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Final Study Guide

by: Katelynn Jones

Final Study Guide CH 101

Katelynn Jones
GPA 3.4

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What I put together from old study guides and tests
General Chemistry
Jared Allred
Study Guide
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This 13 page Study Guide was uploaded by Katelynn Jones on Monday May 2, 2016. The Study Guide belongs to CH 101 at University of Alabama - Tuscaloosa taught by Jared Allred in Spring 2016. Since its upload, it has received 142 views.


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Date Created: 05/02/16
Chapter 1 1. Elements, Compounds, and Mixtures  A pure substance is made up of only one type of particle (one component) and its composition is invariant (it does not vary from one sample to another). A mixture is a substance that cannot be chemically broken down into simpler substances. A compound is composed of two or more elements in fixed, definite proportions. Water and sand is a heterogeneous mixture, one in which the composition varies from one region of the mixture to another- the different particles that compose water and sand do not mix uniformly. Sweetened tea is a homogeneous mixture, one with the same composition throughout- the particles that compose sweetened tea mix uniformly. An element is a substance that cannot be chemically broken down into simpler substances 2. What is a molecule  Two or more atoms joined chemically in a specific geometrical arrangement. 3. States of matter  The particles that compose liquid matter pack about as closely as particles do in solid matter, but slightly weaker attractions between the particles allow them to move relative to each other, giving liquids a fixed volume but not a fixed shape.  The particles that compose gaseous matter attract each other only very weakly- so weakly that they do not clump together as particles do in a liquid or solid. Instead the particles are free to move large distances before colliding with one another.  The particles that compose solid matter attract one another strongly and therefore pack close to each other in fixed locations. Although the particles vibrate, they do not move around or past each other. Consequently, a solid has a fixed volume and rigid shape. 4. Laws, theories, observations, and measurements  A law is a brief statement that summarizes past observations and predicts future ones. A theory is a model for the way nature is ad tries to explain not merely what nature does but why. 5. Law of conservation of mass  In a given process, the total mass of the system stays the same. Summarized by Lavoisier. In a chemical reaction, matter is neither created nor destroyed. Mass of reactants=mass of product 6. Law of definite proportions  All samples of a given compound, regardless of their source or how they were prepared, have the same proportions of their constituent elements. Example on page 11. 7. Law of multiple proportions  Fixed ratios of elements go into making two pure compounds made from the same elements. An example of this would be: two different compounds formed from carbon and oxygen have the following mass ratios: 1.33g O: 1g C and 2.66g O: 1g C. Example on page 12. 8. John Dalton’s Atomic Theory  The atomic theory states that, each element is composed of tiny, indestructible particles called atoms. All atoms of a given element have the same mass and other properties that distinguish them from the atoms of other elements. Atoms combine in simple, whole- number ratios to form compounds. Atoms of one element cannot change into atoms of another element. In a chemical reaction, atoms only change the way that they are bound together with other atoms. 9. J.J. Thomson’s Cathode Ray Experiment  Thomson applied a high electrical voltage between two electrodes at either end of the tube. He found that a beam of particles, called cathode rays, traveled from the negatively charged electrode (called the cathode) to the positively charged one (called the anode). Thomson observed that the particles that compose the cathode ray have the following properties: they travel in straight lines, they are independent of the composition of the material from which they originate and they carry a negative electoral charge. Thomson had discovered the electron, a negatively charged, low mass particle present within all atoms. 10. Millikan’s oil- drop experiment  Millikan sprayed oil into fine droplets using an atomizer. The droplets were allowed to fall under the influence of gravity through a small hole into the lower portion of the apparatus where Millikan viewed them with the aid of a light source and a viewing microscope. During their fall, the drops acquired electrons that had been produced by bombarding the air in the chamber with ionizing radiation. The electrons imparted a negative charge to the drops. In the lower portion of the apparatus, Millikan created an electric field between two metal plates. Since the lower plate was negatively charged, and since Millikan could very the strength of the electric field, he could slow or even reverse the free fall of the negatively charged drops. 11. Rutherford’s gold-foil experiment  Demonstrated the existence of a tiny, charged nucleus in atoms that contains most of the mass. Came up with the nuclear theory with three basic parts: 1. Most of the atoms mass and all of its positive charge are contained in a small core called the nucleus. 2. Most of the volume of the atom is empty space, throughout which tiny, negatively charged electrons are dispersed. 3. There are as many negatively charged electrons outside the nucleus as there are positively charged particles (protons) within the nucleus, so that the atom is electrically neutral. 12. Sub atomic particles and their properties  All atoms are composed of the same subatomic particles: protons, neutrons, and electrons. Protons and neutrons have nearly identical masses. Recall that the proton and the electron both have electrical charge. The electron is negative and the proton is positive. The neutron has no charge and is found in the nucleus of most elements. 13. Atomic Symbols  The most important number to the identity of an atom is the number of protons in its nucleus. The number of protons defines the element. The number of protons in an atom’s nucleus is its atomic number and is given the symbol Z. The number of electrons and protons is equal to a neutral atom. 14. Isotopes  Atoms with the same number of protons but different numbers of neutrons are called isotopes. They have different masses, but the same properties. 15. Atomic masses a. 1 mole= 6.02214 X 10^23. Example of converting to moles on page 52-54. b. Find an atomic mass from an isotopic mass or vice versa- Examples on page 23-25 and in notes. Chapter 2 1. Significant figures  2 rules of sig figs: 1. Addition and subtraction you truncate at last significant digit. 2. Multiplication and division keep smallest number. 2. Precision vs. Accuracy  Accuracy refers to how close the measured value is to the actual value. Precision refers to how close a series of measurements are to one another or how reproducible they are. Measurements are precise if they are consistent with one another, but they are accurate only if they are close to the actual value. Random error is an error that has equal probability of being too high or too low. Almost all measurements have some degree of random error. Systematic error is an error that tends toward being either too high or too low. 3. Density  Density is a characteristic physical property of a substance. Density=mass/volume. 4. Energy  The total energy of an object is a sum of its kinetic energy, the energy associated with its motion, and its potential energy, the energy associated with its position or composition. Thermal energy is the energy associated with the temperature of an object. Thermal energy is actually a type of kinetic energy. KE=1/2 X mv^2. 1kg X m^2/s^2=1J. Energy is conserved in any process but can be transferred between a system and its surrounding. An exothermic process involves the transfer of energy from the system to the surroundings and carries a negative sign. An endothermic process involves the transfer of energy to the system from the surroundings and carries a positive sign. 5. Unit conversion  1 inch= 2.54 cm  1 kJ= 1000 J  1 cal= 4.184 J  1 mL= 1 cm^3  Given unit X desired unit/given unit= desired unit 6. The following metric prefixes a. T- tera 10^12 b. G- giga 10^9 c. M- mega 10^6 d. k- kilo 10^3 e. c- centi 10^-2 f. m- milli 10^-3 g. μ- micro 10^-6 h. n- nano 10^-9 Chapter 3 1. Wave particle duality  Particles can behave as waves and vice versa. Electrons, neutrons all have a wavelength when they’re moving. Certain properties of light are best described by thinking of it as a wave, while other properties are best described by thinking of it as a particle. 2. The photoelectric effect  The observation that many metals emit electrons when light shines upon them. A high frequency, low intensity light produces electrons without the predicted lag time. The light used to dislodge electrons in the photoelectric effect exhibits a threshold frequency, below which no electrons are emitted from the metal, no matter how long the light shines on the metal. Planck’s constant is h= 6.626 X 10^-34 J. 3. Atomic spectra  When an atom absorbs energy- in the form of heat, light, or electricity- it often reemits that energy as light. Atomic spectroscopy is the study of the electromagnetic radiation absorbed and emitted by atoms. 4. The Bohr model  Bohr attempted to develop a model for the atom that explained atomic spectra. In his model, electrons travel around the nucleus in circular orbits. Bohr’s orbits exist only at specific, fixed distances from the nucleus. The energy for each Bohr orbit is also fixed, or quantized. Bohr called these orbits stationary states and suggested that, although they obey the laws of classical mechanics, they also possess a peculiar, mechanically unexplainable stability. It is only when an electron jumps, or makes a transition, from one stationary state to another that radiation is emitted or absorbed. 5. The deBroglie wavelength  A single electron traveling through space has a wave nature; its wavelength is related to its kinetic energy. The faster the electron is moving, the higher its kinetic energy and the shorter its wavelength. The de Broglie relation is wavelength= Planck’s constant over mass times velocity. Example given on page 78. The double slit experiment shows us that matter has a wave like part to it. The electrons are diffracted by the two slits, which means they interfere with each other. The double slit experiment proves the wave nature of matter. Since quantum mechanical theory is universal, it applies to all objects, regardless of size. λ = h/mv 6. The Schrodinger equation  HY=EY can be used to find the energy of the electronic state in a hydrogen atom. It can be used to find the spatial distribution (shape) of an orbital. 7. The wavefunctions of hydrogen  Quantum numbers- the principle quantum number is an integer that determines the overall size and energy of an orbital.  Orbitals and their shapes  Atomic transitions  Relationship between number of nodes and n quantum number  Relationship between average radius of orbitals and n quantum number Equations that wont be given: The behavior of light: E=hv ν = c /λ The de Broglie relation: λ =h/mv Chapter 4 1. Pauli Exclusion Principle  2 electrons cannot have the same 4 quantum numbers on the same atom. 2. Shielding/Penetration  Effective nuclear charge- the actual nuclear charge experienced by an electron, defined as the charge of the nucleus plus the charge of the shielding electrons. The trend in atomic radius as we move to the right across a row in the periodic table is determined by the inward pull of the nucleus on the electrons in the outermost principal energy level.  We can think of the repulsion of one electron by other electrons as screening or shielding that electron from the full effects of the nuclear charge. As the outer electron undergoes penetration into the region occupied by the inner electrons, it experiences a greater nuclear charge and therefore a lower energy. 3. Aufbau Principle  The principle that indicates the pattern of orbital filling in an atom. 4. Hund’s Rule  States that when filling degenerate orbitals, electrons fill them singly first, with parallel spins. 5. Valence electrons, determining the number and configuration  For main group elements, the valence electrons are those in the outermost principal energy level. For transition elements, we also count the outermost d electrons among the valence electrons. 6. Types of elements: metals, metalloids, and non metals  Metals lie on the lower left side and middle of the periodic table and share some common properties: they are good conductors of heat and electricity; they can be pounded into flat sheets (malleability); they can be drawn into wires (ductility); they are often shiny; and most importantly they tend to lose electrons when they undergo chemical changes.  Several metalloids are classified as semiconductors because of their intermediate electrical conductivity. Many of the elements that lie along the zigzag line are metalloids.  Nonmetals lie on the upper right side of the periodic table. The division between metals and nonmetals is the zigzag line from boron to astatine. Nonmetals have varied properties: tend to be poor conductors of heat and electricity; they all tend to gain electrons when they undergo chemical changes. Chlorine is one of the most reactive nonmetals. 7. Periodic trends  Size of atoms- in general, atomic radii increase as we move down a column and decrease as we move to the right across a period in the periodic table.  Size of ions  Effective nuclear charge- as we move to the right across a row in the periodic table, the effective nuclear charge experienced by the electrons in the outermost principal energy level increase, resulting in a stronger attraction between the outermost electrons and the nucleus, and smaller atomic radii.  Ionization energies- ionization energy increases as we move to the right across a period and decreases as we move down a column in the periodic table. st i. 1 ionization energy trends- peak at each noble gas and bottom at each alkali metal. Ionization starts at a minimum with each alkali metal and rises to a peak with each noble gas. 8. Electron affinity. What does it mean?  The electron affinity of an atom or ion is the energy change associated with the gaining of an electron by the atom in the gaseous state. Electron affinity does not have any definite trend in the columns or groups. Among the group 1A electron affinity becomes more positive as we move down the column. Electron affinity generally becomes more negative as we move to the right across a period (row) in the periodic table. As we move to the right across a row (or period) in the periodic table, metallic character decreases. As we move down a column (or family) in the periodic table, metallic character increases. Chapter 5 1. Ionic vs. covalent bonds  A bond that forms between a metal and a nonmetal is an ionic bond. A bond that forms between two or more nonmetals is a covalent bond. 2. Empirical vs. molecular formulas  An empirical formula gives the relative number of atoms of each element in a compound. A molecular formula gives the actual number of atoms of each element in a molecule of a compound. 3. Lattice energies- what chemical step does this refer to and what does it mean?  Lattice energy is the energy associated with the formation of a crystalline lattice of alternating cations and anions from the gaseous ions. 4. Naming compounds  Binary ionic compounds- contain only two different elements. i. Metal forms 1 kind of ion- To name these compounds you take the name of the metal and then the base name of the nonmetal + -ide. ii. Metal forms more than 1 kind of ion- for these types of metals the name of the metal is followed by a roman numeral in parenthesis which indicates the charge of the metal in that particular compound. To name these compounds you take the name of the metal (charge of metal in roman numerals) base name of nonmetal + -ide. iii. Polyatomic ions- an ion composed of two or more atoms. We name ionic compounds that contain a polyatomic ion in the same way that we name other ionic compounds except that we use the name of the polyatomic ion whenever it occurs. Most common polyatomic ions are oxyanions, anions containing oxygen and another element. We name them systematically according to the number of oxygen atoms in the ion. If there are more than two ions in the series, we use the prefixes hypo- and per. iv. See table 5.5(pg. 159) you should know these names!  Hydrates- contain a specific number of water molecules associated with each formula unit. We name hydrates like we name other ionic compounds, but we give them the additional name “prefixhydrate”, where the prefix indicates the number of water molecules associated with each formula unit.  Binary molecular compounds- prefix-first element-prefix-base name of 2 ndelement + -ide. 5. Molar mass of a compound  Page 166-169 Chapter 6 1. Electronegativity  The ability of an atom to attract electrons to itself in a chemical bond.  For main group elements the periodic trends are: generally increases across a period in the periodic table; generally decreases down a column; fluorine is the most electronegative element; francium is the least electronegative element. 2. Lewis structures. Drawing the “best” ones  Formal charges- fictitious charge assigned to each atom in a Lewis structure that helps us to distinguish among competing Lewis structures. The formal charge of an atom is the charge it would have if all boding electrons were shared equally between the bonded atoms. Formal charge= # of valence electrons- (#of nonbonding electrons+1/2 # if bonding electrons). In general there are 4 rules: the sum of all formal charges in a neutral molecule must be zero; the sum of all formal charges in an ion must equal the charge of the ion; small (or zero) formal charges on individual atoms are better than large ones; when formal charge cannot be avoided, negative formal charge should reside on the most electronegative atom.  Resonance structures- is one of two or more Lewis structures that have the same skeletal formula, but different electron arrangements. These are represented with a double-headed arrow. 3. VSEPR  Two electron groups: linear  Three electron groups: trigonal planar  Four electron groups: tetrahedral. One lone pair: trigonal pyramidal. Two lone pairs: bent  Five electron groups: trigonal bipyramidal. When two of the five electron groups around the central atom are lone pairs it is T shaped. When three of the five electron groups around the central atom are lone pairs it is linear.  Six electron groups: octahedral. One lone pair: square pyramidal. Two lone pairs: square planar. 4. Bond strengths and bond lengths  In general, for a particular pair of atoms, triple bonds are shorter than double bonds, which are shorter than single bonds. Chapter 7 1. Valence band theory  Two orbitals making a sigma bond can easily rotate with respect to each other. Two orbitals making a pi bond cannot rotate with respect to each other without breaking the bond. 2. Molecular orbital theory- when two atomic orbitals come together to form two molecular orbitals, one molecular orbital will be lower energy than the two separate atomic orbitals and one molecular orbital will be higher in energy than the separate atomic orbitals.  Antibonding orbitals contain a *  Bond order= (# of electrons in bonding MOs)-(# of electrons in antibonding MOs)/2  Paramagnetic contain unpaired electrons and diamagnetic are paired. 3. Band theory  Could you distinguish the picture of bands of a metal from those of an insulator? - Metal has no energy gap, a semiconductor has a small energy gap and an insulator has a large energy gap. 4. True for molecular orbital theory, but not true for the valence bond or Lewis dot models: when atoms bond they form a set of bonding and antibonding orbitals. Chapter 8 1. Distinguish between chemical and physical properties/changes  In a chemical change, atoms rearrange, transforming the original substances into different substances (rusting of iron, burning of sugar, fossil fuel burning). In a physical change the substances do not change their composition. A chemical property is a property that a substance displays only by changing its composition. A physical property is a property that a substance displays without changing its composition. 2. Reaction stoichiometry  The coefficients in a chemical equation specify the relative amounts in moles of each of the substances involved in the reaction.  Combustion reactions (react with oxygen) i. Reaction that represents the combustion of C6H12O2= C6H12O2 (l) + 8O2(g) 6CO2(g)+6H2O(g) 3. Can you figure out what the limiting reagent is?  A limiting reagent is the reactant that limits the amount of product in a chemical reaction. Notice that the limiting reactant is the reactant that makes the least amount of product. 4. Calculate theoretical yield using the same sets of info  Theoretical yield is the amount of product that can be made in a chemical reaction based on the amount of limiting reactant. 5. Calculate percent yield  Percent yield= actual yield/theoretical yield X 100%  Actual yield is the amount of product actually produced by a chemical reaction. Chapter 9 1. Molarity  Molarity= amount of solute (in mol)/ volume of solution (in L)  Build on reaction conditions from chapter 8 2. Know the solubility rules we covered in class and that are tabulated in the book  Page 312! 3. Strong/weak/non-electrolytes, and their properties  Substances that dissolve in water to form solutions that conduct electricity. Substances such as sodium chloride that completely dissociate into ions when they dissolve in water are strong electrolytes. Compounds such as sugar that do not dissociate into ions when dissolved in water are nonelectrolytes. Acids are molecular compounds that ionize to form H+ ions when they dissolve in water. HCl= strong acid and strong electrolyte. A solution of a weak acid is composed mostly of the nonionized acid only a small percentage of the acid molecules ionize. 4. Precipitation reactions  A reaction will only occur if an insoluble compound is present. A reaction in which a solid forms upon the mixing of two solutions. Precipitation reactions are predictable if you understand that only insoluble compounds form precipitates. 5. Ionic solutions  Complete (total) ionic equations are chemical equations that show all the ions in a reaction.  Net ionic equations are equations that show only the species that change during the reaction.  Example of a net ionic equation would be 2 H+(aq)+ S2-(aq) H2S(g)  Example of complete ionic equation would be:  2Li+(aq)+S2-(aq)+Cu2+(aq)+2NO3-(aq)CuS(s)+2Li+(aq) +2NO3-(aq) 6. Recognize the following reactions  In an acid base reaction an acid reacts with a base and the two neutralize each other, producing water. HCl (aq) is an acid because it produces H+ ions in a solution. Bases produce OH- in a solution. Sodium hydroxide (NaOH) is a base. Binary acids contain only two elements and oxyacids contain oxygen. Binary acids are composed of hydrogen and a nonmetal. To name- hydro base name +ic acid. Oxyanions ending with – ate= base name +ic acid. Oxyanions ending with ite= base name +ous acid. In titration a substance in a solution of known concentration is reacted with another substance in a solution of unknown concentration. At the equivalence point- the point in titration when the number of moles of OH- equals the number of moles of H+ in a solution titration is complete. In a gas evolution reaction a gas forms resulting in bubbling. Many gas evolution reactions are also acid base reactions. Oxidation- reduction reactions or redox reactions are reactions in which electrons transfer from one reactant to the other. These types of reactions occur both in and out of solution. The rusting of iron, the bleaching of hair, and the production of electricity in batteries involve redox reactions. Many redox reactions involve the reaction of a substance with oxygen. Oxidation is the loss of electrons and reduction is the gain of electrons. Rules for assigning oxidation states: 1. The oxidation state of an atom in a free element is 0. 2. The oxidation state of a monoatomic ion is equal to its charge. 3. The sum of the oxidation states of all atoms in a neutral molecule or formula unit is 0 and an ion is equal to the charge of the ion. 4. In their compounds, metals have positive oxidation states. Group 1A metals always have an oxidation state of +1. Group 2A metals always have an oxidation state of +2. 5. In their compounds, we assign nonmetals oxidation states according to the table on page 330. Oxidation state of S in S03 2-= +4. Chapter 10 1. Specific heat capacity  q- heat  Cs- specific heat capacity  ΔT- change in temperature  C= heat capacity  C=q/ ΔT  q= C X ΔT  q= m X Cs X ΔT 2. Calorimetry  In calorimetry, we measure the thermal energy the reaction (defined as the system) and the surroundings exchange by observing the change in temperature of the surroundings. If a reaction is carried out at a constant volume, V=0 and w=0. The heat evolved (or given off), called the heat at constant volume (qv), is then equal to Erxn=qv. When a chemical reaction occurs open to the atmosphere under conditions of constant pressure the energy can evolve as both heat and work. Under conditions of constant pressure, a thermodynamic quantity called enthalpy represents this. Enthalpy (H) of a system as the sum of its internal energy and the product of its pressure and volume. In a bomb calorimeter, reactions are carried out at fixed volume.  Esystem= -Esurroundings. State functions do not depend on the path taken to arrive at a particular state. Hrxn can be determined using constant pressure calorimetry. Erxn can be determined using constant volume calorimetry. 3. Calculate ΔH 0 rxn using various methods:  Page 358 and on


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