|Ch.1 - Intro to General Chemistry||2hrs & 53mins||0% complete||WorksheetStart|
|Ch.2 - Atoms & Elements||2hrs & 49mins||0% complete||WorksheetStart|
|Ch.3 - Chemical Reactions||3hrs & 25mins||0% complete||WorksheetStart|
|BONUS: Lab Techniques and Procedures||1hr & 38mins||0% complete||WorksheetStart|
|BONUS: Mathematical Operations and Functions||47mins||0% complete||WorksheetStart|
|Ch.4 - Chemical Quantities & Aqueous Reactions||3hrs & 55mins||0% complete||WorksheetStart|
|Ch.5 - Gases||3hrs & 47mins||0% complete||WorksheetStart|
|Ch.6 - Thermochemistry||2hrs & 28mins||0% complete||WorksheetStart|
|Ch.7 - Quantum Mechanics||2hrs & 35mins||0% complete||WorksheetStart|
|Ch.8 - Periodic Properties of the Elements||1hr & 57mins||0% complete||WorksheetStart|
|Ch.9 - Bonding & Molecular Structure||2hrs & 5mins||0% complete||WorksheetStart|
|Ch.10 - Molecular Shapes & Valence Bond Theory||1hr & 31mins||0% complete||WorksheetStart|
|Ch.11 - Liquids, Solids & Intermolecular Forces||3hrs & 40mins||0% complete||WorksheetStart|
|Ch.12 - Solutions||2hrs & 17mins||0% complete||WorksheetStart|
|Ch.13 - Chemical Kinetics||2hrs & 22mins||0% complete||WorksheetStart|
|Ch.14 - Chemical Equilibrium||2hrs & 26mins||0% complete||WorksheetStart|
|Ch.15 - Acid and Base Equilibrium||4hrs & 42mins||0% complete||WorksheetStart|
|Ch.16 - Aqueous Equilibrium||3hrs & 48mins||0% complete||WorksheetStart|
|Ch. 17 - Chemical Thermodynamics||1hr & 44mins||0% complete||WorksheetStart|
|Ch.18 - Electrochemistry||2hrs & 58mins||0% complete||WorksheetStart|
|Ch.19 - Nuclear Chemistry||1hr & 33mins||0% complete||WorksheetStart|
|Ch.20 - Organic Chemistry||3hrs||0% complete||WorksheetStart|
|Ch.22 - Chemistry of the Nonmetals||2hrs & 1min||0% complete||WorksheetStart|
|Ch.23 - Transition Metals and Coordination Compounds||1hr & 54mins||0% complete||WorksheetStart|
|Rate of Reaction||11 mins||0 completed|
|Average Rate of Reaction||18 mins||0 completed|
|Arrhenius Equation||16 mins||0 completed|
|Rate Law||38 mins||0 completed|
|Integrated Rate Law||52 mins||0 completed|
|Collision Theory||9 mins||0 completed|
|Instantaneous Rate of Change|
|Identifying Reaction Order|
Heat or thermal energy play a pivotal role in the rate of a chemical reaction. The Arrhenius equation represents the formula that shows the dependence of the rate and rate constant k on the absolute temperature, as well as other variables such as activation energy, and the pre-exponential factor.
Rate of Reaction
Generally, as we increase the temperature of a chemical reaction, molecules will absorb this new source of thermal energy. The result is that molecules will move more quickly, striking one another more frequently and with greater force. If molecule A can hit molecule B with enough energy and the correct orientation then they will successful combine and form a new molecule AB.
From the chart provided, as the temperature increases the rate of formation for molecule AB increases. The rule of thumb is that a rise of 10 °C will double the rate of reaction.
Although it was well known that increases in temperature led to increases in rate, it wasn’t until the Arrhenius Equation developed by Svante Arrhenius that this temperature-rate relationship could be accurately explained.
When examining the variables of the Arrhenius Equation: k is the rate constant and, Ea is the activation energy normally in J/mol or kJ/mol, T is the absolute Temperature in Kelvin, R is the universal gas constant equal to 8.314 J/mol K, and A is the pre-exponential factor, frequency factor or Arrhenius constant.
Since the pre-exponential factor is a constant based on the given chemical reaction and the universal gas constant is always the same value, the way to affect k is based on changing Ea or T. Based on the Arrhenius equation, increasing the rate constant k can be accomplished by either increasing the temperature or decreasing the activation energy, usually through the use of a catalyst.
Let’s approach a question using the Arrhenius equation.
PRACTICE: The rate constant of a reaction at 33.0 °C was measured to be 5.2 × 10−2 s−1. If the frequency factor is 1.2 × 1013 s−1, what is the activation barrier?
STEP 1: Identify the given variables.
STEP 2: Convert the temperature given in Celsius to Kelvin.
STEP 3: Multiply both sides of the equation by the natural logarithm, ln.
When multiplying the Arrhenius equation by ln we obtain:
Rearranging this new formula gives us the straight-line form:
STEP 4: Plug in the given values into straight-line form of the Arrhenius Equation.
Solving for some of the portions simplifies it to:
STEP 5: Subtract 30.1158 from both sides.
This simplifies the equation to:
STEP 6: Multiply both sides by 8.314.
The equation is now:
STEP 6: Divide both sides by 0.003266.
This gives the value of Ea as:
Arrhenius Equation: Plot-wise approach
As we’ve shown before, the Arrhenius equation can be written in a non-exponential form or straight-line form.
This form of the Arrhenius equation allows us to interpret our values graphically.
By setting our y-axis as ln k and our x-axis as the inverse temperature we can correlate the non-exponential form of the Arrhenius equation to a straight-line.
By using the straight-line or non-exponential form of the Arrhenius equation we can determine either the frequency factor or the activation energy of a chemical reaction.
PRACTICE: The rate constant for a reaction was measured as a function of temperature. Based on the plot of ln k versus 1/T, what is the activation energy for the reaction?
STEP 1: Using the equation for the straight line from the chart provided, apply it to the non-exponential form of the Arrhenius equation.
STEP 2: Isolate the slope of the straight line and set it equal to – Ea/R.
Since the universal gas constant R is equal to 8.314 J/mol K, solve for the activation energy of the chemical reaction.
STEP 3: Multiply both sides by – 8.314 to isolate the activation energy.
This gives Ea as:
Arrhenius Equation: 2-Point Form
The 2-point form of the Arrhenius Equation is a way mathematically calculating the changes in the rate constant k as the temperature changes. Recall that an increase in temperature causes an increase in the rate constant.
Depending on the placement of T1 and T2 the 2-point form of the Arrhenius equation can be written in two different ways. Algebraically either form can be used for solving.
From the 2-point form of the Arrhenius Equation: k1 and k2 are the rate constants and, Ea is the activation energy normally given in J/mol or kJ/mol, T1 and T2 are the absolute Temperatures in Kelvin, R is the universal gas constant equal to 8.314 J/mol K.
PRACTICE: The activation energy for a reaction is 37.6 kJ/mol. The rate constant for the reaction is 5.4 x 10-3 s-1 at 45 °C. Calculate the rate constant at 145 °C.
STEP 1: Identify the given variables.
STEP 2: Convert the temperatures given in Celsius to Kelvin.
STEP 3: Since the universal gas constant R uses joules in its units then the activation energy Ea must be converted from kJ/mol to J/mol.
STEP 4: Plug the given variables into the equation.
STEP 5: Solve the portions of the equation with values into order to simplify.
STEP 6: Combine the integers on the right side by multiplying them together.
STEP 7: Take the inverse e of the natural logarithm to both sides of the equation to further isolate k2.
When you take the inverse of the natural logarithm you eliminate the ln on the left side.
This simplifies the expression to:
STEP 8: Solve for the right side of the equation.
STEP 9: Multiply both sides by 5.4 x 10-3 s-1.
STEP 10: Isolate the value for k2.
The Arrhenius equations form an integral part in our study of Chemical Kinetics and give a mathematical view of how molecules successful react to form products. Other important concepts dealing with the rates of chemical reactions include general rate, rate law, the integrated rate laws, energy diagrams and reaction mechanisms.
Jules felt a void in his life after his English degree from Duke, so he started tutoring in 2007 and got a B.S. in Chemistry from FIU. He’s exceptionally skilled at making concepts dead simple and helping students in covalent bonds of knowledge.
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