state)

A

an unstable arrangement of atoms containing partially formed and unformed bonds that represents the maximum potential energy point in the change

48
Q

factors affecting reactant rates

A

chemical nature of reactants: for any reactant, the bond type, strength, and number determine the activation energy required for successful collisions
concentration of reactants: if the concentration of a reactant is increased, the probability of collisions between reactant molecules increases and a great number of effective collisions is likely to occur
surface area: for any reaction in which the reactants are in more than one state of matter (solid and liquid), the greater the surface area of the solid, the faster the reaction rate
temperature of the reaction system: in general, the rate of a chemical reaction increases with the temperature at which it is carried out (the reaction rate of many reactions doubles for every 10°C temperature increase)
presence of a catalyst: a catalyst provides an alternative pathway for a reaction, which has a lower activation energy

49
Q

catalyst

A

a substance that alters the rate of a chemical reaction, without itself being permanently changed

50
Q

heterogenous catalyst

A

a catalyst in a reaction, in which the reactants and the catalysts are in different physical states

51
Q

homogeneous catalyst

A

a catalyst in a reaction, in which the reactants and the catalysts are in the same physical states

52
Q

elementary step

A

a step involving one-, two-, or three-entity collisions that cannot be explained by simpler reactions

53
Q

reaction mechanism

A

a series of elementary steps by which a chemical reaction occurs

54
Q

reaction intermediate

A

an entity that is neither a reactant, nor a product, but is formed and consumed during the reaction sequence

55
Q

molecularity

A

refers to the number of reactant particles that are involved in an elementary step (i.e. unimolecular, bimolecular, termolecular)

56
Q

rate law equations

A

A → products: rate = k[A]
A + A → products: rate = k[A]2
A + B → products: rate = k[A][B]

Chemistry - SCH4U1 (5 decks)

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