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Chemistry Notes for Week of 4-12 and 4-14

by: Kelly Johnson

Chemistry Notes for Week of 4-12 and 4-14 Chem 107

Marketplace > West Chester University of Pennsylvania > Chemistry > Chem 107 > Chemistry Notes for Week of 4 12 and 4 14
Kelly Johnson
GPA 3.63

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Second half of Chapter 7 and first half of chapter 8
General Chemistry for Health Science
Jacqueline Butler
Class Notes
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This 6 page Class Notes was uploaded by Kelly Johnson on Friday April 15, 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 17 views. For similar materials see General Chemistry for Health Science in Chemistry at West Chester University of Pennsylvania.


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Date Created: 04/15/16
Chapter 7 Continued 1. Thermodynamics (in first set of ntoes) 2. Experimental Determination of Energy Change in Reactions (in first set of notes) 3. Kinetics (in first set of notes) 4. Equilibrium a. Rate and Reversibility i. Equilibrium reactions are chemical reactions that do not go to completion.  AKA all reactants do not convert to products. Also known as incomplete  reactions. b. Physical Equilibrium i. Reversible reactions. SO they can occur in either direction (products to  reactants or reactants to products) ii. Represented by double arrows, one facing each way 1. Double arrow serves as an indicator for a reversible process, an  equilibrium process, or the dynamic nature of a process iii. Dynamic Equilibrium­ when the rate of the forward process of the reaction is exactly equal to the rate of the backward process.  Rate = rate  1.  solute settles 2. Ex. Sugar (s)sugar(aq) = sugar (aq) sugar (s) iv. Chemical Equilibrium 1. As a reaction begins, its forward reaction rate is quick, but as time  goes on that rate decreases.  a. To start­ [reactants] are high and [products] is low b. At end­ [reactants] is decreasing and [products] increasing 2. K  eq constant at a constant temperature v. Equilibrium Constant 1. At equilibrium Rate=ratef r a. Rate=rafe forward b. Rate= reverse  r [kf ] 2. K =eq [kr] a. [k]=fconcentration of products b. [k]=rconcentration of reactants c. K is unitless d. In the equilibrium constant, any exponents are numerically  equal to the coefficients in the reaction equation [C] [D] d i. If aA + bB = cC + dD, then K = a b eq [A] [B] 3. Rules for writing an equilibrium constant a. Can only be written with a balanced chemical equation b. Every reaction has a unique constant value at a specified  temperature c. Brackets represent molarity of that substance d. All equilibrium constants are unitless e. Concentrations of gases and substances in solution are  shown, liquids and solids are not shown vi. Interpreting Equilibrium constants 1. If a double arrow exists, then an equilibrium constant exists 2. Numerical value tells you the extent to which reactants have turned to products. a. If K eqs greater than 100, than there is mostly product  present b. If K  is less than .01, than there is mostly reactant present eq c. If K eqs between .01 and 100, than products and reactants  are very close in concentration vii. Calculating Equilibrium Constants 1. A reversible reaction will proceed until the system reaches  equilibrium, which means reactant and product concentrations no  longer change 2. Steps to calculate a. Look to see if double arrows are present b. Cross out and solids or liquids in the reaction equation c. Plug in products and reactants into the K eqeq ion d. Substitute each substance with its concentration i. Add a exponents if needed e. Solve c. LeChatleir’s Principle i. Definition 1. If stress is added to a system, that system will respond by altering  the equilibrium composition to minimize stress 2. Systems are stressed in a reactant or product is added to the system ii. Changes 1. If reactants are added, then the system will move toward product  formation to reduce the reactants 2. If products are added, then the system will move toward reactant  formation to balance it out 3. If added or removed, concentration of those materials also changes accordingly a. System will alter to realign concentrations iii. Effects of Heat  1. Heat is treated as a product or reactant a. Exothermic reaction i. Heat is treated as a product ii. Production of product will decrease and production  of reactants will increase iii. This is a shift of equilibrium to the left b. Endothermic reaction i. Heat is treated as a reactant ii. Production of product will increase and production  of reactants will decrease iii. This is a shift of equilibrium to the right iv. Effects of Pressure 1. Pressure only effects equilibrium if one or more of the substances  are gases 2. The number of moles of gas on either side must differ a. If they differ and pressure goes up, the equation will shift to the side with a lower number of moles b. If they differ and pressure goes down, the equation will  shift to the side with a larger number of moles v. Effects of a catalyst 1. A catalyst will have no effect on equilibrium because it effects  both the forward and backwards reactions 2. Equilibrium will however, be reached in a shorter time  Chapter 8: Acids and Bases 1. Acids and Bases a. Characteristics i. Acids are sour and can dissolve in some metals ii. Bases taste bitter, are slippery, and are corrosive b. Theories i. Arrhenius Theory 1. Acid­ a substance, when dissolved in water, dissociates to produce  Hydrogen ions (protons) a. There are 6 strong acids  2. Base­ a substance, when dissolved in water, dissociates to produce  hydroxide ions  a. Strong bases end in OH 3. NH is3 n exception ii. Bronsted­Lowry Theory 1. Acid­ a proton donor 2. Base­ a proton acceptor 3. Water can be an acid or a base in each equation c. Acid­Base Properties of Water i. Water has both acid and base properties (amphiprotic) ii. Most commonly used solvent for acids and bases iii. Solute­solvent interactions promote both dissociation and solubility d. Acid and Base strength i. Measured based on degree of dissociation ii. To be strong acids or bases, they must dissociate completely iii. Weak acids and bases do not dissociate completely, and have a double  arrow in their reaction equation 1. Weak acids can be used as buffers iv. Strong Acids 1. HCl 4. HNO3 2. HBr 5. H2SO4 3. HI 6. HClO4 v. Strong Bases 1. NaOH 2. KOH 3. Ba(OH) 2 4. All metal hydroxides e. Conjugate Acids and Bases ­ + i. HA + B  A + HB 1. Where HA is the acid and B is the base  ­  + 2. A is the conjugate base and HB  is the conjugate acid ii. A conjugate acid is what the base becomes after it accepts a proton iii. A conjugate base is what the acid becomes after it donates a proton iv. Each conjugate pairs with a reactant  f. Acid­Base Dissociation i. If an acid or base dissociates 100%, then only a single arrow is used g. Dissociation of Water i. Pure water is basically 100% molecular ii. It is a very weak electrolyte, so very few molecules dissociate h. Hydronium Ion + i. H O3 is the hydronium ion ii. Pure water has a 1 x 10  [H O ] and [OH] ­ 3 iii. K =eqH O ]3OH] ­ iv. Ion Product of Water 1. K =w[H O ]3OH] ­ + ­ ­7 2. We know the values of [H O ] and3[OH] are 1.0 x 10 3. So, K = w.0 x 10 ­14 2. pH a. Definition i. The pH scale indicates the acidity or basicity of a solution 1. Range from 0 (very acidic) to 14 (very basic) + 2. Defined as pH= ­log[H O ] 3 ii. If you add an acid, [H O ]3increases and [OH] decreases + ­ iii. If you add a base, [H O ]3decreases and [OH] increases iv. An acid and base are present if [H O ] =3[OH ] ­ b. Measuring pH i. pH can be calculated if the concentration is known of either OH  or H O ­ 3 + ii. pH paper indicators iii. pH sensor c. Calculating pH i. Equations + 1. pH= ­log[H O ] 3 2. pOH= ­log[OH] ­ 3. pH + pOH = 14.00 4. 10 ­pH= [H O3]+ ­pOH  ­ 5. 10 = [OH] 6. K = w.0 x 10  = [H O ][OH3 + ­ ii. pH scale 1. Represents the concentration of H and OH +  ­ 2. Ranges from 0 to 14 3. Each increase in 1 pH unit is a 10x increase in concentration d. Importance of pH and pH Control i. Agriculture­ plants can only grow in soil of a certain pH ii. Physiology­ a shift of 1pH in the blood is fatal iii. Acid Rain­ lowers pH of aquatic systems and kills fish iv. Municipal­ sewage treatment and water purification require certain pH’s v. Industry­ strict pH control for cost effective production 3. Reactions Between Acids and Bases a. Neutralization Reactions i. When a strong acid and a strong base combine to form a salt and water 1. ACID + BASE = SALT + H O + 2 ­  ii. The net ionic equation will be H  + OH  H O 2 1. More accurately H O (aq) + OH(aq)  2H O(l) 3 2 b. Titration i. Is used to neutralize solutions and determine concentration ii. A standard solution (one of known concentration) is placed in a buret and  slowly added to the unknown substance in a flask. A indicator is also  added to notify when the pH changes c. Determining the Concentration of a Solution i. M VacidMacid baseV base ii. This only works if 1 OH is being accepted or donated d. Polyprotic Substances i. Donate or accept more than one proton per formula unit ii.


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