University:
High School
Course:
High School Physics
Academic year:

2024
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235

Pages:

67
Author:

larbiwxcvbn

Thermochemistry, Reaction Rates, & Equillibrium Reaction Rates The rate at which chemical reactions occur Reaction Rates • RXN rate = rate at which reactants change into products over time. – This tells you how fast a reaction is going • Some reactions are fast: TNT exploding • Some are slow: nail rusting Reaction Rate ex: CH4 + O2  CO2 + H2O - At the start there are 100 % reactants & 0% products. - As the reaction proceeds, amount of reactants ↓ & products ↑ time Collision Theory In order for a reaction to occur 3 things must happen: 1. reactants must collide 2. while favorably oriented/ positioned 3. must combine with enough energy to disrupt current bonds (activation energy) Reaction Rates are affected by: 1) Temperature 2) Surface Area 3) Concentration 4) Pressure (gases only) 5) Catalysts 6) Nature of reactants Temperature • ↑ temperature = ↑ the rate of reaction • ↑ temperature means particles move faster and are more likely to collide. • Ex: Egg cooks faster with higher heat • Ex: refrigerators slow down bacterial growth Surface Area • ↑ the surface area = ↑ reaction rate. • It ↑ the exposure of the reactants to each other = ↑ chance for collisions. Demo: Coffee Creamer and Bunsen Burner Concentration • The more particles that are around in a given volume, the more likely they are to collide and react. Pressure (on gases only) • molecules in a gas are very spread out. • By ↑ the pressure, you squeeze the molecules together so you will ↑ the frequency of collisions between them. Catalysts- speed up chemical reactions by lowering activation energy TiO2 How they are written: CO2 + H2O  H2CO3 • They do not get used up in the reaction & can be used over and over. • Ex: enzymes in your digestive track help break down food. Chemical Equilibrium Chemical Equilibrium Ex: Cl2(g) + H2 (g)   2 HCl (g) + heat Chemical equilibrium the rate of the forward reaction is equal to the rate of the reverse reaction. • The concentration of reactants and products is constant. Equilibrium Example 500 cars/ hour 500 cars/ hour Equilibrium Ex: -if 500 cars/hour cross the Golden Gate bridge in each direction then the rate of reaction is equal. -The # of cars in each city remains constant, but isn’t the same. Le Chatelier’s Principle • if a stress is applied to a system at equilibrium, the reaction will shift one direction to relieve the stress. • Stresses: – Change in concentration – Change in temperature – Change in pressure Change in Concentration & Equilibrium Ex: N2 (g) + 3 H2 (g)   2 NH3 + heat Stress Relief • increase N2 Use excess N2 How • Reaction will shift , (right) Change in Temperature & Equilibrium N2 (g) + 3 H2 (g)   2 NH3 + heat Stress • add heat Relief Use up excess heat How • Reaction will shift to the  left • because that direction is endothermic (heat absorbing). Going right would increase the heat. Change in Pressure & Equilibrium (Gases only!) • Increase of pressure causes shift in the direction that results in the fewer number of moles of a gas • if pressure is relieved, the reaction will shift in the direction that produces more moles of a gas N2 (g) + 3 H2 (g)   2 NH3 + heat (4 moles) (2 moles) Stress Relief • add pressure How lower pressure – Reaction will shift to the , (right), forward rxn to reduce the pressure (less moles of gas in products) Thermochemistry Study of energy transfers that occur during chemical & physical changes Energy is absorbed or released in all chemical reactions & physical changes • Chemical reactions: - Energy absorbed to break bonds - Energy given off when bonds form • Physical (aka phase) changes: Energy absorbed or released – when molecules move faster/ slower – break/ form attractions between molecules Endothermic & Exothermic Endothermic & Exothermicdirection that heat flows • Exothermic= process that releases energy to the environment. Ex: 2 H2 + O2  2 H2O + 483.6 kJ (heat) • Endothermic= process where energy is absorbed from the environment. Ex: 2 H2O + 483.6 kJ (heat)  2 H2 + O2 Endothermic Processes- heat is absorbed from environment •In cold packs, ammonium nitrate, NH4NO3 , is often used because it absorbs a lot of heat when it dissolves in water. Heat is absorbed by the ice pack and reduces swelling Heat is taken from the environment into the cold pack. Endothermic Reaction (a graph view) • Endothermic: Reaction gains heat from the environment • Products have more heat energy than the reactants Exothermic Processes- releases heat into the environment •In hot packs, O2 reacts with Fe inside the hot pack. •This chemical reaction releases heat. Heat is being released from the chemical reaction into the environment Exothermic Reaction (graph view) • Exothermic: reaction loses heat to the environment • Products have less heat energy than the reactants. Heat Heat = energy that transfers from a higher temp. object to a lower temp. object • Heat always moves from a hotter object  colder object • Heat moves until thermal equilibrium is established and the temperatures are the same. • Ex: getting used to the swimming pool Heat will enter the ice cube from the environment and cause the ice cube to melt. Temperature vs. Heat • Temp (°C, K) = avg kinetic energy in molecules (how fast they are moving) • Heat (Joules/ kJ or Calories) = energy that can be transferred between objects – based on TOTAL amount of energy (kinetic & potential) Ex: Heat Vs. Temperature • Temps are same b/c both contain molecules moving at the same avg. speed But… • Bath contains more heat because it has more molecules that can transfer more heat Which contains more heat? 40° C Beaker A 40° C Example of heat vs. Temp. Which contains more heat? 40° C TEA!!!! 40° C • Both contain molecules moving at the same avg. speed (temp.) •Heat content is much more for the tea because it has more molecules that can transfer more heat How your body senses hot/ cold:  Thermoreceptors= special sensory cells that detect gain or loss of heat COLD!!  Heat is leaving your foot and going to the cold pack, thereby making your skin “feel cold” “Feel hot” when warm object transfers heat to your cells. Hot/ Cold Vs. Temperature • Temperature= is absolute, measurement of average speed of molecules. • Hot/ Cold= terms are relative and reflect heat loss or gain. • Ex: same temperature, yet 1 person might say “hot” other might say “cold” How Thermometers Work • Does not measure temp. directly • Measures the effect of temperature on physical property of material used in the thermometer. Mercury expands when it absorbs heat, so it rises up. Calculations involving Specific Heat Units for Measuring Heat (q) • Joule = SI system unit for measuring heat: • Calorie = heat required to raise the temperature of 1 gram of water by 1 Celsius degree 1calorie 4.18Joules Calculation involving Heat q=m·c·ΔT q = amount of energy lost or gained during a chemical reaction Mass (grams) “delta” T is change in temperature Specific Heat Reported in Joules or calories Final Temp – Initial Temp Amount of heat it takes to raise the temperature of 1 gram of a substance by 1 C. Specific Heat • Is unique to every substance due to the way they absorb energy. – Energy can make molecules move faster (temp. increases) – Energy can break attraction between molecules (does not raise temperature) •The plastic slide has a higher “specific heat” (aka heat tolerance) than the metal slide. •It takes more heat to raise the temperature of 1 g of the plastic by 1 °C than it does the metal. Which slide would you rather go down in summer? Specific Heat of some common subtances Water has an extremely high specific heat When you heat 1 gram of water it take 4.18 joules of energy to get the temperature to increase 1° C. Water (solid) 2.1 Water (gas) 1.9 Pools cost a lot of money to heat because of this. Sample Calculation • Gold has a specific heat of 0.129 J/(g×°C). How many joules of heat energy are required to raise the temperature of 15 grams of gold from 22 °C to 85 °C? Answer: q = m · c · Δ T q = (15 g) 0.129 J/(g×°C) (85 °C – 22 °C) q= 120 Joules of heat is absorbed by gold This is an endothermic reaction because energy is entering from outside the system. Bomb Calorimeter • Measures the heat absorbed or released in a physical or chemical change • Amount of heat lost/gained is calculated using the specific heat of water (4.18 J/g ºC) Homemade calorimeter •Determines how many joules of heat are released by a burning peanut (or any food item). •The heat from the peanut is transferred to water and that amount of heat is calculated for the water using q = m · c · Δ T Phase Changes Phase changes in water Phase Change Names “Fusion” “Vaporization” Phase Diagram (water) Why phase changes require heat energy Heat breaks/ weakens the intermolecular attraction between molecules. Calculations involving Phase changes Calculations with heat & phase changes 2 things can happen when you heat a substance: 1) heat can be used to speed up the molecules of the substance (temp. change) q = m · c · Δ T 2) energy is used to break the H- bonding between the molecules (substance can change phase). q = m ·Δ HFus (constant) or q = m ·Δ Hvap (constant) Δ H fus H2O = + 334 Joules/ g ΔH vap water= + 2260 Joules/ g Δ H fus H2O = - 334 Joules/ g ΔH vap water= - 2260 Joules/ g See notes from class… I did #9 on heating curve Wksht. Sample problem How much heat is required to melt 2 grams of ice at O C? • This is a phase change (no temp. change) • Use formula: q = m ·Δ HFus q= (2 g) (334 J/ g) q= 668 Joules to melt ice Heat energy being used to disrupt hydrogen bonds between water molecules. Sample problem How much heat is required to vaporize 2 grams of water at 10O C? • This is a phase change (no temp. change) • Use formula: q = m ·Δ Hvap q= (2 g) (2260 J/ g) q= 4520 Joules to vaporize Heat energy being used to disrupt hydrogen bonds between water molecules…all have to be broken to turn into gas Summary: Fusion and Vaporization for Water Gas ΔH ΔH vap water= + 2260 Joules/ g vap water= - 2260 Joules/ g Liquid H fus water = + 334 Joules/ g H fus water = - 334 Joules/ g Δ Δ Solid Sample Problem: How much heat in KJ is required to change 50 grams of ice at -10 ° C to a liquid at 40 ° C? 3 2 1 Step 1: Find heat it takes to raise ice from -10 ° C to 0 ° C (q = m · c · Δ T) q= (50g) (2.1 J/g ° C ) (10 ° C ) = 1050 J Step 2: Find heat to do a phase change from solid- liquid (q = m Δ H Fus) q = (50 g) (334 J/ g) = 16700 J Step 3: Find heat it takes to raise temp from 0 ° C to 40° C (q = m · c · Δ T) q= (50 g) (4.18 J/g ° C ) (40 ° C )= 8360 J Total (add steps 1+2+ 3)= 27150 J or 27.150 KJ Hydrogen bonding in Ethanol (alcohol) @ -186° C • Notice how there is less hydrogen bonding than water. • Melting point= -114 ° C • Boiling Point = 78.4 ° C Applications Purpose of Sweat When sweat evaporates from your skin, it removes excess heat and cools you 1 gram sweat uses 2260 Joules of body heat to evaporate HeatGear- pulls perspiration away from the skin and push it through the clothing's surface, where it evaporates Polar Phase Change Cooling Garments • “Passive Cooling Garments that release long lasting, 58 °F temperature specific, cooling energy.” How misters work -Cools patios by up to 35 ° F -Evaporation uses heat from the air, thus dropping the temperature. Why does warm air rise? This creates our wind/ weather patterns Extras Bond Energies Energy is used to break bonds & is released when new bonds form. + 185 KJ 436 KJ + 243 = + 679 KJ (absorbed) 2 (432)= -864 KJ (released) = -185 KJ lost FYI: 2 ways Vaporization can occur: Evaporation Vs. Boiling -Occurs at temperatures less than boiling (100° C) -Occurs at 100° C, but never goes above. -occurs on surface of the liquid -Can heat up faster or slower, but same temp. Phases of Matter Phases of Matter (review) Solid Liquid Gas •Motion of particles is restricted (only vibrational) •Particles are more closely packed •Lowest kinetic energy •Strong intermolecular forces •Have definite volume •Particles are in constant motion •Lower kinetic energy than gas particles •Attractive forces between particles effect their behavior •Particles are not fixed, but are closer together than in a gas •Particles are in constant, rapid motion •Highest kinetic energy state of matter •Attractive forces do not effect their behavior •Particles are far apart Salt in Boiling water • The answer to the question why do bubbles form when water is boiled? explains how, as the temperature of water rises, its molecules move around faster, collide more often and release more water vapor gas molecules. When the temperature reaches boiling point - about 100°C (212°F) - the pressure from the release of these molecules (the vapor pressure) becomes greater than atmospheric pressure and water vapor starts to escape as bubbles. In salty water, Na+ and Cl- ions occupy some of the space between the water molecules. As temperature increases, although the water molecules are moving faster, there are less of them, so there are fewer collisions, less release of water vapor molecules and lower vapor pressure compared to pure water at the same temperature. It takes more energy (temperature) for the vapor pressure of salt water to reach and exceed atmospheric pressure and start to boil. To raise the boiling point of one liter (34 ounces) of water by 1°C (1.8°F) requires about 58 grams (2 ounces) of salt. This is much more that the amount of salt typically added to boiling vegetables, which is done primarily for taste. If you ever watch salt melting ice, you can see the dissolving process happen -- the ice immediately around the grain of salt melts, and the melting spreads out from that point. If the temperature of the roadway is lower than 15 F or so, then the salt really won't have any effect -- the solid salt cannot get into the structure of the solid water to start the dissolving process. In that case, spreading sand over the top of the ice to provide traction is a better option. • Put another way, if a solute is dissolved in a solvent, then the number of solvent molecules at the surface of the solution is less than for pure solvent. The surface molecules can thus be considered “diluted” by the less volatile particles of solute. The rate of exchange between solvent in the solution and in the air above the solution is lower (vapour pressure of the solvent is reduced). A lower vapour pressure means that a higher temperature is necessary to boil the water in the solution, hence boiling-point elevation. Joke • Two scientists walk into a bar • The first one says “I’ll have some H2O.” The second one says, “I’ll have some H2O too.” Then he dies