Honors Text: 8.1-8.6, 17.1-17.4, 9.1-9.2 CP Text: NA

Honors Text: 8.1-8.6, 17.1-17.4, 9.1-9.2 CP Text: NA

Honors text: 8.1-8.6, 17.1-17.4, 9.1-9.2 CP text: NA

Thermochemistry

  • Thermochemistry: the study of energy (in the from of heat) changes that accompany physical & chemical changes
  • heat flows from high to low (hot cool)
  • endothermic reactions: absorb energy in the form of heat; show a positive value for quantity of heat (q > 0)
  • exothermic reactions: release energy in the form of heat; show a negative value for quantity of heat (q < 0)

Magnitude of Heat Flow:

  • units of heat energy:
  • the relationship between magnitude of heat flow, q, and temperature change, t, is:

q =

C =

t =

  • for a pure substance of mass m, the expression of q can be written as:

q =

m =

c =

t =

  • specific heat = the amount of heat that must be added to a stated mass of liquid to raise its temp. by 1C, with no change in state.

specific heat values (in J / g C):

Examples:

  1. How much heat is given off by a 50.0 g sample of copper when it cools from 80.0 to 50.0C?
  1. Iron has a specific heat of 0.446 J/gC. When a 7.55 g piece of iron absorbs 10.33 J of heat, what is the change in temperature? If it was originally at room temp. (22.0C), what is the final temperature?
  1. The specific heat of copper is 0.382 J/gC. How much heat is absorbed by a copper plate with a mass of 135.5 g to raise its temperature from 25.0C to oven temperature (420F)?

Calorimetry

  • to measure the heat flow in a reaction, it is carries out in a calorimeter.
  • it is possible to calculate the amt. Of heat absorbed or evolved in a reaction if you know the heat capacity, Ccal, and the temp. change, t, of the calorimeter:

Coffee Cup Calorimeter

  • the cup is filled with water, which absorbs the heat evolved by the reaction, so:

Example: When 1.00 g of ammonium nitrate, NH4NO3, is added to 50.0 g of water in a coffee cup calorimeter, it dissolves, NH4NO3 (s)NH4+(aq) + NO3-(aq), and the temperature of the water drops from 25.00C to 23.32C. Calculate q for the reaction system.

Bomb Calorimeter:

Examples:

  1. The reaction between hydrogen and chlorine, H2 + Cl22HCl, can be studied in a bomb calorimeter. It is found that when a 1.00 g sample of H2 reacts completely, the temp. rises from 20.00C to 29.82C. Taking the heat capacity of the calorimeter to be 9.33 kJ/C, calculate the amount of heat evolved in the reaction.
  1. When 1.00 mol of caffeine (C8H10N4O2) is burned in air, 4.96 x 103 kJ of heat is evolved. Five grams of caffeine is burned in a bomb calorimeter. The temperature is observed to increase by 11.37C. What is the heat capacity of the calorimeter in J/C?
  1. When twenty milliliters of ethyl ether, C4H10O. (d=0.714 g/mL) is burned in a bomb calorimeter, the temperature rises from 24.7C to 88.9C. The calorimeter heat capacity is 10.34 kJ/C.
  1. What is q for the calorimeter?
  1. What is q when 20.0 mL of ether is burned?
  1. What is q for the combustion of one mole of ethyl ether?

Enthalpy

  • Enthalpy = a type of chemical energy, sometimes referred to as “hear content”
  • exothermic reactions:
  • endothermic reactions:

Thermochemical Equations (8.4)

  • a chemical equation that shows the enthalpy (H) is a thermochemical equation.

Rules of Themochemistry:

Rule #1) The magnitude of H is directly proportional to the amount of reactant of product.

Example: H2 + Cl22HclH = - 185 kJ

Calculate H when 1.00 g of Cl2 reacts.

Example: When an ice cube weighing 24.6 g of ice melts, it absorbs 8.19 kJ of heat. Calculate H when 1.00 mol of solid water melts.

Example: Methanol burns to produce carbon dioxide and water:

2CH3OH + 3O2 2CO2 + 4H2O + 1454 kJ

What mass of methanol is needed to produce 1820 kJ?

Example: How much heat is produced when 58.0 liters of hydrogen (at STP) are also produced?

Zn + 2HCl  ZnCl2 + H2 + 1250 kJ

Rule #2) H for a reaction is equal in the magnitude but opposite in sign to H for the reverse reaction. (If 6.00 kJ of heat absorbed when a mole of ice melts, then 6.00 kJ of heat is given off when 1.00 mol of liquid water freezes)

Ex: Given:

H2 + ½ O2 H2OH = -285.8 kJ

Calculate H for the equation:

2H2O 2H2 + O2

Rule #3) The value of H for a reaction is the same whether it occurs in one step or in a series of steps.

for the overall equation is the sum of the H’s for the individual equations:

Hess’s Law:H = H1 + H2 + …

Example: Given:

C + O2 CO2H = -393.5 kJ

2CO + O2 2CO2H = -566.0 kJ

Calculate H for the reaction:

C + ½ O2 CO

Example: Find the heat of reaction (enthalpy) for the following reaction

NO + ½ O2 NO2H = ?

Given the following equations….

½ N2 + ½ O2 NOH = +90.4 kJ

½ N2 + O2NO2H = +33.6

Enthalpies of Formation

Hf = enthalpy of formation

  • usually exothermic
  • see table 8.3 for Hf value
  • enthalpy of formation of an element in its stable state =
  • these can be used to calculate H for a reaction
  • standard enthalpy change, H, for a given thermochemical equation is = to the sum of the standard enthalpies of formation of the product – the standard enthalpies of formation of the reactants.
  • elements in their standard states can be omitted:

2 Al(s) + Fe2O3(s)  2 Fe(s) + Al2O3(s)

  • the coefficient of the products and reactants in the thermochemical equation must be taken into account:

2 Al(s) + 3 Cu2+(aq)  2 Al3+(aq) + 3 Cu(s)

Example: Calculate H for the combustion of one mole of propane:

C3H8 (g) + 5O2 (g)  3CO2 (g) + 4H2O (l)

Example: The thermochemical equation for the combustion of benzene,

C6H6, is:

C6H6(l) + 15/2 O2(g)  6CO2 (g) + 3H2O(l)

H = -3267.4 kJ

Example: When hydrochloric acid is added to a solution of sodium carbonate, carbon dioxide gas is formed. The equation for the reaction is:

2H+ (aq) + CO32-(aq)  Co2 (aq) + H2O (l)

Calculate H for this thermochemical equation.

Thermodynamics

• THERMODYNAMICS = the study of energy changes that accompany physical and chemical changes.

• Enthalpy (H):

• Enthalpy change (H):

• Exothermic reactions/changes: release energy in the form of heat; have negative H values.

• Endothermic reactions/changes: absorb energy in the form of ehat; have positive H values.

• Changes that involve a decrease in enthalpy are favored!

Reaction pathways:

• Entropy (S): the measure of the degree of disorder in a system; in nature, things tend to increase in entropy, or disorder.

• all physical & chemical changes involve a change in entropy, or S. (remember that high entropy is favorable)

• enthalpy and entropy are DRIVING FORCES for spontaneous reactions (rxns that happen at normal conditions)

• it is the interplay of these 2 driving forces that determines whether or not a physical or chemical change will actually happen.

• Free Energy (G): relates enthalpy and entropy in a way that indicates which predominates; the quantity of energy that is available or stored to do work or cause change.

where:G =

H =

T =

S =

G: pos value means change is NOT spon.

G: neg value means change IS spon.

Relating Enthalpy and Entropy to Spontaneity

Example of reaction / H / S / Spontaneity
H2O(g) H2O(l)
H2O(s) H2O(l)

Examples:

1) For the decomposition of O3 (g) to O2(g): 2O3(g)3O2(g)

H = -285.4 kJ/mol and S = 137.55 J/mol•K at 25°C.

a) Calculate G for the reaction.

b) Is the reaction spontaneous?

c) Is H or S (or both) favorable for the reaction?

2) What is the minimum temperature (in °C) necessary for the following reaction to occur spontaneously?

Fe2O3 (s) + 3CO(g)  2Fe(s) + 3CO2 (g)

H = +144.5 kJ/mol; S = +24.3 J/K•mol

(Hint: assume G = -0.1 kJ/mol)

Liquids and Solids

  • Intermolecular forces: the forces between (among) individual particles (atoms, ions, molecules); they are weak relative to intramolecular forces (i.e. covalent and ionic bonds within a compound)
  • intermolecular forces are increasingly significant in this order:

Main Types of Intermolecular Attraction:

  1. Ion-ion interactions: the force of attraction between 2 oppositely charges ions, ionic bonding/attraction is quite strong, so ionic compounds have high melting points.
  • containing highly-charged ions have higher melting points than compounds containing univalent (1+ or 1-) ions.

Example: arrange the following ionic compounds in the expected order increasing melting and boiling pts:

NaF, CaO, CaF2

  1. Dipole-dipole interactions: occur between the + end of one polar molecular and the - end of another. (see figure 9.9)
  1. Hydrogen bonding: occurs in polar molecules that contain hydrogen that is bonded to one of the very electronegative elements O, N, or F.
  • a + hydrogen atom is attracted to an unshared pair of electrons on an O, N, or F atom on an adjacent molecule:

LIQUIDS and their properties:

  • surface tension: a measure of the inward forces that must be overcome to expand the surface of a liquid; molecules on the surface are attracted only toward the interior molecules.
  • cohesive forces: the forces that hold a liquid together.
  • adhesive forces: the forces between a liquid and another surface.
  • evaporation: the process in which molecules escape from the surface of a liquid; occurs more rapidly as temperature increases.
  • vapor pressure: the pressure exerted by the vapor of the liquid on its surface at equilibrium in a closed container; increases as temperature increases

***liquids with high boiling points have low vapor pressures and relatively strong intermolecular attractions.

  • boiling point:the temperature at which the vapor pressure of a liquid equals the applied (usually atmospheric) pressure

***the normal boiling pt = the temp at which a liquid’s vapor pressure = 760 mm Hg or 1 atm

Heat Transfer in Liquids and Solids

  • specific heat = the amt of heat that must be added to a stated mass of a substance to raise its temp by 1C, with no change in state.

Ex: How much heat is released by 250.0 g of H2O as it cools from 85.0C to 40.0C? (specific heat of water = 4.18 J/gC)

  • heat of vaporization: the amt of heat that must be added to 1 g of a liquid at its boiling point to convert it to vapor with no change in temp.
  • heat of vaporization of water = 2260 J/g

Ex: How much heat energy is required to bring 135.5 g of water at 55.0C to its boiling point (100C) and then vaporize it?

  • heat of fusion: the amt of heat needed to melt 1 g of a solid at its melting point.
  • heat of fusion of ice = 334 J/g
  • When substances change state, they often have different specific heats:

specific heat of ice = 2.09 J/gC

specific heat of water = 4.18 J/gC

specific heat of steam = 2.03 J/gC

Ex: How much heat energy is required to convert 15.0 g of ice at –12.5C to steam at 123.0C?

Phase Diagrams

  • a phase diagram shows the equilibria pressure-temperature relationship among the different phases of a given substance.

H2O CO2

AB =

AC =

AD =

  • triple point = Point ______

The point at which all 3 phases of a substance (solid, liquid, gas) can coexist at equilibrium.

  • critical point = Point ______

The combination of critical temperature and critical pressure.

  • critical temp =

The temperature above which a gas cannot be liquefied. (H2O=374ºC)

  • critical pressure =

The pressure required to liquefy a gas at its critical temperature. (H2O=218 atm)

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