(CHEM 1041) Exam 1 Study Guide
(CHEM 1041) Exam 1 Study Guide CHEM1041
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This 6 page Study Guide was uploaded by Macen Notetaker on Sunday February 7, 2016. The Study Guide belongs to CHEM1041 at University of Cincinnati taught by Dr. Waddell in Spring 2016. Since its upload, it has received 231 views. For similar materials see General Chemistry 2 in Chemistry at University of Cincinnati.
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Date Created: 02/07/16
CHEM 1041 General Chemistry 2 Exam 1 Study Guide 1. Chapter 12 a. Intermolecular Forces i. Inter vs Intra 1. Intra is within molecules. a. They are bonding forces, and influence chemical properties b. Ex. Ionic, Covalent, Metallic 2. Inter is between molecules. a. They are nonbonding forces, and influence physical properties b. Ex. H Bonding, Dipole-Dipole, Dispersion, etc. ii. Differences in states of matter 1. Compressibility 2. Ability to flow 3. Definite or Indefinite Volume and Shape iii. Kinetic Energies vs. Interaction Energies 1. Higher IEs lead solids to have fixed shapes 2. Lower IEs lead gases to lack fixed shape or volume iv. Phase Changes 1. Names of each phase change 2. Enthalpy Changes a. Reverse property. Enthalpy of melting is opposite of freezing, for example. 3. Quantitative Aspects a. Within a phase: q = amount * molar heat capacity * ΔT b. During a phase change: q = amount * enthalpy of phase change 4. Liquid – Gas Equilibrium & Vapor Pressure a. Vapor Pressure increases with temperature b. Clausius-Clapeyron Equation -∆vap 1 i. ln P = R *T+C P -∆vap 1 1 ii. ln( )= *( - ) P1 R T2 T1 c. Boiling Point i. Normal Boiling Point is at 760 torr ii. As external pressure on a liquid increases, so does boiling point 5. Phase Diagrams v. Polarizability 1. Polarizability is the ease with which an electron cloud is distorted a. More polarizable = stronger intermolecular forces b. Polarizability trends down to the left, same as atomic radius vi. Effect of IMFs on other properties 1. Surface tension increases with stronger IMFs 2. Boiling point increases with stronger IMFS 3. Viscosity increases with stronger IMFs 4. Vapor Pressure decreases with stronger IMFs vii. Properties of Water 1. Caused by Water’s ability to form H bonds to other molecules 2. High Heat Capacity 3. High Heat of Vaporization 4. High Surface Tension 5. Universal Solvent b. Solids & Crystal Structure i. Unit Cells 1. Simple Cubic 2. Body Centered Cubic 3. Face Centered Cubic ii. Edge Length for Each Cell 1. Simple: A=2 r 4 2. Body Centered: A= r √3 3. Face Centered: A= 8 √ iii. Types of Crystalline Solids 1. Atomic 2. Molecular 3. Ionic 4. Metallic Network Covalent iv. Molecular Orbital Band Theory 1. Energy gaps in relation to conductors vs. insulators 2. Chapter 13 a. Solutions & Solubility i. Like Dissolves Like 1. Polar dissolves polar 2. Similar IMFs dissolve in one another ii. Dual Polarity and Alcohols 1. Longer chains are more soluble in nonpolar solvents 2. Shorter chains are more soluble in polar solvents 3. There is a happy medium around ethanol and propanol b. Why Substances Dissolve i. Entropy 1. Gases have the highest entropy 2. Solutions have higher entropy than pure solvents and solutes 3. Higher entropy is favored ii. Heat of Solution 1. Sum of enthalpies for solute, solvent, and mixing 2. Solvation and Hydration a. Solvation is surrounding a particle with solvent particles b. If surrounding with water, this process is hydration c. Equal to sum of enthalpies of solvent and mixing 3. Heat of Solution can be written as the sum of the enthalpies of solvation and solute iii. Charge Density 1. Measure of charge per area a. Larger charges and smaller radii yields larger charge densities 2. Determines the trend of enthalpy of Hydration values 3. Charge density trends up and to the right, with electronegativity iv. Solubility and Equilibrium 1. Saturated solutions contain some undissolved solute 2. This undissolved solute is in equilibrium with dissolved solute v. Factors Affecting Solubility 1. Temperature a. Most solids are more soluble at higher temperatures 2. Pressure a. Gases are more soluble at a higher pressure b. Solids and liquids are unaffected vi. Henry’s Law 1. Solubility of a gas is proportional to the partial pressure of the gas above the solution a. S gas= kH* P gas vii. Concentration Definitions 1. Molarity 2. Parts by Volume 3. Molality 4. Parts by Mass 5. Mole Fraction c. Colligative Properties i. Depend on number of solute particles, not the chemical identity of the solute ii. Electrolytes 1. Electrolytes separate into ions when dissolves in water 2. Strong electrolytes dissociate completely a. Weak electrolytes don’t 3. Nonelectrolytes will not dissociate at all when in water iii. Types of Colligative Properties 1. Be sure to also know the associated equation 2. Vapor Pressure lowering 3. Boiling Point elevation 4. Freezing Point depression 5. Osmotic Pressure iv. Van’t Hoff factor 1. Determines the “effective” number of ions 2. Ideal van’t Hoff Factor a. Given by the number of ions that are in one molecule of substance i. NaCl has an ideal factor of 2 ii. Sugar has an ideal factor of 1, as it does not dissociate 3. Calculated van’t Hoff Factor a. Given by ration between measured value for electrolyte solution and the expected value for nonelectrolyte solution 3. Chapter 16 a. Reaction Rate i. Factors Affecting Reaction Rate 1. Rate increases with concentration 2. Physical state affects concentration 3. Rate increases with temperature ii. Expressing Rate 1. For a standard reaction, A yields B ∆[A] a. Rate=- ∆t b. The negative sign makes the value positive, as it would otherwise be negative with A being a reactant that decreases in concentration as the reaction moves forward iii. Three Types of Rate 1. Average Rate a. Slope of the concentration vs time plot 2. Instantaneous Rate a. Slope of the tangent line 3. Initial Rate a. Rate when time = 0 iv. Rates Relative to One Another 1. For the Reaction aA + bB → cC + dD a. - * ∆ A=- * ∆ B]= * ∆ C= *1 ∆ D] a ∆t b ∆t c ∆t d ∆t b. Rate Law i. Rate as a function of concentration ii. General Form 1. For the reaction aA + bB → cC + dD m n a. Rate=k[A] [B] i. K is the rate constant ii. M and N are determined through experiments iii. Reaction Orders 1. First Order ([A] ) a. Rate doubles when [A] doubles 2. Second Order ([A] ) 2 a. Rate quadruples when [A] doubles 3. Zero Order ([A] )0 a. Rate doesn’t change when [A] doubles 4. Adding the individual reaction orders for each reactant will yield the overall reaction order iv. Determining Reaction Orders 1. You will be given experimental data 2. To determine the order of reactant A: a. Pick two trials where everything but [A] and the rate change b. Find the ratio between the two concentrations of A c. Find the ration between the two rates d. The number that you need to raise your answer to b to in order to get your answer to c is your reaction order 3. For example: a. If the rate changes by a factor of 4 when the reactant changes by an order of 2, the order is 2. i. Because 2 must be raised to the power of 2 to yield 4 b. If the rate changes by a factor of 4 when the reactant changes by an order of 4, the order is 1 i. Because 4 to the power of 1 is 4 c. If the rate changes by a factor of 1, the order is 0 i. Anything to the power of 0 is 1 4. Determine the rate constant by plugging in the concentrations for one trial, along with the values found for order. Set this equal to the appropriate rate, and do algebra v. Integrated Rate Law 1. Concentration as a function of time st nd th 2. Equations for 1 , 2 , and 0 order will be given a. But you must know how to manipulate each equation to solve for any variable vi. Units for Rate Constant 1. Can be found by remembering that rate is in units of M/s or molarity per second 2. Plug in M for each concentration and raise them to the proper exponent 3. Then do algebra to find k 4. Units can be memorized for orders of 0, 1, and 2 a. First Order: 1/s b. Second Order: 1/(M*s) c. Zero Order: M/s c. Half Life i. The amount of time required for the concentration of a reactant to reach half of its initial value ii. Equations will be given 1. Again, be familiar with manipulation
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