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CHM 113 Exam 2 Study Guide

by: Andrew Notetaker

CHM 113 Exam 2 Study Guide Bio 113

Marketplace > Arizona State University > Science > Bio 113 > CHM 113 Exam 2 Study Guide
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About this Document

This is a full study guide covering all topics that will be on Exam 2.
Udo Savalli
Study Guide
Redox reactions, hess's law heats of reactions general chemistry janet degrazzia engineering engineers thermochemistry, Thermodynamics, electrons, Electromagnetic Waves, atomic orbitals
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This 9 page Study Guide was uploaded by Andrew Notetaker on Sunday October 9, 2016. The Study Guide belongs to Bio 113 at Arizona State University taught by Udo Savalli in Spring 2015. Since its upload, it has received 113 views. For similar materials see Dinosaurs in Science at Arizona State University.


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Date Created: 10/09/16
Chapter 4 (Reactions in Aqueous Solution) What you should know for this test: The characteristics of and be able to identify the following types of reactions (some are from Ch. 3):  Single Replacement One atom replaces another atom in a reaction, typically a metal replacing a metal. AB + C -> CB + A  Double Replacement (metathesis) The cations of two molecules replace each other. AB + CD -> AD + BC o Precipitation o Acid/Base neutralization HCl + NaOH -> H O2+ NaCl ( And acid and a base form a salt and water) o Gas forming neutralization (watch for CO3 2- , SO3 2- , & S2- ions combining with H+ from an acid) AB +CD -> Salt & Gas Bases Capable of forming a gas: Sulfite Ion - SO3- Na S2 (a3) + 2HCl(aq)  SO (g2 + H O(l) 2 2NaCl(aq) Sulfide Ion - S- Na S2aq) + H SO (2q) 4H S(g) + Na S2 (aq) 2 4  Decomposition AB -> A + B  Combination (Synthesis) A + B -> AB  Combustion A hydrocarbon and oxygen to react to release energy, carbon dioxide and water. The definition of an electrolyte and how to identify strong, weak, and non-electrolytes The meaning of “dissociation” and “solubility” Electrolytes dissolve in water to form ions. Strong electrolytes completely dissolve in water and form ions. Weak electrolytes do not fully dissolve in water, but form ions. Non-electrolytes dissolve in water but do not form ions. Group I and group II containing molecules are typically strong electrolytes. How to use solubility rules (a table of solubility rules will be provided on the exam - see practice exam) How to compare the relative strengths of two strong electrolytes (by # of ions formed) The more ions produced, the stronger the electrolyte. How to write a balanced molecular, complete ionic, and net ionic equations for single/double replacement and acid/base reactions Balanced molecular- Write out full equation and make sure there are equal amounts of atoms on both sides by altering coefficients Complete ionic- Write out the full equation with ions broken apart with their charges Net ionic- Remove ions that appear on both sides of the equation (spectator ions) and rewrite How to predict products when ionic compounds are mixed in aqueous solution Using the activity table and solubility rules it can be determined what products will be formed. The definition of spectator ions and how they relate to a net ionic equation Ions that appear on both sides of an equation, reactants and products. These are removed from the final net ionic equation. The definition of Arrhenius, Bronsted-Lowry, and Lewis Acids and Bases Acid/base reactions; know how to write net ionic equations for strong acid/base, weak acid/strong base, weak base/strong acid reactions. Arrhenius defined acids as substances that increase the concentration of H ion when dissolved in water. - He defined bases as substances that increase the concentration of OH ion when dissolved in water. Bronsted-lowry definition of an acid is a proton donor. Bases are proton acceptors. Strong acid/base reaction: net ionic only water is left Weak acid/strong base: acid and water is left over Weak acid, weak base: Both are left over Strong acid/weak base: base is left over How to assign oxidation numbers to elements in ionic or molecular compounds Refer to periodic table Determine if oxidation-reduction takes place in a chemical reaction; be able to ID oxidizing/reducing agents (see Mastering assignment question: Oxidation-Reduction Reactions. Reduced: Gains electrons Oxidized: Loses electrons What an activity series is and how to use it; be able to predict the reactivity of a metal with other metal cations; metals with water or acids. Look at activity series table, the higher up a metal is on the table, the elements below it can be reduced by it. Li is the strongest reducing metal. Using molarity as a conversion factor to calculate moles and volume (don’t worry about mass %, mole fraction, or molality for this exam) Molarity is the moles of the solute divided by the moles of a solution. How to perform stoichiometry calculations that involve molarity (many types of problems represented by this statement) Titration calculations – standard equivalence point calculation (watch out for diprotic acids and group II metal hydroxides). Moles of acidic hydrogen = moles base Concentration is mol of H / original volume of acid Chapter 5 (Thermochemistry) What you should know for this test: The definition of energy and the difference between kinetic and potential. Energy is the capacity to do work or transfer heat, kinetic energy is the energy of motion, potential energy is stored energy from attractions/repulsions. The definition of heat and some common units of heat energy. Heat is the transfer of energy from a hotter to cooler object, common units of heat energy are Joules, J, or Kilojoules, kJ. The definition of Internal Energy (E) and ∆E Internal energy is the sum of the kinetic energy and potential energy of a system. The change in internal energy is the sum of final energy minus the sum of the initial energy. The meaning of system and surroundings and what sign (+ or -) accompanies changes in heat and work from system and surroundings System is the part of the universe we are interested in studying (a chemical reaction), surroundings are the rest of the universe or anything but the system. ΔE = q + w to determine internal energy For q + means gains heat  Means loses heat For w + means work done  Work is done by on system system For ΔE + means net gain of  Means net loss of energy by system energy by system The definition of Enthalpy (H) and ∆H That ∆H = qP and that ∆E ≈ ∆H and therefore for chemical reactions, ∆E ≈ qP ΔE= q + w = q - PΔV For most chemical reactions:ΔE= q p Which means the change in energy during a chemical reaction under constant pressure is approximately equal to the heat involved in the reaction. Q ps called Enthalpy (H) When does ΔH =ΔE? The enthalpy change for a reaction equals the change in internal energy when work is negligible. Therefore ΔE approximately equals ΔH The meaning of exothermic and endothermic reactions and what the sign of ∆H is for each Exothermic releases energy/heat, the change in H is negative Endothermic absorbs energy/heat, the change in H is positive Relative energy-releasing and energy-requiring processes (bond-breaking and bond-forming) for exothermic and endothermic reactions. To form bonds (anabolic) energy is absorbed, endothermic To break bonds (catabolic) energy is released, exothermic How to relate specific heat, mass, energy and temperature change. Energy of a reaction, Q can be determined through mass x specific heat x change in temperature. Q = m x s x ΔT q= heat (J) M = mass (g) S = specific heat (J/g*°C) ΔT=(T finalTinitial How to use the specific heat equation to calculate q, m, Cs, or ∆T Use algebra to solve for each specific variable. How to use calorimetry data to calculate q for a reaction or process, e.g., mixing of two solutions. Heat released/consumed by a reaction = heat gained/lost by a calorimeter solution 3.99 of ammonium nitrate NH NO 480.3g/mol) dissolves in 60.0g water in a calorimeter. The temp drops from 23.0°C to 18.4°C (ΔT = -4.6°C) q= 63.88 x 4.184 x(-4.6) q=1,229 1,229/(3.88g/80.1g/mol)= -25.37 kJ/mol How to use Hess’s Law and standard enthalpies of formation (∆Hºf) to calculate ∆Hºrxn Hess's Law states: If a reaction is carried out in a series of steps, ΔH for the overall reaction will be equal to the sum of the enthalpy changes for the individual steps. 1. ΔH for a reaction forwards is equal in size, but opposite in sign to ΔH of the reverse reaction 2. Enthalpy is extensive, if the reaction is multiplied, ΔH is multiplied 3. ΔH for a reaction depends on the state of the products and reactants That tabulated ∆Hºf values are: for standard states, for 1 mole of formation, extensive, reversible by changing the sign An element in its normal state has an enthalpy of formation of 0, 1 mol of the product is made by the enthalpy of formation. How to use ∆H values along with balanced equations to determine enthalpy changes for specific amounts (mols or grams) of reactants or products 2NI 3(s)→N 2(g)+3I 2(g), ΔH rx=−290.0 kJ When 2 mol of NI 3decomposes, the enthalpy of the reaction is −290.0 kJ. If 5mol ofNI 3decomposes, the enthalpy of the reaction will be (5/2×−290.0) = −725.0 kJ . What will not be on this test: • Section 5.8 (Food and Fuels) Chapter 6 (Electronic Structure of Atoms) What you should know for this test: The definitions of frequency and wavelength for electromagnetic (EM) radiation. Frequency- The number of waves passing a given point per unit of time. Wavelength is the distance between corresponding points on adjacent waves. The appearance of the EM spectrum and what regions are on it (Gamma rays, X-Rays, UV, Visible, IR, microwave, TV waves, Radio Waves) The order of visible light (and other forms of EM radiation) by increasing or decreasing energy, frequency, and wavelength. Gamma rays have most energy, frequency and shortest wavelength. Radio waves have the least energy, frequency and highest wavelength. Conceptual relationships between frequency, energy, and wavelength (how one changes if you change the others—can be predicted by formulas) For waves traveling at the same velocity, the speed of light c is shorter than the wave length, the greater the frequency. (Inverse) C =V How to use formulas to calculate energy, frequency, and wavelength (formulas and constants will be provided) E = hν = hc/λ h is Planck’s constant, 6.63 x 103J•s What a photon is and how does it relate to the wave nature of light. Particles of light that are ejected which carry the electromagnetic radiation. Know about the Photoelectric Effect, Blackbody Radiation, Continuous and Line Spectra When light is shined on metals, particles are ejected by the surface of the material at a minimum frequency. Blackbody radiation is light emitted from heated metal atoms which is not continuous. Diffraction produces a continuous spectrum. Diffraction of light from pure elements results in a line spectrum, not a continuous spectrum. What happens when regular white light is passed through a prism and when an atom of a single element is excited and what this tells us about energy levels within an atom When white light is passed through a prism, this produces a continuous spectrum of all wavelengths of visible light. When an atom of a single element is excited, its energy levels increase with potential energy. The ordering of energy levels in an atom and what happens to the spacing between levels as the shell # increases As energy levels and shell numbers increase, electrons get further away from an atom. The meaning of atomic absorption and emission Specific wavelengths of electromagnetic radiation are absorbed by an element The release of energy as light is known as emission at discrete, specific wavelengths. What happens to the wavelength/energy of emitted or absorbed light as an electron travels between energy levels Energy increases as electrons transfer from higher energy levels to lower energy levels. How to use Bohr’s equation to calculate hydrogen absorption/emission energy/wavelength/frequency. de Broglie formula of mass, velocity, and wavelength as it pertains to the wave-matter relationship. m = h / c and = h / mc The meaning of the Heisenburg Uncertainty Principle The position/ velocity can never be determined simultaneously The meaning of the Schrödinger wave equations as they relate to an electron probability density map. Solving the wave equations gave n, the energy or distance from nucleus, l the shape and M t1e direction. How to write a ground state electron configuration for elements #1-36 on the periodic table (with or without a noble gas core condensed configuration), including ions of those elements. Oxygen- 1s 2s 3p or [He]2s 3p 4 Also be able to create orbital diagrams (boxes with up/down arrows). Oxygen- Why elements like Cr (3d5 4s1) and Cu (3d10 4s1) fill and half-fill the 3d level before the 4s level. Cabriac didn’t explain this but its just the way it is. The elements are more stable this way. The order of electron stacking into s, p, d, f sublevels (using the 1s 2s 2p, etc. notation or the orbital boxes) 1s 2s, 2p 3s, 3p, 3d 4s, 4p, 4d, 4f 5s, 5p, 5d, 5f, (5g) 6s, 6p, 6d, 6f, (6g, 6h) Diagonal rule!! Using probability density plots to determine electron location Refer to graphs The shapes of s, p, d-orbitals S orbitals are sphere shaped. P orbitals are dumbbell shaped, d orbitals have four lobes, (cloverleaf) the last d orbital resembles a p orbital with a doughnut around the center. The definition of Hund’s Rule, the Pauli Exclusion Principle, and the Aufbau Principle and how to recognize when they are followed or violated Hund’s rule- electrons fill singly into different orbitals of with the same spin before pairing. Pauli exclusion- two electrons can occupy the same orbital but they must have opposite sign Aufbau- electrons fill into the lowest energy orbitals first. What will not be on this test: Writing electron configurations for elements above #36 (Kr) How to write quantum numbers (n, l, m). However, you ARE responsible for the underlying concepts (shells, subshells, electron spin, etc.)


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