chapter 5

Question 1

when thinking about phase changes, we should be thinking about?

Answer 1

• intermolecular forces
  
  • stronger intermolecular forces tend to be associated with solides
  • weaker intermolecular forves with gases
  • liquids are in between

Question 2

what is a solid?

Answer 2

• a structure with a tightly packed organization of atoms, such as ice, NaCl, and most metals
• they are characterized by a fixed volume, meaning that they do not expand and are not compressible and have a fixed shape and do not flow but their particles do vibrate in place
  
  • a solid can be crystalline or amorphous

Question 3

what is a crystalline vs amorphous solid?

Answer 3

• crystalline solids enhibit a regular arrangement of atoms (lattice structure like NaCl) which is extremely difficult to disrupt
  
  • the lattice energt of an ionic solid refers to the amount of energy required to separate the solid into its component cations and anions (large amount of energy)
• amorphous solids are solids that do not have a regular crystal structure (ex. glass)

Question 4

what are liquids?

Answer 4

• have a fixed volume, meaning they are not compressible
• do not have a fixed shape
• have viscosity- the resistance of the liquid to deformation by certain forces

Question 5

what are gases?

Answer 5

• not in a condensed phase
• lack a fixed shape or volume so density of a given gas is not constant
  
  • if the gas is forced into a smaller container, its density will increase as its particles pack more closely together

Question 6

what is it called when a substance goes from solid to liquid?

Answer 6

• melting or fusion

Question 7

What is it called when a substance goes from liquid to solid?

Answer 7

• freezing

Question 8

What is it called when a substance goes from liquid to gas?

Answer 8

• evaporation or boiling

Question 9

What is it called when a substance goes from gas to liquid?

Answer 9

• condensation

Question 10

What is it called when a substance goes from solid directly to gas?

Answer 10

• sublimation

Question 11

What is it called when a substance goes from gas directly to solid?

Answer 11

• deposition

Question 12

what are the 6 phase changes?

Answer 12

Question 13

what are the 3 endothermic phase changes that require input of heat?

Answer 13

• melting, evaporation, and sublimation
  
  • breaking of bonds or intermolecular forces

Question 14

what are the 3 exothermic phase changes that release heat into the environment?

Answer 14

• freezing, condensation, and deposition
  
  • bond formation or an increase in intermolecular forces

Question 15

what is the heat of fusion?

Answer 15

• the amount of heat that was used solely to disrupt these interactions and melt the ice

Question 16

the heat required to convert liquid to gas at constant temperature is?

Answer 16

• the heat of vaporization

Question 17

both the specific heat of vaporization, which describes how much heat energy is necessary to boil a substance and the specific heat of fusion, which describes how much heat energy is necessary to melt a substance are given in units of?

Answer 17

• energy per unit of mass or per mole

Question 18

what is specific heat capacity?

Answer 18

• the amount of heat required to raise the temperature of one unit mass (typically one gram or kg) of a substance by one degree
  
  • the specific heat capacity of water is 4.184J/g•°C

Question 19

specific heat capacity can be used in the following equation that related the heat applied to (or released by) a system to the temperature chage:

Answer 19

Q = mc deltaT

• Q is the heat (usually in J) applied or released by the system
• m is the mass
• c is the specific heat capacity of a substance
• deltaT is the change in temperature in ºC

Question 20

• at phase changes, use heat of fusion or vaporization
• between phase changes, heat will be related to temperature change so use Q = mc deltaT
  
  • try practice problem:

Answer 20

Imagine that we want to take 250g of ice from a -20°C freezer and turn it into water at 30°C under standard conditions. how much heat will this process equire?

• 3 steps
  
  1. heating the ice. we can use Q = mc delta T and approximate c_ice as 2 J/g•°C so Q = (250g)(2 J/g•°C)(20°C) = 10,000J
  2. melting the ice. we need the heat of fusion of water. Q = (250g)(334 J/g) = 83,500J
  3. heating the water. Q = mc deltaT and we can estimate c_water as 4 J/g•°C so Q = (250g)(4 J/g•°C)(30°C) = 30,000
  4. add all the joules up to gice 123,500J

Question 21

the phase of a substance is affected by temperature and pressure. these relationships are illustrated using?

Answer 21

• phase diagrams, which are usually drawn with pressure on the y-axis and the temperature on the x-axis
  
  • in phase diagrams, phase changes occur across the solid lines
  • an increase in pressure or decrease in temperature can convert gas to solid (deposition) gas to liquid (condensation), or liquid to solid (freezing) depending on the temperature
  • a decrease in pressure or increase in temperature can promote solid to gas (sublimation), liquid to gas (boiling), and solif to liquid (melting)
  • solid to liwuid changes are depicted by green lines. for most substances, the solid phase is denser than the liquid phase so an increase in pressure causes liquids to form solids, ice however is less dense than liquid water due to its unique crystal structure so an increasing pressure can melt ice into liquid water so the phase change line between solid and liquid tilts to the left for water

Question 22

phase diagrams also have a triple point and a critical point which mean what?

Answer 22

• the triple point is the point at which solid, liquid and gas are in equilibrium
• the critical point is the point that represents the end of the liquid-gas interface. above this point, matter exists as supercritical fluid, which possesses properties of both liquid and gas

Question 23

what is the concept of vapor pressure?

Answer 23

• describes what happens when a vapor, or gas, is in thermodynamic equilibrium with the liquid phase (or theoretically, the solid phase too)
  
  • there is always some interchange between the liquid and gas phase at the boundaries of a liquid
    
    • defined as the pressure exerted by the molecules of that substance that are in gas form in a closed system at a given temperature
    • vapor pressure increases with temperature, and the point at which the vapor pressure of a liquid is equal to that of the surrounding atmospheric pressure corresponds to the boiling point of a substance

Question 24

since gases are compressible, their density is affected by pressure and temperature, this means we have to specify certain pressure and temperature conditions in order to be able to compare density in a meaningful way. in general gases will specify conditions of standard temperature and pressure (STP) of:

Answer 24

• temperature = 0°C (273K)
• presure is 1 atm or 760 mmHg or 760 torr

Question 25

at STP one mole of any gas has a volume of?

Answer 25

22. 4 L
    * in general, one mole of any gas at a given temperture and pressure will have a constant volume thus equal volumes of any 2 gases under the same conditions will always contain the same number of gas molecules. (Avogardo’s law which states that the volume divided by the number of moles is constant)

Question 26

ideal (or theoretical) gas follows the parameters of kinetic molecular theory which states that:

Answer 26

• the average kinetic energy of the gas molecules is directly proportional to temperature
• the has particles have no volume
• gas particles do not exert forces on each other, although they do exert a force on the walls of the container. All collisions with the walls of the container are elastic

Question 27

gases deviate from ideal behaviour at?

Answer 27

• high pressures and low temperatures
  
  • so thet behave most ideally at low pressures and high temperatures

Question 28

always express temperature in Kelvin for?

Answer 28

• gas equations

Question 29

freezing of water occurs at?

Boiling of water occurs at?

what is body temperature?

what is room temperature?

Answer 29

• 0°C and 273K and 32°F
• 100°C and 373K and 212°F
• 37°C and 98.6°F
• 20-25°C and 68-77°F

Question 30

the relationship between Celsius and Farenheit is:

Answer 30

F = (9/5)C + 32

Question 31

what does Boyle’s law state?

Answer 31

• that pressure and volume of a gas are inversely proportional at a constant temperature
  
  • P₁V₁ = P₂V₂
  • PV = constant (pressure of a gas multiplied by its volume forms a constant at a certain temperature)

Question 32

what does Charles’ law state?

Answer 32

• that the volume and temperature of a gas are directly proportional under constant pressure
  
  • V₁/T₁ = V₂/T₂
  • V/T = constant
    
    • increase temperature increases volume, decreasing temperature will decrease volume

Question 33

What is the ideal gas law?

Answer 33

• PV = nRT
  
  • n is the number of moles of a gas
  • R is the universal gas constant (0.08206 L•atm/K•mol or 8.314 J/K•mol)

Question 34

The ideal gas law can be modified to model the behaviour of non-ideal gases in a modified form known as?

Answer 34

• Van der Waal’s equation
  
  • we need to account for:
    
    • the volume taken up by the molecule of gas (V_m - b) where V_m is the molar colume of the gas at a given temp and pressure and b is the volume occupied by the molecules per mole
    • the attractive forces experienced among the gas molecules, we add a/V²m to the P term where a is a constant that is different fir every gas and V_m is the molar volume of the gas
      
      • (P + a/V_m²)(V_m - b) = RT

Question 35

what is the mole fraction?

Answer 35

• X_gas whcih refers to the number of moles of given gas divided by the total moles of gas in the mixture:
  
  • X_gas = n_gas/n_total

Question 36

what is the partial pressure?

Answer 36

• the pressure that a gas in a mixture would exert if it took up the same volume by itself; this corresponds to Pgas_​
  
  • P_gas = X_gas • P_total
    
    • this equation states that the partial pressure of a gas is equal to the total pressure of the mixture multiplied by the mole fraction of that gas. Implicit in this is a relationship known as Dalton’s law which states that the total pressure of a mixture of gases is equal to the sum of the partial pressure of its components

Question 37

Solutions are?

Answer 37

• homogenous mixtures containing particles that are evenly distibutred throughout the solution

Question 38

when does dissolution occur?

Answer 38

• when particles, called solutes, are dissolved in another substance, usually fluid, called a solvent

Question 39

As solute is added to a solvent, the solution is considered saturated when the maximum amount of solute that can be dissolved has been added. upon heating, more solute can be dissolved in the solution, and then upon slowly cooling the solution, the same concentration of solute will remain dissolved in what is now considered?

Answer 39

• a supersaturated solution
  
  • crystals can form in supersaturated solutions with the addition of a small amount of solute, which creates a nucleation site for solute to precipitate and form a crystal. this process is called crystallization

Question 40

the concentration of a solution can be expressed in multiple ways. Solute concentrations are frequently expressed in terms of molarity (M), which is the number of moles of solute per liter of solution:

Answer 40

• molarity (M) = moles of solute/ litres of solution = mol/L

Question 41

what is molality (m)?

Answer 41

• is a measure of concentration that represents the number of moles of solute per kilogram of solvent:
  
  • molality (m) = moles of solute/kg of solvent (mol/kg)

Question 42

what is the normality (N)?

Answer 42

• the N of a solution is the number of equivalents of reactive species per litre of solution, for which we must define the reactive species
• normalirt is often used to express the concentration of H⁺ or OH^- ions produced in acid-base reactions
  
  • normality (N) = equivalents of solute/litres of solution

Question 43

an easy way to remember how to deal with normality for acids (where it most often shows up) is to use the simple equation:

Answer 43

• normality (N) = molarity (M) x # of protons in acid

Question 44

there are some important properties of solutions that are related to the total number of solute molecules present in the solution (regardless of whether they are the same or different compounds). these are known as?

Answer 44

• colligiative properties
  
  • there are 4 basic colligiative properties:
    
    • vapor pressure reduction
    • boiling point elevation
    • freezing point reduction
    • osmotic pressure

Question 45

the addition of solutes reduces the vapor pressure of a solvent in a relationship proportional to molal solute concnetration. In other words, the presence of solute decreases the concentration of solvent that exists in the gas phase and reduces evaporation. Vapor pressure reduction is expressed by?

Answer 45

• Raoult’s law
  
  • P = X_AP_A°
    
    • P is the vapor pressure of the solution
    • X_A is the mole fraction of the solvent
    • P_A° is the vapor pressure of the pure solvent

Question 46

Reducing the vapor pressure of a solution is equivalent to increasing the boiling point, since the boiling point is defined as the temperature at which the vapor pressure equals to the ambient pressure and vapor point increases with temperature. Boiling point elevation can therefore be calculated as follows:

Answer 46

• deltaT = iK_bm
  
  • deltaT_b is the boiling point elevation
  • i is the ionization factor
  • K_b is the boiling elevation constant
  • m is the molal solute concentration

Question 47

Noncolatile solutes also decrease the freezing point of solutions. freezing point depression can calculated as follows:

Answer 47

• deltaT_f - iK_fm
  
  • deltaT_f is the freezing point depression
  • i is the Van’t Hoff ionization factor
  • K_f is the freezing point depression constant
  • m is the molal solute concentration

Question 48

boiling point is elevated and freezing point is?

Answer 48

• depressed

Question 49

osmotic pressure relates to the principle of osmosis, the flow of solvent through a semipermeable membrane

osmostic pressure describes the pressure required to prevent osmosis, and is given by:

Answer 49

Π = MRT

Question 50

what is solubility?

Answer 50

• the extent to which various substances will dissolve in a solvent
  
  • solubility is an equilibrium process between the non-dissolved form of a substance and the dissolved form

Question 51

A high K_sp means that?

Answer 51

• a substance will readily dissolve in a solution while a low K_sp value means that it is insoluble

Question 52

What is the common-ion effect?

Answer 52

• The common ion effect describes the effect on ​equilibrium that occurs when a common ion (an ion that is already contained in the solution) is added to a solution. The common ion effect generally decreases ​solubility of a solute.

Question 53

higher temperatures generally favour the solubility of ionic compounds, exceptions has to do with whether the dissolution reactions in water are endothermic or exothermic:

Answer 53

• most dissolution reactions in water are endothermic, meaining that increased heat favours them
• some are exothermic, in which case the opposite effect can be expected

Question 54

gases in contrast are:

Answer 54

• more solubke at lower temperatures
  
  • this is because higher temperatures provide gases with more kinetic energy that they can use to escape the solution
  • additionally, pressure favours the solubility of gases

Question 55

soluble ionic compounds:

Answer 55

• alkali metals (Li⁺, Na⁺, Rb⁺, Cs⁺, Fr⁺) and NH₄⁺
• nitrates (NO_3^-) and chlorates (ClO₃^-)
• halides (Cl^-, Br^-, I^-) except compounds containinf Ag⁺, Pb²⁺, or Hg²⁺
• sulfates (SO₄^2-) except compounds containing Ca²⁺, Sr²⁺, Ba²⁺ or Pb²⁺

Question 56

insoluble ionic compounds:

Answer 56

• carbonates (CO₃^2-), phosphates (PO₄^3-), sulfides (S^2-) except compounds containing alkali metals or NH₄+
• hydroxides (OH^-), and metal oxides (O^-) except compounds containing alkali metals Ca²⁺, Sr²⁺, or Ba²⁺