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AP Chemistry > Unit 7- Liquids & Solids > Flashcards

Unit 7- Liquids & Solids Flashcards

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1
Q

Explain the difference between an intermolecular force and an intramolecular force. Give specific examples of each.

A

Intramolecular Force: forces occurring inside molecules between atoms; ex: Polar covalent bonds in H2O
Intermolecular Forces: forces between molecules; ex: One CO molecule forming a dipole-dipole force to another CO molecule.

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2
Q

Describe everything about London Dispersion

A

The random movement of electrons creates an instantaneous or temporary dipole in a molecule. This temporary dipole creates an induced dipole in a neighboring molecule thus producing an intermolecular attraction called London Dispersion Forces (LDFs); all molecules have electrons, so all molecules have LDFs; the more the electrons, the stronger the LDF- Also, a chain of atoms has stronger LDFs than branched atoms.

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3
Q

Describe everything about Dipole-Dipole

A

Electrostatic attraction of polar molecules: be sure to draw molecule to determine polarity. Dipole-Dipole forces are stronger than LDFs; only polar molecules have dipole-dipole forces; greater the dipole magnitude (difference in electronegativity) gives rise to stronger dipole-dipole forces.

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4
Q

Describe everything about Hydrogen Bonding

A

A type of dipole-dipole force that occurs between hydrogen and a very electronegative atom (F,O,N), Small size of hydrogen allows a strong electrostatic force. Strongest IMF; molecules must be polar with hydrogen DIRECTLY bonded to F,O, or N; the greater the difference in electronegativity, the stronger the H-Bond. The more hydrogen bonding sites, the stronger the H-Bond.

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5
Q

Define Evaporation and its trend related to IMFs

A

Liquid to Gas (Evaporation is a cooling process); Inverse Relationship

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6
Q

Define Heat of Vaporization and its trend related to IMFs

A

Heat energy required to vaporize one mole of a sample (kJ/mol-Endothermic); Direct Relationship

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7
Q

Define Condensation and its trend related to IMFs

A

Gas to Liquid (Exothermic- Heat must be released); Direct Relationship

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8
Q

Define Vapor Pressure and its trend related to IMFs

A

Pressure exerted by the vapor above a liquid once the vapor-liquid equilibrium has been established; Inverse Relationship

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9
Q

Define Boiling Point and its trend related to IMFs

A

Temperature at which the vapor pressure is equal to the external pressure; Direct Relationship

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10
Q

Define Surface Tension and its trend related to IMFs

A

Energy required to break through the surface of a liquid or to spread out a film- resistance of a surface to increase its surface area; Direct relationship

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11
Q

Define Capillary Action and its trend related to IMFs

A

Spontaneous rising of a liquid in a narrow tube (caused by cohesive and adhesive forces); Polar liquids tend to display capillary action

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12
Q

Define Viscosity and its trend related to IMFs

A

Measure of a liquid’s resistance to flow; Direct relationship

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13
Q

At room temperature, Cl2 is a gas, Br2 is a liquid, and I2 is a solid. Explain the trend.

A

As the mass increases, so does the electron count and the strength of the IMFs (LDFs). Thus, I2 with the most mass and electrons would have the strongest IMFs.

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14
Q

At room temperature, CO2 is a gas and CS2 is a liquid. Why is this reasonable?

A

Both molecules are nonpolar and only have LDFs. Since CS2 has the most number of electrons, it will have stronger LDFs.

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15
Q

If H2O were a linear molecule, could it have hydrogen-bonding interactions?

A

No, since it would be nonpolar.

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16
Q

What is the difference between a covalent bond and hydrogen bonding?

A

Covalent bonds arise from a sharing of electrons between nonmetals to have filled orbitals. These are 100x stronger than H-Bonds which are only electrostatic attractions.

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17
Q

Which should have stronger hydrogen bonding, NH3 or H2O?

A

H2O because of its greater difference in electronegativity.

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18
Q

Why is the motor oil more viscous than water? Does motor oil have a greater surface tension than water?

A

Stronger IMFs. Motor oil would have a higher surface tension.

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19
Q

Explain how molecules of a liquid can go into the vapor state if the temperature is below the boiling point.

A

The pressure of the atmosphere is low or some molecules have a high enough Kinetic Energy.

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20
Q

Why does a summer rainstorm lower the temperature?

A

The water evaporates quickly by absorbing heat.

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21
Q

Ethyl Chloride boils at 12 degrees Celsius. When it is sprayed on the skin, it freezes a small part of the skin and thus serves as a local anesthetic. Explain how it cools the skin.

A

The alcohol absorbs energy as it quickly evaporates.

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22
Q

Given a sample of water at 90 degrees Celsius and a sample of water at 30 degrees Celsius, in which one does the temperature change at a faster rate when both are allowed to evaporate?

A

The 90 degrees H2O changes at a faster rate.

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23
Q

A beaker of a liquid with a vapor pressure of 350 torr at 25 degrees Celsius is set alongside a beaker of water, and both are allowed to evaporate. In which liquid does the temperature change at a faster rate? Why?

A

Since its vapor pressure is higher than the water, the liquid evaporates at a faster rates because its IMFs are more easily broken.

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24
Q

What is implied by the word “equilibrium” in equilibrium vapor pressure?

A

At equilibrium vapor pressure, molecules are condensing and vaporizing at the same rate.

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25
Q

What is the difference between boiling point and normal boiling point?

A

Boiling point can be at any temperature and pressure while the normal boiling point is the temperature when the vapor pressure is equivalent to 1 atm.

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26
Q

A liquid has a vapor pressure of 850 torr at 75 degrees Celsius, is the substance a gas or a liquid at 75 degrees Celsius and 1 atm pressure?

A

It is a gas because the vapor pressure is greater than the pressure of the atmosphere.

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27
Q

How can the boiling point of a pure liquid be raised?

A

The boiling point can be raised by increasing the atmospheric pressure.

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28
Q

The boiling point of water in Death Valley, California, is about 100.2 degrees Celsius. Why is the actual boiling point higher than the normal boiling point?

A

Since the pressure of the atmosphere increases, the boiling point increases.

29
Q

Propane is used as a fuel to heat rural homes where natural gas pipelines are not available. It is stored as a liquid under normal temperature conditions, although its normal boiling point is -42 degrees Celsius. How can propane remain a liquid in a tank when the temperature is well above its normal boiling point.

A

The tank is pressurized high enough so that Propane remains a liquid.

30
Q

The normal boiling point of Neon is -246 degrees Celsius and that of Argon is -186 degrees Celsius. What accounts for this order of boiling points.

A

Since it has more electrons, Argon has stronger LDFs.

31
Q

The normal boiling point of HCl is -84 degrees Celsius and that of HBr is -70 degrees Celsius. Why is this order reasonable on one account but opposite from what one would expect from a consideration of polarity? Which trend is more important in this case?

A

HCl has a greater difference in electronegativity so one would think its greater polarity would give rise to stronger IMFs. Since HBr has much more electrons, it LDFs are however stronger than HCl’s dipole-dipole interactions.

32
Q

Describe everything about Spectroscopy/Spectrophotometry

A

It is the study of absorption and emission of light in matter; Compounds can absorb and emit light at different wavelengths;

33
Q

Define Transmittance

A

The amount of photons passing through a sample

34
Q

Define Absorbance

A

The amount of photons absorbed by the sample

35
Q

Describe everything about Beer’s Law

A

It gives us meaning for absorbance; A= elc where A=Absorbance (Unitless) L= Path Length (usually 1cm) C= Concentration in Molarity E= Molar extinction coefficient (1/Mcm); Every compound has a unique molar absorptivity constant at a given wavelength; Beer’s Law allows us to determine concentration from absorbance; Concentration and Absorbance show a linear relationship with the slope of eL

36
Q

What information does the phase diagram provide?

A

Temperature and pressure at each phase (solid, liquid, and gas) and also boiling points, freezing points, density, and etc.

37
Q

Define Triple Point

A

the temperature and pressure where the solid, liquid & gaseous phases of a substance exist at the same time (often in equilibrium)

38
Q

Define Critical Point

A

the last point where the gas can be compressed to form a liquid. Beyond this point, the boundaries between the liquid and gaseous phase disappear.

39
Q

Define Critical T&P

A

the last point where the liquid & gas exist at separate phases.

40
Q

Define Supercritical Fluid

A

Above this point (Critical Point), liquid and gas are indistinguishable from each other. Many useful properties result like fine tuning density or dissolving substances.

41
Q

What are all of the phase changes?

A

MFVCSD: Melting- Freezing- Vaporization- Condensation- Sublimation- Deposition; Melting (Solid to Liquid)- Freezing (Liquid to Solid)- Vaporization (Liquid to Gas)- Condensation (Gas to Liquid)- Sublimation (Solid to Gas)- Deposition (Gas to Solid)

42
Q

What does the negative slope along the interface between solid and liquid tell us?

A

It tells us that the liquid form of water has a higher density than the solid form.

43
Q

How can a phase diagram be used to justify the ease with which an ice skater glides across the ice?

A

It can be explained by the fact that the pressure exerted by their skates melts a small portion of the ice that lies beneath the blades.

44
Q

How is the greater density of liquid water life saving for aquatic life in the winter in cold regions?

A

Because if the solid form was more dense than the liquid form, the bottom of the ocean would be solid water an eventually the oceans would freeze over and everything would die.

45
Q

How are critical temperature and room temperature related?

A

If the critical temperature is higher than the room temperature, the substance can be liquified; if the critical temperature is lower than the room temperature, the substance can’t be liquified.

46
Q

What is the slope and intercept of the Clausius-Clapeyron Equation?

A

-DeltaHvap/R and intercept C

47
Q

What are the steps for calculating the vapor pressure of an unknown using the Clausius- Clapeyron Equation?

A

1) Plug right side in the calculator and solve 2) Multiply the known pressure by “e” to the power of the right side.

48
Q

Define Alloys

A

An alloy is a blend of metallic elements prepared by mixing the molten components and then cooling the mixture to produce a metallic solid.

49
Q

What are solid solutions classified as?

A

Substitutional and Interstitial.

50
Q

Define a substitutional solid.

A

A solid solution in which atoms of the solute metal occupy some of the locations of the solvent metal atoms.

51
Q

Define an interstitial solid.

A

A solid solution in which the solute atoms occupy the interstices (the holes) between solvent atoms.

52
Q

What three criteria need to be met for a substitutional solid to be formed?

A

1) The atomic radii of the elements are within about 15% of each other. 2) The crystal structures of two pure metals are the same; this similarity indicates that the directional forces between the two types of atom are compatible with each other. 3) The electropositive characteristics of the two components are similar; otherwise compound formation, where electrons are transferred between species, would be more likely.

53
Q

Explain the process of the formation of interstitial solids.

A

Interstitial solid solutions are often formed between metals and small atoms that can inhabit the interstices in the structure. The small atoms enter the host solid with preservation of the crystal structure of the original metal and without the transfer of electrons and formation of ionic species.

54
Q

Distinguish between crystalline and amorphous solids. Give some properties and examples of each.

A

Crystalline solids are regular repeating structures- they have a sharp MP/BP- they cleave along specific angles- examples are CuSO4, diamond, and sugar.

Amorphous solids have no regular array (arranged randomly)- they have no sharp MP- they cleave at irregular angles- examples are coal, rubber, plastic, and glass

55
Q

Define Lattice Point

A

Intersection of two or more grid lines where atoms and molecules are formed

56
Q

Define Unit Cell

A

the most basic repeating structure of a crystalline sold.

57
Q

Define Lattice Energy

A

the energy released when a crystal forms from atoms.

58
Q

What is the ranking of the strength of IMF, ionic bonds, and covalent bonds?

A

Both ionic and covalent bonds are stronger than IMFs.

59
Q

What are the types of crystalline solids?

A

IMMNG (Ionic-Molecular-Metallic-Network-Group 18)

60
Q

What are the properties of ionic solids?

A

Lattice points are occupied by ions- strong attraction between oppositely charged ions- High MP, High Hardness, rigid, conducts electricity when dissolved/molten.

61
Q

What are the properties of molecular solids?

A

Lattice points are occupied by molecules- IMF bonding- Low MP, Low Hardness, Nonconducting

62
Q

What are the properties of metallic solids?

A

Lattice points are occupied by metal nuclei- mobile sea of electrons that nuclei attract- High MP, Malleable, Ductile, Conductive

63
Q

What are the properties of network solids?

A

(Know SiO2, diamond, quartz, and graphite) Lattice points are occupied by nonmetals- covalent bonds between nonmetals- Very Hard, High MP, Generally nonconductive

64
Q

What are the properties of Group 18 solids?

A

Lattice points are occupied by nonmetal noble gas- extremely weak dispersion forces- Low MP, Nonconducting

65
Q

Why is ice less dense than water?

A

The crystal lattice structure of H2O causes water molecules to spread out in comparison to the liquid form.

66
Q

A solid is hard, brittle, and electrically nonconducting. Its liquid and aqueous solution form conducts electricity. Classify this solid.

A

Ionic.

67
Q

A solid is soft and has a low melting point. The solid, liquid, and aqueous solution do not conduct. Classify this solid.

A

Molecular

68
Q

A solid is very hard and has a high melting point. Neither the solid nor liquid conducts. Classify the solid.

A

Network

AP Chemistry (8 decks)

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