A student is heating water to cook pasta. She notices it begins to boil at 100°C. What is the temperature of the water if she keeps boiling it for 15 more minutes?
Write potential or kinetic beside each of the following. A stretched rubber band___________ A cookie ________ A glass on a shelf ___________
A falling leaf _______
What are endothermic and exothermic processes?
What has a higher heat capacity, water or Al metal? Explain your answer.
What is the formula to calculate heat energy?
Thermochemistry- Chapter 11
11.1 Thermochemistry
• Thermochemistry - is concerned with the ____heat changes____ that occur during chemical reactions.
• Energy – The capacity ___for doing work or supplying heat__. It comes in many forms. • __Kinetic_______ Energy: motion of particles
• ___Potential______ Energy: stored energy
EXAMPLE: When gasoline burns, the potential energy stored in its chemical bonds is released as KE to do work, such as moving car.
• Heat (q)- energy that transfers from one object to another because of a temperature difference between them___. Heat always flows from a warmer object to a cooler object. • System – the part of the universe on which you focus your attention
• _Surroundings____ – everything else in the universe
• Law of Conservation of Energy – in any chemical or physical process, energy is neither _created nor destroyed________.
EXAMPLE: Heat may be lost by the system, but it is not destroyed. It is transferred to the surroundings.
• Endothermic process- system __gains____________ heat from the surroundings • Exothermic process – system ___loses____________ heat to the surroundings • Endothermic or Exothermic process?
– Evaporating alcohol: endothermic – Leaves burning: exothermic – Boiling water: endothermic – Water cooling: exothermic – Melting ice: endothermic – Freezing water: exothermic
• Energy units: ___calorie (cal) and joule (J)_______
• Calorie – quantity of heat needed to raise the temperature of _1 gram of water 1° Celcius__ • A food Calorie is used in nutrition and is capitalized.
• 1 Calorie = __1000____ cal = _____1____ kcal
Example: If a label on a candy bar indicated it contains 180 Calories, that is really 180 kcal, or 180,000 calories! If “burned” the sugar and fat in the candy bar release 180,000 cal of energy.
• Heat capacity- amount of heat needed to increase the temperature of an object 1oC. It
depends on ___the mass of composition of the substance_____________________.
EXAMPLE: It takes more heat to increase the temperature of a large pot of water than a small cup of water. It takes more heat to raise the temperature of water than metal.
• C = q (cal or J)
[mass (g)][T(°C)]
• Specific heat units ___ J/g°C______
• A low specific heat is matter that loses or gains heat quickly (Tiles on space shuttle) • A high specific heat is matter that loses or gains heat slowly (water)
• Water has a uniquely __high_____ specific heat compared to other substances.
• To calculate the heat energy required for a temperature change, use the following formula: ____q=mC T __________________
– q = energy – m = mass – C= specific heat
– ∆T = change in temperature
EXAMPLE: How much energy is required to heat an iron nail with a mass of 7.0g from 25oC
until it becomes red hot at 750oC?
equation: q = mC∆T
First, write out your givens. T = 752°-25°=727°C
q = mCT
q = (7.0g)(0.46J/g°C)(727°C) q = 2340.94 or 2300 J
EXAMPLE: If 5750 J of energy is added to a 455g piece of glass at 24.0ºC, what is the final temperature of the glass?
q = mC(Tf-Ti)
5750J = (455g)(0.50J/g°C)(Tf – 24.0°C)
Tf = 49°C
EXAMPLE: A 30.0g sample of an unknown metal is heated from 22.0 ºC to 59.2 ºC. During the process, 1.00 KJ of energy is absorbed by the metal. What is the specific heat of the metal?
q = mC(Tf-Ti)
1,000J = (30.0g)(C)(37.2°C) C = 0.896 J/g°C
EXAMPLE: If it takes 3590 calories to heat up a sample of water by 12.2 ºC, what is the mass of the water?
q = mC(Tf-Ti)
11.2 Calorimetry - the measurement of heat change for chemical and physical processes. • Heat released by system = ___heat absorbed__ by surroundings
• __Calorimeter__________- insulated device used to measure the absorption or release heat in a chemical or physical process.
• A bomb calorimeter is used to measure heat at a constant volume
EXAMPLE: A 25g sample of a metal at 75.0 ºC is placed in a calorimeter containing 25g of H2O at 20 ºC. The temperature stopped changing at 29.4 ºC. What is the specific heat of the
metal? 2 Steps to solve:
1. Find the q value for the water
2. Use the value from step one to find the C value of the metal.
You will have 2 sets of givens. Remember the final temperature and energy for both will be the same.
qwater = mC(Tf-Ti) qmetal = mC(Tf-Ti)
qwater = (45gH2O)(4.18J/g°C)(6.4°C) 1203.84J = (25g)(C)(75.0°-29.4°)
qwater = 1203.84J Cmetal = 1.1J/g°C
qwater = qmetal
EXAMPLE: A 25g sample of a metal at 75.0 ºC is placed in a calorimeter containing 25g of H2O at 20 ºC. The temperature stopped changing at 29.4 ºC. What is the specific heat of the
metal?
Enthalpy
Enthalpy (H) is the amount of _____heat in a system at a given temperature_____. Enthalpy Change: (same as energy)
______ H = q = mC∆T ________________________________ Exothermic reaction has ____-q___
Endothermic reaction has _____+q_______
__Thermochemical_______ equation includes heat changes. Physical state must be included!!!
– CH4(g) + 2O2(g) à CO2(g) + 2H2O(g) + 890KJ
• ∆H= -890 kJ
• Exothermic (energy released)
– 2H20 + 241.8 KJ à 2H2(g) + O2(g)
• ∆H= 241.8 kJ
• Endothermic (energy absorbed) What is
EXAMPLE
– CH4(g) + 2O2(g)
• Exothermic – energy is a ______product____ • Endothermic – energy is a _____reactant________ • ∆H is also called the ____heat of reaction_____.
11. 3 Heat in Change of State
The specific heat of water is ____ or 1.0cal/gºC The specific heat of ice and steam is _____ or 0.5cal/gºC q = mc∆T
Heat of __Fusion__ - heat required to melt 1 gram of solid
Heat of ___Solidification___ - heat released as 1 gram of liquid freezes Heat of fusion of water = _____6.01kJ/mol____ = heat of solidification
q = mHf
Heat of vaporization- heat required to vaporize ____1 mole of a liquid_____
Heat of ____Condensation_____- heat released as 1 gram of a gas condenses to a liquid. Heat of vaporization for water = _____40.7 kJ/mol___ = Heat of Condensation
q = mHv
q= ____moles x heat of phase change____
*Temperature is constant during a phase change
Heat
Heating/Cooling Curve shows energy changes. All matter follows curve when energy is added or lost.
• Horizontal portions of curve indicate a physical state change. Notice __that the temperature does not change___. However, there is a change in particle position resulting in a change in potential energy.
• Slope portions show ___temperature_ change which indicates a change in _____kinetic_____ energy as well.
EXAMPLE: How much heat, in calories, is needed to melt 150g of ice at 0 ºC? How many joules of
energy would it take to melt 23 mol water?
Which value is higher, the heat of fusion or the heat of vaporization?
EXAMPLE: How much heat, in calories, is needed to heat the liquid water in the above problem to 20. ºC?
EXAMPLE: A 50g sample of ice is held at -10 ºC. Will 270 cal of heat be sufficient to raise the temperature of the ice to 0 ºC?
EXAMPLE: How many calories are released when 36g of steam at 100 ºC condenses to water at 100 ºC?
EXAMPLE: How many calories are needed to convert 5.0g of ice at -15ºC to steam at 130ºC? (five problems)
11. 4 Standard Heat of Formation
• Standard heat of ____formation___ of a compound (∆Hfº)
• ∆Hfº of a free element in its ____standard state___ is zero.
• This is another way to calculate ∆H for a reaction.
• ∆H =___Hf°products - Hf°reactants__-use table of “Standard Heats of Formation”.
EXAMPLE: Calculate ∆H for the following reaction: CaCO3(s) à CaO(s) + CO2(g)
*First, make sure equations are balanced. *You will multiply the ∆H by the coefficient.
(all coefficients for this problem are 1, so do not need to worry about that)
EXAMPLE: Calculate the heat of reaction for the following reaction: 2H2(g) + O2(g) à 2H2O (g)
