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Chemistry 1151k Exam 2 Study Guide

by: Stephanie Argueta

Chemistry 1151k Exam 2 Study Guide CHEM 1151K

Marketplace > Georgia State University > CHEM 1151K > Chemistry 1151k Exam 2 Study Guide
Stephanie Argueta

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Everything from chapter 5 to chapter 7 and some work problems written out if you need up. THERE WILL BE AN UPDATE IN A FEW DAYS: I have some confusion on some principles that i don't want to publis...
Dr. Nilmi Fernando
Study Guide
radiation, Chemistry, covalent bonds, ionic bonds, Lewis Structure, resonance, compounds, Oxidation-Reduction
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This 12 page Study Guide was uploaded by Stephanie Argueta on Sunday October 16, 2016. The Study Guide belongs to CHEM 1151K at Georgia State University taught by Dr. Nilmi Fernando in Fall 2016. Since its upload, it has received 66 views.


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Date Created: 10/16/16
Chemistry 1151k Fernando Exam 2 Study Guide Highlight = Important Principle Highlight = Important Concept Highlight = Key Term Chapter 5: Nuclear Chemistry Radioactive elements emit radioactive waves - Elements with an atomic number less than 20 are radioactive b/c there will be more repulsion than attraction, therefore atom will become energetically unstable - Types of radiation: - Alpha: two protons & two neutrons; mass number of 4 and atomic number of 2 AND a charge of 2+ - - Beta: charge of 1- and mass of 0 - - Positron has positive 1+ charge w/mass of 0. It is produced by an unstable nucleus when a proton is transformed into a neutron and positron - Gamma rays: actually high energies of radiation; when an unstable nucleus is undergoing a rearrangement of its particles to make the nucleus more stable - - Biological effects of radiation : - Radiation goes through your: bone marrow, skin, reproductive organs, and intestinal lining, which WILL GIVE YOU CANCER OR BODY DAMAGE - Radiation - Air travel - Tissue - Protection depth - Alpha - 2-4 cm - 0.05mm - skin/clothing - Beta - 200-300 - 4-5mm - Thick lab cm coat, mask, gloves - Gamma - 500cm - >50cm - Inches of lead & a foot of concrete - Gamma will do enough damage however alpha particles will do the most damage if there is a lot of it - Radiation measureme nt - Curie (Ci)- original unit of activity, the number of disintegration that occur in 1 second for 1 gram of radium (3.7x10^10 disintegrations per second) - Bequerel (Bq)- similar to Ci but is only 1 disintegration per second - Rad (radiation absorbed dose ): measures the amount of radiation absorbed by a gram of material such as body tissue à SI unit: Gray (Gy) - Rem (Radiation equivalent in humans): measures the biological effects of different kinds of radiation à to determine this: absorbed dose is multiplied by a factor that adjusts for biological damages caused by a particular form of radiation - Rem= rad x factor - 1 rem = 1000mrem 1 Sv (Sievert) = 100 rem Medical applications using radioactivity - Scams with radioisotopes - Positron emission tomography (PET) à these produce 3D images of organs - Computed tomography- monitors absorption of 30,000 x-ray beams directed at layers of the organ - Magnetic resonance imaging (MRI) à absorption of energy when protons in hy drogen atoms are excited by a strong magnetic field (this process does not use x -ray radiation) Nuclear Energy Generation: - Nuclear fission: breaking the nucleus - Nuclear fusion: combining the nucleus ( this is an experimental technique b/c it is expensive and time consuming) Half Life of Radioactive elements: time it takes for a substance to be reduced to ½ of the orginal mass Here is how to solve them: Chapter 6: Compounds Ionic Compounds: - Positive ions: mostly metals (solids) that remove an electron - Across the periodic table, ion size decreases - Positive ions are SMALLER than their or iginal atom size - Negative Ions: non metals that GAIN electrons - Repulsion is greater than attraction, so negative ions get BIGGER across the table. - So, negative ions are BIGGER t han their original atom size - Naming these beauties: - Qualities of ionic compounds: - Very high melting point - Crystalline 3D lattices - Strong electrostatic attraction - Mostly SOLIDS @ room temperature - Some positive ions have variable charge, meaning they can either a + or 2+ charge. Please refer to Fernando’s table on icollege to find all these charges (since it ’s not my work I can’t publish them here) - YOUR FAVORITES: POLYATOMIC IONS ! PLEASE memorize these AND their charges: Covalent bonds: sharing electrons between atoms (AKA molecules) - Prefixes: mono (1); di (2); tri (3); tetro (4); penta (5); hexa (6) - Represented by Lewis structur es: valence electrons of all atoms are arranged to give octets (8 atoms) EXCEPT for hydrogen that only takes in 2 electrons - Other exceptions you should remember: - Any elements in energy level 3 (or in period 3 in the periodic table) can have more than an octet due to empty D orbitals **THIS IS ONLY FOR ENERGY 3** - Be can only have 2 electrons (a duet) - When drawing your Lewis structure, a molecule can also have resonance, which is basically alternating your bonds. You add this symbol to show resonance: ßà - Exceptions: - Hydrogen will always have one bond w/no lone pairs - Halogens only form one bond and lone pairs - Shaping of electrons - Linear: two electron groups and two bond groups OR one bond group & lone pairs - Linear bonds have 180 degree angles, can be polar (different electronegativity & unequal sharing of electrons) or non polar (identical electronegativity) based on the direction of dipoles - Trigonal planar: 3 electron groups & 3 bonding groups; Bond angle is 120 degrees, can be polar/ non-polar - Tetrahedral: 4 electron groups w/ 109 degree angle; if you have the same 4 elements in the group, the bond will be overall non polar -Valence Shell Electron -Pair Repulsion (VSEPR): electron groups are arranged as far apart as possible around the central atom to minimize the repulsion between negative charges. Structures will adopt to the minimum energy structure to avoid repulsion. Electronegativity - Ability to attract the shared electrons in a chemical bond - Bonds based on polarity (i n increasing order of polarization): - Nonpolar covalent bonds (weakest): must have polarity between 0 and 0.4 - Polar Covalent Bonds: polarity must be between 0.4 and 1.8 - Ionic bonds (strongest): polarity is above 1.8 - REMEMBER: - Electropositive elemtns GIVE their electrons (found far left of the periodic table) - Electronegative elements TAKE IN electrons (found in the far right of the table) - Therefore, electronegativity INCREASE as you go across the table but DECREASES when you go down a group - A bond has two electron s; a double bond has four; a triple bond has 6 - When writing your Lewis structure, you put the most electronegative element in the middle of your structure! Types of interactions between molecules: - Intermolecules: in between molecules - Ex: Ionic bonds exist in lattices and no individual units are isolated. ALL IONCIS INTRA AND INTER - Intramolecules: inside molecules - Outside covalent bonds, there are three bonds than could happen: - Hydrogen Bonding - Dipole-Dipole interactions - Van der waal forces Chapter 7: Reactions & Chemical Equations Structure and balancing Chemical Equations: Types of Reactions: - Combination reaction - Decomposition Reactions Displacement/Replacement reactions - Combustion Reaction: - Oxidation reductions: OXIDATION-REDUCTION REACTION: v Remember this: Oxidation is the REMOVAL of electrons o One species (molecule) OXIDIZES when you remove electrons, such as ANY METAL ION/ANY POLYATOMIC ION o One species REDUCES when you GAIN electrons, such as ANY NONMETAL ION/POLYATOMIC ION. v Before you attempt ANY problem to look for oxidation numbers, memorize these rules o 1) An element that is by itself in an equation has an oxidation number of zero. § Fe + O2 à FeO 2 In this case, the Fe and O on the REACTANT side has an oxidation number of 0. Ignore the products for n w o Sometimes group elements have the same oxidation number as their CHARGES. So, Group 1A has a charge of 1+ AND oxidation number of 1+. o Same for Groups 2A: Oxidation number and charge number is 2+ o F (Fluorine) always has an oxidation number of 1 o EXCEPTIONS FOR WHEN GROUP # DOES NOT EQUAL OXIDATION NUMBER: § Cl (Chlorine) will have an oxidation number of 1 § O will have an oxidation number of +2 ONLY WHEN it is with F (Fluorine § Hydrogen (H) is mostly +1, however in meta hydroids (MgH2, NaH), H will have a charge of 1 o Total oxidation number in a neutral compound will always be ! § Ex: Fe2O3 = 0 o And sometimes oxidation number in POLYATOMIC ions will be the same as their charge (just in case, make sure you memorize the POLYATOMIC ions chart) v Got everything memorized? Good. Now, here is how you tackle a problem involving how to find oxidation numbers o Let’s take the equation 2à 2Fe 32O3 o Since Fe and O2 are by themselves ON THE REACTANT SIDE, they have an oxidation number of 0 o Now take a look at the product side 2O3 Fe § Since O3 is not by F, it has an oxidation number of -2 (same as its ionic charg ) § Since you have 3 O, your total charge for O is -6 (-2 3) o Remember the 5 rule of all compound the oxidation number is 0/neutral? It applies to 2FO 3 § Think it as like this: 2O 3 = 0 § Second step: Fe2 + (-6) = 0 • The -6 is from the total charge for O § Third step: Fe2 6 = 0 Fe2 = +6 § Fourth step: divide by 2 Fe = +3 § And now you know for one Fe has an oxidation number of +3 o Okay, so now we know that in the PRODUCT side, O has a charge of -6 and since we have two Fe, it has an oxidation number of +6. How do we know which is oxidized and which is reduced? § Remember, in oxidization we REMOVE electrons , so let’s see which element REMOVED electrons: • Fe was originally 0 on the REACTANT side, but on the PRODUCT side it became +6, how did it do that? o In order for become a POSTIVE ION, it must REMOVE electrons. So, Fe LOST 6 ELECTRONS ( 6 e-) and became Fe +6. So Fe OXIDIZED. • As for O, it was originally 0 on the REACTANT side, but it became -6 on the PRODUCT side. Guess how it did it? o YES! It GAINED electrons (6 e-) to become a NEGATIVE ION. So, O became REDUCED. o Yay! Not so bad, was it?? But before we can try another problem, I have to tell you about oxidizing agents and reducing agents. § Basically oxidizing agents HELPS reducing species (molecules), so whichever element REDUCES is the oxidizing agent. In our example, O would be our oxidizing agent. § Reducing agents HELPS oxidizing species, so which ever element OXIDIZES is the reducing agent. In our last example, Fe would be the reducing species o These are weird wordings but if you remember this you won ’t forget it! v Okay, last one and I hope you can try this first w/o checking the answer below. Remember, check the first step-by-step problem to see if you’re following it correct: C2H 2 + O 2 à CO 2 + H 2O Tip: don’t worry about H 2O, focus on CO 2 o On the product side, O has a total charge of -4 (2x-2), so in order to NEUTRALIZE (make it 0), C has to be +4 o On the REACTANT side, C has a charge of -1 b/c : § C2H2 = 0 H has a charge of +1 always UNLESS it’s a meta hydroid § C2 + (+2) = 0 § C2 = -2 § C= -1 o So, C -1 changed into C +4, so it REMOVED 5 electrons to become POSITIVE. So this is your oxidized species AND your reducing agent o As for O, it was originally 0 (remember, single elements always have an oxidation number of 0) and changed in to O - 4, so it GAINED electrons to become NEGATIVE. So O is your reduced species AND your oxidized agent. v Phew! Yeah, it’s strange and if you’re still confused, try these tips: o Remember how ions are formed o Don’t worry if you still don’t get it, practice makes perfect! o If it’s killing you, take a break. Eat ice cream. Watch one episode on Netflix. You will get this! Acid Base Reaction: v Memorize these equations, she ’ll mostly ask what type of reaction it is (they’re all acid ba e): o HCl + NH4OH à H 2O + NH 4Cl o H+ + OH- à H2O o HNO 3 + NaCH- à NaNO 3 + HO o H2SO 4 + NaOH à H O + Na 2SO


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