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Chem 107 Notes 4/5 and 4/7

by: Kelly Johnson

Chem 107 Notes 4/5 and 4/7 Chem 107

Kelly Johnson
GPA 3.63

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General Chemistry for Health Science
Jacqueline Butler
Class Notes
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This 5 page Class Notes was uploaded by Kelly Johnson on Thursday April 7, 2016. The Class Notes belongs to Chem 107 at West Chester University of Pennsylvania taught by Jacqueline Butler in Winter 2016. Since its upload, it has received 32 views. For similar materials see General Chemistry for Health Science in Chemistry at West Chester University of Pennsylvania.


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Date Created: 04/07/16
Chapter 7 Energy, Rate, and Equilibrium 1. Thermodynamics a. Definition i. Thermodynamics is the study of energy, work, and heat b. Basic Concepts i. Molecules are in constant and random motion ii. Molecules frequently collide with each other iii. Collisions with sufficient enough energy will break molecular bonds iv. If bonds are broken, new bonds may form causing new product formation c. Changes in Energy i. Energy stored cannot be measured, but you can measure the change in  energy ii. Energy may be gained by a system as it is lost in the surroundings  1. System­ contains the process under study 2. Surroundings­ the rest of the universe d. First Law of Thermodynamics (Law of the Conservation of Energy) i. Energy in the universe is constant  e. Chemical Energy Changes i. What 1. AB + CD = AD + CD 2. Each bond in this equation stores energy 3. In a reaction, bonds may break and that requires energy ii. Types of Reactions 1. Exothermic a. Energy required to break the bond is less than what is  released when the bonds are formed b. There is a net release of energy in the form of heat c. AB+CDAD+CB+ HEAT 2. Endothermic a. Energy required to break the bond is greater than what is  released when bonds are formed b. Energy/heat must be added to the reaction c. HEAT+AB+CD=AD+CB iii. Enthalpy 1. Represents heat energy 2. Change in Enthalpy (ΔH°) a. Represents the energy difference between the products and  reactants of a chemical reaction b. In an exothermic reaction, energy is released so ΔH° is  negative c. In an endothermic reaction, energy is absorbed so ΔH° is  positive d. Can be spontaneous or nonspontaneous i. Spontaneous occur without external input ii. Exothermic reactions are almost always  spontaneous 3. Entropy a. The second law of thermodynamics states that the universe  tends toward increasing disorder or randomness b. Entropy (S°) is the measure of randomness in a chemical  system i. High entropy= highly disordered ii. Low entropy= well disorganized c. You cannot have a negative S° value but you can have a  negative ΔS° d. ΔS° = S°(products) ­ S°(reactants) i. If ΔS° is positive, there is an increase in disorder 1. This is spontaneous ii. If ΔS° is negative, there is a decrease in disorder 1. This is nonspontaneous e. Melting, vaporization, & dissolution become more  disordered (+ΔS°) f. Entropy values per state i. As a gas becomes a solid, it gains order  ii. As solids change to gas, they lose order 4. Free Energy a. Represents the combined contribution of the enthalpy and  entropy values for a chemical reaction (ΔG°). Predicts the  spontaneity of a chemical reaction. b. ΔG° = ΔH° ­ TΔS° i. A negative ΔG° is always spontaneous 1. If exothermic and positive ΔS° ii. A positive ΔG° is never spontaneous 1. If endothermic and negative ΔS° c. Need to know ΔH and ΔS  to predict the sign of ΔG d. Temperature also may determine direction i. +ΔH and – ΔS  ΔG always + regardless of T ii. – ΔH and + ΔS  ΔG always – regardless of T iii. + ΔH and + ΔS  ΔG depends on T iv. – ΔH and – ΔH  ΔG depends on T 2. Experimental Determination of Energy Change in Reactions a. Calorimetry i. The measurement of heat energy changes in a chemical reaction ii. Calorimeter is used to measure changes in calories iii. Change in temperature is used to measure loss or gain of heat b. Heat Energy in Reactions i. Changes in temp caused by a reaction can help calculate gain or loss of  heat energy 1. Endothermic­ heat released is absorbed 2. Exothermic­ reactants absorb heat from the solution ii. Specific Heat is the number of calories of heat needed to raise the  temperature of 1g of the substance 1°C iii. Q = m ΔT SH 1. M= mass of the solution 2. ΔT= change in initial to final temperatures 1cal 3. SH= specific heat (for water it is  g°C ) c. Bomb of Calorimeter and Measurement of Calories in Foods i. Measures nutritional calories (Cal) 1. 1 Cal=1 kilocalorie (kcal)=1000 calories 2. Used for fuel value of food ii. Uses same formula (Q = m ΔT SH) 3. Kinetics a. Definition i. The study of the rate or speed of chemical reactions ii. This also hints at the mechanism of a reaction b. Change over time i. Measures the disappearance of A and the appearance of B ii. The color change over time can be used to monitor the progress of a  chemical reaction  c. Reactions i. Energy from collisions is required to break bonds ii. When collisions occur they must be effective or bonds will not break iii. Activation energy is the minimum amount of energy required to initiate a  chemical reaction iv. An activated complex is extremely unstable and short­lived. Requires  energy to overcome the barrier to start the reaction. d. Factors That Affect Reaction Rate i. Structure of the Reacting Species 1. Opposite charged species react more rapidly 2. Ions of the same charge do not react 3. Bond strength has an effect a. The more bonds/the stronger the bond= harder to break ii. Molecular Shape and Orientation 1. Only collisions with correct orientation can lead to productive  formation 2. Large molecules may obstruct the reactive part of the molecule iii. Concentration of Reactants  1. Rate will generally increase as concentration increases a. More molecules means more collisions iv. Temperature of Reactants 1. Rate increases as temperature increases a. Increased average kinetic energy means molecules move  faster and collide more often and a higher percentage of  these collisions will result in product formation v. Physical state of Reactants 1. In order of fastest to slowest reactions a. Liquid­ particles are free to move and lie close together so  they collide often b. Gas­ particles are free to move but lie far apart which make collisions infrequent c. Solid­ particles are close together but are restricted in  movement so they can’t collide as often vi. Presence of a Catalyst 1. Catalyst is a substance that increases the reaction rate. They  undergo no net change and do not alter.  e. Mathematical Representation of Reaction Rate i. There is a generic rate equation (rate law) 1. Rate= k [A] [B] n m a. [ ] represents molarity or concentration b. Rate is normally in mol/s ii. Rate only depends on reactants  iii. Exponent represents the order of the reaction. It is determined through  experiments. IT IS NOT THE SAME AS THE COEFFIECIENT. 1. For example, if n=2 then it is in the 2  orderd iv. If concentration is changed, then place that number in the bracket the letter is in and calculate. v. Examples 1.  In the equation CH (g) +42O (g)  22 O(l) + CO2(g) 2 a. Rate=k[CH ][O ] 4 st2 b. CH is 4  the 1  order because it is not raised to a power c. O  i2 of the first order because it is not raises to a power nd d. The overall order is the 2 2. In the equation 2NO(g) + O (g)  22O (g) 2 2 a. Rate= k[O ][NO2 b. O is2 f the 1  order nd c. NO is of the 2  order d. The overall order is the 3 rd 3. In the equation 2NO(g) + 2H (g) N (g) + 2H O(g) 2 2 2 2 a. Rate= k[NO] [H ] 2 b. If the concentration of H  is 2oubled, the rate will double.  2 1 (Rate= k[NO] [2H ]2 )2 c. If the concentration of NO is doubled, the rate will increase 2 2 4x. (Rate= k[2NO] [H ]2 ) 2 d. If the concentration of NO is tripled, the rate will increase  2 2 9x. (Rate= k[3NO] [H ]  3 2


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