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by: Gabrielle Notetaker

Chemistry CHM103

Gabrielle Notetaker
GPA 3.95

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Notes and study guides for Chem 103
Intro to chemistry
Study Guide
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This 6 page Study Guide was uploaded by Gabrielle Notetaker on Saturday April 2, 2016. The Study Guide belongs to CHM103 at University of Rhode Island taught by Dombi in Spring 2016. Since its upload, it has received 56 views. For similar materials see Intro to chemistry in Chemistry at University of Rhode Island.


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Date Created: 04/02/16
Chemistry Test #3 Study Guide  Chapter 5 Chemical Equations:   Chemical Equations are written in terms of reactants and products o Reactants are the substances on the left side of the arrow  o Products are the substances on the right side of the arrow   Balanced chemical equations are equations in which the number of atoms of each  element in the reactants is equal to the number of atoms of the same element in the  products (This is achieved by conservation of matter)  o An equation is balanced by inserting coefficients o 2H 2+ O 2 2 H O2   Because both the reactant and product sides have 4 H atoms and 2 O  atoms Types of Chemical Reactions      Redox is the combination of the words reduction and oxidation.   Oxidation means to combine with oxygen to lose, to lose hydrogen, to  lose electrons and to increase oxidation number (becomes reduced)  Reduction means to lose oxygen, combine with hydrogen, gains  electrons and to decrease the oxidation number. (Became oxidized)  In a redox reaction the substance that contains an element that is oxidized during the reaction is called the reducing agent  In a redox reaction the substance that contains the element that is reduced during the reaction is the oxidizing agent.   Oxidation Numbers are positive or negative numbers assigned to elements in chemical  formulas according to a set of rules  o Rule 1: The oxidation number of any uncombined element is 0 o Rule 2: The oxidation number of simple ions is equal to the charge on the ion  o Rule 3:  The oxidation number of group IA and IIA elements when they are  compounds are always +1 and +2  o Rule 4: The oxidation number of hydrogen is always +1 o Rule 5: The oxidation number of oxygen is always ­2 (except in peroxides when  its ­1) o Rule 6: The algebraic sum of the oxidation numbers of all atoms in a complete  compound equals 0 o Rule 7:  The algebraic sum of the oxidation numbers of all the atoms in a  polyatomic ion is equal to the charge of the ion  Decomposition Reactions­ one substance is broken down into two or more similar  substances o A B+C  Combination Reactions­ or more substances react to form a single substance. o A+B C  Single Replacement Reaction­ is always redox reactions because they occur when one  element reacts with a compound, displaces one of the elements from the compound and  becomes a part of a new compound. o A+BXAX+B  Double Replacement Reaction­ is never redox reactions. These reactions often take place  when something is dissolved in water. In typical reactions, two dissolved compounds  react and exchange partners to form tow new compounds.  o AX+BYAY+BX o A subset of Double Replacement Reactions is Ionic Equations  Ionic compounds and some polar covalent compounds break apart  (dissociate) when they dissolve in water and form ions  In a total ionic equation, all soluble ionic substances are represented by  the ions they form in solution. Substances that do not dissolve or that  dissolve but do not dissociate into ions are represented by their formulas.  In a net ionic equation, only unionized or insoluble materials and ions  that undergo changes as the reaction proceeds are represented.  Any ions that appeared on both the left and right side of the total  ionic equation are called spectator ions and are not included in  the net ionic equation.  Energy and reactions­ all reactions either absorb or give up energy. o Most often, all or most of the energy takes the form of heat o Exothermic Reaction: chemical reactions that release heat o Endothermic Reaction: chemical reaction in which heat is absorbed  Moles­ The mole concept can be applied to balanced chemical equations and used to  calculate mass relationships in chemical reactions o  Steps for the Factor Unit Method o Step 1:  Write down the known or given quantity.  Include both the numerical  value and units of the quantity. o Step 2:  Leave some working space and set the known quantity equal to the units  of the unknown quantity. o Step 3: Multiply the known quantity by one or more factors, such that the units  of the factor cancel the units of the known quantity and generate the units of the  unknown quantity. o Step 4: After you generate the desired units, do the necessary arithmetic to  produce the final answer.  Reaction Yields: The amount of product calculated in the examples is called the  theoretical yield.  The amount of product actually produced is called the actual yield o  % Yield = (actual yield / theoretical yield) x 100 Chapter 6 Physical Properties of Matter  Physical: density, shape, compressibility and thermal expansion  o Density= Mass/Volume  o Shape: shape matter takes dependent on physical state o Compressibility: volume change of a sample that results from a pressure change  o Thermal expansion: change in volume of sample resulting from temperature  change  Kinetic Molecular Theory of Matter: KE=mv / 2 (m=mass, v=velocity) o Postulate 1:  Matter is made up of tiny particles called molecules, which includes  atoms. o Postulate 2:  The particles of matter are in constant motion and therefore possess  kinetic energy. o Postulate 3:  The particles possess potential energy as a result of repelling or  attracting each other. o Postulate 4:  The average particle speed increases as the temperature increases. o Postulate 5:  The particles transfer energy from one to another during collisions  in which no net energy is lost from the system.  Potential Energy: stored energy o Potential energy is the energy a particle has as a result of being attracted to or  repelled by other particles. All the molecules in one corner of the container have  high potential energy. o Cohesive Force: an attractive force between particles  All particles have very weak Induced dipole attractions.  Some particles have stronger dipole interactions or hydrogen bonds. o Disruptive Force:   A disruptive force results from particle motion.  It is associated with  kinetic energy.   To melt or evaporate requires added energy, usually heat o States of Matter  Solid: high density, definite shape, small compressibility and very small thermal  expansion  o Types of Solids  Covalent Network of Solid­ High melting points, extremely hard  Metallic Solids­ Middle melting points, pliable, loose electrons  Ionic Solids­ High Melting points, brittle, some water soluble  Molecular Solids­ Low melting points, often soft, many hydrophobic  Liquid: high density, an indefinite shape that depends on the shape of its container, a  small compressibility, and a small thermal expansion.  Gaseous state­low density, an indefinite shape that depends on the shape of its container,  a large compressibility, and a moderate thermal expansion.  The Gas Laws­ mathematical equations that describe the behavior of gases as they are  mixed, subjected to pressure or temperature changes, or allowed to diffuse o o Combined Gas Law:  Avogadro’s Law: V/n=k (v is volume, n is the amount of gas in moles and k is a  proportionality constant (STP= 0 C, and 1.00 atm)  o 1 mole of any gas molecules has a volume of 22.4 L at STP   Ideal Gas Law: PV=nRT (p=pressure, V=volume, T is the gas temperature in Kelvin, n is the number of moles of gas and R= 0.821 L atm/ mol K  Ideal Gases vs. Real Gases o Ideal gases don’t exist, but if they did they would behave exactly as predicted by  the gas laws at all temperatures and pressures unlike real gases, which deviate  from the gas laws especially at low temperatures.  Dalton’s Law of Partial PressureP total  P  individgases  o As more gas is added the pressure increases, whether it is a new gas or more of  the first gas   Grahams Law o Effusion is a process in which a gas escapes from a container through a small  hole in the container o Diffusion is a process that causes gases to spontaneously mix when they are  brought together.   Increasing size of molecules increases the boiling point. ‘ o H bonding has the highest bp  o Dipole has the medium bp  o IDDI has the lowest bp   Changes in State are caused by the adding or removal of heat  o Evaporation: liquid to gas; endothermic o Sublimation: solid to gas; endothermic o Melting/Fusion: solid to liquid; endothermic o Condensation: gas to liquid; exothermic o Deposition/Condensation: gas to liquid; exothermic o Freezing/Crystallization: liquid to solid; exothermic  Vapor Pressure is the pressure exerted by a vapor that is in equilibrium with its liquid  Boiling point of a liquid is the temperature at which the vapor pressure of the liquid is  equal to the prevailing atmospheric pressure o As elevation increases, boiling point decreases  Specific heat of a substance is the amount of heat required to raise the temperature of  exactly 1 g of a substance exactly 1°C.  o The specific heat of ice is 0.51 g / (°C cal). o The specific heat of water is 1.00 g / (°C cal). o The specific heat of steam is 0.48 g / (°C cal). o Heat = q=mc∆T    m = mass (g) c = specific heat    Heat of Fusion of a substance is the amount of heat required to melt exactly 1g of a solid  substance at constant temperature. Ice to water = 80 cal /g = water to ice. o There is no temp symbol in constant, since no temp change. o How much heat is necessary to raise the temperature of 100 g of ice from 0 °C to  liquid water at 0 °C?  Cal = 100 g (80 cal)/g  = 8000 cal  Heat of vaporization of a substance is the amount of heat required to vaporize exactly 1g  of a liquid substance at constant temperature.  540 cal /g water to steam   Chapter 7 Solutions   Solutes vs. Solvents o Solvents: larger quantity, usually the solution is in the same phase as the solvent o Solutes: smaller quantity  Solubility of the solute is the maximum amount f the solute that can be dissolved in a  specific amount of solvent under specific conditions of temperature and pressure  o Solubility always goes up with temperature increase, except in the case of gases  in liquids (then it will decrease)  Immiscible: liquids that are not soluble in other liquids. The liquids will not mix to form  solutions   Saturated solutions: unstable solution that contains more solute than solvent under  specific circumstances  Solution process: involves interaction between solvent molecules (often water) and the  particles of the solute. An example of solution process for an ionic solute in water.   Relative Solubility in covalent molecules: o Nonpolar covalent (IDDI)­ least soluble o Polar Covalent (Dipole)­ some solubility  o Polar Covalent (H Bonds)­ most soluble


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