CHANGES IN MATTER & ENERGY

1. Chapter 5.1 CHANGES IN MATTER & ENERGY

2. Changes in Matter (v.1) • 3 types – Physical, Chemical and Nuclear • During any of these changes energy changes also occurs. • The study of the energy changes that occurs during changes in matter is called THERMOCHEMISTRY. It is a branch of Thermodynamics. • Complete Q#1 (a) – (f) on page 300 of the textbook.

3. Discuss answers • Gas barbecue operating • An ice cube melting in someone’s hand • White gas burning in a camping lantern • Wax melting on a hot stove • Zinc added to an acid solution in a beaker • Ice applied to an athletic injury

4. Changes in Matter (v.2) • From energy perspective, changes in matter are of 2 types – Exothermic & Endothermic • EXOTHERMIC: change in which heat is lost from the system • Example: reaction of zinc and sulphur • Reaction of bromine with aluminum • ENDOTHERMIC: change in which heat is absorbed into the system • Example: using a cold pack • Mixing ammonium thiocyanate and barium hydroxide Answer Q#5 (a) – (c) on page 300 of the textbook.

5. System and Surroundings • When studying transfers of energy, it is important to distinguish between the substance undergoing change, called the chemical system, and the system’s environment, called the surroundings. • A system is usually represented by a chemical equation. • Identify the system and surroundingsin each of the examples shown in the next slide.

6. Identify system and surroundings in the following examples: • Gas barbecue operating • An ice cube melting in someone’s hand • White gas burning in a camping lantern • Wax melting on a hot stove • Zinc added to an acid solution in a beaker • Ice applied to an athletic injury

7. Exothermic and Endothermic revisited Exothermic: • system  surroundings heat - causes an increase in temperature of surroundings Endothermic: • system  surroundings heat • causes a decrease in temperature of surroundings

8. System – Open, Closed or Isolated? • OPEN System is one in which both matter and energy can move in or out. • CLOSED System is one in which energy can move in or out, but not matter. • ISOLATED System is an ideal system in which neither matter nor energy can move in or out. • Answer Q#3 on page 300 of textbook.

9. Discuss Answers (Q#3 on pg#300) • Gasoline burning in an automobile engine • Snow melting on a lawn in spring • A candle burning on a restaurant table • Addition of baking soda to vinegar • A gas barbeque operating

10. Some other terms that you should know……. • Heat (Q) refers to the transfer of energy, and is measured in joules (J). • Temperature (T) – measures the average kinetic energy of particles that make up a system or substance; measured in ℃ or K. K= ℃+273.15 • Thermal Energy 0r Heat Energy – is random kinetic energy; all particles have heat energy by virtue of their motion; particles stop moving at 0 K (zero Kelvin)

11. ENERGY and its UNITS…… • Energy refers to the capacity of doing work or producing heat • Joule (J) is the SI unit of energy. • 1 Joule = Work done when a force of 1N acts over a distance of 1 m. Heat quantities in a chemical reaction can be quite large so chemists use kilojoule (kJ) • Older unit is calorie (cal) or kilocalorie (kcal) 1 cal = 4.184 J

12. How is heat related to mass? • Consider two samples of water: • A bath tub • A small bottle of water Which sample will take longer to get hot? the tub! • more water means it requires more energy to get hot! • Heat (q)absorbed is directly related to the mass of the object

13. How is heat related to temperature change? • Consider two samples of water: a saucepan filled with – • Cold water (10 ºC) • Warm water (30 ºC) Which sample will take longer to boil? Cold water sample! • Since more heat is required to raise the temperature of cold water to 100 ºC! • Heat absorbed is directly related to initial temperature of the sample.

14. Heat Capacity • Heat Capacity is the measure of how much heat is required to increase the temperature of an object by 1K or 1ºC. • The higher the heat capacity of the substance, the more energy is required to change the temperature. • Measured in J/ºC or J. ºC-1 • Heat capacity of a solid is generally less than that for a liquid, which is less than that for a gas. • Heat capacity depends both on the amount of the substance as well as the type of the substance.

15. Specific Heat Capacity • Specific heat capacity (c) refers to the amount of heat required to change the temperature of 1g of substance by 1 ºC. c = heat absorbed/mass x change in temperature • Units are J/g. ºC or J.g-1. ºC-1 • Specific heat capacity depends on the nature and state of the substance rather than on the quantity of the substance. • Refer to the table #1 given on page 301

16. Molar Heat Capacity • Molar heat capacity (C) refers to the amount of heat required to change the temperature of 1 mole of substance by 1 ºC. • Units are J.mol-1.K-1 or J.mol-1ºC-1 • A 3.1 g ring was heated using 10 J. The temperature of the ring rose by 17.9°C. Calculate specific heat and heat capacity. Is the ring pure gold? c=q/mT= 10 J /(3.1 g)(17.9C) = 0.18 J/gC (gold is 0.129 J/gC ; it is not pure) heat capacity = c.m = 0.18 J/gC x 3.1 g = 0.558 J/ C

17. HOMEWORK • Practice Q#8-13 on page 302 of the textbook.