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Honors Chemistry > Module 07: Energy in Reactions > Flashcards

Module 07: Energy in Reactions Flashcards

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1
Q

07.01 Endothermic and Exothermic

07.01 Endothermic and Exothermic

A
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2
Q

07.01 Endothermic and Exothermic

What is thermochemistry?

A

a branch of chemistry concerned with the quantities of heat released or absorbed during chemical reactions

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3
Q

07.01 Endothermic and Exothermic

What are the three energies of a system?

A
  1. Temperature
  2. Thermal Energy
  3. Heat
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4
Q

07.01 Endothermic and Exothermic

What is “temperature?”

A

Measure average kinetic energy in particle

direct relationship kinetic energy and temperature

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5
Q

07.01 Endothermic and Exothermic

What is thermal energy?

A

Measure total kinetic energy in sample

  • difference temperature:
    • Thermal: represents kinetic energy of a sample together
    • Temperature: average kinetic energy of individual particles
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6
Q

07.01 Endothermic and Exothermic

What is heat?

A

Transfer of thermal energy from one substance to another due to temperature differences

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7
Q

07.01 Endothermic and Exothermic

How does heat transfer?

A

From matter at high temperature to low temperature

  • depends temperature differences (not in total thermal energy of substances)
  • continue until both have the same temperature
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8
Q

07.01 Endothermic and Exothermic

How is heat and thermal energy measured?

A

SJ units: joule (J)

  • unit kilojoules (kJ) → calories
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9
Q

07.01 Endothermic and Exothermic

Define enthalpy.

A

total heat content of a system equal to the internal energy of the system plus the product of pressure and volume

  • denoted “q”
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10
Q

07.01 Endothermic and Exothermic

What is the difference between endothermic & exothermic?

A

ENDOTHERMIC:

reactions absorb energy from surroundings → increase net energy (+q)

  • absorb kinetic energy > convert potential energy
  • amount energy absorbed: dependent amount of reactants and difference of potential energy

EXOTHERMIC:

  • reactions releases energy to surroundings → decrease net energy (-q)*
  • stored potential energy released into kinetic energy
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11
Q

07.01 Endothermic and Exothermic

What is a change in enthalpy (△H) ?

A

Equal heat gained or lost between system and surroundings

  1. gain: +q
  2. lost: -q
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12
Q

07.01 Endothermic and Exothermic

What is a potential energy diagram?

A
  • tracks potential energy
  • y-axis: measure potential energy & x-axis: progress of reaction
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13
Q

07.01 Endothermic and Exothermic

What is th potential energy diagram of an exothermic reaction?

A
  • products less PE than reacts
  • starts higher and ends lower

Law of Conservation of Energy: obeyed extra energy not end up is released into the surroundings

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14
Q

07.01 Endothermic and Exothermic

What is the diagram for potential energy of a endothermic reaction?

A
  • products more potential energy than reactants
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15
Q

What is activation energy? What is it on a potential energy diagram?

A

the minimum amount of energy colliding particles must have for a chemical reaction to occur

Diagrams: representated positive slope that follows reactatns

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16
Q

07.02 Enthalpy Values

07.02 Enthalpy Values

A
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17
Q

07.02 Enthalpy Values

What is the property of specific heat capacity?

A

the quantity of heat required to raise the temperature of one gram of a substance by one degree Celsius

  • intensive property
  • impacted by mass (how much thermal energy can be absorbed)
  • J/g °C
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18
Q

07.02 Enthalpy Values

What is the specific heat formula?

A

C = q/m x △T

  • C: specific heat
  • q: heat generated by temperature change
  • m: mass of substance
  • △T: change in temperature
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19
Q

07.02 Enthalpy Values

How can energy transfers be calculated?

A

q = m x C x △T

  • q: the heat gained (positive) or lost (negative) by the system, in joules
  • m: the mass of the sample, in grams
  • C: the specific heat capacity of the sample, in joules per gram per degree Celsius Jg °C
  • △T: change in temperature (°C)
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20
Q

07.02 Enthalpy Values

How does enthalpy apply to phase change?

A
  • Deposition/Condensation/Freezing: Exothermic (−△H or −q)
  • Sublimation/Vaporization/Melting: Endothermic (+△H or +q)
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21
Q

07.02 Enthalpy Values

How can the overall change of enthalpy be determined?

A
  • comparing the sum of the enthalpies of reactants to sun of products
  • Heat of Reaction
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22
Q

07.02 Enthalpy Values

What is Heat of Reaction?

A

△H = Hproducts − Hreactants

The heat of reaction (overall change in enthalpy) is found by taking the sum of enthalpies of the products and subtracting the sum of enthalpies of the reactants.

  • Endothermic:
    • +△H
    • energy required break bonds is more than amount of energy released
  • Exothermic:
    • -△H
    • energy required break less amount energy released in forming bonds in product
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23
Q

07.02 Enthalpy Values

The chemical reaction and enthalpy change for burning methane (CH4) is given below. If 38.5 grams of methane are burned, how many kilojoules of energy are transferred, and is this reaction endothermic or exothermic?

CH4 + O2 → CO2 + 2H2O

△H = −890.0 kJ/mol

A

The enthalpy change value of this reaction tells us that for every mole of methane (CH4) burned, 890.0 kilojoules are given off because the enthalpy change is negative. This is an exothermic reaction. Use the enthalpy change as a conversion factor between moles of methane and kilojoules of energy.

38.5 g CH4÷1 × 1 mol CH4 ÷ 16.05 g CH4 × −890.0 kJ ÷ 1 mol CH4 = −2,135 kJ

This shows that as 38.5 grams of methane are burned, 2,140 kilojoules of energy are released.

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24
Q

07.02 Enthalpy Values

If the amount of energy required to break bonds in the reactants is more than the amount of energy released in forming bonds in the products, then the reaction will have a negative change in enthalpy (−ΔH). (2 points)

True

False

A

False

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25
Q

07.02 Enthalpy Values

The substances below are listed by increasing specific heat capacity value. Starting at 30.0 °C, they each absorb 100 kJ of thermal energy. Which one do you expect to increase in temperature the least? (3 points)

  1. Cadmium, 0.230 J/(g °C)
  2. Sodium, 1.21 J/(g °C)
  3. Water, 4.184 J/(g °C)
  4. Hydrogen, 14.267 J/(g °C)
A

4. Hydrogen, 14.267 J/(g °C)

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26
Q

07.02 Enthalpy Values

CaO(s) + H2O(l) → Ca(OH)2(s): ΔH = −65.2 kJ

Which statement about the reaction between calcium oxide and water is correct? (3 points)

  1. 65.2 kJ of heat are released for every mole of CaO that reacts.
  2. 130 kJ of heat are released for every mole of H2O that reacts.
  3. 130 kJ of heat are absorbed for every mole of CaO that reacts.
  4. 65.2 kJ of heat are absorbed for every mole of H2O that reacts.
A

1. 65.2 kJ of heat are released for every mole of CaO that reacts

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27
Q

07.02 Enthalpy Values

A 52-gram sample of water that has an initial temperature of 10.0 °C absorbs 4,130 joules. If the specific heat of water is 4.184 J/(g °C), what is the final temperature of the water? (4 points)

  1. 11 °C
  2. 19 °C
  3. 29 °C
  4. 51 °C
A

3. 29 °C

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28
Q

07.02 Enthalpy Values

The specific heats of two natural substances are shown in the table.

Specific Heat:

Wet Mud: 2.5 J/g °C

Sandy Clay: 1.4 J/g °C

A pile of wet mud and a pile of sandy clay are placed next to each other on a sunny day.

If the temperature of each pile is equal and allowed to cool in the shade for a few minutes, which substance is likely to show a lower temperature and why? (4 points)

  1. Wet mud, because it released more heat while placed in the sun.
  2. Sandy clay, because it absorbed more heat while placed in the sun.
  3. Wet mud, because it requires less energy to increase its temperature.
  4. Sandy clay, because it needs to release less energy to decrease its temperature.
A

4. Sandy clay, because it needs to release less energy to decrease its temperature

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29
Q

07.02 Enthalpy Values

Some thermodynamic properties of ethanol are listed in the table.

Thermodynamic Properties

c (solid): 0.5 J/g °C

c (liquid): 1.0 J/g °C

c (gas): 2.0 J/g °C

Melting Point: −114 °C

Boiling Point: 78 °C

How much heat is released when 40.0 g of ethanol cools from −120 °C to −136 °C? (4 points)

  1. 640 J
  2. 580 J
  3. 320 J
  4. 290 J
A

3. 320 J

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30
Q

07.03 Honors Entropy

07.03 Honors Entropy

A
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31
Q

07.03 Honors Entropy

What is a “system” in chemistry and what are the types of systems?

A

System is the container or entity of interest

  1. Isolated System
  2. Closed System
  3. Open System
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32
Q

07.03 Honors Entropy

What is a Isolated System?

A
  • neither energy or matter is premitted exchange with surroundings
  • total amount of energy = constant
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33
Q

07.03 Honors Entropy

What is a closed system?

A
  • Energy can or leave, but matter cannot*
  • boundary between systems
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34
Q

07.03 Honors Entropy

What is an open system?

A

Both matter and energy are exchanged freely between the system and the surroundings

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35
Q

07.03 Honors Entropy

What is a spontaneous change?

A

Change in a system that proceeds without a net input of energy from an outside source.

  • not instantaneous
  • needs some activation energy to get started → without any continuing outside influence (heating, cooling, stirring)
  • fast/slow
  • Factor affects spntaneity: change in entropy (△S) - reaction or process takes place
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36
Q

07.03 Honors Entropy

How does entropy compare between exothermic and endothermic reactions?

A

In general, an exothermic reaction will increase the entropy of its surroundings because it releases energy.

An endothermic reaction will decrease the entropy of the surroundings because it absorbs energy.

37
Q

07.03 Honors Entropy

What is entropy?

A

the property that describes the disorder of a system (S)

38
Q

07.03 Honors Entropy

How does entropy change in relation to the phases of matter?

A

Rises from Solid → Gas

Gas: highest

Solid: lowest

  • Entropy increases as a substance changes to a less condensed state (solid to liquid or liquid to gas). This is because an increase in the motion of particles means an increase in disorder. Gas particles have many more positions and arrangements available to them than the same particles in the solid state.
  • Entropy increases when substances mix or something is dissolved because mixing particles together increases the disorder.
  • Entropy increases when the total number of moles of products is greater than the number of moles of reactants because a greater number of particles leads to a greater chance of disorder.
  • The entropy of any sample increases when temperature increases because of the increase in kinetic energy and movement.
  • Entropy is greater for particles with a greater mass. For example, one mole of Cl2 has a greater mass than one mole of F2, so the entropy of the chlorine sample is greater.
  • Entropy is greater for weakly bonded compounds than for strongly bonded compounds. For example, the carbon atoms in graphite have weaker bonds than the carbon atoms in diamonds, so the entropy is greater for the sample of graphite.
  • Entropy is greater for compounds with greater complexity (greater number of atoms, greater mass, etc.).
40
Q

07.03 Honors Entropy

What is free energy?

A

the energy in a physical system that can be converted to do work

41
Q

07.03 Honors Entropy

What is Gibbs’s Free Energy Equation?

A

△G = △H − (T × △S)

  • △G: free energy change
  • △H: enthalpy change
  • T: temperature
  • △S: entropy change

At a constant temperature and pressure, a process is spontaneous in the direction that gives a negative change in free energy (−△G).

42
Q

07.03 Honors Entropy

A metal pot is filled with water and placed on top of a stove. As the stove heats up, the water begins to boil, and steam can be seen escaping from the pot. What type of system is this an example of? (4 points)

  1. A closed system, because energy can enter or leave the system, but matter cannot.
  2. An open system, because heat and matter are both able to enter or leave the system.
  3. A closed system, because matter can enter or leave the system, but energy cannot.
  4. An isolated system, because heat and matter can enter but cannot leave the system.
A

2. An open system, because heat and matter are both able to enter or leave the system

43
Q

07.03 Honors Entropy

Which of the following changes will always be true for a spontaneous reaction? (4 points)

  1. +ΔG
  2. –- ΔG
  3. +ΔS
  4. –- ΔH
A

2. - ΔG

44
Q

07.03 Honors Entropy

Which of the following changes has a decrease in entropy? (4 points)

  1. CO2(s) → CO2(g)
  2. 2H2O(l) → 2H2(g) + O2(g)
  3. H2O(l) → H2O(s)
  4. 2KClO3(s) → 3O2(g) + 2KCl(s)
A

3. H2O(l) → H2O(s)

45
Q

07.03 Honors Entropy

Which of the following combinations will result in a reaction that is spontaneous at all temperatures? (4 points)

  1. Negative enthalpy change and negative entropy change
  2. Negative enthalpy change and positive entropy change
  3. Positive enthalpy change and negative entropy change
  4. Positive enthalpy change and positive entropy change
A

2. Negative enthalpy change and positive entropy change

46
Q

07.03 Honors Entropy

Combustion of ethane gas is an exothermic reaction. Which of the following best describes the temperature conditions that are likely to make the combustion of ethane gas a spontaneous change? (4 points)

  1. Any temperature, because combustion of ethane leads to an increase in entropy.
  2. Any temperature, because combustion of ethane leads to a decrease in entropy.
  3. Low temperature only, because combustion of ethane leads to an increase in entropy.
  4. High temperature only, because combustion of ethane leads to a decrease in entropy.
A

1. Any temperature, because combustion of ethane leads to an increase in entropy

47
Q

07.03 Honors Entropy

A sealed, insulated calorimeter contains water at 310 K. The surrounding air temperature is 298 K, and the water inside the calorimeter remains at 310 K two hours later. What type of system does the calorimeter attempt to model? (4 points)

  1. A closed system, because energy can enter or leave the system, but matter cannot.
  2. A closed system, because neither heat nor matter is entering or leaving the container.
  3. An isolated system, because energy can enter or leave the system, but matter cannot.
  4. An isolated system, because neither heat nor matter is entering or leaving the container.
A

4. An isolated system, because neither heat nor matter is entering or leaving the container

48
Q

07.03 Honors Entropy

If a reaction is not spontaneous at any temperature, what is true of the reaction? (4 points)

  1. It has a negative enthalpy change and positive entropy change.
  2. It has a negative enthalpy change and negative entropy change.
  3. It has a positive enthalpy change and positive entropy change.
  4. It has a positive enthalpy change and negative entropy change.
A

4. It has a positive enthalpy change and negative entropy change

49
Q

07.04 Honors Calorimetry

07.04 Honors Calorimetry

A
50
Q

07.04 Honors Calorimetry

What is calrimetry?

A

Calorimeter: used measure heat absorbed or released by a chemical or physical change

Calorimetry: measurement of change

51
Q

07.04 Honors Calorimetry

How can the equation q = m × c × △T be applied to calorimeter?

A
  • determine energy gained or lost by the water in the calorimeter
  • q → + (gain energy) & - (loss energy)
52
Q

07.04 Honors Calorimetry

What is the Calorimetry equation?

A

qsurroundings = −qsystem

m × c × △T = −(m × c × △T)

* negative sign: designate flow of every

53
Q

07.04 Honors Calorimetry

How is specific heat equations applied to calorimetry (steps)?

A

Step 01: Indenfity known variables

Step 02: Determine unknown variables

Step 03: Calculate qsurroundings and qsystem.

Step 04: solve for c (specific heat)

54
Q

07.04 Honors Calorimetry

Q: A 150.0-gram sample of metal at 75.0 °C is added to 150.0 g of water at 15.0 °C. The temperature of the water rises to 18.3 °C.

What is the specific heat of the metal?

A

qsurroundings = m × c × △T

qsurroundings = 150.0 g × 4.184 × (18.3 °C − 15.0 °C)

qsurroundings = 2,100 J

qsurroundings = −qsystem

−2,100 J = 150.0 g × c × (18.3 °C − 75.0 °C)

c = 0.25

55
Q

07.05 Reaction Rates

♣ 07.05 Reaction Rates ♣

A
56
Q

07.05 Reaction Rates

What is the Collision Theory?

A

“For particles to react successfully, they must collide”

  • states rate of reaction is determined by the number of successful collisions
  • occur over a given amount of time

Successful Collision: enough energy for reactants to intersect valence shells, create new bonds, make product

Two Factors:

  1. adequate speed of particles
  2. correct orientation
57
Q

07.05 Reaction Rates

How does complexity and alignment relate?

A

More complex → more difficult align with one another

  • must collide energy equal/greater
58
Q

07.05 Reaction Rates

What is the Maxwell-Boltzmann distribution curve?

A

Represent distribution of particle energy within a sample at a given temperature

  • activation energy requirement → requirement for successful collisions
    • under particles under this area: not enough energy to react
    • only particles represented on upper part: enough energy to react

When analyzing a Maxwell-Boltzmann distribution curve, a common mistake is thinking that the curve with the higher hump is the one with the greater temperature. Remember, the curve shows the number of particles with a given energy. But because temperature is a measurement of the average kinetic energy, the curve that has more particles with higher energy (or shifted right) is the one with the higher temperature. Therefore, in the graph above, curve T2 reflects a higher temperature than curve T1.

59
Q

07.05 Reaction Rates

What factors influence reaction rates?

A
  1. Concentration and Rates
  2. Temperature and Rates
  3. Catalysts and Rates
60
Q

07.05 Reaction Rates

What is “Concentration and Rate” ?

A
  • Increae concentation: increase rate of a reaction
  • Concentration higher: chances of particles coming in contact increases

Increase pressure or decreasing volume: increase reaction rate

61
Q

07.05 Reaction Rates

How are reaction rates related to “temperature and rates?”

A
  • slight increase temperature: changes distribution of energy in a sample
    • kinetic increase little amount → large increase number of particles meet and exceed activation eneryg
  • Reation not happen fast enough: raised temperature help
62
Q

07.05 Reaction Rates

How do cataylsts and rates infuence reaction rates?

A
  • temperature not increased meet activation energy needed
  • Catalyst: alternative path → less activation energy required
    • results larger percentage of successful collisions

Maxwell-Boltzman distribution: lower activation energy requirements allows more partciles to meet energy requiremetns for successful collision without increase in temperature

63
Q

07.05 Reaction Rates

how do you calculate the average rate of a reaction?

A

Rates: measures change in something over a given amount of time

Average rate of change: change in concentration of a reaction over a time interval

  • depends concentration of reactants present
    • decreases as reactants are used up
  • Concentration measured molarity (M)
64
Q

07.05 Reaction Rates

What is the formula for the average rate?

A

(Final Concentration - Initial Concentration)

÷

(Final Time - Initial Time)

65
Q

07.06 Equilibrium

♣

07.06

Equilibrium

♣

A
66
Q

07.06 Equilibrium

What is a dynamic equilibrium?

A
  • Dynamic → reaction occurs in both directions
  • Dynamic equilibrium: rates of forwards and reserve reaction continues to occur*

A + B ⇄ C + D

  • Forward reactions follow the right-pointing arrow in the equation and yield products.
  • Reverse or backward reactions follow the left-pointing arrow in the equation and yield reactants.
67
Q

07.06 Equilibrium

When does a reaction reach dynamic equilibrim in a closed system?

A
  1. Rate of forward reaction = rate of reverse reaction
  2. Concentrations of the reactants and products remain constant
68
Q

07.06 Equilibrium

What are the rtaes of opposite reactions in a dynamic equilibrium?

A
  • equal*
  • reaction not stop when in equilibrium
69
Q

07.06 Equilibrium

Describe dynamic equilibrium in relation to evaporation and condensation.

A

The evaporating and condensing phases of water continue back and forth in dynamic equilibrium.

70
Q

07.06 Equilibrium

Describe dynamic equilibrium in relation to dissolving and crysallization.

A

Dynamic equilibrium is also possible during dissolving and crystallization, which are both physical changes. For instance, when salt dissolves in water, it doesn’t change chemically; it just breaks apart. As more salt is dissolved in the water, it reaches a saturation point, and the salt begins to crystalize again. At this point, a balance has been reached in this physical process; salt is continuously dissolving at the same rate as it is crystallizing. The salt and water solution is in dynamic equilibrium.

71
Q

07.06 Equilibrium

What is the Law of Mass Action?

A
  • General description of the equilibrium conditions, including an equation defined the equilibrium constant*
  • Equilibrium Constant: (K)

the value obtained when equilibrium concentrations are plugged into the law of mass action equilibrium expression; represented by the letter K, it is constant for a given system at a given temperature

72
Q

07.06 Equilibrium

What is the equilibrium constant?

A
  • In the forward reaction, w moles of A react with x moles of B to produce y moles of C and z moles of D.
73
Q

07.06 Equilibrium

What is the equilibrium position?

A

a given set of equilibrium concentration values

  • depends on intial concentration of each substance
  • predicted based on value of equilibrium constant
74
Q

07.06 Equilibrium

What is the rates of Reactions Graph?

A

Compares the reaction rates over time

  • see when equilibrium happens
  • The point on the graph when the forward and reverse reactions reach equal rates, is when the reaction has reached dynamic equilibrium.
75
Q

07.06 Equilibrium

What is the Concentration Graph?

A

describes the change in concentrations of reactants and products over time

  • direct relationship between concentration and rates of reverse reaction
  • The point on the graph when both concentrations plateau and remain constant when the system has reached dynamic equilibrium.
76
Q

07.06 Equilibrium

What is the significance of the K (equilibrium constant)?

A

K > 1

The forward reaction proceeds to the right, giving more products than reactants when the system reaches equilibrium.

K < 1

The forward reaction will not proceed very far before equilibrium is established, and the system will have more reactants than products at equilibrium.

77
Q

07.07 Le Châtelier’s Principle

♣

07.07

Le Châtelier’s Principle

♣

A
78
Q

07.07 Le Châtelier’s Principle

What is Le Châtelier’s Principle?

A
  • equilibrium reactions want to stay in equilibrium → distrurbed through stress → change concentration to shift to a new equilibrium.
    • discovered by Henri Louis Le Châtelier

Equilibrium reactions show the forward and reverse reactions occurring simultaneously using a double arrow pointed in both directions in the chemical equation.

A + B ⇄ C + D

79
Q

07.07 Le Châtelier’s Principle

How is stress on equilibrium reaction described?

A
  • A shift right means the reaction moves in the forward direction and yields more products.
  • A shift left means the reaction moves in reverse (backward) and yields more reactants.
80
Q

07.07 Le Châtelier’s Principle

How do catalysts affect equilibrium?

A
  • increases rate of reaction (providing alternative reaction path)
  • Lowers activation eneryg for both forwards and reverse reactions
    • no effect equilibrium
81
Q

07.07 Le Châtelier’s Principle

What is the Le Châtelier’s Principle for Concentration?

A

If a component (reactant or product) is added to a reaction system at equilibrium, the equilibrium position will shift away from what was added to lower its concentration.

If a component is removed from a reaction system at equilibrium, the equilibrium system will shift to replace some of what was removed.

82
Q

07.07 Le Châtelier’s Principle

What is

Le Châtelier’s Principle for Pressure?

A

When the pressure of an equilibrium system is increased, the system will shift in the direction that decreases the total moles of gas in the system to minimize the stress.

The opposite is true if pressure is decreased; the system will shift toward the side of the equation that has a greater total number of moles of gas.

83
Q

07.07 Le Châtelier’s Principle

Le Châtelier’s Principle for Temperature:

A

When an equilibrium system is heated, the temperature increases in the system. It will shift forward (right) in an endothermic reaction and backward (left) in an exothermic reaction.

When an equilibrium system is cooled, the temperature decreases. The system will shift backward (left) in an endothermic reaction and forward (right) in an exothermic reaction.

84
Q

07.07 Le Châtelier’s Principle

A catalyst lowers the activation energy for both the forward and the reverse reactions in an equilibrium system, so it has no effect on the equilibrium position of a system. (2 points)

True

False

A

True

85
Q

07.07 Le Châtelier’s Principle

According to Le Châtelier’s principle, how will an increase in pressure affect a gaseous equilibrium system? (3 points)

  1. Shift it toward the products
  2. Shift it toward the reactants
  3. Shift it toward the side with higher total mole concentration
  4. Shift it toward the side with lower total mole concentration
A

4. Shift it toward the side with lower total mole concentration

86
Q

07.07 Le Châtelier’s Principle

What change would shift the equilibrium system to the left?

A(g) + B(s) + Energy ⇌ 3C(g) (3 points)

  1. Adding more of gas C to the system
  2. Heating the system
  3. Increasing the volume
  4. Removing some of gas C from the system
A

1. Adding more of gas C to the system

87
Q

07.07 Le Châtelier’s Principle

What stress will shift the following equilibrium system to the left?

2SO2(g) + O2(g) ⇌ 2SO3(g); ΔH= −98.8 kJ/mol

(3 points)

  1. Decreasing concentration of SO3
  2. Decreasing volume
  3. Increasing concentration of SO2
  4. Increasing temperature
A

4. Increasing temperature

88
Q

07.07 Le Châtelier’s Principle

A chemical equilibrium between gaseous reactants and products is shown.

N2(g) + 3H2(g) ⇌ 2NH3(g)

How will the reaction be affected if the pressure on the system is decreased? (3 points)

  1. It will shift toward the reactant side as there is lower pressure on the reactant side.
  2. It will shift toward the product side as there is higher pressure on the product side.
  3. It will shift toward the reactant side as there are a greater number of moles of gas on the reactant side.
  4. It will shift toward the product side as there are a fewer number of moles of gas on the product side.
A

3. It will shift toward the reactant side as there are a greater number of moles of gas on the reactant side

89
Q

Nitrogen dioxide gas is dark brown in color and remains in equilibrium with dinitrogen tetroxide gas, which is colorless.

2NO2(g) ⇌ N2O4(g)

When a light brown colored mixture of the two gases at equilibrium was moved from room temperature to a higher temperature, the mixture turned dark brown in color. Which of the following conclusions about this equilibrium mixture is true? (3 points)

  1. This reaction is exothermic because the system shifted to the left on heating.
  2. This reaction is exothermic because the system shifted to the right on heating.
  3. This reaction is endothermic because the system shifted to the left on heating.
  4. This reaction is endothermic because the system shifted to the right on heating.
A

1. This reaction is exothermic because the system shifted to the left on heating

Honors Chemistry (8 decks)

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