Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

Chemistry > Reactivity 1. 1: Energetics > Flashcards

Reactivity 1. 1: Energetics Flashcards

You may prefer our related Brainscape-certified flashcards:
Periodic Table flashcards Chemistry 101 flashcards AP® Chemistry flashcards Organic Chemistry 101 flashcards
1
Q

a) Bond enthalpy
b) write an bond enthalpy in CH3Cl

A

a) the energy needed to break one mole of bonds in gaseous molecules under standard conditions
b) Cl3Cl (g) ➯ CH3 (g) + Cl (g), HBR (g) ➯ H(g) + BR (g)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

The average bond enthalpy

A

the average bond enthalpy is the energy needed to break one mole of bonds in gaseous molecules under standard conditions averaged over similar compounds

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

Exothermic

A

the energy released from the system to the surrounding loosing free energy with no external heat therefore, products are more stable than reactants . The surrounding becomes hotter (∆H is negative)

  • eg: Gas -> Liquid -> Solid, combustion, bond formation
  • Σ𝐸(𝑏𝑜𝑛𝑑𝑠 𝑏𝑟𝑜𝑘𝑒𝑛) > Σ𝐸(𝑏𝑜𝑛𝑑𝑠 𝑓𝑜𝑟𝑚𝑒𝑑)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

Endothermic

A

energy gained to the system from the surrounding. The surrounding becomes cooler, to form products less stable thatn reactants. (∆H is positive)

  • eg, bond breaking, photosyntehsis, solid->liquid->gas
  • Σ𝐸(𝑏𝑜𝑛𝑑𝑠 𝑏𝑟𝑜𝑘𝑒𝑛) < Σ𝐸(𝑏𝑜𝑛𝑑𝑠 𝑓𝑜𝑟𝑚𝑒𝑑)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

Temperature

A

Is the measure of average kinetic energy (motion Ek) of the particles in the substance

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

Heat

A

transfer of energy (Ep) as result of temperature difference, increasing its temperature/ quantity of energy transfer

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

energy

A

ability to do work

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

standard conditions for enthalpy changes ∆H𝜃

A

𝜃: the pressure of 100 kPa, concentrations of 1 moldm^3 for all solutions, and all the substances in their standard states at 298K (25celsius)

consistent intermolecular forces of the reactants and product

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

open system

A

energy and matter

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

a) Specific heat capacity
b) specific heat capacity of water

A
  • The energy required to raise the temperature of 1g of substance by 1K
  • 4.18 J-g-1k-1

heat capacity depends on the number of particles presemt in unit mass sample, therefore mass of the indiviaul particles
* the specific heat capacity of water is 4.18 J-g-1k-1
* the unit for specific heat capacity can be Jg-1k-1 or Jg-1C-1
* highest temperature happens at the lowest specific heat capacity

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Closed system

A

only energy is transferred to the surroundings

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

isolated system

A

matter nor energy can be transferred

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

standard bond enthalpy change of combustion ∆Hºc

A

the enthalpy change that occurs when one mol of a compound reacts with exess oxygen under standard condition of at 298ºK and 1.00 * 10^5 Pa

All combustion reactions are exothermic

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

What is the reason to difference in experimental and theoretical value of enthalpy change of combustion from the databooklet?

A

1) Decrease: heat loss form the sides of copper calorimeter
2) Decrease: incomplete combustion
2)Decrease: evaporation of alchol during experiment
3)Increase: the thermometer touches the calorimeter
4)random measurement error

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

enthalpy

A

The amount of heat energy contained in a substance. Stored in chemical bonds and intermolecular forces as potential energy.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

The standard enthalpy change of formation ∆𝐻𝑓𝜃

A

the energy change when one mole of a compound is formed from its elemts in their standard states under standard conditions (298 K (25 °C) and 1.00 × 105 Pa)

∆𝐻 𝜃 = ∑ ∆𝐻𝑓𝜃(𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠) − ∑ ∆𝐻𝑓𝜃(𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠)

  • (Elements have a ∆𝐻𝑓𝜃 value of zero) as it dosent take much energy to trun element into element/ most stable
  • x(solid) + y (gas) ➟ xy (gas)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

Method for calculating the energy released, in kJ, when two reactant react together given the thermodynamic equation and ∆𝐻𝜃

A

Step 1: Calculate how many moles of reactant there are
Step 2: compare 1 mole of ∆𝐻𝜃 in the thermodynamic equation to moles of reactant.
Step 3: Use ratios to find how many KJ are produced comapred to 1 mol.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

What are three factors that temperature change of an object depends on

A
  1. the mass of the object
  2. the heat added
  3. the nature of the substance
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

a) Equation for enthalpy change of combustion
b) what are assumption made on the combustion experiment?
c) How can you improve the calorimetry experiment
d) how can you reduce percentage uncertainty of an experiemnt?

A

1.
* all heat is transferred from the beaker to the water/surrounding → no heat loss (some transferd to metal or in the system)
* all solution has same specific heat capacity as pure water
* experiment not done in standard condition
* incomplete combustion (soot)
* the specific heat capcity of the beaker is negligble

    • insulate/ put a lid on
    • bomb calorimeter
    • accurate balance
    • large temperature different
    • large quanitity of water/sulfur
    • more precise division per degree on the therometer

The unit q: J
The unit of water: Jg-1k-1 or Jg-1c-1
The unit of temperature: kelvin or celsius
* The mass of water in grams can be placed by cm3

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

1) Define calorimetry
2) what are the 5 steps to finding the enthalpy change from calorimetry?

A

a) technique used to measure the enthalpy change of a reaction. Used for combustion, neutralization and displacement reaction.

b)
1. inital mass - final mass = m (H2O)
2. Tfinal - Tinitial = ∆T (H2O)
3. exploriate the line, T(recorded maximum temp/maxium temperature with heat loss)- initial temperature
4. calculate Q of water using Q = mc∆T
5. calculate ∆H using ∆H = - Q/n

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

1) Equation for standard enthalpy of neutralization ∆𝑯𝒏𝒆𝒖𝒕
2) What are assumptions to neutrallization reaction

A
  1. no heat loss from the system
  2. all the heat produced from neturalization reaction passes into the water
  3. the volume of solution is equal to the volume of water ( solution cm^3 ➟ g ➟ m(H2O) g)
  4. water has a density of 1gcm-3
  5. at standard temperature

  • neutralization reaction is always exothermic since more stable water is formed
  • temperature found from: (maximum tempeature allowing heat loss) - (recorded maximum temperature)
  • In neutralization reaction, the sytem is the compound and the surrounding is the water.
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

1) Define Hesses law
2) Draw the Hesses law reaction for 3C +4H2 +5O2

A

total change in chemical potential energy (enthalpy change) must be equal to the energy lost or gained by the reaction system due to the law of conservation of energy

must write its states

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

a) find ∆Hºf and ∆Hºc
b) why is ∆Hºf and ∆Hºc different?

A

b)
Formation Enthalpy is the energy it takes to form a compound from elements
Bond Enthalpy: an approximation of the energy it takes to form a compound from anything by calculating the bond energies (bond energies of the reactants - bond energies of the products)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

1) define average bond enthalpy
2) What is the equation for avergae bond enthalpy?

A

1) energy needed to break one mol of bonds in gaseous molecules under standard conditions averaged over similar compounds

2) 𝚫𝑯 = 𝚺𝑬 (𝒃𝒐𝒏𝒅𝒔 𝒃𝒓𝒐𝒌𝒆𝒏/endo/reactants)−𝚺𝑬 (𝒃𝒐𝒏𝒅𝒔 𝒇𝒐𝒓𝒎𝒆𝒅/exo/products) = 𝒌𝑱∗𝒎𝒐𝒍−𝟏

  • single bonds < double bonds < triple bonds
  • bond forming and bond breaking
  • ignores intermolecular forces since its in gaseous state
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Q

1) What are the products of incomplete combustion and complete combustion?
2) how do you distinguish different carbon chain compounds through combustion analysis
3) Why is carbon monoxide an dangerous pollutant?

A

1) complete: CO2 +2H2O
incomplete: CO + 4H2O (+c if asked)
2) find carbon percentage content in the hydrocabon chain. The more soities flame has the largest carbon content.
3) CO affects oxygen uptake in the blood.
It is absorbed by the lungs and binds to haemoglobin in red blood cells more effectively than oxygen and si bound permanently. This prevents oxygen from being transported around the body.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
26
Q

Which reaction has higher bond enthalpy?

A

bond enthalpy of HCl is higher than HBr. Cl is higher in group than Br therefore has higher polarity/electronegativity/reactivity therefore harder to break the bonds

  • related to displacement reaction
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
27
Q
A
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
28
Q

Draw hesses diagram for

A
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
29
Q

Define born harber cycle

A

a) the lattice enthalpy of an ionic soid is the energy required to completely seperate one mole of a solid ionic compound into its gaseous ions

  • compound can be made from energy change of formation or the born harber cycle
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
30
Q

a) Define standard enthalpy of ionization Δ𝐻 I.E
b) write an equation for first and second ionization energy of Ca

A

a) the enthalpy change when one mol of electrons are mremoved from one mole of gaseous atoms at standard condition
b)
Ca(g) ➱ Ca+(g) + e-
Ca+(g) ➱ Ca2+(g) + e-

31
Q

a) Define standard enthalpy of atomization Δ𝐻at
b) wirte an equation Δ𝐻at of Br2(l)

A

a)The enthalpy change when one mole of gaseous atoms is formed from an elements in its standard state

b) 1/2 Br 2 (l) ➱ Br (g)

32
Q

a) Define standard enthalpy of ionization 1st E.A
b) write an equation for first and second affinity energy of O

A

a) The enthalpy change when one moles of electrons is gained by one mole of gaseous atoms at its standard state

b)
O(g) +e- ➱ O-(g)
O(g) +e- ➱O2-(g)

33
Q

a)Define lattice enthalpy Δ𝐻lat
c) write Δ𝐻latt for K2O

A

a) The enthalpy change when one mole of an ionic compound is broken into its constituients gaseous ions at its standard state
b) K2O (s) ➱ 2K+(g) + O2- (g)

has an unit of KJmol-1

34
Q

Draw born-Harber cycle for NaCl

A

  • if there are two charges being gain, you need to write first electron affinity (exothermix) + second electron affinity (endothermic)
  • if there are two chraged being lost, you need to times the lattice enthalpy by 2
  • ionization energy is always edothermic +ve charge
  • the unit for lattice enthalpy is KJmol-1
35
Q

1) What are two factros that affects the lattice enthalpy in a ionic strucure.
2) explain why the experimental lattice enthalpies are always larger than theoritical values

A
  • higher charge (more cation) and smaller radius ➨ higher charge density ➨ higher lattice enthalpy ➨ higher meltig point
    2)
  • ions with larger radius show more covalent tendency as they can more eaisly polarize.
  • Theoretical values are based on an ionic
    model. It does not take into account any additional covalent contributions to the
    bonding. The smallest lattice enathlpy value has the lowest electronegative value.
36
Q

formula for lattice enthalpy

A

(constant x charge on cation and anion)/sum of ionic radii

37
Q

What does it mean by high entropy and low entropy

A
  • spontaneous processes always occur with an increase of entropy in the universe
  • high enthalpy ➟ disordered state ➟ more ways
    distributing/dispsersing energy and matter (gas) ➟ endothermic
  • low enthalpy ➟ordered state vless ways of dispersing energy and matter (solid) ➟ exothermic
38
Q

1) Define entropy ∆𝑆𝜃 of system
2) Equation for the total entropy changes
3) Equation for the entropy changes of reaction

A

measure of how number of ways of distributing of total available energy among the particles in a system in standard state. measured in K/K*mol

  • unit of JK-1mol-2
  • A perfect lattice strucutre at absolute zero has zero entropy, therefore indivisual atoms’s antropy can never be 0 (impossible).
  • DIfference between reactant entropy and product entropy can be -ve or +ve.
  • If the overall entropy differece is -ve ➟ reactant entropy>product’s entropy ➟ more disordered (towards gasesous state) ➟ more spontaneous
  • If the overall entropy is difference is +ve ➟ reatanct entropy< product’s entropy ➟ more ordered (towards solid state) ➟ less spontaneous
39
Q

3 things to look if the equation’s entropy is positive or negative

A

1)
* States: gas ➟ solid to gas (+ve entropy)
* gas to solid (-ve entropy)

2)
* number of mols: decrease in mol ➟ -ve entropy as mass is divided into bigger chunks and moves more slowly to behave like solid
* number of mols: increase in mol ➟ +ve entropy as mass is divided into smaller chucks to behave like gas

3)
* state and mols same: entropy of 0

  • can determine if an reaction has postive or negative entropy value
40
Q

How does ∆𝐻𝜃, 𝑇 and ∆𝑆𝜃 affects spontaenous reactions?

A

  • Some reactions will always be spontaneous depending on ∆𝐻𝜃, 𝑇 and ∆𝑆𝜃
41
Q

1) define Gibbs free energy
2) equation for Gibbs free energy
3) predict the gibbs free energy at high temperature and low temperture

A

1)energy free to do work rather than leave the system as heat. It measures the spontaneous reaction.
3)
* high temperature of the system: gibbes -ve therefore spontanous
* low temperature of the system: gibbes +ve therefore non spontaneous

  • the gibbs equation accounts for only the gibbs energy of the system
  • increase in entropy increases makes a reaction spontaenous
  • increases in temperature makes a reaction spontaenous
  • decreases in enthalpy makes the reaction spontaenous
  • a. If ∆𝐺 𝜃 < 0 the reaction is spontaneous (reaction will happen) ➟ foward reaction: spontaneous, reverse reaction: non spontaneous
  • b. If ∆𝐺 𝜃 = 0 the reaction is at equilibrium ➟ rate of foward reaction is equal to the rate of reverse reaction
  • a. If ∆𝐺 𝜃 > 0 the reaction is not spontaneous (reaction will not happen) ➟ foward reaction: non-spontaneous, reverse reaction: spontaenous reaction

∆Gº: change in Gibbs free energy (kJ/mol-1)
∆H: enthalpy change (kJ/mol-1)
T: temperature (K)
∆Sº: entropy (convert to kJ/Kmol -1 from JKmol-1)

42
Q

entropy ∆𝑆𝜃 of surrounding

A
  • High temperature: Entropy of system increases with temperture due to higher kinetic energy of particles
  • this decreases the entropy of the system, as heat is released
  • the entropy of the surrounding increases as the hat given by the reaction inceases the dispersal of the surrounding
  • ∆H in the gibbs energy formula indirecrly measure the entropy change of ∆𝑆𝜃
43
Q

∆G reaction equation

A

  • calculated from ∆G formation values
  • elements have gibbs free energy of 0
44
Q

Why does low temperature favours exothermic and high temperature endothermic?

A
  • Low temperature: exothermic - If the temperature is decreased for a system in equilibrium then the equilibrium will shift to increase the temperature
  • High temperature: endothermic- readjusts itself to counteract (partially) the effect of the applied change by “trying” to lower the temperature by absorbing heat and a new equilibrium is established.
45
Q

reaction quotient/ equillbrilium constant

A

measures the relative amounts of prodict and reactant present during reaction at particular point of time. Increase in Q values means there are more products.

46
Q

What does the graph indicate?

A

*the rate of reaction from products to reactants increases
* the rate of reaction from reactants to products decreases
* The rate reaches equal and there is no longer any change in the concentration of products and the reactantsn

47
Q

Draw a graph when pure substance

A
  • The gibbs free enrgy decreases (becomes spontaneous) as the reaction moves towards the equillbrium
  • equal amount of reactants and products at the beginning of the reaction
  • Q=1: equillbrium: gibbs free energy =0, K=1
48
Q

Draw a graph when the mixture when gibbs energy is negative

A
  • negative gibbs energy therefore the reactant is spontaenous
  • more pure reactant with high gibbs energy than pure product, higher spontaneous reaction from reactants into product (big increase in Q value) until it reaches minimum gibbs energy
  • gibbs energy < 0, q>1, mainly products, K>1
49
Q

Draw a graph when the mixture when gibbs energy is positive

A
  • negative gibbs energy therefore the pure reactant are spontaenous
  • more pure products with high gibbs energy than reactants, therefore products are turned into reactants (decrease in Q) until it reaches minimum gibbs energy
  • ther is equillbrium micture containing mostly reactant
  • gibbs energy > 0, q<1, more reactants, K<1
50
Q

Equation of equillbrium constant

A
51
Q

What is the disadvantage of coal

A
  • Coals contain 90% of carbon content by mass therefore little energy is released per unit of mass
  • cheap and abundant, long lifespare, converted to liquid feuls and gases, safer than nuclear power
  • high carbon composition producing particulates (soot,Co2,long carbon chain), acid rain from SO2 & CO2, difficult transport, diffult mining, formation of photochemical smog

  • gases fuel such as hydrogen and methane have low energy density
  • hydrogen has higher heat capacity than methane
  • liquid hydrocarbon fuel have high energy density
52
Q

How are fossils fuels formed

A
  • energy from the sunlight
  • absence of oxygen/reduction from trapped animals and plants making hydrogen carbons (CHNS) under pressure and heat
53
Q

What is the trend of energy as number of carbon number in alkane increases
What are the impacts to increasing carbon contents?

A

Trend
* longer carbon chain means more carbon added to the chain, therefore carbon contents increases and hydrogen content decreases by percentage.

  • Longer the carbon contents, more carbon reacts with oxygen increasing the mass of Co2 produced when 1g of feul is burnt.
  • Less energy is produced as carbon contents increases as longer alkanes are likely to experience greater incomplete combustion and energy is lost as heat.
  • low carbon content (H2, CH4➤natural gas) are gases, higher carbon contents are liquids (petroleum, coal)
  • gases with low carbon contents have lower energy density and releases more energy, liquid high carbon contents and have higher energy density and releases less energy.
  • overall, higher carbon number alkanes (petroleum, coal) has higher carbon dioixde pollution (mass) and low energy released where as lowc carbon number like natural gas has less carbon dioixide pollution and releases more energy per mass.
  • : natural gas (high specific heat) > protroleum & coal (low specific heat) > biofuel (lowest specific heat)
    Impacts
  • high percentage of carbon content is likely to experience incomplete combustion so produces many carbon monixide
  • natural gas (clean/mostly methane/small carbon conentent)< coal(dirty/high carbon content
54
Q

1)what 3 types of fossil fuel

A
  1. coal, natural gas, petroleum
55
Q

What is the disadvantage of crudeoil/petroleum

A

cudeoil:
*unsaturated alkanees with N,O,S, from plantss and sea animals, high carbon contents
* crude oil useful for raw material such as chemical feedstock, used in cars with high enthalpy density. Sulfur impurities can be easily removed.
* limited lifespan, acid rain, global warming, carbonmonoxide pollutants from gasoline in combustion engines as incomplete combustion,

56
Q

What is the disadvantage and advntatges of natural gas

A

natural gas
* decomposition of crude oil and coals.
* natural gas is mostly methane gas (S+N), less impurities (lowe carbon contents
* easily transported in pipelines and pressurized contaners, no acid rain
* limited, global warming, can explode due to leaking

57
Q

1) Explain the trend between fossil fuels and greenhouse effect

A
  • number of carbon dioxide increases as greenhouse effect therefore global warming
58
Q

1) How does greenhouse effect work
2) what are the effects of greenhouse effect

A

1)
* solar shortwave UV light from the sun not abosorbed but reflected by the atmosphere containing carbon dioxide to to be rediated as infraed light
* Carbon dioxide IR ative. Stretch oppsite and same direction, bends up and down. Dipole changes storing infraed light, releases and returns to carbon dioxide.
2) agriculture, rising sea level, melting polar ice caps, changes in desertifaction and fish habitats

59
Q

1)What is a example of renewable Biofuel
2)write chemical formula that provides energy for the biofuel

A

1) wood, ethanol (liquid biofuel)
2) photosyntheis reaction: 6Co2 + 6H2O = C6H12O6 + 6O2

60
Q

1) write equation for complete combustion of ethanol
2) state diadvtagnes and advantages of ethanol
3) What are other gases of biofuel that is made with limited supply of oxygen

A

1) C6H12O6 →2C2H5OH +2CO2
2)
Pro: renewable, lower CO, NO emission, country uses less fossil fuels
cons: corrosion, makes hydrogen bond with water seperating from rest of hydrocarbon in the fuels
3) methanol, mathane

61
Q

What are advantages and disadvatages of using biofuel

A

a) pros: cheap, replanted, sustainable
Cons: high cost of harvesting and transportation, fertilizers, low specific energy than fossil fuels, photosythesis is not efficient

62
Q

1) Draw the fuel cell
2)Deduce the two half-reactions that occur at the anode and the cathode ni the hydrogen-oxygen fuel cel with aqueous potassium hydroxide as the electrolyte.
3)Deduce the two half-reactin for methane fuel cell

A
63
Q

what are the steps to oxidation and recution hlaf reaction

A

1) mol
2) add H2O ( for oxygen)
3) add H+ (for H)
4) add electrons (for charge)

  • oxidation half equation: loss of electorns
  • reduction hald eqaution: gain of electrons
  • oxidation half equation + reduction hald eqaution= balanced redox half equation (number of mols must be multiplied so that only # of electrons must cancle out)
64
Q

find the percentage error

A
65
Q

a) heat chage formula
b) enthalpy change formula

A

b) o If the temperature of the system/compound increased but temperature of the surrounding/water decreased, then the reaction is endothermic and 𝑞 must be positive

o If the temperature of the system/compound decreased but temperature of the surrounding/water DECREASED, then the reaction is exothermic and 𝑞 must be negative

66
Q

What temperature on a temperature time graph would you use?

A

Maximum temperature, not recorded maximum temperature

67
Q

3 rules to determining if the equation has positive or negative entropy

A
  • Look at the state: liquid to gas
68
Q

Define electron affinity

A

The amount of energy released when one mole of electrons is gained by one mole of atoms of an element in the gaseous state to form one mole of gaseous ions

69
Q

Deduce the equation for complete combustion of hydrocarbon CxHy

A

  • law of conservation of mass, the number of atoms need to be equal before and after the chemical equation
70
Q

Find the constant for equillbirum constant

A
71
Q

Formula from finding temperature of an spontaneous reaction/ boiling point from enthalpy and entropy values

A
72
Q

Which conditions corresponds to a system at equillbirum

A

The system will be at equillbrium, k=0, when the entropy of the stem is at a maximum and the gibbs energy of the system is at minmum.

  • Gibbs free energy is a measure of the energy put into the system for a chemical reaction to occur. At equilibrium, the Gibbs free energy is at a minimum because the system has reached a state of balance where the forward and reverse reactions occur at equal rates. This means that the system is at its most stable state, and the no further reaction needs to proceed.
  • Entropy is a measure of the amount of disorder or randomness in a system. At maxiumum entropy, the system has reached a state of maximum disorder and randomness, and there is no further potential for change in the system without an external influence.
73
Q

calculate the percentage uncertainty

A

(0.5+0.5/15) x100 =6.7%

Chemistry (10 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions