What is Energy?
What is Energy? 1. Energy – The ability to do work. 2. Potential energy – stored energy in chemical bonds. 3. Average Kinetic energy – energy of motion. Forms of energy 1) Thermal 2) Chemical 3) Light 4) Nuclear 5) Electromagnetic Activation Energy Activation energy – energy needed to start a chemical reaction. Examples: Heat – flame to start barbecue Mechanical – squeezing ice pack to mix chemicals Physical Behavior of Matter 1) Matter – anything that has mass and occupies some volume. 2) Phases of matter: a) Solid(s) – Definite volume, definite shape, rigid, fixed, regular, crystalline structure. b) Liquid(l) – Definite volume, no definite shape. Takes the shape of its container. c) Gas(g) – No definite shape and no definite volume. Takes the shape of the container uniformly. Gas molecules move in constant, random, straight-line motion. Heat and Temperature Heat Temperature MEASUREMENT of Average KE energy of a substance. Units for measuring temp. 1. Kelvin (K) 2. Celsius (0C) 3. Absolute zero = zero k 1. Kinetic energy of Temp. at which particles stop molecules. moving 2. Heat comes from the Conversion scale (table T) random movement of molecules in a substance. K = 0C + 273 3. Heat flows from source to a 0C = K - 273 sink. (high to low) Convert the following Temperatures Table T 1. 110oC K 383 K 2. -35oC K 238 K 3. 300K oC 27 oC 4. 75K oC 5. 500K 0C -198 oC 227 oC K = oC + 273 oC = K - 273 Convert the following: use table T 500 C to K K = 50 +273 = 323K 238 K to 0 C 0C = 238 – 273 = -350C Which cup of coffee has the greatest average Kinetic Energy? Explain 80 oC 90 oC 70 oC 85 oC How can we describe matter changing phase when energy is added or removed? CONSERVATION of ENERGY - In any chemical reaction or physical change, energy is conserved. 1) Energy can be changed from one form to another. (PE KE) (KE PE) 2) Energy can be transferred from one body or system to another. Heat is transferred from a source to a sink. 3) BUT……THE TOTAL AMOUNT OF ENERGY REMAINS CONSTANT!!!!!!! Units of Energy Joules 1. The joule is the SI unit for energy symbolized (J). 2. The amount of heat or energy needed to increase the temperature of a substance. Example: 1. The amount of heat required to raise 1 gram of water by 1 Celsius degree is 4.18 J 2. One KJ = 1000 Joules Heating and Cooling Curves System + Surroundings = Universe 1) System – what we want to study or examine. The reaction in a flask: NH4Cl(s) + H2O(l) 2) Surroundings – everything outside the system. 3) Universe – System + Surroundings. Surroundings Surroundings System System Surroundings Heat Transfer Exothermic 1. Temperature of the surroundings is measured 2. Increase in Temperature 3. Heat is released from the system to the surroundings 1. Reactions that release energy. 2. Heat is a product. Endothermic Sulfuric acid and sugar Phases of Matter (water) 1. Temperature of the surroundings is measured 2. Decrease in Temperature 3. Heat is absorbed from the surroundings to the system 1. Reactions that absorb energy. 2. Heat is a reactant. Exothermic or Endothermic? Why? Endothermic or exothermic? How do you know? 1) KNO3(s) + 34.89 kJ H2O K+ (aq) + NO3- (aq) •Endothermic •Heat is a reactant (left side of arrow) •Bonds Broken •Temperature of surroundings decreases 2) 4Al(s) + 3O2(g) 2Al2O3(s) + 3351 kJ •Exothermic •Heat is a product (right side of arrow) •New Bonds Formed •Temperature of surroundings increases Exothermic or Endothermic? Why? Exothermic or Endothermic? Why? Increase in energy Solid Liquid Gas Summary: PHASE CHANGES Energy is absorbed Particles spread out and move faster Melting SOLID Vaporizing SUBLIMATION: SOLID TO GAS LIQUID Freezing Condensing Energy is released Particle speed slows down and the particles move closer Gas Exothermic or Endothermic? Why? Aim: Does the Temperature of Water change during a Phase Change? BBC - GCSE Bitesize: Exothermic and endothermic reactions 1. Exothermic reactions release energy from the system (reaction occurs here) to the surroundings. INCREASE in TEMPERATURE 2. Endothermic reactions have the system absorbing energy from the surroundings. DECREASE in TEMPERATURE. Concept Check: 1) Which sample of water contains particles having the highest average kinetic energy? (1) 25 mL of water at 95°C (2) 45 mL of water at 75°C (3) 75 mL of water at 75°C (4) 95 mL of water at 25°C 2) Which change is exothermic? a)Freezing of water b)Melting of iron c)Vaporization of ethanol d)Sublimation of iodine 3) As the temperature of a substance decreases, the average KE of its particles a)Increases b)Decreases c)Remains the same 4) In a laboratory where the air temperature is 22°C, a steel cylinder at 100.°C is submerged in a sample of water at 40.°C. In this system, heat flows from (1) both the air and the water to the cylinder (2) both the cylinder and the air to the water (3) the air to the water and from the water to the cylinder (4) the cylinder to the water and from the water to the air Units of Energy Joules 1. The joule is the unit for energy symbolized (J). 2. The amount of heat or energy needed to increase the temperature of a substance. Example: 1. The amount of heat required to raise 1 gram of water by 1 Celsius degree is 4.18 J 2. One KJ = 1000 Joules Heating and Cooling Curves Phase Changes at Constant Temperature Heat of Fusion Q=mHf 1. Solid Liquid (endothermic) 2. Solid Liquid (exothermic) Heat of Vaporization Q=mHv 1. Liquid Gas (endothermic) 2. Liquid gas (exothermic) Specific Heat Heat capacity of a substance 1. Amount of heat needed to increase the temperature of one gram of a substance by 1 degree Celsius. 2. Because substances have different compositions, each substance has its own specific heat. Specific Heat of Water 4.18 J/g 0C or 4.18 J/gK Which has a higher specific heat, pure water or concrete? Water = 4.18 joules /g 0C Concrete = 0.84 joules /g 0C Specific heat formula Table T Q = mcΔT 1) q is the heat gained or lost (J) 2)m is the mass of the heated substance (grams) 3)c is the specific heat capacity (J/g 0C) 4)ΔT (pronounced delta T) is the change in temperature. 5)ΔT = (Tfinal - Tinitial) in 0C or K How many joules are absorbed by 30.0g of water when its temperature is raised from 20.0 0C to 40.0 0C? Formula for measuring specific heat Table T 1. Unit for q = 2. Unit for mass = 3. Unit for specific heat = 4. Unit for change in temperature = 1) Heat is added to a 200. gram sample of H2O(s) to melt the sample at 0°C. Then the resulting H2O(ℓ) is heated to a final temperature of 65°C. a) Determine the total amount of heat required to completely melt the sample. b) Calculate the total amount of heat required to raise the temperature of the H2O(ℓ) from 0°C to its final temperature. c) Compare the amount of heat required to vaporize a 200. g sample of H2O(ℓ) at its boiling point to the amount of heat required to melt a 200. g sample of H2O(s) at its melting point. 1) The temperature of a sample of water changes from 10°C to 20°C when the sample absorbs 418 joules of heat. What is the mass of the sample? (1) 1 g (3) 100 g (2) 10 g (4) 1000 g 2) A 36-gram sample of water has an initial temperature of 22°C. After the sample absorbs 1200 joules of heat energy, the final temperature of the sample is (1) 8.0°C (3) 30.°C (2) 14°C (4) 55°C Change in State - YouTube Time
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