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FINAL EXAM Study Guide

by: Briana Marcy

FINAL EXAM Study Guide CHEM 211-002

Briana Marcy
GPA 3.8

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This study guide is mostly conceptual, but should be helpful when studying for the final, along with the ACS study guide.
General Chemistry 1
Paul Cooper
Study Guide
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This 23 page Study Guide was uploaded by Briana Marcy on Friday December 11, 2015. The Study Guide belongs to CHEM 211-002 at George Mason University taught by Paul Cooper in Summer 2015. Since its upload, it has received 124 views. For similar materials see General Chemistry 1 in Chemistry at George Mason University.


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Date Created: 12/11/15
CHEM 211 Dr. Cooper Final Exam Study Guide Briana Marcy Chapter 1 Know the differences between physical and chemical properties and changes -Basically, if the composition changes, it’s chemical, if not- physical *see textbook for examples of physical vs. chem. Changes Total energy equals potential plus kinetic energy LOWER states of energy are the most stable and are preferred over high energy states “Energy is neither created nor destroyed” It is always conserved, and can be converted to different forms Antoine Lavoisier, 1743-1749 He was the “father” of modern chemistry The basic fundamental of chemistry is the existence of atoms and molecules, which is really just a very good and well-tested theory Make sure you know how to do chemical problem solving Remember: units are manipulated in the same way as numbers See textbook pages 32-39 for practice/ review How to measure the volume of a solid? Measure the volume of water before and after placing the solid in the container Density=mass/volume or D=M/V If the temperature and pressures are known, the density of a substance is a physical property with a specified value The unit for density is g/cm^3 1 All phrases in quotations are taken directly from the textbook or Dr. Cooper’s powerpoints, as they cannot be reworded or the wording is ideal. Temperature- Kelvin is universally used in science and all values are above zero, and zero is “absolute zero”- nothing can be colder than that Make sure to know how to convert between Kelvin and Celsius (273.15 degree difference) Significant Figures (sig figs) All measurements have some amt of uncertainty, so the digit on the far right is always estimated. But the more sig figs there are, the more certain your answer is So which digits are significant?  NOT zeros that are only used to position the decimal point  Zeros that END a number are significant (whether before or after decimal point, both places count) but there HAS TO BE a decimal So for example, 700. has 3 sig figs but 700 only has one If you want practice for sig figs, see the sample problem on page of the textbook Rules for Sig Figs in various calculations 1.)In multiplicaton and division- answer will have same # of sig figs as the measurement with the fewest SIG FIGS 2.)In addition and subtraction- the answer will have the same # of DECIMAL places as the smallest measurement in the problem Rounding? Typically you should round at the end to avoid errors Rounding Rules  For digits more than 5, the preceding # goes up 1, if less than 5 it doesn’t change  If the 5 is removed but nothing (or zeros) are after it, even numbers stay the same but odd #s go up 1 Note: exact #s don’t limit the # of sig figs in a calculation Precision, Accuracy, and Error Precision means that the data are all close together Accuracy means the data are close to the correct value Random Error when values are both above and below the actual value Chapter 2 The Components of Matter Element- the most basic part of a substance, has only one type of atom and CANNOT be broken down any further Molecule- contains 2 or more atoms that are bound together chemically- so it works as an independent unit Compound- a substance made up of 2 or more chemically combined elements Mixture- a grouping of 2 or more elements that are physically combined Law of Mass Conservation “The total mass of substances present does not change during a chemical reaction” *See powerpoint slide 2-7 for a review Law of Definite, or Constant, Composition “No matter the source, a particular compound is composed of the same elements in the same parts (fractions) by mass” Law of Multiple Proportions “If elements A and B react to form two compounds, the different masses of B that combine with a fixed mass of A can be expressed as a ratio of small whole numbers” Dalton’s Atomic Theory He said that:  All matter is made of atoms  Atoms of one type of element can’t be converted to completely different atoms  Atoms in a certain element are unique from atoms in other elements  “Compounds result from the chemical combination of a specific ratio of atoms of different elements” Mass conservation Atoms can’t be created or destroyed Every atom has a fixed, specific mass that “constitutes a fixed fraction of the total mass in a compound” Mass of an electron= mass/charge x charge Names of interest: Millikan and J.J. Thomas, worked to find the mass of an electron Millikan conducted the oil-drop experiment Rutherford is known for his -scattering experiment and for discovering the nucleus of the atom Make sure to know general characteristics of an atom (protons, neutrons, electrons, their charges, etc.) When naming binary ionic compounds, the metal comes first Atomic Symbol, Number and Mass X= atomic number of the element A= mass number (A=Z+N) Z= atomic number (#protons) N= neutrons in the nucleus ISOTOPES are atoms of the same element w/ the same number of everything EXCEPT nuetrons, which in turn ups the mass number ***Know where metals, metalloids, and non-metals are on the periodic table You should also be familiar with the position of noble gases and often-used elements Molecules consist of two or more atoms bonded by the sharing of electrons (most covalent substances consist of molecules Ions are a single atom or covalently bonded group of atoms that have an overall electrical charge (the are NO molecules in an ionic compound) ELEMENTS THAT OCCUR AS MOLECULES: Diatomic- H2, N2, 02, F2, Cl2, Br2, i2 Tetratomic- P4 Octatomic- S8, Se8 *A polyatomic ion consists of two or more atoms that are covalently bonded together and have an overall charge In many reactions the polyatomic ion will remain together as a unit Naming Binary Ionic Compounds st For all ionic compounds, the name and formula lists the cation 1 , then the anion In a binary ionic compound, both the cation and anion are monatomic The name of the cation is the same as the name of the metal- HINT: many metal names end in –ium The anion is named by adding the suffix –ide to the root of the nonmetal name Ex. Calcium and bromine makes calcium bromide MUST MEMORIZE FIGURE 2.17 on ch.2 PP ***Know prefixes for hydrates and binary covalent compounds. Mono-, di-, etc. Naming acids: see PP slide 50 Naming Binary Covalent Compounds: slide 52 Naming straight-chain alkanes: Hydrocarbons are compounds that contain only carbon and hydrogen atoms Alkanes are the simplest type of hydrocarbon, usually named using a root name followed by the suffix –ane See slide 56 for molecular masses from chemical formulas *Know the models used for representing molecules Mixtures Heterogenous mixture- one or more visible boundaries between the compounds Homogenous mixture- no visible boundaries bc the components are mixed as ind. Atoms, ions, and molecules ANOTHER NAME for a homogenous mixture is a SOLUTION Solutions in water are called AQUEOUS SOLUTIONS Chapter 3 A mole (mol) is the amt. of a substance that contains the same number of “entities” as there are atoms in exactly 12g of carbon- 12 Entities: atoms, ions, molecules, formula units, or electrons- any type of particle One mole contains 6.022x10^23 entities- to 4 sig figs This number is called Avogadro’s number and is abbrev. N EXAMPLES: 1 atom of Fe has the mass 55.85amu, 1 mol of Fe has mass 55.85g And 1 molecule of H2O has mass 18.02 amu, 1 mol of H20 has mass 18.02g Molar Mass is the mass per mole of its entities For monatomic elements, the molar mass is the same as the atomic mass in grams/mole (the atomic mass is found on the Per. Table) Ex. Molar mass of Ne=20.18 g/mol Mass Percent from the chemical formula See slide 3-10 ***Many of the formula problems will need to just be practiced, practice, practiced! See PP slides for some, the book for others, and YouTube is helpful for some concepts as well! Mass Percent and Mass of an Element Mass percent can be used to calc. the mass of a particular element in any mass of a compound ***see slide 3-13 for formula Empirical and Molecular Formulas Purpose: to show the ratios The empirical formula is the simplest formula for a compound that agrees w/ the elemental analysis. It shows the lowest whole # of moles and gives the relative # of atoms for each element present Ex. Emp. Formula for hydrogen peroxide is HO The molecular formula shows the ACTUAL # of atoms for each element in a molecule of the compound Ex. Molec. formula for hy. perox. Is H202 (#s in right subscripts) Determining the Molecular Formula The molecular formula gives the actual numbers of moles of each element present in 1 mol of compound The molecular formula is a whole-number multiple of the empirical formula Chemical Equations A chemical equation uses formulas to express the identities and quantities of substances involved in a physical or chemical change Stoichiometric Calculations The coefficients in a balanced chemical equation Represent the relative number of reactant and product particles and the relative number of moles of each Since moles are related to mass The equation can be used to calculate masses of reactants and/or products for a given reaction The mole ratios from the balanced equation are used as conversion factors Reactions in Sequence Reactions often occur in sequence The product of one reaction becomes a reactant in the next An overall reaction is written by combining the reactions Any substance that forms in one reaction and reacts in the next can be eliminated Limiting Reactants So far we’ve assumed that reactants are present in the correct amounts to react completely In reality, one reactant may limit the amt of product that can form The limiting reactant will be completely used up in the reaction The reactant that Is not limiting is in excess- some of this reactant will be left over Reaction Yields The theoretical yield is the amount of product calculated using the molar ratios from the balanced equation The actual yield is the amt of product actually obtained The actual yield is usually less than the theoretical yield Solution Stoichiometry- Many reactions occur in solution A solution consists of one or more solutes dissolved in a solvent The concentration of a solution is given by the quantity of solute present is a given quantity of solution MOLARITY (M) is often used to express concentration CHAPTER 4 Water as a solvent -Water is a polar molecule --since it has uneven electron distribution --and a bent molecular shape Water readily dissolves a variety of substances Water interacts strongly with its solutes and often plays an active role in aqueous reactions Electron charge distribution in H2 is symmetrical, in H20 it’s asymmetrical Writing equations for Aqueous Ionic Reactions The molecular equation shows all reactants and products as if they were intact, undissociated compounds This gives the least information about the species in solution The total ionic equation shows all soluble ionic substances dissociated into ions This gives the most accurate information about species in solution Spectator ions are ions that are not involved in the actual chemical change. Spectator ions appear unchanged on both sides of the total ionic equation. Determine and eliminate the spectator ions to make the net ionic equation, which shows only the actual chemical change Precipitation Reactions  In a precipitation reaction 2 soluble ionic compounds react to give an insoluble product, called a precipitate  The precipitate forms through the net removal of ions from solution  Is it possible for more than one precipitate to form in such a reaction Predicting Whether a Precipitate Will Form  Note the ions present in the reactants  Consider all possible cation-anion combinations  Use the solubility rules to decide whether any of the ion combinations is insoluble o An insoluble combination identifies the precipitate that will form STUDY solubility rules (slide 14) Acid-Base Reactions  An acid is a substance that produces H+ ions when dissolved in H2O  A base is a substance that produces OH- ions when dissolved in H20  An acid-base reaction is also called a neutralization reaction Acid-Base Titrations  In a titration, the concentration of one solution is used to determine the concentration of another  In an acid-base titration, a standard solution of base is usually added to a sample of acid of unknown molarity  An acid-base indicator has different colors in acid and base, and is used to monitor the reaction progress  At the equivalence point, the mol of H+ from the acid equals the mol of OH- ion produced by the base o Amount of H+ ion in flask= amount of OH- ion added  The end point occurs when there is a slight excess of base and the indicator changes color permanently STUDY OXIDATION NUMBER RULES slide 32 Oxidation-Reduction (Redox) Reactions OXIDATION is the LOSS of electrons -The reducing agent loses electrons and is oxidized REDUCTION is the GAIN of electrons -The oxidizing agent gains electrons and is reduced A redox reaction involves electron transfer -Oxidation and reduction occur together Elements in Redox Reactions- Types of Reactions Combination Reactions  2 or more reactants combine to form a new compound:  X+YZ Decomposition Reactions  A single compound decomposes to form 2 or more products:  ZX+Y Displacement Reactions  Double displacement: AB+CDAC+BD  Single displacement: X+YXXZ+Y Combustion  The process of combining w/ O 2 CHAPTER 5 Pressure= force/area Atmospheric pressure decreases w/ altitude Pressure arises from countless collisions between gas particles and walls **know what a manometer is and how it works (closed-end and open-end) KNOW common units of pressure conversions IDEAL GAS LAW Typically gases behave most like an ideal gas at high temp. and low pressure Boyle’s Law Charles’s Law Avogadro’s Law Gases mix homogenously in any proportions - And each gas present behaves as if it were the only gas present Dalton’s Law EFFUSION- lightest gases move most quickly, fastest rate of effusion Chapter 6 q= heat w= work c= specific heat capacity total energy of the universe, energy of system and surroundings remains constant 1 cal= 4.184 J 1 Btu= 1055 J 1000J= 1 kJ Know Enthalpy and equations dealing with it Know how coffee-cup calorimeters are different from bomb calorimeters One thing to note: coffee-cup: measures heat transferred at constant PRESSURE Bomb calorimeter: measures heat released at constant VOLUME HESS’ LAW Enthalpy change of an overall process is the sum of the enthalpy changes of its individual steps CHAPTER 7 Quantum Theory and Atomic Structure The Wave Nature of Light The wave properties of electromagnetic radiation are described by three variables: -frequency- cycles per second -wavelength- the distance a wave travels in one cycle -amplitude- the height of a wave crest or depth of a trough *the speed of light is a constant: =3.00*10^8 m/s in a vacuum Energy and Frequency A solid object emits visible light when it is heated to about 1000 K. This is called blackbody radiation The color (and the intensity) of the light changes as the temp. changes. Color is related to wavelength and frequency, while temperature is related to energy Planck proposed that one certain quantities of energy can be emitted (or absorbed) The Quantum Theory of Energy Any object (including atoms) can emit or absorb only certain quantities of energy (hv) Energy is quantized, it occurs in fixed quantities, rather than being continuous. Each fixed quantity of energy is called a quantum (hv) An atom changes its energy state by emitting or absorbing one or more quanta of energy E= nhv where n can only be a whole number- the smallest change being where n=1 Eatom= hv= Ephoton The Bohr Model of the Hydrogen Atom Bohr’s atomic model postulated the following:  The H atom had only certain energy levels, which Bohr called stationary states o Each state is associated w/ a fixed circular orbit of the electron around the nucleus o The higher the energy level, the farther the orbit is from the nucleus o When the H electron is in the first orbit, the atom is in its lowest energy state, called the ground state o The atom does not radiate energy while in one of its stationary states o The atom changes to another stationary state only by absorbing or emitting a photon  The energy of the photon (hv) equals the difference between the energies of the 2 energy states  When the E electron is in any orbit higher than n=1, the atom is in an excited state Matter and energy are alternate forms of the same entity ALL MATTER exhibits properties of both particles AND waves Heisenberg’s Uncertainty Principle Heisenberg’s Uncertainty Principle states that it is not possible to to know the position and momentum of a moving particle at the same time -The more accurately we know the speed, the less accurately we know the position, and vice versa QUANTUM NUMBERS AND ATOMIC ORBITALS An atomic orbital is specified by three quantum numbers -The principal quantum number (n) is a positive integer The value of n indicates the relative size of the orbital and therefore its relative distance from the nucleus The angular momentum quantum number (l) is an integer from 0 *The value of l indicates the shape of the orbital The magnetic quantum number (m1) is an integer with values from –l to +l The value of m1 indicates the spatial orientation of the orbital CHAPTER 8 Shielding and Orbital Energy  Electrons in the same energy level shield each other to some extent  Electrons in inner energy levels shield the outer electrons very effectively o The further from the nucleus an electron is, the lower the Zeff for that particular electron Splitting of Levels into Sublevels Each energy level is split into sublevels of differing energy Splitting is caused by penetration and its effect on shielding For a given n value, a lower I value indicates a lower energy sublevel Order of sublevel energies: s< p< d< f Building Orbital Diagrams -The aufbau principle is applied- electrons are always placed in the lowest energy sublevel available -The exclusion principle states that each orbital may contain a maximum of 2 electrons, which must have opposite spins Hund’s Rule specifies that when orbitals of equal energy are available, the lowest energy electron configuration has the maximum number of unpaired electrons with parallel spins CHAPTER 9 CHAPTER 10 CHAPTER 11 Valence Bond (VB) Theory: Basic principle: a covalent bond forms when the orbitals of two atoms overlap and a pair of electrons occupy the overlap region The space formed by the overlap can hold a MAX of 2 electrons, and they must have opposite (paired) spins The greater the overlap, the stronger the bond VB Theory and Orbital Hybridization The orbitals that form when bonding occurs are different from the atomic orbitals in the isolated atoms -Atomic orbitals “mix” or hybridize when bonding occurs to form hybrid orbitals -The spatial orientation of these hybrid orbitals correspond with observed molecular shapes Mixing one s and two p orbitals gives three sp3 hybrid orbitals. The third 2p orbital remains un-hybridized See slide 11-16 Types of Covalent Bonds A sigma  bond is formed by end-to-end overlap of orbitals All single bonds are  bonds A pi is formed by sideways overlap of orbitals A  bond is weaker than a sigma bond because sideways overlap is less effective than end-to-end overlap ***A double bond consists of one sigma and one pi bond Sources: Class textbook: Chemistry, Silberberg and Amateis, 2015) Dr. Cooper’s powerpoints and verbal lectures


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