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Chem 116 (Final exam study guide)

by: Jamisha Evans

Chem 116 (Final exam study guide) CHEM 116

Marketplace > Western Kentucky University > Chemistry > CHEM 116 > Chem 116 Final exam study guide
Jamisha Evans
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These are the notes for the final exam. i apologize ahead of time if they are not as good as my other study guides. i was not able to spend a lot of time on it of it because my laptop stopped worki...
Bangbo Yan
Study Guide
Chem, 116, study, guide
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This 8 page Study Guide was uploaded by Jamisha Evans on Friday May 6, 2016. The Study Guide belongs to CHEM 116 at Western Kentucky University taught by Bangbo Yan in Spring 2016. Since its upload, it has received 57 views. For similar materials see INTRO TO COLLEGE CHEMISTRY in Chemistry at Western Kentucky University.


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Date Created: 05/06/16
Chemistry 116: Chapter 1 *Week 2 Notes* Measurement & Significant Figures • Unit: Comparison of a physical quantity with a fixed standard of measurement • Accuracy and Precision repeated measurements for four trials are Trial 1: 32kg, Trial 2: 34kg, Trial 3: 33kg, Trial 4: 34kg then your average is 33.25 and this means your results are precise but NOT accurate. • Significant Figures: meaningful digits plus a final digit with some uncertainty (AKA: SigFigs) EXAMPLE: 25,425 or 23.932 *Both have 5 SigFigs* EXAMPLE: 0.0903 *3 SigFigs* cant EXAMPLE: 0.00034 *2 SigFigs* EXAMPLE: 0.200 *3 SigFigs* significant • Scientific Notation: Representation of a numbers 1-10 multiplied by 10 (A x 10^n) • Rounding: Dropping nonsignificant numbers There are only 3 SigFigs in each measurement so there can only be 3 SigFigs in your answer or your answer would be more accurate than your measurements! If the number after 3 were greater than five then you would round the 3 to a 4. (.3936 .394) r answer should have the same amount of SigFigs as your measurement with the least number of SigFigs EXAMPLE: 1.51g/ 3.845L=.3927178153 .393 decimal places as the measurement with the least number of decimal places. EXAMPLE: 8.7781 - 0.42589= 8.35221 (Since we have a measurements with 4 decimal places and one with 5 decimal places our answer should only have 4 decimal places). • Exact number: Counted number or defined number : There are exactly 2.54cm in an inch so 2.54cm is the exact number and SHOULDN’T be used to determine the amount of SigFigs in your answer. 2.54 1000 in. X 1 = 2540 cm ( So 1000 should be used to determine the amount of SigFigs in your answer which is why our answer has 4 SigFigs and not 3 or 1). SI Units • SI units are made up of a particular choice of metric units • Base units Length Meter (m) Mass Kilogram (kg) Time second (s) Temperature Kelvin (K) Amount of substance Mole (mol) Electric current Ampere (A) Luminous intensity Candela (cd) Area: length x length (m^2) Volume: length x length x length (m^3) Density: mass/unit volume (kg/m^3) Speed: distance/unit time (m/s) Acceleration: change in speed/ unit time (m/s^2) • Temperature Equation (TK-32˚F) x (100/180) Equation (T˚C) x (9/5) + 32 Equation TK-273 Equation ˚C+273 Equation (T°F + 459.67) x 5/9 Equation (TK ) x (9/5) – 4 Chapter 2: Atoms, molecules and ions Lecture 6 Questions to consider 1. What concepts were apart of John Daltons Atomic theory of matter? 2. Distinguish between the law of conservation of matter and the law of multiple Questions to consider: 1. What does the atom consist of? 2. Who discovered the electron? 3. Who measured the charge of electrons? 4. Who is Ernest Rutherford? 5. What is the charge of the electron? Questions to consider: 1. Understand what the nuclear structure and what it consists of. 2. What is an isotope? 3. Who is James Chadwick? 4. What is the charge of protons and neutral? 5. Understand the relationship protons and electrons share. ________________________________________________________________________ Lecture 7 Questions to consider: 1. Understand how to calculate average atomic mass 2. What is the importance of Dimitri Mendeleev 3. Understand how to properly read the periodic table 4. What is the difference between a metal nonmetal and metalloid 5. What are the two types of ions 6. Understand what an ionic compound is and how to create a proper ionic compound formula. ________________________________________________________________________ ___ Lecture 8 Questions to consider: 1. Know the difference between organic and inorganic compounds. 2. What type of organic compound contains only hydrogen and carbon? Know a few examples. 3. What are functional groups and types of functional groups? 4. Understand how to name ionic compounds. 5. Understand how to predict the charge and name monatomic ions, transition metal ions, and nonmetal main group element monatomic ions. 6. What is a polyatomic ion and know example. Lecture 9: Questions to consider: 1. What is a binary compound? Give examples 2. What is a hydrate? 3. What are the 10 prefixes 4. Understand how to balance a chemical equation  Ionic theory of solution and solubility rules • Ionic theory of solution ◊ Ionic theory: some substances produce freely moving ions when dissolved in water and can conduct an electric current in an aqueous solution. ◊ Electrolyte: substance that conducts electricity when dissolved in water EXAMPLE: NaCl ◊ Nonelectrolyte: substance that doesn’t conduct energy when dissolved in water. EXAMPLE: Sugar ◊ Weak electrolyte: substance that produces only a minimum amount of energy when dissolved in water. EXAMPLE: CH3COOH • Ionization of acetic acid CH3COOH is a reversible reaction which means the reaction can occur in both ways. • Solubility rules SOLUBLE ◊ Group 1A an ammonium compounds are soluble ◊ Acetates and nitrates are soluble ◊ Most chlorides, bromides and iodides are soluble • Exceptions: AgCl, AgBr, AgI Hg2Cl2, Hg2Br2, Hg2I2 HgBr2, HgI2 PbCl2, PbBr2, PbI2 ◊ Most sulfates are soluble •Exceptions: CaSO4 Ag2SO4 SrSO4 Hg2SO4 BaSO4 PbSO4 INSOLUBLE ◊ Most carbonates areinsoluble • Exceptions: Group 1A carbonates and (NH4)2CO3 ◊ Most phosphates are insoluble • Exceptions: Group 1A phosphates and (NH4)3PO4 ◊ Most sulfides are insoluble • Exceptions: Group 1A sulfides and (NH4)2S ◊ Most hydroxides are insoluble • Exceptions: Group 1A hydroxides Ca(OH)2, Sr(OH)2, Ba(OH)2 (3/31)  Molecular and ionic equation • Molecular equation: a chemical equation I which reactants and products are written as molecular substances. They may actually exist in solutions as ions. EXAMPLE: Pb(NO3)2 (aq) + 2Kl (aq)>>> PbI2 (s) + 2KNO3 (aq) • Complete ionic equation: a chemical equation in which strong electrolytes are written as separate ions in the solution and other products are written in molecular form • Spectator ions: an ion in an ionic equation that does not take part in the reaction. It appears as a reactant and product in complete ionic equations. • Net ionic equation: a chemical equation in which spectator ions have been omitted. It shows the reaction that actually occurs at the ionic level.  Types of chemical equations • Precipitation ◊ Precipitate: insoluble solid, and insoluble solid compound formed during a chemical reaction in solution. EXAMPLE: The above reaction is precipitation reaction • Acid base reactions ◊ Arrhenius acid: A substance that produces (H+) in H2O ◊ Arrhenius base: A substance that produces ◊ Bronsted acid: Proton donor ◊ Bronsted base: Proton acceptor ** Two arrows are used for things that don’t ionize completely** ◊ How to complete acid base reactions EXAMPLE: Ba(OH)2 (aq) +H3PO4 (aq) 1. Identify the acid and base, determine if the base is ammonia ( NH3) 2. H+ and OH- to form H2O and other ions from salt 3. Determine the solubility of the salt 4. Balance the equations If the base is NH3 I acid-base reaction 1. Identify the acid and base, determine if the base is ammonia 2. Identify anions and cations of acid 3. H+ and NH3 to form NH4+ then form salt with the anion and acid 4. Balance the equation ◊ How to write net ionic equations for acid base reaction EXAMPLE: HCl (aq) + NaOH (aq) 1. Complete the reaction 2. Write complete ion equation 3. Cancel spectator ions balance the equation  Strong and weak acid • Strong acid: acid that ionizes completely in water. ◊ Present entirely as ions ◊ Strong electrolyte ◊ Common strong acids: HNO3, H2SO4, HCLO4, HCl, HBr, HI ◊ One arrow • Weak acids: acid that only party ionizes in water. ◊ Present primarily as molecules and partly as ions ◊ Weak electrolyte ◊ Common weak acids: HF, CH3COOH, HCN, H2CO3, H2SO3, H3PO4 ◊ Two arrows  Strong and weak bases • Strong Base: Ionizes completely in water ◊ Present entirely s ions ◊ Strong electrolyte ◊ Common strong bases: LiOH, NAOH, KOH, CA(OH)2, Sr(OH)2, Ba(OH)2 • Weak bases: partly ionizes in water ◊ Present primarily as molecules ad partly as ions ◊ Weak electrolyte ◊ Example: NH3  Neutralization reaction • acid + base salt + water EXAMPLE: NaOH (aq) + HCl (aq) NaCl (aq) +H2O (l)  Gas forming reactions • Gas forming acid base reactions EXAMPLE: NA2S (aq) + 2HCl (aq) 1. Identify salt and acid 2. Write salt and acid into ions recombine ions to form two new compounds (a salt and an acid) (H2S (g), H2CO3 ( written as H2O (l) and CO2(g)), H2SO3 (written as H2O (l) and SO2 (g))  Molarity concentration • Concentration of solution: amount of solute present in a given quantity of solvent or solution • M( molarity)= moles of solute/ liter of solute( n/v) EXAMPLE: a sample of NaNO3 weighing 0.38 g is placed in a 50.0 mL volumetric flask. The flask is then filled with water tot eh mark on the neck, dissolving the solid. What is the molarity of the resulting solution? 1. Identify known and unknown measurements ( if known is In grams change it to moles and if known is in mL change it to L 2. Calculate the unknown • Dilution: procedure for preparing a less concentrated solution form a more concentrates solution (stock solution) ◊ Diluting a solution EXAMPLE: You are given a solution of 14.8 M NH3. How many milliliters of this solution do you require to give 100.0mL of 1.00 M NH3 when diluted? 1. Identify known and unknown measurements M initial= 14.8M V initial=? mL M finial= 1.00 mL V final= 100.0 mL 2. Calculate the unknown MiVi= MfVf Vi= (Mf)(Vf)/Mi= 6.76 mL


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