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| 8137544514 | 1. Math of Chemistry: | 0 | ||
| 8137545630 | What are significant figures and how do we determine them? | sig figs are digits that carry meaning contributing to its measurement resolution. -All non zero digits are significant (1,2,3,4,5...). -Zeros between non-zero digits (102, 2005, etc). -Leading zeros are NEVER significant (0.02, 0.000515) -In a number with a decimal point, trailing zeros, those to the right of the last non zero digit, are significant (2.02000, 5.400, etc). -In a number without a decimal point, trailing zeros are usually not significant. (unless a decimal point is places after the number-- ex: 100. has 3 sig figs) | 1 | |
| 8137547491 | How do we use significant figures? | to convey the precision and accuracy of a measurement (bc a measurement always has some degree of uncertainty). Accuracy refers to the agreement of a particular value with the true value. Precision refers to the degree of the agreement among several measurements of the same quantity. Precision reflects the producibility of a given measurement. | 2 | |
| 8137547492 | Why are significant figures important? | Sig figs are important because they indicate the degree of accuracy. The more sig figs the more accurate the data will be. | 3 | |
| 8137550035 | Multiplication/Division vs Addition and Subtraction (significant figures) | Multiplication and Division: The LEAST number of sig figs in any number of the problem determines the number of sig figs in the answer. ex: 2.5 x 3.42 -- the answer is 8.6 (rounded from 8.55) because 2.5 has only two sig figs while 3.42 has three. Two sig figs is less precise than three, so the answer has two sig figs. Addition and Subtraction: 1.Count the number of sig figs in the decimal portion of each number in the problem (digits to the left of the decimal are not used to determine number of decimal places in final answer) 2.Add or subtract in normal fashion 3. Round the answer to the LEAST number of places in the decimal portion of any number in the problem. | 4 | |
| 8137550036 | Why are conversions important? | It's important because we need answers that are universally accepted and meaningful. Or, if you have a certain unit of measure given in the problem and it wants a different one in the answer. ex: If it gives you 1 ft but wants the answer in cm you could use the conversion 1 ft=12 in= 30.48 cm. | 5 | |
| 8137551366 | How do we use box problems to convert between units? | 6 | ||
| 8137551945 | What are the SI base units for volume, energy, temperature, pressure, mass, moles? | Volume: cubic meter (m^3) Energy: Joule (J) Temperature: Kelvin (K) Pressure: Pascal (Pa) Mass: Kilogram (kg) Moles: g/mol?? | 7 | |
| 8137553284 | 2. The Mole | 8 | ||
| 8137553933 | Molar Masses of Elements and Compounds | The molar mass of an element is found on the periodic table (its atomic weight) The molar mass of a compound is found by using the chemical formula to determine the number of each type of atom present in the compound. Multiply each molar mass of the element by the number of atoms of that element present in the compound. Add it all together and give units of grams/mole. | 9 | |
| 8137559200 | Percent Composition: | Mass of Element in 1 Mole of Compound/ Mass of 1 Mole of Compound | 10 | |
| 8137559201 | Empirical vs Molecular Formula | Empirical Formula is the base ratio of the elements in the compound. The molecular formula of a compound can be determined from the empirical formula if the molecular weight is known. To find Empirical Formula: -find mol amount of all elements used in the compound -determine the smallest mol amount of all the elements and divide the other mol amounts by that number (they should all end up as ratios, although not exact should be like 2.999 is 3, or 0.999 is 1) -Those ratios are the base amounts for each element and therefore gives Empirical Formula To find Molecular Formula: (molecular weight/empirical weight) molecular weight= g of molecule/mol of molecule empirical weight= atomic weight of that element multiplied by their base ratio. Do for all the elements and add up. Ex: The compound dioxane contains only carbon, hydrogen, and oxygen. When 0.956 g dioxane is burned, 1.91 g carbon dioxide and 0.782 g water are formed. In another experiment, it was determined that 6.04x10^-3 mol dioxane weighs 0.532 g. What is the molecular formula of dioxane. Empirical Formula: ______ mol C= 1.91 g CO2/44.01 g CO2 x 1 mol CO2/ 1 mol CO2 x 1 mol C= 0.0433 mol C ______ mol H= .782 g H2O/18.016 g H2O x 1 mol H2O/ 1 mol H2O x 2 mol H= 0.0875 mol H ______ mol O= .956 g dioxane- .521 g C- 0.087 g H= .348 g O--> .02195 mol O .0433/.02195= about 2 .0875/.02195= about 4 .02195/.02195= 1 So Empirical Formula = C2H4O Molecular Formula--> molecular weight= .532 g dioxane/ 5.04x10^-3 mol dioxane= 88.1g empirical weight= (12.01)2+(1.008)4+(16.00)1= 44.052 88.1/ 44.052= about 2 Use that as your ratio and multiply that to the empirical--> molecular formula= C4H8O2 | 11 | |
| 8137560487 | 3. Reactions | 12 | ||
| 8137560488 | Balancing Equations | 1. Write down given equation 2. Write down the number of atoms per each element that you have on each side of the equation 3. Always leave hydrogen and oxygen for last. 4. Start by choosing one of the elements and attempt to get the same number of that element on each side. 5. Balance the hydrogen 6.Balance the oxygen | 13 | |
| 8137563402 | Predicting Products from Reaction Types | Synthesis: two or more substances combine to form a new compound (A+X --> AX) Decomposition: Usually has heat as a reactant, and causes the compound to 'decompose' into two different compounds, usually to CO2 and a metal oxide. (AX-->A+ X) Combustion: When reactive elements combine with oxygen, releasing a large amount of energy in the form of light and heat. (AX+O2--> AO +XO) Double Replacement: The anion and cation from the two different compounds switch places (AX + BY --> AY + BX) Single Replacement: when a substance in a compound is replaces by a substance outside of the compound (A+BX --> B+AX) Redox: or reduction-oxidation reaction, in which elements are either oxidized or reduced (OIL-RIG) | 14 | |
| 8137566025 | 4. Stoichiometry | 15 | ||
| 8137570553 | Limiting Reagent vs. Excess Reagent | Limiting Reagent: Balance the chemical equation for the chemical reaction. Convert the given information into moles. Use stoichiometry for each individual reactant to find the mass of product produced. The reactant that produces a lesser amount of product is the limiting reagent. The reactant that produces a larger amount of product is the excess reagent. | 16 | |
| 8137571292 | Determining Theoretical Yield given a reaction and an excess left over | Theoretical Yield: The smallest yield of product is called the theoretical yield (same amount usually as the limiting reagent. Excess Reagent Leftover: Determine the mass used of both the excess and limiting reagent and then subtract the excess amount from the limiting amount and that is how much is left over. | 17 | |
| 8137572206 | Percent Yield vs. Percent Error | Percent Yield: Actual Yield/Theoretical Yield Percent Error: | ( measured - accepted) / accepted | * 100 | 18 | |
| 8137572764 | 5. Solutions | 19 | ||
| 8137579881 | Polar vs. nonpolar solutes in polar vs. nonpolar solvents | Polar solvents can dissolve polar solutes. Non-polar solvents can't dissolve polar solutes. Non-polar solvents can dissolve non-polar solutes. Polar solvents can dissolve non-polar solutes Non-polar solvents can't dissolve a polar substance because there is no charge separation in the solvent. Non-polar substances don't have enough force to rip apart a polar substance, but they do have enough to take in other non-polar materials. | 20 | |
| 8137593274 | The "Big 4" of solubility rules | Solids tend to dissolve best when: Heated Stirred Ground into small particles Gases tend to dissolve best when: Solution is cold Pressure is high Nitrates, Halogens, Alkali Metals are very soluble. Ammonium is VERY soluble. | 21 | |
| 8137594134 | Electrolytes and conductivity of solutions | Electrolytes are substances which, when dissolved in water, break up into cations and anions. Strong Electrolytes ionize completely (100%). Fall into three categories: strong acids, strong bases, and salts. while weak electrolytes ionize only partially. Usually include weak acids and weak bases. Solution is a good conductor when there are a lot of free electrons in it. (so strong electrolytes are good conductors). | 22 | |
| 8137594135 | Strong vs. Weak Acids and Bases | Strong Acids:All halogens except for fluorine create strong acids (binary acids). Ionizable hydrogens are hydrogens that come off in solution. Some acids can only release one hydrogen (monoprotic), while others release more (di, tri, etc. -polyprotic). Protic=proton [H+] Weak Acids: such as lactic acid, usually ionize less than 5% of the time. Many of these weaker acids are "organic" acids that contain a "carboxyl" acid. Strong base showing dissociation. Strong bases fully dissociate to give ions in solution. Weak bases only partially dissociate. | 23 | |
| 8137595673 | Molarity as a conversion factor: | Molarity (M) is moles of solute per volume of solution in liters. Gives concentration of a solution. Molarity= moles of solute/ Liters of solution A 1 M solution has 1 mol of solvent for every one liter of solution | 24 | |
| 8137596503 | Molarity as a concept: | 25 | ||
| 8137643759 | Solving for moles given a volume | from volume of pure liquid or solid: 1.Multiply the volume by the density to get the mass 2.Divide mass by molar mass to get moles Moles from volume of solution: n= M x V ex: How many moles of NaCl are contained in 0.300 L of 0.400 M NaCl solution n= 0.300 L x 0.400 mol NaCl/ 1 L= 0.120 mol NaCl | 26 | |
| 8137644471 | Preparing Solutions (molarity) | You prepare a solution by dissolving a known mass of solute into a specific amount of solvent. Ex: Prepare 1 L of 0.5 M NaCl solution First calculate molar mass of NaCl (58.44 g/mol). 0.5 x 58.44 = 29.22 g NaCl dissolve 29.22 g NaCl in 1 L of solution. | 27 | |
| 8137645063 | Diluting Solutions (molarity) | use C1V1=C2V2 C- concentration V- volume plug known values into the equation. ex: dilute a 5 M solution with water to make 1 L of a 1 mM solution. We know the initial and final concentration and final volume, but not the initial volume needed. 1mM= 0.001 M (5M)V1= (0.001 M)(1L) --> V1= (0.001 M)(1 L)/ (5 M)= .0002 or 0.2 L V1= 0.2 mL we would first measure 0.2 mL of our 5 M solution. Next, we would add enough water to increase the volume to 1 L. 1 L- 0.2 mL= 0.9998 L or 999.8 mL. So we would add 999.8 mL of water to the 0.2 mL of 5M solution. | 28 | |
| 8137646218 | Complete and net ionic equations (Precipitation Reactions) | Complete Ionic Equation: Includes everything, even if it isn't involved in the chemical reaction. ex: Na + Cl + Ag+ NO3 --> Na+ NO3+AgCl Net Ionic Equation: Includes only the species involved in the chemical reaction: Cl+Ag--> AgCl | 29 | |
| 8137647536 | Formation of Ionic Solids (Precipitation Reactions) | Solids composed of oppositely charged ions. They consists of positively charged cations and negatively charged anion. When they dissolve in water, the cation and anion separate and allow the solution to conduct an electrical current. | 30 | |
| 8137648534 | Bronsted-Lowry Model (Acid-Base Reactions) | The Bronsted-Lowry Model and theory of acids and bases is where the acid is a proton (hydrogen ion) donor, and the base is a proton acceptor. This then turns the base into the new acid, and the old acid into the new base. These acid and base could technically continuously exchange the proton. | 31 | |
| 8137650985 | Titrations of Acids and Bases: Procedural Set-Up | Before beginning the experiment, obtain all necessary materials and clean all necessary items with distilled water. Measure out a precise amount of analyte; this will make up the solution of unknown concentration. Quantitatively transfer the analyte into a beaker or Erlenmeyer flask. Make sure to rinse all of solid analyte into the beaker or Erlenmeyer flask with distilled water. Add additional distilled water until the anlayte is fully dissolved. Measure and record volume of aqueous solution, the process of titration will solve for concentration of this solution. Add four to five drops of the appropriate color indicator into the beaker. Swirl the beaker in order to mix the aqueous solution of the analyte and the drops of indicator. Fill the burette with an excess amount of titrant. The titrant is the standard solution of known concentration and should be in aqueous form. Clamp the burette carefully to a burette stand. The tip of the burette should not be touching any surfaces. Place the beaker or Erlenmeyer flask containing the aqueous solution of unknown concentration under the burette. Record the initial volume of the burette. Make sure to measure at the bottom of the meniscus. Turn on the stopcock (tap) of the burette, so that standard solution is added to the beaker. This should cause a color change so be sure to swirl the beaker or Erlenmeyer flask until the color disappears. Repeat the above step until the color does not disappear. This means you have reached the endpoint Stop when you've reached endpoint, which is the point when the reactant within the solution of unknown concentration has been completely neutralized. You can tell you've reached the endpoint because the color will change. Measure and record your final volume of the burette. Calculate the volume of standard solution used by subtracting the initial volume measurement from the final volume measurement of the burette. Now perform the necessary calculations in order to obtain the concentration of the unknown solution | 32 | |
| 8138484195 | Titrations of Acids and Bases: Graphical Analysis |  | 33 | |
| 8138483622 | Titrations of Acids and Bases: Concept and Endpoint vs. Equivalence Point | Titration is used for determining how much of an analyte in moles is in a solution. this is done by slowly adding a standard solution, or a reagent of known concentration, until the titration is determined to be complete. This typically occurs after the titration has passed an equivalence point, or when the amount of reagent equals, chemically, the amount of analyte. The equivalence point is not something that is typically observed however because around the equivalence point, one drop before and there is no change in pH and the next drop changes the pH sometimes by 3-4 units. The equivalence point is between these two drops and the closer these two drop are to each other the better the quantification of the analyte. The endpoint is when the indicator first changes in appearance, it is always slightly after the equivalence point. | 34 | |
| 8137651581 | Assigning Oxidation States | Oxidation state of an element is 0. Oxidation state of monatomic ion= charge of ion Oxygen= 2 in covalent compounds (except peroxide=-1) Hydrogen= +1 in covalent compounds Fluorine= -1 Sum of oxidation states= 0 Sum of oxidation states= charge of the ions. | 35 | |
| 8137651582 | Determining if something is an oxidation reduction reaction or not. | If charges don't change on the products side of the equation then it is not a redox. | 36 | |
| 8137653546 | 6. Gas Laws | 37 | ||
| 8137683926 | Avogadro's Law | Volume and number of moles are directly related. n1/V1=n2/V2 | 38 | |
| 8137683927 | Boyle's Law | Pressure and Volume are inversely related (constant temperature, moles of gas). P1V1=P2V2 | 39 | |
| 8137684674 | Charles' Law | Volume and temperature (in Kelvin) are directly related (constant pressure and amount) V1/T1=V2/T2 | 40 | |
| 8137685374 | What are assumptions of the ideal gas law? | 1.All gas particles are in constant motion and collisions between the gas molecules and the walls of the container cause the pressure of the gas. 2.The particles are so small that their volume is negligible compared with the volume occupied by the gas. 3.The particles don't interact. There are no attractive or repulsive forces between them 4. The average kinetic energy is proportional to temperature. | 41 | |
| 8137686270 | What causes the ideal gas law to fail? | When the pressure is high, the volume is low, the temperature is low, or there are significant intermolecular forces. | 42 | |
| 8137686271 | What are real gases? | Non-hypothetical gases whose molecules occupy space and have interaction. The Ideal Gas equation works well enough for most gases at ordinary pressure, as long as the temperature is reasonably high. | 43 | |
| 8137686957 | Molar Volume of a gas at STP | 22.4 L/mol | 44 | |
| 8137687421 | Converting from a gas to another gas using a balanced equation | Ex: 6 L of H2(g) reacts with excess nitrogen. How many liters of NH3 are produced balanced equation: 3 H2+ N2--> 2 NH3 6L H2/ 22.4 L H2 x 1 mol H2/ 3 mol H2 x 2mol NH3/ 1 mol NH3 x 22.4 L NH3= 4 L NH3 | 45 | |
| 8137688092 | Converting from a gas to a solid or visa versa. | How many L O2 at STP with decomposition of 50 g of solid potassium chlorate. balanced equation: 2 KClO3 (s)-->2 KCl+ 3 O2 50 g KClO3/ 122.55 g KClO3 x 1 mol KClO3/ 2 mol KClO3 x 3 mol O2/ 1 mol O2 x 22.4 L O2= 13.7 L O2 | 46 | |
| 8137689606 | Determining Pressures of gases already mixed (Dalton's Law) | For a mixture of gases in a container, the total pressure exerted is the sum of the pressures that each gas would exert if it were alone. Ptotal= P1+P2+P3+... or ntotal(RT/V) Assuming that each gas behaves ideally, the partial pressure of each can be calculated with ideal gas law: p1= n1RT/V or P1V1/Vtotal For a mixture of ideal gases, is the total number of moles that is important, not the identity or composition of the involved particles. | 47 | |
| 8137691887 | Mixing two gases into different containers and finding new pressures (Dalton's Law) | 48 | ||
| 8137693016 | How are mole fractions related to pressure? | (Mole of Compound/ Total Moles Used)= mole fraction Partial Pressure of Compound= Mole Fraction x Total Pressure | 49 | |
| 8137693017 | How can you compare moles of different substances? | 50 | ||
| 8137694202 | Postulates and Deviations of the Kinetic Molecular Theory | The particles are so small compared to the distances between them that the volume of the individual particles can be assumed to be negligible. These particles are in constant motion. The collisions of the particles with the walls of the container are cause of the pressure exerted by the gas Particles are assumed to exert no forces on each other;they are assumed neither to attract nor to repel each other. The average kinetic energy of a collection of gas particles is assumed to be directly proportional to the Kelvin temp of the gas Ideal Gases: Random, fast, travel in a straight line Collide— elastic Volume is negligible No attractions Real Gases: Not Random, curved patterns Collisions occur— are inelastic Volume exists (heavier) | 51 | |
| 8137695161 | Temperature vs Energy (Kinetic Molecular Theory) | Temperature measures the average kinetic speed of molecules. Usually as temperature increases so does energy. | 52 | |
| 8137695162 | Mass vs rate of diffusion | As mass decreases, the rate of diffusion increases. (diffusion also increases with a temperature increases) | 53 | |
| 8137695752 | Graham's Law | Rate of effusion for gas 1/ Rate of effusion for gas2 = √M2/√ M1 M1 and M2 represent molar masses of the gases the rate of effusion of a gas is inversely proportional to the square root of the mass of its particles | 54 | |
| 8137696668 | 7. Law of Conservation of Energy | 55 | ||
| 8137697394 | Potential vs Kinetic Energy | Potential Energy is energy due to position or composition. For example: Skiers at the top of a hill Ammonium nitrates— fertilizers and explosives TNT— trinitrotoluene (Three separate nitrogen atoms each connected to three oxygen atoms, one of which, in each group, is double bonded. Connected by carbon atoms). Kinetic Energy is the energy due to motion of the object and depends on the mass of the object and its velocity. | 56 | |
| 8137697395 | Internal Energy of a System | Internal Energy (E) of s system is the sum of the kinetic and potential energies of all the "particles" in the system. To change the internal energy of a system: E=q+w q— heat; w— work Q negative— heat lost Q positive— heat gained W negative— work done BY the system W is positive— work done ON the system | 57 | |
| 8137698212 | System vs. Surrounding vs. Universe | A system is part of the universe on which we wish to focus attention. The surroundings include everything else in the universe. Therefore, system+surroundings= universe | 58 | |
| 8137698213 | State Functions | Energy is a state function; work and heat are not. State function-- property that does not depend in any way on the system's past or future (only dependent on present state). Ex: Airplane trip. What you eat during the flight doesn't matter, a flight last week doesn't effect it, just that it ends up in the destination. | 59 | |
| 8137700476 | Definition of Heat | is a form of energy transfer among particles in a substance by means of kinetic energy of those particles. A form of energy that flows between two samples of matter because of their differences in temperature. | 60 | |
| 8137700477 | How does heat transfer | It is the exchange of thermal energy between physical systems. Involves the transfer of energy between two objects due to temperature difference. Washing your hands on a cold day at the mountain Touching a hot stove The rate of heat transfer is dependent on the temperatures of the systems and the properties of the intervening medium through which heat is transferred. | 61 | |
| 8137701047 | Endothermic vs. Exothermic | Endothermic reactions are when heat flows into a system or absorbs energy from surroundings. In an equation sign of values, it is an endothermic process when q is positive (sign reflects the SYSTEMS point of view) In a balanced equation, it is endothermic if energy must be added to the reactants, so on the left side of the equation. Exothermic reactions are when energy flows out of the system. In an equation sign of values, it is an exothermic process when q is negative. In a balanced equation, it is exothermic if energy is released with the products, so on the right side of the equation. | 62 | |
| 8137703187 | Where can you find out a reaction is exothermic: Energy Diagrams | Exothermic |  | 63 |
| 8138814962 | Where can you find out a reaction is endothermic: Energy Diagrams |  | 64 | |
| 8137704331 | Work's relation to energy and heat | Energy is a state function; work and heat are not. Energy is the ability to do work and the transfer of heat between the system and the surroundings. | 65 | |
| 8137705397 | How is work done? | Assuming the only work done is due to volume changes in the system at a constant pressure, the work done by a system is given by ∆W= P∆V | 66 | |
| 8137705404 | Expansion vs. Compression | Expansion is when there is a positive delta v, increase, and has a -w result. The energy of the system decreases. Work has been done BY the system on the surroundings and loses energy. Compression is when there is a negative delta v, decrease, and has a positive w results. The energy of the system increases. Work has been done ON the system by the surroundings and gains energy. For expanding gas, delta V is a positive quantity because volume is increasing so we have to have opposite signs. To convert between L·atm and joules is 101.3 joules per L·atm. | 67 | |
| 8137706560 | What is Enthalpy | Enthalpy is a state function. It is the "heat content"of a system or the potential of a system to create heat. Since enthalpy is a state function, the change in enthalpy in going from some initial state to some final state is independent of the pathway. This means that in going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes place in one step or a series of steps (Hess's Law). △H= q at constant pressure ∆H= Hproducts- Hreactants | 68 | |
| 8137708211 | How is enthalpy related to heat? | Enthalpy is the amount of energy in a system and when this changes (when a reaction happens), the energy is either released or absorbed and this energy is usually released or absorbed as heat. | 69 | |
| 8137710811 | What is a calorimeter | an object used for calorimetry, or the process of measuring the heat of chemical reactions or physical changes as well as heat capacity. They're what we made in class | 70 | |
| 8137712082 | Heat capacity of a substance | Is the heat needed for a substance's temperature to change by one degree. | 71 | |
| 8137712606 | Heat capacity vs. specific heat capacity vs. molar heat capacity | Specific heat capacity is the energy required to raise the temperature of one gram of a substance by one °C. Molar heat capacity is the energy required to raise the temperature of one mole of a substance by one °C. Heat capacity is the heat needed for a substance's temperature to change by one degree. | 72 | |
| 8137712607 | q= mc∆T | q= heat energy m=mass c=specific heat ∆T: change in temperature if q is negative it is exothermic reaction. If q is positive it is endothermic. | 73 | |
| 8137713666 | Conceptual Implications of Hess' Law | This law means that in going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes place in one step or a series of steps. | 74 | |
| 8137713667 | Mathematical Implications of Hess' Law | Calculations involving Hess's law typically require that several reactions be manipulated and combined to finally give the reaction of interest. In doing this, you should: Work backward Reverse any reactions as needed Multiply reactions If a reaction is reversed, the sign of ∆H is also reversed. The magnitude of ∆H is directly proportional to the quantities of reactants and products in a reaction. If the coefficients in a balanced reaction are multiplied by an integer, the value of deltaH is multiplied by the same integer, | 75 | |
| 8137715929 | How to solve Standard Enthalpies of Formation | Defined as the change in enthalpy that accompanies the formation of 1 mole of a compound from its elements with all substances in their standard states. A degree symbol on a thermodynamic function indicates that the corresponding process has been carried out under standard conditions. The standard state of a substance is a precisely defined reference. We must use a common reference state to properly compare the thermodynamic properties of two substances. Steps steps for this are the same as Hess's law and change in enthalpy calculated from enthalpies of formation or reactants+products. | 76 | |
| 8137715930 | Understanding the idea of what standard state is | The standard state for a substance is a precisely defined reference state. We must use a common reference state to properly compare the thermodynamic properties of two substances. For a compound the state of a gaseous substance is a pressure of 1 atm. For a pure substance in a condensed state (liquid or solid), the standard state is the pure liquid or solid. For a substance present in a solution, the standard state is a concentration of exactly 1M. For an element the standard state of an elemnt is the form in which the lement exists under conditions of 1 atm and 25°C. | 77 | |
| 8137716736 | Solving Standard Enthalpy of Formation problems using the equation | 78 | ||
| 8137730445 | 8. Electromagnetic Radiation | 79 | ||
| 8137731751 | How light is emitted from heated elements | The atomic spectrum is produced when heated elements are passed through a prism. Consist of lines of different colors separated by dark areas. Series of lines specific for each element produced by photons emitted by electrons dropping to lower energy levels. When the element is heated, electrons go up and down (relax) and then go up again in the orbitals. It gives up energy in the form of light when they rest. | 80 | |
| 8137731752 | Bohr Model vs. Quantum Mechanical Model | this model gave a hydrogen atom energy levels consistent with the hydrogen emission spectrum. Says that energy (a photon) must be given off when an electron moves toward the nucleus and that energy must be absorbed from a photon to move an electron away from the nucleus. However, this model is technically incorrect because it only holds true for hydrogen. The second model is more correct and began with the discovery that electrons act like waves, not particles. Only certain circular orbits have a circumference into which a whole number of wavelengths of the standing electron wave will "fit." The probability of finding an electron at a particular position is the greatest close to the nucleus and drops off rapidly as the distance from the nucleus increases. | 81 | |
| 8137732748 | Energy Released using Bohr Model | 82 | ||
| 8137735070 | Characteristics of a Wave: Wavelength Frequency Speed Energy | Wavelength (λ)— the distance (m) between two consecutive peaks of troughs in a wave. Frequency (v) --number of wave (cycles) per second (Hz or s⁻1) that pass a given point in space Speed (c) — Speed of light (2.9979·10⁸ m/s) Energy can be gained or lost only in whole number multiples or "hv." A system can transfer energy only in whole quanta (or "packets") Energy has mass... E=mc² | 83 | |
| 8137735626 | Relationships between all characteristics of a wave | Waves carry energy. There is a relationship between the frequency of a wave and the amount of energy it carries. A wave with a large wavelength carries less energy than the same kind of wave with a smaller wavelength. The speed of a wave is the time it takes for one part of the wave to travel a certain distance. | 84 | |
| 8137735627 | Calculating for frequency, wavelength, or energy | Frequency: v= λ/c Wavelength: λ= c/v Energy: E= hv | 85 | |
| 8137736468 | Quantums of Energy/Photons | Energy can be released (or absorbed) by atoms only in "packets" of some minimum size. This minimum energy packet is called a quantum. The energy (E) of a quantum is related to frequency (v) by planck's constant. The light in the photoelectric effect was a stream of tiny energy packets called photons. Each photon has an energy proportional to its frequency | 86 | |
| 8137736469 | Photoelectric Effect | The photoelectric effect is the observation that light shining on a metallic surface can cause the surface to emit electrons. For each metal there is a minimum frequency of light below which no electrons are emitted, regardless of the intensity of the light. The higher the light's frequency above this minimum value, the greater the kinetic energy of the released electron(s). | 87 | |
| 8137737280 | Duality of Light Wave Particle | The duality of light is unique. Under certain conditions, such as when we shine it through narrow slits and look at the result, it behaves as only a wave can. Under other conditions, such as when we shine it on a metal and examine the spray of electrons that comes off, light behaves as only particles can. Light behaves as a wave, or as particles, depending on what we do with it, and what we try to observe. This is what lies at the heart of the Heisenberg uncertainty principle. | 88 | |
| 8137737281 | 9. Electrons | 89 | ||
| 8137737986 | Wave functions and Quantum Mechanical Model | Wave function in quantum mechanics is a function that encodes the state of a quantum-mechanical system. Typically the wave function obeys a wave equation or modified wave equation that has wave-like solutions. The most well known example of such a wave equation is the Schrodinger equation, The Quantum Mechanical model of the atom states that the location of the electrons around the atom cannot be precisely determined. The region where the electron can probably be found is known as the electron cloud. | 90 | |
| 8137740809 | Concepts and Implications of the Heisenberg Uncertainty Principle | The Heisenberg Uncertainty Principle states that there is a fundamental limitation to just how precisely we can know both position and momentum of a practice at a given time. | 91 | |
| 8137742624 | Mathematical Model of Heisenberg Uncertainty Principle | ∆X · ∆ (M · V) is ≥ h/4π ∆X = uncertainty in a particle's position ∆(m · v) = uncertainty in a particle's momentum H = planck's constant | 92 | |
| 8137742625 | Electron Density Map |  | 93 | |
| 8137744056 | Radial Probability vs. Distance from Nucleus Graph (For all orbitals and different energy levels) | Look up picture | 94 | |
| 8137745413 | Quantum Numbers Concepts | Principal quantum number (n) — size and energy of the orbital Angular momentum quantum number (l) — shape of atomic orbitals (sometimes called a subshell- s, p, d, f). Magnetic quantum number (ml) — orientation of the orbital in space relative to the other orbitals in the atom. Electron spin quantum number (ms) — can be +½ or -½ | 95 | |
| 8137745969 | Subshells with quantum numbers | 96 | ||
| 8137745970 | Orbital Shapes | S= spherical shaped orbital P = bowtie shaped orbital D = clover shaped orbital (dx2 is upward bowtie with donut shape around the middle) F = two intersected clover shaped orbital | 97 | |
| 8137746894 | Aufbau Principle | states that, hypothetically, electrons orbiting one or more atoms fill the lowest available energy levels before filling higher levels (Ex: 1s before 2s). In this way, the electrons of an atom, molecule, or ion harmonize into the most stable electron configuration possible. | 98 | |
| 8137746895 | Pauli Exclusion Principle (Electron Spin and Magnetic Fields) | in a given atom no two electrons can have the same set of four quantum numbers (n, l, ml, ms). Since electrons in the same orbital have the same values of n, l, and ml, this postulate says that they must have different values of ms. Then, since only two different ms, an orbital can only hold two electrons and they must have opposite spins. | 99 | |
| 8137750010 | Hund's Rule | states that the lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by the pauli principle in a particular set of degenerate (same energy) orbitals. | 100 | |
| 8137750662 | Orbital Diagrams according to Aufbau, Pauli, Hund's rules/principles | Fill all of the orbital (all the way across) before going back to add the second (with the opposite spins) |  | 101 |
| 8137752048 | Concept/Analogy of Electron Configurations | 102 | ||
| 8137752049 | Complete ground state electron configurations | Example for ground state (this is Mg): 1s^2 2s^2 2p^6 3s^2 | 103 | |
| 8137752817 | Noble Gas Configuration | Example for Noble Gas (this is Mg): [Ne]3s2 | 104 | |
| 8137752818 | Non-ground state configurations and ions. | Non-ground state configurations and ions follow basic principles and rules for the element, you just add or take away electrons in the sublevels according to the ion charges. Ex: Mg2+ has two less electrons than normal, so instead of the electron configuration being 1s^2 2s^2 2p^6 3s^2 it would now be 1s^2 2s^2 2p^6 | 105 | |
| 8137753626 | Valence Electrons of the representative (main group elements). | The elements in the same group on the periodic table have the same valence electron configuration Ex: 1s^2 2s^2 2p^6 shows an 8 valence electron configuration. | 106 | |
| 8137753627 | 10. PES— Photoelectron Spectroscopy | 107 | ||
| 8137754430 | Photoelectric effect | the observation that light shining on a metallic surface can cause the surface to emit electrons. For each metal there is a minimum frequency of light below which no electrons are emitted, regardless of the intensity of the light. The higher the light's frequency above this minimum value, the greater the kinetic energy of the released electron(s). | 108 | |
| 8137755173 | Ionization Energy | is the energy required to remove an electron from a gaseous ion where the atom or ion is assumed to be in ground state. The first ionization energy is the energy required to remove the highest energy electron of an atom. The first ionization energy is considerably smaller than the second ionization energy because as more electrons are removed, the more unstable the atom can become causing it to hold onto its remaining electrons stronger (therefore, requiring more energy to remove the next electron. Think of this like the someone stealing your money example in class) | 109 | |
| 8137755174 | Describing electron structure from PES |  | 110 | |
| 8137755763 | Calculations using PES | 111 | ||
| 8137755764 | 11. Periodic Trends | 112 | ||
| 8137756893 | Ionization Energy Periodic Trend | Ionization energy decreases going down columns (more orbitals, electrons increased distance from the nucleus is greater than the increase of protons in the nucleus). It increases going left to right (because in rows the atom has the same number of orbitals, but has an increased number of protons that holds tighter onto the electrons making them harder to remove. ) | 113 | |
| 8137757713 | First vs second/third/fourth/etc ionization energies | Ionization Energy increases as number of ionization energies does. This is because when the first electron is removed, it is the furthest possible out. However, there are still the same number of protons pulling on a smaller amount of electrons, so when you try to take away yet another electron, it requires more energy. This increase continues as you keep trying to take away more electrons. | 114 | |
| 8137758622 | Electron Affinity Concept | Electron Affinity is the energy change associated with the addition of an electron to a gaseous atom. | 115 | |
| 8137758623 | Atomic Radius Concept | Atomic radius can be obtained by measuring the distance between atoms in a chemical compound. These radii are often called covalent atomic radii because of the way they are determined. | 116 | |
| 8137759573 | 12. Alkali Metals Properties and Trends | Density increases going down group 1A because atomic mass generally increases more rapidly than atomic size. They are very reactive, have low ionization energies, and are very reactive with water. They form ionic solids with non-metals. | 117 | |
| 8137759574 | 13. Bonding | 118 | ||
| 8137761439 | Why do chemical bonds occur? | To lower the energy of the system. To fill their valence orbital. Because of the octet (or in hydrogen's case the duet) rule. | 119 | |
| 8137761440 | How do atoms bond to form compounds | Atoms form compounds by sharing electrons to create full valence orbitals which keep them balanced and uniform. To cause them to be more stable. | 120 | |
| 8137762709 | What are the types of chemical bonds and what types of elements are involved in these bonds | Ionic Bonds: bonding that creates thermal stability due to the electrostatic attraction of the closely packed, oppositely charged ions. Usually between an a metal and nonmetal. Covalent Bond: when electrons are shared equally. Usually formed between two or more nonmetals Polar Covalent Bond: when the atoms electrons aren't completely taken, but instead just unequally shared. | 121 | |
| 8137763351 | 14. Interactions of Atoms | 122 | ||
| 8137763949 | What happens when atoms get close and can they get too close? | If they can bond, they will. And if the electronegativity difference is strong enough, sometimes it will take an electron completely. It a rare case, if they are pushed too close together, they will become very unstable and become another atom. But more times than not they will repel from each other due to electron shielding and repelling. | 123 | |
| 8137763950 | Graphical Representations of Interaction of Atoms | 124 | ||
| 8137765216 | 15. What is the relationship between bond length (internuclear distance), bond type (single/double/triple), and bond energy? | As bond type increases (more bonds), the bond length decreases and bond energy increases. | 125 | |
| 8137765217 | 16.Electronegativity | 126 | ||
| 8137766159 | Defintion of Electronegativity | The ability of an atom in a molecule to attract shared electrons to itself. The greater the difference in electronegativities of the atoms, the greater the ionic component of the bond and the greater the value of ∆. | 127 | |
| 8137766160 | Electronegativity Trend | Increases left to right and decreases going down a column. | 128 | |
| 8137766824 | Electronegativity Concept | The concept of electronegativity was put on a quantitative footing by Linus Pauling. He came up with Pauling's model (which we don't need to know math of, however) to understand this model, consider a hypothetical molecule HX. The relative electronegativities of the H and X atoms are determined by comparing the H-X bond energy with the "expected" H-X bond energy (which is an average of the H-H and X-X bond energies). | 129 | |
| 8137766825 | Electronegativity relation to bond character | 0-0.3 electronegativity difference is non-polar covalent. 0.3-1.7 difference is a polar covalent bond. Difference of 1.7 or higher is an ionic bond. | 130 | |
| 8137767815 | 17. Polarity of Molecules | 131 | ||
| 8137767816 | What affects Polarity? | The electronegativity difference between atoms affects polarity and their shape. This then affects physical properties such as melting and boiling points or solubility. The three polarities are ionic, polar or nonpolar. | 132 | |
| 8137768924 | What is a dipole moment and how do you represent one in a molecule? | A dipole moment is a property of a molecule whose charge distribution can be represented by a center of positive charge and a center of negative charge (said to be dipolar). The dipole character of a molecule is often represented by an arrow pointing towards the negative charge with the tail of the arrow indicating positive charge. |  | 133 |
| 8137768925 | 18. Isoelectronic Series | 134 | ||
| 8137770707 | definition and concept of an Isoelectronic Series and how do they compare in terms of size, number of electrons, and number of protons? | Ions containing the same number of electrons. They all have the same electron configuration but have different protons. Ex: K+, Ca2+, Ar, Cl- In general, for a series of isoelectronic ions, the size decreases as the nuclear charge increases. (Ca2+ would be the smallest and Cl- would be the largest). | 135 | |
| 8137771853 | 19. Ionic Compounds | 136 | ||
| 8137772649 | Lattice Energy and Formation How is it like Coulomb's Law How is it different | Is the energy change occurring when separated gaseous ions are packed together to form an ionic solid. This is often defined as the energy released when ionic solid forms from its ions. LE= k (Q1Q2/r) is exactly like coloumb's law except instead of multiplied by the size? of en electron (2.31x10^-19 Jxnm) it is multiplied by the constant k. Q1/Q2= charge of atoms. r= shortest distance between anion and cation. | 137 | |
| 8137773386 | Born-Haber Cycle What is it? What does it show? What is it very similar to, but different? Which processes are endothermic and which are exothermic? | It is the cycle concerned with the formation of an ionic compound from the reaction of a metal (often from alkali or alkali-earth metals) with a halogen. The values used in this cycle are all predetermined charges in enthalpy. It is similar to Hess' Law. How to use this cycle: 1. Determine energy of the metal and nonmetal in their elemental forms (elements in their natural state have an energy level of zero.) Subtract from this the heat of formation of the ionic solid that would be formed from combining this element. This is the energy of the ionic solid, and will be used at the end of the process to determine lattice energy. 2.Make sure involved elements are in their gaseous form. Add these changes enthalpy. 3. Turn all polyatomic species into single atoms. (ex: Cl2--> 2Cl). Add the energy required for this to the value from step 2. 4. Turn them into their ionic forms (ex: Cl- or Na+). To do this, ionization energy will be added to the value from step 3. Then the electron affinity of the nonmetal will be subtracted from the previous value (bc it is a release of energy associated with the addition of an electron) 5. Now the metal and nonmetal will be combined to form the ionic solid. This will cause a release of energy, called the lattice energy. The value for lattice energy is the difference between the value from Step 1 and Step 4. |  | 138 |
| 8137773387 | 20. Models | 139 | ||
| 8137773881 | Assumptions of Models | Models are attempts to explain how nature operates on the MICROscopic level based on experiences in the MACROscopic world. | 140 | |
| 8137773882 | Limitations of Models | A model does not equal reality Models are oversimplifications, and therefore often wrong. Models become more complicated and are modified as they age We must understand that underlying assumptions in a model so that we don't misuse them When a model is wrong, we often learn much more than when it is right. | 141 | |
| 8137774296 | 21. Bond Energies | 142 | ||
| 8137775059 | What energy is associated with breaking bonds? | To break bonds, energy must be added to the system (endothermic, energy term carries a positive sign). | 143 | |
| 8137775060 | What energy is associated with forming bonds? | To form bonds, energy is released (exothermic, energy term carries a negative sign). | 144 | |
| 8137775751 | What equation can be used to describe the enthalpy formation of a compound and how is it different from most equations | ∆H= sum of energies required to break old bonds (positive or endothermic) - sum of the energies released in formation of new bonds (negative or exothermic) ∆H = Σn×D(bonds broken) - Σn×D(bonds formed) D represents the bond energy per mole of bonds (always has a positive sign). It is different from most equations because the ∆ is reactants-products (∆ is usually products - reactants) | 145 | |
| 8137775752 | 22. Localized Electron Model | 146 | ||
| 8137776456 | Lone pairs | Pairs of electrons localized ON an atom. | 147 | |
| 8137777027 | Bonding pairs | pairs of electrons found in the space between the atoms. | 148 | |
| 8137777873 | How many lone and bonding pairs can be given to an atom and are there exceptions. | The group number of the element in the periodic table will tell about the number of bonds it will form to complete its octet or duet rule. For example, N has 5 valence electrons and needs three more to complete the octet rule. It will form three bonding pairs to complete this rule (in most cases) and then have one lone pair. Carbon however will form four bonding pairs and zero lone pairs. There will be an exception if the element is in group 3 or further, because it has a d block it can fill it can take on extra electrons if necessary. ex: SF4 | 149 | |
| 8137779694 | What are the steps for writing a lewis structure? | 1. Sum the valence electrons from all the atoms 2.Use a pair of electrons to form a bond between each pair of bound atoms. 3. Atoms usually have noble gas configurations. Arrange the remaining electrons to satisfy the octet rule (or due rule for hydrogen) | 150 | |
| 8137781425 | Octet vs Duet Rules | Describes the tendency for main group atoms to preferentially combine into molecules that optimizes having eight electrons in their valence shell. and The adaption of the octet rule that applies to hydrogen and lithium. They only possess S orbitals, two electrons can create a filled energy state like the nearest noble gas, helium. | 151 | |
| 8137782370 | What is resonance and do the lewis structures have to be equivalent? | is invoked when more than one valid lewis structure can be written for a particular molecule. The resulting electron structure of the molecule is given by the average of these structure. So no, the lewis structures don't have to be equivalent, it is an average of all of them. | 152 | |
| 8137783506 | Formal Charges Why use them? What do they help show How do you calculate them for an atom? Can you calculate it for a compound? | It is used to evaluate nonequivalent lewis structures. Atoms in molecules try to achieve formal charges as close to zero as possible. Any negative formal charges are expected to reside on the most electronegative atoms. FC= (# of valence electrons when neutral)-(lone pair electrons*)-(total bonded electrons divided by 2) For a compound you would find the formal charge for each atom. *count all lone pair electrons. ex-2 lone pairs is four electrons | 153 | |
| 8137783507 | What is the VSEPR Model? | The valence shell electron pair repulsion (VSEPR) theory is a model used to predict the geometry of individual molecules from the number of electron pairs surrounding their central atoms. | 154 | |
| 8137785001 | How does the VSEPR model explain how molecules are shaped? | the premise of this model is that the valence electron pairs surrounding an atom tend to repel each other and will adopt an arrangement that minimizes the repulsion, thus determining the molecule's geometry. Lone pairs tend to take more space than bonded pairs leading into somewhat distorted structures. Bonded pairs count as one area of electron density. It is based on observable electron density. | 155 | |
| 8137785510 | How can you apply the VSEPR model to a lewis structure | VSEPR can be used in a lewis structure to display the orientation and geometry of a molecule. A solid wedge-shaped line represents bonds that project up out of the paper towards you. A hatched wedge-shaped line are for bond that project down into the paper away from you. Helps create more visually accurate shapes of molecules. *Q in the picture |  | 156 |
| 8137786092 | What is the shape of a molecule based on its bonded and lone pairs? | It's molecular geometry/lewis structure?? | 157 |
