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Energy and Chemical Change Measuring and calculating the energy involved in chemical changes

Energy Energy is defined as the ability to do work or produce heat. Kinetic Energy Potential Energy Energy is defined as the ability to do work or produce heat. SI unit for energy is Joules (J). There are 2 types of energy

Energy 1. Potential Energy (PE) Energy that is based on an objects composition (chemical PE) or position; associated with attractions and repulsions. Chemical potential energy the energy stored in a substance because of its composition (ex: plants, gasoline) plays an important role in chemical reactions

Energy 2. Kinetic Energy (KE) - Energy of motion. Kinetic Energy Can be measured as temperature or heat. KE of a substance is directly related to the motion of its atoms and temperature.

Energy * Every object has either potential or kinetic energy!!!

Energy Law of Conservation of Energy: energy can be converted from one form to another, but it is neither created nor destroyed.

Heat Form of energy associated with changing the temperature of an object. An object’s temperature increases because energy is transferred into it Temperature is a measure of the average kinetic energy of the particles in a sample of matter; it does not depend on the amount of matter in the sample.

Heat The heat (energy) transfer occurs when 2 objects of different temperatures are brought into contact to reach equilibrium. Depends on the mass of the sample. Symbol = m Unit = grams Reactions can be endothermic (heat/energy in) or exothermic (heat/energy out).

How Can I Measure Energy/Heat? There are many units that can express an amount of energy. calorie (cal) - amount of heat required to raise the temperature of 1.00 g of pure water by 1oC.

How Can I Measure Energy/Heat? kilocalorie (kcal) - 1000 calories or one Calorie – this is one food Calorie! Joule - SI unit of heat and energy. (What we will mainly use) Important conversions: 1 Calorie = 1 kcal 1 kcal = 1000 cal 1 cal = 4.184 J

How Can I Measure Energy/Heat? heat changes that occur during chemical and physical processes can be measured accurately and precisely using a calorimeter a calorimeter is an insulated device used for measuring the amount of heat absorbed or released during a chemical or physical process

How Can I Measure Energy/Heat? Inside most calorimeters you will need: a known mass of water is placed in an insulated chamber in the calorimeter it is designed to absorb the energy released from the reacting system (unknown/known substance) or to provide the energy absorbed by the system the data collected is the change in temperature of the water inside the calorimeter

How Can I Measure Energy/Heat? most calorimetric reactions occur at constant pressure can also be used to determine the specific heat of an unknown metal assuming that no heat is lost to the surroundings, the heat lost by the metal is equal to the heat gained by the water

What is Specific Heat? (Or, How do I Calculate Amount of Energy Transferred?) Specific Heat is the amount of heat required to raise the temperature of one gram of any substance by 1oC. Units = J/g x OC of kJ/g x oC Each substance has its own specific heat. Knowing the specific heat of a substance allows you to calculate the heat transferred by measuring the mass of the substance and the temperature change.

What is Specific Heat? (Or, How do I Calculate Amount of Energy Transferred?) Specific heat of water = 4.184 J/ g x oC Formula: q = c x m x ∆T c = specific heat m = mass ∆T = Tfinal – Tinitial Heat transferred = (specific heat)* (mass)* (change in T measured in OC ) q = (g)*(J/g x OC)*(∆T)

What is Specific Heat? (Or, How do I Calculate Amount of Energy Transferred?) --This equation also indicates which direction the heat is traveling. **If ∆T is (+) then heat is being transferred to or absorbed by the substance (endothermic) **If ∆T (-) then heat is being released during the reaction (exothermic)

Keep in mind… the specific heat of a substance is a measure of how efficiently that substance absorbs heat

PRACTICE!!! Problem- How much heat is lost when solid Al with a mass of 4110 g cools from 660.0 oC to 25 oC? (Specific Heat of Al is 0.902 J/g x oC) First determine the ΔT = Tfinal -Tinitial (25oC-660oC = -635oC) Remember the equation is q= cxmxΔT c = specific heat m = mass

PRACTICE!!! Plug in the values q = (0.902 J/gxoC)*(4110g)*(-635oC) Cancel like units Multiply all values -2354084.7J or divide by 1000 and get -2354.0847kJ Significant figures gives us -2354kJ or -2.35 x 106J Type of reaction: EXOTHERMIC

PRACTICE!!! Problem- How much heat is required to raise the temperature of 854 g of H2O from 23.5 oC to 85.0 oC? ΔT = Tfinal – Tinitial ΔT = 85.0oC – 23.5oC = 61.5oC q = c x m x ΔT q = (4.184 J/g x oC)*(854g)*(61.5oC)= 219747.9J or 219.7kJ or 2.20 x 105J Type of reaction: ENDOTHERMIC

So, What is Thermochemistry? Thermochemistry is the study of heat changes that accompany chemical reactions and phase changes. ***Remember: Energy often takes the form of HEAT As a chemical reaction progresses some of the energy in the bonds may be released as heat. These reactions are called EXOTHERMIC.

So, What is Thermochemistry? In an exothermic reaction the products have less potential energy than the reactants. Exothermic reactions release energy; more energy is released than is required to break bonds in the initial reaction.

So, What is Thermochemistry? Examples: burning wood, any reaction in which the test tube or beaker gets warm, formation of water (explosion of hydrogen and oxygen), an athletic hot pack. the contents of the athletic hot pack is the system which is the specific part of the universe that contains the reaction or process you wish to study the universe is defined as the system plus the surroundings (which is everything in the universe other than the system)

So, What is Thermochemistry? Some chemical reactions absorb heat as they progress. These reactions are called ENDOTHERMIC. In an endothermic reaction the products have more heat energy (enthalpy) than the reactants.

So, What is Thermochemistry? Reaction that needs energy to complete the reaction; a greater amount of energy is required to break the initial bonds than is released from the new bonds forming Examples: cooking eggs, any reaction in which the beaker or test tube gets cooler, an athletic cold pack, decomposition of water (electrolysis).

Quick Quiz!!! Identify each of the following as endothermic or exothermic: a. burning natural gas (methane) exothermic b. baking a cake endothermic c. melting ice

What is Enthalpy? Enthalpy is the chemistry word for heat content (remember that chemical compounds have POTENTIAL energy in their bonds – this is the heat content). Enthalpy is represented by the symbol H You cannot measure the actual enthalpy of a substance…instead you measure the change in enthalpy (the heat absorbed or released in a reaction)

What is Enthalpy? The change in enthalpy is called the enthalpy heat of reaction and is represented by ∆H or ∆Hrxn. ΔHrxn = Hproducts – Hreactants

Enthalpy of Exothermic Reactions Enthalpy changes for exothermic reactions are always negative. ∆H is (-) In an exothermic reaction the products have less potential energy than the reactants Hprod. < Hreact.

Enthalpy of Exothermic Reactions Energy is a product Example: CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) + 802 kJ product CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) ΔH = -802 kJ (because it is on the product side)

Enthalpy of Endothermic Reactions Enthalpy changes for endothermic reactions are always positive. ∆H is (+) In an endothermic reaction the products have more heat energy (enthalpy) than the reactants. Hprod > Hreact

Enthalpy of Endothermic Reactions Energy is a reactant Example: 2H2O(l) + 572 kJ  2H2(g) + O2(g) 2H2O(l)  2H2(g) + O2(g) ΔH = 572 kJ (because it is on the reactant side)

Thermochemical Equations Thermochemical equations are balanced chemical equations that include the physical states of all reactants and products and the energy change. 2H2O(l)  2H2(g) + O2(g) ΔH = 572 kJ CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) ΔH = -802 kJ

Enthalpy of Combustion Enthalpy heat of combustion (ΔHcomb) is the enthalpy change for the complete burning of one mole of the substance. -- carried out under standard conditions which are one atmospheric pressure (1 atm) and 298K (250C) - C6H1206(s) + 6O2(g)6CO2(g)+ 6H2O(l) ΔHcomb = -2808kJ

Changes of State Molar enthalpy (heat) of vaporization (ΔHvap) is the heat required to vaporize one mole of liquid. -- think of water vaporizing from your skin after you take a hot shower. Your skin provides the heat needed to vaporize the water and as the water absorbs the heat you feel cool (shiver) ΔHvap = -ΔHcond (condensation)

Changes of State Molar enthalpy (heat) of fusion (ΔHfus) is the heat required to melt one mole of a solid substance. --think of ice in a drink. The drink cools as it provides the heat for the ice to melt ΔHfus = - ΔHsolid (solidification—freezing)

Phase Changes Occurs when energy is added or removed from a system and the substance can go from one physical phase to another

Endothermic Phase Changes Melting The energy absorbed to melt a solid is not used to raise the temperature of that solid The energy instead disrupts the bonds holding the solid’s molecules together and cause the molecules to move into the liquid phase

Endothermic Phase Changes The amount of energy required to melt one mole of a solid depends on the strength of the forces that hold the solid together The melting point of a crystalline solid is the temperature at which the forces holding its crystal lattice together are broken and it becomes a liquid

Endothermic Phase Changes Vaporization Particle that escape from the liquid enter the gas phase and those liquids at room temperature the gas phase is called vapor Vaporization is the process by which a liquid changes into a gas or vapor Once the solid becomes a liquid then and only then does the temperature of the substance begin to increase

Endothermic Phase Changes When vaporization takes place only at the surface of the liquid it is called evaporation Evaporation is the method by which the human body maintains and controls its temperature

Endothermic Phase Changes The pressure exerted by a vapor over a liquid is called the vapor pressure The temperature at which the vapor pressure of a liquid equals the external or atmospheric pressure is called the boiling point

Endothermic Phase Changes Sublimation Is the process by which a solid changes directly to a gas without first becoming a liquid Dry ice (CO2) and snow are the most common examples

Endothermic Phase Changes If ice cubes are left in the freezer for extended periods of time, they will eventually sublime and become smaller This process is also helpful in freeze drying foods for hikers and astronauts

Exothermic Phase Changes Condensation When a vapor molecule loses energy its velocity is reduced therefore colliding more with other molecules to form a liquid Condensation is the process by which a gas or vapor becomes a liquid and it is the reverse action of vaporization

Exothermic Phase Changes Deposition Is the process by which a substance changes from a gas or vapor to a solid without first becoming a liquid It is the reverse action of sublimation The formation of snow crystals high up in the atmosphere is an example

Exothermic Phase Changes Freezing Point Is the temperature at which a liquid is converted into a crystalline solid The same temperature as the melting point of a given substance

Phase Change Graph

Phase Change Graph Graph shows the energy required to go from one phase to the other Where the graph inclines, potential energy is at its greatest and temperature is increasing Where the graph plateaus (flat region) kinetic energy is at its greatest but the temperature remains constant

Phase Diagrams A phase diagram is a graph of pressure versus temperature that shows in which phase a substance exists under different conditions of temperature and pressure

Phase Diagrams The triple point is the point on a phase diagram that represents the temperature and pressure at which three phases of a substance can coexist The critical point is the point that indicates critical pressure and temperature above which water cannot exist as a liquid

Phase Diagrams Different for each substance because of the different boiling/freezing points