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Gen. Chem. I > Module 3 > Flashcards

Module 3 Flashcards

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1
Q

What does kinetics tell you/measure?

A

How fast, temperature

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2
Q

What does thermodynamics measure/tell you?

A

How much

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3
Q

Definition of internal energy

A

The energy that’s intrinsic to a material

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4
Q

State functions

A

Tells us the state of the system.

Does NOT depend (mathematically) on how we got from A to B, just uses the initial and final conditions.

Value is dependent only on the state the system is in, never on how that state was reached.

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5
Q

Path functions

A

(Mathematically) depends on how we got from A to B. Measures what happens in between the initial and final conditions.

Value is dependent on how the state of the system was reached.

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6
Q

First Law of Thermodynamics

A

The energy of the universe is constant.

The total change in energy of a system consists of the heat transferred to (delta H or q) and the work done (w) on a system.

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7
Q

If something releases heat, it is…

A

Exothermic

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8
Q

If something absorbs heat, it is…

A

Endothermic

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9
Q

Definition of specific heat capacity

A

The quantity of heat required to increase the temperature of one gram of a substance by one degree C/K

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10
Q

Is delta H (change in enthalpy) a state function or a path function?

A

State function. It uses only the initial and final conditions.

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11
Q

What does calorimetry measure?

A

Heat flow

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12
Q

Is enthalpy (H) a state function or a path function?

A

It’s a state function.

Note that enthalpy values can’t be measured directly. Only changes in enthalpy can be determined.

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13
Q

If delta H is negative…

A

…the rxn is exothermic, indicates that heat has been lost from the system.

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14
Q

If delta H is positive…

A

…the rxn is endothermic, indicates that heat has been gained in the system.

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15
Q

What does delta H predict?

A

Whether reactions will be endothermic or exothermic.

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16
Q

What does entropy predict?

A

Whether reactions will be spontaneous or not.`

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17
Q

Second Law of Thermodynamics

A

The entropy of the universe is increasing to a maximum.

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18
Q

What happens to entropy if the temperature of the system increases while volume remains constant?

A

Entropy increases, delta S is positive

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19
Q

What happens to entropy if the volume of the system increases while temperature remains constant?

A

Entropy increases, delta S is positive

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20
Q

What happens to entropy if both the temperature and the volume of a system increase?

A

Entropy increases, delta S is positive

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21
Q

What three properties do liquids have?

A

Viscosity (resistance to flow)

Surface tension

Vapor pressure

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22
Q

What would make a liquid more viscous?

A

Decreasing the temperature, decreasing the IMFs present

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23
Q

What would increase a liquid’s surface tension?

A

Strengthening the IMFs

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24
Q

What would increase a liquid’s vapor pressure?

A

Weakening the IMFs, increasing temperature

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25
Q

When at low temperatures is a rxn spontaneous?

A

At low temps, delta G is roughly equal to delta H.

Rxn must be endothermic (delta H is negative) for it to be spontaneous.

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26
Q

When at high temperatures is a rxn spontaneous?

A

At high temps, delta G is roughly equal to -TdeltaS.

Rxn must increase in entropy (delta S is positive) for it to be spontaneous.

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27
Q

How does delta G predict spontenaity?

A

When delta G is negative, rxn is spontaneous.

When delta G is positive, rxn is NOT spontaneous, instead the reverse rxn is spontaneous.

When delta G = 0, the system is at equilibrium and neither the forwards nor reverse rxns are spontaneous.

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28
Q

Where does vapor pressure come from?

A

Particles at the surface of a liquid having enough kinetic energy so that they can escape the IMFs holding them in the liquid, vaporizing and entering the gas state above the liquid.

29
Q

What would decrease a liquid’s boiling point?

A

Weakening the IMFs, lowering the external pressure

30
Q

Breaking bonds is…

A

…endothermic.

31
Q

Forming bonds is…

A

…exothermic.

32
Q

When will a liquid boil?

A

When it’s vapor pressure is equal to the external pressure.

33
Q

Why would a rxn be endothermic?

A

More bonds are created than broken.

More moles of stuff are created/present at the end of the rxn.

34
Q

What is Gibbs Free Energy (G)?

A

Combines enthalpy and energy. Whether G is positive or negative allows us to predict a rxn’s spontenaity, and thus whether or not it will occur.

It also tells us how much of a particular type of work a thing/system can do.

35
Q

How do we find at what temperature a rxn will become spontaneous, aka its crossover temp?

A

Set delta G = 0, then solve. When delta G = 0, the system is at equilibrium. It marks the crossover temp/point from spontaneous to non-spontaneous.

deltaG = 0 = deltaH - TdeltaS

36
Q

Why is water a poor ideal gas?

A

Because it has a lot of IMFs holding it together, causing its collisions to be less than perfectly elastic.

37
Q

Why do gases turn into liquids (condense) when the ideal gas law states that they shouldn’t?

A

IMFs

38
Q

Which state of matter has more entropy: solid, liquid, or gas?

A

Gas

39
Q

What are London Dispersion Forces and where do they come from?

A

LDFs are the weakest IMF. They exist in al matter and arise from the transient fluctuating dipole in each electron cloud inducing a similar dipole in neighboring electron clouds.

40
Q

What kinds of IMFs are found in pure noble gas samples?

A

Only London Dispersion Forces

41
Q

What would increase a molecule’s LDF?

A

Increasing surface area and mass

42
Q

What are dipole-dipole interactions and where do they come from?

A

Occurs only in intrinsically polar molecules, molecules with a permanent dipole.

43
Q

What would increase dipole-dipole IMFs?

A

Having a larger individual/intrinsic dipole, caused by a larger electronegativity difference in the molecule

44
Q

What are hydrogen bonding IMFs and where do they come from?

A

Subset of dipole-dipole IMFs that arise because hydrogen is so small it can get really close to the next atom/molecule. Occurs only between hydrogen and either oxygen, nitrogen, or fluorine.

45
Q

What are ion dipole IMFs and where do they come from?

A

Occurs when a molecule with an intrinsic dipole interacts with a dissolved ion, causing the molecules to orient themselves in certain ways

46
Q

What would increase ion dipole IMFs?

A

Increasing the ionic charge of the dissolved ion

47
Q

The heat of condensation describes…

A

…condensing. Gas to liquid.

48
Q

The heat of vaporization describes…

A

…vaporizing. Liquid to gas.

49
Q

The heat of crystallization describes…

A

…freezing. Liquid to solid.

50
Q

The heat of fusion describes…

A

…melting. Solid to liquid.

51
Q

What properties are solids dependent on?

A

Type of particle, the geometry of the particles in arrangement, and the forces holding the particles together

52
Q

What are the four types of crystalline solids?

A

Ionic, metallic, covalent network, and molecular

53
Q

What type of crystalline solid conducts electricity as a liquid but not as a solid, has a very high melting point, and is very brittle?

A

Ionic solid

54
Q

What type of crystalline solid is good at conducting heat and electricity and is made out of electropositive elements?

A

Metallic solid

55
Q

What type of crystalline solid is not conductive, has a very high melting point, and is very hard?

A

Covalent network solid

56
Q

What type of crystalline solid is made of atoms from electronegative elements?

A

Covalent network solid

57
Q

What type of crystalline solid is malleable and ductile and has a variable hardness and melting temp?

A

Metallic solid

58
Q

What type of crystalline solid has a low melting point and is held together by IMFs?

A

Molecular solids

59
Q

What type of crystalline solid is made of nonmetal element compounds, in building blocks of molecules, and is held together by IMFs?

A

Molecular solids

60
Q

What is sublimation?

A

Phase change from solid to gas, without become a liquid in between

61
Q

What is a triple point?

A

The temperature and pressure at which the three phases of a substance coexist in thermodynamic equilibrium.

62
Q

What is a critical point?

A

The end point of a phase equilibrium curve, where phase boundaries vanish.

In the case of the liquid-vapor critical point, its the temp and pressure above which a substance exists as a supercritical fluid

63
Q

Order the four IMFs from strongest to weakest.

A

Ion-dipole, hydrogen bonding, dipole-dipole, LDFs

64
Q

What is deposition?

A

Phase change from a gas to a solid, without becoming a liquid in between.

65
Q

In a unit cell, corner-centered atoms count as…

A

1/8

66
Q

In a unit cell, face centered atoms count as…

A

1/2

67
Q

In a unit cell, edge centered atoms count as…

A

1/4

68
Q

In a unit cell, body centered atoms count as…

A

1

69
Q

Which solid crystalline structure features separate bonding and inter-particle force interactions?

A

Molecular covalent solids

Gen. Chem. I (7 decks)

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