Difference between revisions of "Chemistry Lab"
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==Kinetics: Rate Laws== | ==Kinetics: Rate Laws== | ||
*In an equation with rate constant k and reactants A, B, and C, <math>Rate=k([A]^x)([B]^y)([C]^z)</math>. x, y, and z are whole numbers that indicate the concentrations' effect on reaction rate. They can be determined by using a table of data showing concentrations of reactants and resulting reaction rates. When the concentration of all reactants but one (A) stay the same, the concentration of that one product is multiplied by a factor of p, and the reaction rate is multiplied by a factor of q, the whole number exponent (x) of that product in the rate law is equal to log(q)/log(p). The reaction is considered to be x-order with respect to A and (x+y+z)-order overall. | *In an equation with rate constant k and reactants A, B, and C, <math>Rate=k([A]^x)([B]^y)([C]^z)</math>. x, y, and z are whole numbers that indicate the concentrations' effect on reaction rate. They can be determined by using a table of data showing concentrations of reactants and resulting reaction rates. When the concentration of all reactants but one (A) stay the same, the concentration of that one product is multiplied by a factor of p, and the reaction rate is multiplied by a factor of q, the whole number exponent (x) of that product in the rate law is equal to log(q)/log(p). The reaction is considered to be x-order with respect to A and (x+y+z)-order overall. | ||
| − | *In reactions that are 0th-order overall (see first bullet), the <math>rate=k</math>. The graph of the concentration of a product or a reactant over time yields a straight line, and the absolute value of the slope of this line equals k. The half-life of the reaction equals [A]0/(2k). | + | *In reactions that are 0th-order overall (see first bullet), the <math>rate=k</math>. The graph of the '''''concentration''''' of a product or a reactant '''''over time''''' yields a straight line, and the absolute value of the slope of this line equals k. The half-life of the reaction equals [A]0/(2k). |
| − | *In reactions that are 1st-order overall (see first bullet), the <math>rate=k[A]</math>. The graph of the | + | *In reactions that are 1st-order overall (see first bullet), the <math>rate=k[A]</math>. The graph of the '''''natural log''''' of the concentration of a product or a reactant '''''over time''''' does. The absolute value of the slope of this line equals k. The half-life of the reaction equals <math>ln(2)/k</math>. |
| − | *In reactions that are 2nd-order overall (see first bullet), either <math>rate=k[A]^2</math> or <math>rate=k[A][B]</math>. | + | *In reactions that are 2nd-order overall (see first bullet), either <math>rate=k[A]^2</math> or <math>rate=k[A][B]</math>. The graph of the '''''inverse''''' of the concentration of a product or a reactant '''''over time''''' does, and the absolute value of the slope of this line equals k. The half-life of the reaction equals 1/(k[A]0). |
==Aqueous Solutions: Solution Concentration== | ==Aqueous Solutions: Solution Concentration== | ||
Revision as of 23:05, 8 June 2010
Description
1 or 2 people per team. Eye protection #4. 50 minutes. Non-programmable, non-graphing calculator & pencil (NO reference material).
Focus for 2010: Kinetics and Aqueous Solutions.
Sample Questions: Kinetics
Questions in the kinetics section might involve...
- Reaction Rates
- Reaction Conditions (Temperature, Concentration, Particle Size, Cataylsts)
- Rate Laws (at state and national levels)
- Rate Constants (at state and national levels)
Sample Questions: Aqueous Solutions
Questions in the kinetics section might involve...
- Solution Concentration (Molarity, Molality, Mass Percentage, Parts Per Million)
- Conversion Between Units (at state and national levels)
- Determining Concentration using Density, Beer's Law or Titration
- Freezing Point Depression and Boiling Point Elevation
- Factors Affecting Solution Formation
- Solubility
Stoichiometry
I do not know what kind of questions this will bring but I can explain a little about this $5 term. The Merriam-Webster dictionary defines stoichiometry as "a branch of chemistry that deals with the application of the laws of definite proportions and of the conservation of mass and energy to chemical activity". Stoichiometry deals with calculations about the masses (sometimes volumes) of reactants and products involved in a chemical reaction. It is a very mathematical part of chemistry. The most common stoichiometric problem will present you with a certain amount of a reactant and then ask how much of a product can be formed. Ex:: 2A + 3B ---> 3C, Given 25 grams B and unlimited A how much C will be produced. This is called a mass-mass problem. These problems can be solved in 4 simple steps.
- Make sure the chemical equation is correctly balanced.
- Using the molar mass of the given substance, convert the mass given in the problem to moles.
- Construct a molar proportion (two molar ratios set equal to each other) following the guidelines set out in other files. Use it to convert to moles of the unknown.
- Using the molar mass of the unkown substance, convert the moles just calculated to mass.
Other forms of stoichiometric problems are finding the limiting reactant and finding the percentage composition. You can find out more about these in the links below.
Reactions
This is a very generic area so you will have to wait until sample problems come out in the rule book to see what will be on the test. Most problems will probably involve balancing reactions so that the products equal the reactants. You will definitely want to know the five main types of reactions (single displacement, double displacement, combustion, decomposition, and synthesis). Caution: Because this is a very open topic many test makers, especially at the regional level, might take it upon themselves to use questions that you may never have seen before. Just try your best and understand that if you have been diligently studying and haven't seen it, chances are neither have the other teams.
Kinetics: Reaction Rates
[math]\displaystyle{ rate=[limiting reagent]/time }[/math]. To find reaction rates, you need a table or graph showing the concentration of one of the products or reactants over a period of time. If given:
- a line graph showing concentration of a reactant, you can find reaction rate at a given instant for that reactant. It will be equal to the slope of the line tangent to the point on the graph at that instant.
- a line graph showing concentration of a product, you can find the reaction rate at at given instant for that product. It will be equal to the opposite of the slope of the line tangent to the point on the graph at that instant.
- the reaction rate for one reactant or product and the reaction equation, you can find the reaction rates for another reactant or product. Balance the equation, if necessary. Take the rate you are given, multiply by the coefficient of the reactant or product you want the rate for, and divide by the coefficient of the reactant or product whose rate you were given.
Kinetics: Reaction Conditions
- Increasing temperature increases reaction rate.
- Increasing concentration increases reaction rate.
- Increasing particle size decreases reaction rate.
- Adding catalysts increases reaction rate.
Kinetics: Rate Laws
- In an equation with rate constant k and reactants A, B, and C, [math]\displaystyle{ Rate=k([A]^x)([B]^y)([C]^z) }[/math]. x, y, and z are whole numbers that indicate the concentrations' effect on reaction rate. They can be determined by using a table of data showing concentrations of reactants and resulting reaction rates. When the concentration of all reactants but one (A) stay the same, the concentration of that one product is multiplied by a factor of p, and the reaction rate is multiplied by a factor of q, the whole number exponent (x) of that product in the rate law is equal to log(q)/log(p). The reaction is considered to be x-order with respect to A and (x+y+z)-order overall.
- In reactions that are 0th-order overall (see first bullet), the [math]\displaystyle{ rate=k }[/math]. The graph of the concentration of a product or a reactant over time yields a straight line, and the absolute value of the slope of this line equals k. The half-life of the reaction equals [A]0/(2k).
- In reactions that are 1st-order overall (see first bullet), the [math]\displaystyle{ rate=k[A] }[/math]. The graph of the natural log of the concentration of a product or a reactant over time does. The absolute value of the slope of this line equals k. The half-life of the reaction equals [math]\displaystyle{ ln(2)/k }[/math].
- In reactions that are 2nd-order overall (see first bullet), either [math]\displaystyle{ rate=k[A]^2 }[/math] or [math]\displaystyle{ rate=k[A][B] }[/math]. The graph of the inverse of the concentration of a product or a reactant over time does, and the absolute value of the slope of this line equals k. The half-life of the reaction equals 1/(k[A]0).
Aqueous Solutions: Solution Concentration
- Molarity=(number of moles of solute)/(Liters of solution)
- Molality=(number of moles of solute)/(kilograms of solvent)
- Mass Percentage=massA/(massA+massB+...)
- Parts Per Million=Mass Percentage*10,000
Aqueous Solutions: Conversion Between Units
- Molarity->Molality: Multiply by Liters of solution, divide by kilograms of solvent (approximately equal for dilute solutions).
- Molality->Mass Percentage: Multiply by mass of solute, then divide by moles of solute, then multiply by kilograms of solvent, and divide by kilograms of solution (can be appromimated by multiplying by molar mass).
Acid & Bases (2009)
First of all, an old favorite, acids and bases. Last year in all 4 competitions I participated in (2 invitationals, regional, and states) this meant an acid/base titration lab. If you are not familiar with this lab, you will definitely want to ask a teacher to explain it to you before the competition. Believe me, I have seen a team ask the event coordinator what titrate means in the middle of the test, and he was not a happy person. This is a fairly quick and simple lab to complete and it is more than worth your while to DOUBLE CHECK YOUR LAB if you have enough materials. At the state level, I ran the lab through 4 times and averaged the very close results to come up with a more accurate overall result. In a free-response style lab report, this might also get you some extra points for style and accuracy. In the test, the acid/base questions ranged from difficulty of identifying if a solution was an acid based on its pH to balancing advanced reactions trying to find the acidic constant. In order to excel in this event you must be prepared for all levels.
Titration Race (2009)
Despite the name of this portion of the Chem Lab event, this event can barely be considered a race. In fact, a recent rule clarification states that time will not be considered a tie-breaker at the national competition. However if time is considered, here are a few helpful hints to increase both your speed and accuracy in performing a titration.
-Begin with a microtitration. For example titrate with 2 mL of the base to get a ballpark figure of the concentration. Use this to decide how much of the base you will use in the future trials. Remember that the more you use to titrate with, the more accurate your results will be.
-I would recommend three trials. If the ask to show work for only 2 data points, still do three and show the two that are closest together UNLESS your first two seem so close that a third is unnecessary.
-REMEMBER what you're working with. Understand how sulfuric acid might change the calculations to find its concentration.
-DO NOT OVERTITRATE. This is what will seperate the teams the most. A healthy red color is not what you're looking for at the end of a titration. You are looking for your final solution to be barely tinged pink. If given paper, place it below your flask so that you can easily see ANY tinge of pink. If you do overtitrate, go back to the acid and add a drop or two to get to a good final point.
Thermodynamics (2006)
This was the primary focus in a majority of 2005-2006 chem lab competitions. Thermodynamics is a very broad topic, so a variety of problems were used. Although basic enthalpy problems were sometimes found, entropy and Gibbs free energy were significantly more common, as they were more advanced topics.
Strategy
One last thing I would like to mention is strategy. There is a strategy involved all events and I will reveal some hints on a good strategy for this event. First off, don't try to learn everything by yourself. You have a partner for a reason; use him/her/it. It helps if you divide up the study load so that you can learn more in depth in the topics. Next, if the test is extremely long, divide it up. There is no reason for two people to be working on the same problem that one person could do. Lastly, it often helps to have one person do the lab while another starts the test. This saves you valuable time if there is a long test, or allows you to put extra effort into careful practice knowing you are not pressed for time.