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Do Now:. Why is CO 2 a gas and H 2 O a liquid at room temperature? What can we do to change H 2 O to a gas? How can we change CO 2 to a liquid?. Heat and the States of Matter. Homework: Read pages293 – 299 Answer the following questions on page 299 #4, 6, 7, 9 AND

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  1. Do Now: Why is CO2 a gas and H2O a liquid at room temperature? What can we do to change H2O to a gas? How can we change CO2 to a liquid?

  2. Heat and the States of Matter Homework: Read pages293 – 299 Answer the following questions on page 299 #4, 6, 7, 9 AND The even # questions on the handout

  3. Heat • Def: the total amount of kinetic energy of the particles in a sample of matter • Units: Joule (J) • Heat always flows from objects at higher temps to objects at lower temps, until both objects are at the same temperature

  4. Temperature • Def: a measure of the average kinetic energy of the particles in a sample of matter (not a form of ENERGY)

  5. Measuring Heat • Calorimeter is the device used to measure the quantity of heat lost or gained in a reaction • Endothermic – heat absorbed • Exothermic – heat released

  6. Specific Heat Capacity (c) Def: the amount of heat needed to raise the temp of 1 gram of a substance 1 ⁰C • All substances have different specific heat capacities depending on their composition • Water’s specific heat capacity is listed on Table B

  7. Calculating heat lost or gained q = m c D t Where, q = heat lost or gained (J) c = specific heat capacity (J/g⁰C) D t = change in temperature (⁰C) (final temp - initial temp)

  8. How much heat is required to raise 50.0 g of water from 5.00 ⁰C to 20.0 ⁰C? q = m = c = ti = tf = Q = mcDt = (50.0 g)(4.18J/g C) (20.0 C – 5.00 C) = 3138 J = 3140 J

  9. How much heat is required to raise the temp of 200.0g of water 24.5C?

  10. If 500.0 J of heat are added to 10.0 g of water at 30.0 C. What is the final temp of the water?

  11. Using a Calorimeter • While studying heat exchanges we define the experiment that is being performed as the system, and everything outside of the experiment as the surroundings.

  12. Law of Conservation of Energy • Energy can neither be created nor destroyed in a chemical reaction but may be converted from one form to another. Heat lost = Heat gained

  13. The picture shows a reaction to be carried out in a calorimeter. The heat released by the reactants (system) is absorbed by the surrounding water. Measuring the temperaturechange of the water will allow us to calculate the heat lost during the reaction. We can calculate the heat loss using q = mcDtfor the water.

  14. The following is a description of an experiment used to determine the specific heat of copper. • In a laboratory investigation, a 50.0-gram sample of copper is at 100.0°C in a boiling water bath.

  15. A Styrofoam cup with a lid is used as a calorimeter. • The cup contains 100.0 grams of distilled water at 23.2°C.

  16. The hot copper is poured into the cup of water, and the cup is quickly covered with the lid.

  17. A thermometer is inserted through the lid. • The coppe rand water are gently stirred in the cup. • The temperature is checked periodically. • The highest temperature noted is 26.3°C.

  18. In terms of energy flow, explain why the temperature of the water in the calorimeter increases. • Using the information given, complete the data table provided in your answer booklet.

  19. Calculate the number of joules of heat gained by the water.

  20. Calculate the specific heat capacity of the copper.

  21. Homework Go to Chemthink and complete the Particulate Nature of Matter section under the introduction section. Complete both the tutorial and the question set

  22. Phases of Matter

  23. The Kinetic Molecular Theory • Matter is composed of smaller particles • Physical properties determined by distance between particles, and forces (IMF) they exert on one another • Particles are in constant motion • Average kinetic energy is determined by temperature • Collisions between particles transfer energy without a loss of energy • As substances change from a solid  liquid  gas, the amount of motion of the particles increases  • As substances change from solid  liquid  gas, the particles have enough energy to slowly overcome the forces of attraction between particles

  24. Solids • have a definite shape, definite volume (incompressible) • particles of substance are closely packed together • particles vibrate • 2 types of solids • Crystals – particles of solid arranged in a repeating geometric pattern ex) NaCl, Diamond • Amorphous solids – particles are not arranged in any order ex) wax, glass

  25. Liquids • definite volume, no definite shape (assume the shape of container, relatively incompressible) • particles of a substance not held together tightly • particles are in constant motion (can slide past one another)

  26. Gas No definite shape, no definite volume (expand and contract, compressible) particles are widely separated Gases exert a pressure by striking the walls of a container (each collision exerts a force on the container) Measured using a barometer Air pressure measured torr the air pressure at sea level is 760 torr, 1 atmosphere, or 101.325 kPa Volume of a gas can change under different conditions of temp and pressure  Standard temp and press (STP) have been assigned OC and 101.3 kPa (1 atm)

  27. Phase Changes Solid / Liquid • Freezing – the physical change of state from a liquid to a solid - energy is released liquid  solid + heat energy • normal freezing point – the temperature at which a liquid changes to a solid at 101.3 kPa (normal atmospheric pressure) • temperature does not change during phase change

  28. Solid/Liquid phase changes • Melting – the change of state from a solid to a liquid - energy is added to solid, giving the particles sufficient energy to break free from other particles, no temperature change solid + heat energy  liquid • occurs at the same temp as normal freezing point

  29. Heat changes during melting/freezing • Heat of Fusion (Hfus) – the amount of heat energy needed to change a mass of solid into a certain mass of liquid at constant temp for H2O (s) Hfus = 334 J/g (table B) • when a mass of H2O (l) freezes, 334 J/g are released • when a mass of H2O(s) melts, 334 J/g are absorbed • to find the amount of heat energy lost or gained during melting or freezing use: q = mass x Hfus

  30. If 32 g of ice are melted at 0 C, how much heat has been absorbed by the ice? q = m x Hfus q = 32 g x 334 J/g q = 10688 J q = 11000 J

  31. How many grams of ice at 0C can be melted to water at 0C by absorbing 3347 J of heat? q = m x Hfus q = m Hfus 3347 J = m 334 J/g 10.0 g = m

  32. Why does a balloon stay inflated? (explain in terms of the behavior of the molecules that fill the balloon)

  33. Liquid/Gas • Vaporization– change from a liquid to a gas • Boiling- change of a liquid to bubbles of vapor that appear throughout - temp at which a liquid boils is known as normal boiling point • Evaporation – vaporization that takes place at the surface of a liquid - occurs at any temp - particles at surface with greater average kinetic energy overcome the intermolecular forces of attraction that keep particles in liquid state, and escape into gaseous state

  34. Liquid/Gas • Condensation – reverse of vaporization Gas  liquid + heat

  35. Heat calculations for boiling/condensation Heat of vaporization – amount of heat required to change a mass of a liquid to a mass of gas at constant temp * for water Hvap = 2260 J/g (Table B) at 100C q = mass x Hvap

  36. How many joules of heat are needed to change 20. g of H2O at 100C to steam at 100C? q = m x Hvap q = 20. g x 2260 J/g q = 45000 J

  37. But wait! There’s more Vapor Pressure – the pressure exerted by a gas from the liquid, when confined above the liquid at a given temperature • as temp increases, KE increases, • more particles in gaseous state, • vapor pressure increases • Table H

  38. Vapor pressure and boiling points • ****Boiling point – the temperature at which the vapor pressure of a liquid is equal to the atmospheric pressure • at high altitudes, atmospheric pressure is lower than at sea level  the boiling point is lower ex) Salt Lake City and Denver, the BP of water is 95C

  39. Heating Curve

  40. Cooling curve

  41. Sublimation Phase change from a solid to a gas Occurs in substances with weak intermolecular forces of attractions and high vapor pressures Ex) dry ice CO2(s) + energy → CO2(g) Deposition is the opposite process

  42. What would the heating curve look like for a substance that undergoes sublimation?

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