After Being Ignited In A Bunsen Burner Flame? 300 Most Correct Answers

In an endothermic reaction, the reaction mixture absorbs heat from the environment. Hence the products have higher energy than the reactants and #Delta”H”# is positive.

In an exothermic reaction, the reaction mixture gives off heat to the environment. Hence the products will have lower energy than the reactants and #Delta”H”# will be negative.

When the concentration of a dissolved reactant is increased or the pressure of a reacting gas is increased:

There are more reactant particles in the same volume

There is a greater chance that the particles will collide

The reaction speed increases

Temperature and rate of a chemical reaction

key concepts

Reactants must move fast enough and hit each other hard enough for a chemical reaction to take place.

Increasing the temperature increases the average velocity of the reactant molecules.

The faster more molecules are moving, the more molecules are moving fast enough to react, resulting in faster formation of products.

summary

Students will prepare the same two clear, colorless solutions (baking soda and calcium chloride) from Lesson 3. You will help design an experiment to see if the temperature of the solutions affects their reaction rate. Students will then attempt to explain at the molecular level why temperature affects reaction rate.

objective

Students are able to identify and control variables to design an experiment to see if temperature affects the rate of a chemical reaction. The students can explain at the molecular level why the temperature of the reactants influences the reaction rate.

Evaluation

Download the student activity sheet and distribute one per student if indicated in the activity. The activity sheet serves as the “evaluate” component of any 5-E lesson plan.

security

Make sure you and the students wear well-fitting safety goggles.

materials for the demonstration

Hot water in an insulated mug

Ice water in an insulated mug

2 glowsticks

materials for each group

Factors affecting reaction speed

Jessie A Key

Learning Outcomes Gain an understanding of collision theory.

Gain an understanding of the four main factors that affect reaction speed.

Reaction kinetics is the study of the rate of chemical reactions, and reaction rates can vary widely over a wide range of timescales. Some reactions can occur at explosive rates, such as the detonation of fireworks (Figure 17.1 Fireworks at night over a river), while others can occur sluggishly over many years, such as the rusting of barbed wire exposed to the elements ( Figure 17.2 ” Rusted Barbed Wire”).

collision theory

To understand the kinetics of chemical reactions and the factors that affect the kinetics, we should first examine what happens during a reaction at the molecular level. According to the collision theory of reactivity, reactions occur when reactant molecules “effectively collide”. For an “effective collision” to occur, the reactant molecules must be correctly oriented in space to facilitate the breaking and forming of bonds and the rearrangement of atoms that lead to the formation of product molecules (see Figure 17.3 “Collision visualizations”).

During a molecular collision, molecules must also possess a minimal amount of kinetic energy for an effective collision to occur. This energy varies for each reaction and is called the activation energy (E a ​​) (Figure 17.4 “Potential energy and activation energy”). The reaction rate thus depends on the activation energy; A higher activation energy means fewer molecules have enough energy to undergo an effective collision.

Factors affecting the rate

There are four main factors that can affect the reaction rate of a chemical reaction:

reactant concentration. Increasing the concentration of one or more reactants will often increase the reaction rate. This happens because a higher concentration of a reactant leads to more collisions of that reactant in a given period of time. Physical state of reactants and surface. When reactant molecules are in different phases, such as in a heterogeneous mixture, the reaction rate is limited by the surface area of ​​the phases in contact. For example, when a solid metal reactant and a gas reactant are mixed, only the molecules present on the surface of the metal can collide with the gas molecules. So if you increase the surface area of ​​the metal by smashing it flat or cutting it into many pieces, the speed of reaction will increase. Temperature. Increasing the temperature typically increases the reaction rate. An increase in temperature increases the average kinetic energy of the reactant molecules. Therefore, a larger proportion of the molecules will have the minimum energy required for an effective collision (Fig. 17.5 “Temperature and reaction rate”).

presence of a catalyst. A catalyst is a substance that accelerates a reaction by participating in it without being consumed. Catalysts offer an alternative reaction route to obtain products. They are crucial for many biochemical reactions. They will be explored further in the Catalysis section.

KEY FINDINGS Reactions occur when two reactant molecules effectively collide, each with minimal energy and correct orientation.

The reactant concentration, the physical state of the reactants and the surface area, the temperature and the presence of a catalyst are the four main factors affecting the reaction rate.

media attributions

Reaction Rates (continued)

1. Combustion and activation energy

· As you recall, combustion involves a reaction between a fuel and oxygen.

· The products of a combustion reaction always contain an oxide that binds to a sacrificial atom in the fuel (we’re dramatizing here: the sacrificial atom is simply an atom that has donated electrons to the electron-hungry oxygen).

· If there is hydrogen in the molecules of the fuel, water is also formed.

Example 1: Predict what will form from the combustion of:

Fuel Combustion Products Equation Aluminum ( Al) Al 2 O 3 2 Al + 1.5 O 2 –> Al 2 O 3 Methane (CH 4 ) CO 2 + H 2 O CH 4 + 2O 2 –> CO 2 + 2H 2 O Gasoline ( C 8 H 18 ) CO 2 + H 2 O C 8 H 18 + 12.5 O 2 –> 8 CO 2 + 9 H 2 O Coal (mixture of S + C) SO 2 + CO 2 S+ O 2 –> SO2 ; C+ O 2 –> CO 2 wood (mixture of [C 6 H 10 O 5 ] n and [C 10 H 12 O 3 ] n ) CO 2 + H 2 O ———— ———————–

· Since very stable products are formed during combustion, combustion reactions are exothermic.

Example 2: Draw a reaction profile (enthalpy versus reaction progress) for a combustion reaction.

enthalpy (kJ)

progress of the reaction

· Although combustion reactions eventually release heat, they also need heat to get going. The temperature at which oxygen molecules collide with fuel molecules with enough energy to steal electrons and form products is called the ignition point.

Substance Flash point (oC) Paper 232 Wood (varies by species) 190-266 Cotton 266 Methyl alcohol 464 Natural gas (depending on composition) 482-632

· The small bump in the graph on the previous page represents the amount of energy required to reach ignition point. Many non-combustion reactions also require a starting aid, and it is generally known as the activation energy A e. The activation energy can be calculated by:

A e = H maximum – H reactants

Example 3. Calculate the activation energy for the following:

350kJ

300kJ

200kJ

Activation energy = A e = 350-300 kJ = 50 kJ

2. Burn rates

The first three factors that control burn rate are components of the fire triangle. They are heat, oxygen and fuel.

If any of the three components are missing, the rate is zero; in other words, the fire will stop. Firefighters are taught the fire triangle for obvious reasons:

Water extinguishes most fires because, with its high specific heat capacity, it absorbs a lot of heat and thus cools the fire.

CO 2 fire extinguishers that release carbon dioxide by reacting Na 2 CO 3 with H 2 SO 4 (or pressurizing liquid carbon dioxide)

smother the fire and deprive it of oxygen.

When fighting fires, it is important to prevent a fire from reaching fuel-rich areas such as B. Oil stoves. All gas lines leading to the burning area must be automatically shut off when the alarm goes off. See this link for more details on firefighting.

The following factors affect burn rate, or simply put, how fast something burns. Notice that among the five factors, there are 3 components of the triangle of fire.

· The type of fuel used.

· The oxygen concentration

· The surface of the fuel

· Temperature

· Presence of catalysts

Example 1 Describe how each of the first four factors listed above affects the rate at which wood is burned in a fireplace.

a. The type of fuel used:

· The moisture still present in the wood slows down the rate of combustion.

· Certain types of wood contain more resin or oils that increase the rate of combustion.

· Structurally, different trees synthesize different fibers and they don’t all burn at the same rate.

b. The oxygen concentration:

· The higher the oxygen concentration, the faster wood burns. This applies to all other combustible materials.

c. The surface of the fuel:

· If you take two identical logs and turn one into wood chips, the intact one will burn much slower. Due to its smaller surface area, fewer molecules are exposed to oxygen in the intact tree trunk. This means that oxygen molecules only collide with a tiny fraction of the total available wood molecules. Without collisions, no reaction takes place, and therefore low surface area dramatically slows the burn rate.

i.e. Temperature:

· An increase in temperature increases the kinetic energy of both oxygen and fuel molecules. This leads to more collisions overall and more efficient collisions between the fuel and O 2 . As more molecules react at the same time, the rate of combustion increases.

e. Catalysts:

· A catalyst is a chemical that accelerates a reaction. (The catalyst itself is not consumed in the reaction, as we will explore later.) Any chemical that makes it easier for oxygen to remove electrons from wood lowers the activation energy and increases the rate at which wood burns.

Burn Rate Exercises

1. Write a balanced equation for the combustion of pentane. Pentane is C 5 H 12 . Identify the fuel and oxide and add heat on the appropriate side of the equation.

2. Draw a reaction profile where H reactants = 100 kJ, activation energy is 50 kJ and DH = -25 kJ.

3. What is not essential for incineration?

A. a fuel B. oxygen

C. correspond D. to a sufficiently high temperature

4. Carbon burns in air to produce CO 2 and H 2 O. Under what conditions would you expect the fastest burn rate?

A. Chunk carbon is heated and allowed to burn naturally in air.

B. Forcing hot air over the charcoal while it burns.

C. Pulverizing the carbon and allowing it to burn naturally.

D. Pulverize the carbon and blow hot air over it.

5. How does the foam of a fire extinguisher help put out the fire?

6. A cook started a fire by forgetting the oil he was heating up. Then he pulled the pan off the stove and poured baking soda into the pan.

Explain why the chef acted the way he did. Why didn’t he use water?

7. Olive oil has a lower flash point than corn oil. Which is the more practical cooking oil?

8. According to the manufacturer, the average burn rate is at a certain

Kind of candle made of paraffin, C 25 H 52 , is 8.33 x 10-4 mol/min. This kind of

Candle is sold in four sizes: 25, 50, 75 or 100 g. You only want to use one

Candle of this type to provide 4 hours of uninterrupted light. What is the smallest

one you can buy for this purpose?

9. Molecularly, why does a combustion reaction need to reach an ignition point before a fire starts?

10. From a molecular point of view, why does the surface increase the rate of a reaction?

activation energy

The activation energy of chemical reactions

Only a small fraction of collisions between reactant molecules convert the reactants into the products of the reaction. This can be understood by returning to the reaction between ClNO 2 and NO.

ClNO 2 (g) + NO(g) NO 2 (g) + ClNO (g)

In this reaction, a chlorine atom is transferred from one nitrogen atom to another. For the reaction to take place, the nitrogen atom in NO must collide with the chlorine atom in ClNO 2 .

A reaction does not take place when the oxygen end of the NO molecule collides with the chlorine atom of ClNO 2 .

It also does not occur when one of the oxygen atoms of ClNO 2 collides with the nitrogen atom of NO.

Another factor that influences whether a reaction takes place is the energy that the molecules carry when they collide. Not all molecules have the same kinetic energy as shown in the figure below. This is important because the kinetic energy that molecules carry when they collide is the main source of the energy that must be invested in a reaction to get it going.

The total standard free energy for the reaction between ClNO 2 and NO is favorable.

ClNO 2 (g) + NO(g) NO 2 (g) + ClNO(g) Go = -23.6 kJ/mol

However, before the reactants can be converted into products, the free energy of the system must overcome the activation energy for the reaction, as shown in the figure below. The vertical axis in this diagram represents the free energy of a pair of molecules when a chlorine atom is transferred from one to the other. The horizontal axis represents the sequence of infinitesimal changes that must occur to convert the reactants to the products of that reaction.

To understand why reactions have activation energies, consider what must happen for ClNO 2 to react with NO. In the first place, these two molecules must collide and thus organize the system. Not only must they be brought together, but they must also be held in just the right orientation relative to one another for a reaction to take place. Both factors increase the free energy of the system by reducing entropy. Some energy must also be invested to start breaking the Cl-NO 2 bond to allow the Cl-NO bond to form.

NO and ClNO 2 molecules that collide in the right orientation, with enough kinetic energy to overcome the activation energy barrier, can react to form NO 2 and ClNO. As the temperature of the system increases, so does the number of molecules that carry enough energy to react in a collision. The reaction rate therefore increases with temperature. As a rule, the reaction speed doubles for every 10 °C temperature increase in the system.

Purists may notice that the symbol used to represent the difference between the product and reactant free energies in the figure above is Go, not Go. A lowercase “G” is used to remind us that this diagram represents the free energy of a pair of molecules when they react, not the free energy of a system containing many colliding pairs of molecules. If we average the results of this calculation over the entire array of molecules in the system, we get the free energy change of the system, Go.

Purists might also notice that the symbol used to represent activation energy is written with a capital “E”. This is unfortunate because the students think the activation energy is the change in the internal energy of the system, which is not entirely true. E a measures the change in potential energy of a pair of molecules required to begin the process of converting a pair of reactant molecules into a pair of product molecules.

Catalysts and the rate of chemical reactions

Aqueous solutions of hydrogen peroxide are stable until we add a small amount of the I ion, a piece of platinum metal, a few drops of blood, or a freshly sliced ​​turnip, at which point the hydrogen peroxide quickly decomposes.

2 H 2 O 2 (aqueous) 2 H 2 O (aqueous) + O 2 (g)

This reaction therefore provides the basis for understanding the effect of a catalyst on the rate of a chemical reaction. Four criteria must be met for something to be classified as a catalyst.

Catalysts increase the reaction rate.

Catalysts are not consumed by the reaction.

A small amount of catalyst should be able to affect the reaction rate for a large amount of reactants.

Catalysts do not change the equilibrium constant for the reaction.

The first criterion forms the basis for defining a catalyst as something that increases the rate of a reaction. The second reflects the fact that whatever is consumed in the reaction is a reactant and not a catalyst. The third criterion is a consequence of the second; Because catalysts are not consumed in the reaction, they can catalyze the reaction over and over again. The fourth criterion arises from the fact that catalysts accelerate the forward and reverse reaction rates equally, so the equilibrium constant for the reaction remains the same.

Catalysts increase reaction rate by providing a new mechanism that has lower activation energy as shown in the figure below. A larger proportion of the collisions that occur between reactants now have enough energy to overcome the activation energy for the reaction. This increases the reaction speed.

To illustrate how a catalyst can lower the activation energy for a reaction by providing a different pathway for the reaction, let’s look at the mechanism of the I- ion-catalyzed decomposition of hydrogen peroxide. In the presence of this ion, the decomposition of H 2 O 2 need not occur in a single step. This can be done in two steps, both of which are simpler and therefore faster. In the first step, the I ion is oxidized by H 2 O 2 to form the hypoiodide OI-.

H 2 O 2 (aqueous) + I-(aqueous) H 2 O(aqueous) + OI-(aqueous)

In the second step, the OI ion is reduced to I- by H 2 O 2 .

OI-(aq) + H 2 O 2 (aq) H 2 O(aq) + O 2 (g) + I-(aq)

Because there is no net change in the concentration of I ion as a result of these reactions, I ion meets the criteria for a catalyst. Since both H 2 O 2 and I- participate in the first step of this reaction, and the first step of this reaction is the rate-limiting step, the overall reaction rate is first order for both reagents.

Determination of the activation energy of a reaction

The rate of a reaction depends on the temperature at which it is carried out. As the temperature increases, the molecules move faster and therefore collide more often. The molecules also carry more kinetic energy. Thus, the fraction of collisions that can overcome the activation energy for the reaction increases with temperature.

The relationship between temperature and reaction rate can only be explained by assuming that the rate constant depends on the temperature at which the reaction occurs. In 1889, Svante Arrhenius showed that the relationship between temperature and rate constant for a reaction obeys the following equation.

In this equation, k is the rate constant for the reaction, Z is a constant of proportionality that varies from one reaction to another, E a is the activation energy for the reaction, R is the ideal gas constant in joules per mole Kelvin, and T is the temperature in Kelvin .

The Arrhenius equation can be used to determine the activation energy for a reaction. We start by taking the natural logarithm of both sides of the equation.

We then rearrange this equation to match the equation for a straight line.

y = mx + b

According to this equation, a plot of ln k versus 1/T should yield a straight line with a slope of – E a /R as shown in the figure below.

By carefully paying attention to the mathematics of logarithms, it is possible to derive another form of the Arrhenius equation that can be used to predict the effect of a change in temperature on the rate constant of a reaction.

The Arrhenius equation can also be used to calculate what happens to the rate of a reaction when a catalyst lowers the activation energy.

Learning Outcomes Describe the conditions for successful collisions that elicit reactions.

Describe the rate in terms of successful collision conditions.

Describe how changing the temperature, concentration of a reactant, or surface area of ​​a reaction affects the reaction rate.

Define a catalyst and how a catalyst affects the rate of a reaction.

We know that a chemical system composed of atoms (\(\ce{H_2}\), \(\ce{N_2}\), \(\ce{K}\), etc.), ions (\(\ce {NO_3^-}\), \(\ce{Cl^-}\), \(\ce{Na^+}\), etc.), or molecules (\(\ce{H_2O}\ ), \ (\ce{C_{12}H_{22}O_{11}}\) etc.). We also know that these particles move in random motion in a chemical system. The collision theory explains why reactions between these atoms, ions and/or molecules take place at this particle level. It also explains how it is possible to speed up or slow down reactions that occur.

Collision Theory Collision theory gives us the ability to predict what conditions are necessary for a successful reaction to take place. These conditions include: The particles must collide with each other. The particles must collide with enough energy to break the old bonds. The particles must have the correct orientation. A chemical reaction involves breaking bonds in the reactants, rearranging the atoms into new groupings (the products), and forming new bonds in the products. Therefore, not only must a collision take place between reactant particles, but the collision must also have enough energy to break all reactant bonds that need to be broken to form the products. Some reactions require less collision energy than others. The amount of energy that the reaction partners must have to break the old bonds so that a reaction can take place is called the activation energy, abbreviated \(\text{E}_a\). Another way to think of it is to look at an energy diagram as shown in the figure. Particles must be able to overcome the “bump” – the activation energy – if they are to react. If the reactant particles collide with less than the activation energy, the particles will recoil (bounce off each other) and no reaction will take place.

Reaction Rate Chemists use reactions to create a product that they have a use for. For the most part, the reactions that produce a desired compound are only useful if the reaction occurs at a reasonable rate. For example, using a reaction to make brake fluid wouldn’t make sense if the reaction would take 8,000 years to complete the product. Such a reaction would also not make sense if the reaction was so fast that it would be explosive. For these reasons, chemists want to be able to control reaction rates. In some cases, chemists want to speed up reactions that are too slow or slow down reactions that are too fast. In order to control reaction rates, we need to know the factors that affect reaction rates. Chemists have identified many factors that affect the rate of a reaction. The rate or speed at which a reaction occurs depends on the frequency of successful collisions. Remember that a successful collision occurs when two reactants collide with enough energy and with the right orientation. That is, as the number of collisions increases, the number of particles that have enough energy to react increases, and/or the number of particles with the correct orientation increases, the rate of reaction increases.

Effect of Temperature on Reaction Rate Reaction rate was discussed in terms of three factors: collision frequency, collision energy, and geometric orientation. Remember that collision frequency is the number of collisions per second. The impact frequency depends, among other things, on the temperature of the reaction. As the temperature increases, the average speed of the particles increases. The average kinetic energy of these particles is also increased. The result is that the particles collide more often because the particles are moving faster and encountering more reaction particles. However, this is only a small part of the reason why the rate is increasing. Just because the particles collide more often doesn’t mean the reaction will definitely happen. The main effect of increasing the temperature is that more of the colliding particles have the amount of energy needed for an effective collision. In other words, more particles have the necessary activation energy. At room temperature, the hydrogen and oxygen in the atmosphere do not have enough energy to reach the activation energy needed to produce water: \[\ce{O_2} \left( g \right) + \ce{H_2} \ left( g \right) \rightarrow \text{No reaction}

onumber \] At any point in time in the atmosphere, there are many collisions between these two reactants. However, what we find is that water is not formed from the oxygen and hydrogen molecules colliding in the atmosphere because the activation energy barrier is simply too high and all collisions result in rebound. If we raise the temperature of the reactants or add energy to them in some other way, the molecules have the necessary activation energy and can react to form water: \[\ce{O_2} \left( g \right) + \ce {H_2} \left( g \right) \rightarrow \ce{H_2O} \left( l \right)

onumber \] There are times when the speed of a response needs to be slowed down. Lowering the temperature could also be used to reduce the number of collisions that occur, and lowering the temperature would also reduce the kinetic energy available for activation energy. If the particles have insufficient activation energy, the collisions will result in rebound rather than reaction. If the rate of a reaction needs to be slower, keeping the particles of sufficient activation energy will definitely keep the reaction at a slower rate. Society uses the effects of temperature on reaction speed every day. Food storage is an excellent example of how the effect of temperature on reaction speed is being exploited by society. Consumers store food in freezers and refrigerators to slow down the processes that lead to spoilage. The decrease in temperature reduces the rate at which food breaks down or is broken down by bacteria. In the early 20th century, explorers were fascinated by being the first to reach the South Pole. In order to accomplish such a difficult task in an era without most of the technology we take for granted today, they have devised various ways of surviving. One method was to store their food in the snow for later use during their advance to the Pole. On some explorations, they buried so much food that they didn’t have to use it all, and some was left behind. When this food was found and thawed many years later, it turned out to still be edible. For example, when milk is stored in the fridge, the molecules in the milk have less energy. This means that while molecules will still collide with other molecules, few of them will react (meaning “spoil” in this case) because the molecules don’t have enough energy to overcome the activation energy barrier. However, the molecules have energy and collide, so over time the milk will spoil even in the fridge. Eventually, the higher-energy molecules gain the energy needed to react, and when enough of these reactions occur, the milk becomes “sour.” However, if the same carton of milk were at room temperature, the milk would react (in other words, “spoil”) much faster. Most molecules would have enough energy to overcome the room temperature energy barrier and many more collisions would occur. As a result, the milk spoils in a relatively short time. This is also why most fruits and vegetables ripen in summer when the temperature is much warmer. You might have experienced this firsthand if you’ve ever bitten into an unripe banana – it probably tasted sour and might even feel like biting into a piece of wood! When a banana ripens, numerous reactions take place, producing all of the compounds we expect to taste in a banana. However, this can only happen when the temperature is high enough for these reactions to produce these products.

Effect of Concentration on Reaction Rate If you had an enclosed space like a classroom and a red and a green ball were flying through the room in random motion and making perfectly elastic collisions with the walls and with each other, in a given time the balls would be a collide with each other a certain number of times, which is determined by the probability. If you put two red and one green ball in the room under the same conditions, the probability of a collision between a red and a green ball would be exactly doubled. The green ball would have twice the chance of hitting a red ball in the same amount of time. A similar situation occurs in chemical reactions. Particles of two gaseous reactants or two reactants in solution have a certain probability of colliding with each other in a reaction vessel. If you double the concentration of any of the reactants, the chance of a collision doubles. The reaction speed is proportional to the number of collisions per unit of time. If a concentration is doubled, the number of collisions is also doubled. Assuming the percentage of successful collisions does not change, doubling the number of collisions results in twice as many successful collisions. Reaction speed is proportional to the number of collisions over time; Increasing the concentration of any of the reactants increases the number of collisions and therefore increases the number of successful collisions and the speed of the reaction. For example, the chemical test used to identify a gas as oxygen or not relies on the fact that increasing the concentration of a reactant increases the reaction rate. The reaction we call combustion refers to a reaction in which a combustible substance reacts with oxygen. If we light a splint of wood (a thin piece of wood) and then blow out the fire, the splint will continue to glow in the air for a while. If we insert this glowing bar into a gas that doesn’t contain oxygen, the bar immediately stops glowing – that is, the reaction stops. Oxygen is the only gas that supports combustion, air is about \(20\%\) oxygen gas. If we take this glowing splint and insert it into pure oxygen gas, the reaction speed increases by a factor of five, since pure oxygen has 5 times the oxygen concentration of air. If the reaction occurring on the glowing rail increases its speed by a factor of five, the glowing rail will suddenly erupt in full flame again. This test, in which a glowing bar is thrust into a gas, is used to identify the gas as oxygen. Only when the oxygen concentration is higher than in the air does the glowing track burst into flames.

Effect of Surface Area on Reaction Rate The very first requirement for a reaction between reactant particles to occur is that the particles must collide with one another. In the previous section it was pointed out how increasing the concentration of the reactants increases the rate of the reaction as it increases the frequency of collisions between particles. It can be shown that the number of collisions that occur between reactant particles also depends on the surface area of ​​solid reactants. Consider a reaction between Reactant RED and Reactant BLUE where Reactant Blue is in the form of a single blob. Then compare this to the same reaction where the blue reactant was broken into many smaller pieces. In the diagram, only the blue particles on the clump’s outer surface are available for collision with the red reactant. The blue particles inside the clump are protected by the blue particles on the surface. In Figure A, if you count the number of blue particles available for collision, you will find that only 20 blue particles could be hit by one red reactant particle. In Figure A, there are a number of blue particles inside the clump that cannot be hit. In Figure B, however, the lump has been broken into smaller pieces and all of the inner blue particles are now on a surface and available for collisions. In panel B, more collisions occur between blue and red, and therefore the reaction in panel B occurs faster than the same reaction in panel A. Increasing the surface area of ​​a reactant increases the frequency of collisions and increases the reaction rate. Several smaller particles have more surface area than one large particle. The more surface area available for particles to collide with, the faster the reaction occurs. You can see an example of this in everyday life, if you have ever tried to start a fire in the fireplace. If you hold a match against a large tree trunk to start the wood burning, you will find that it was an unsuccessful attempt. Holding a match against a large log will not elicit enough reactions to keep the fire going by providing enough activation energy for further reactions. To start a wood fire, it is common to cut a log into many small, thin sticks called kindling. These thinner wooden sticks offer many times the surface area of ​​a single log. The match will successfully induce enough reactions in the kindling so that sufficient heat is given off to provide activation energy for further reactions. Unfortunately, there have been cases where serious accidents have been caused because the relationship between surface area and reaction speed has not been understood. One such example occurred in grain mills. A grain of wheat is not very combustible. It takes considerable effort to get a grain of wheat to burn. However, when the grain of wheat is pulverized and scattered through the air, a spark is enough to set off an explosion. When the wheat is ground into flour, it is pulverized into a fine powder and some of the powder is scattered in the air. A small spark is then enough to trigger a very rapid reaction that can destroy the entire grain mill. In a 10-year period from 1988 to 1998, there were 129 grain dust explosions at mills in the United States. Efforts are being made in grain mills today to use huge fans to circulate the air in the mill through filters to remove most of the flour dust particles. Another example is the operation of coal mines. Charcoal will burn naturally, but it takes effort to get the charcoal going; Once it burns, it burns slowly as only the surface particles are available to collide with oxygen particles. The inner char particles must wait for the outer surface of the lump of char to burn off before they can collide with oxygen. In coal mines, huge blocks of coal must be crushed before the coal can be taken out of the mine. When breaking up the huge blocks of coal, drills are used to drill into the coal walls. This drilling produces fine coal dust which mixes with the air; Then a spark from a tool can cause a massive explosion in the mine. Coal mine explosions occur for other reasons, but coal dust explosions have contributed to the deaths of many miners. In modern coal mines, lawn sprinklers are used to spray water through the air in the mine and this reduces airborne coal dust and eliminates coal dust explosions.

Effect of a Catalyst on Reaction Rate The final factor affecting reaction rate is the effect of a catalyst. A catalyst is a substance that accelerates the rate of a reaction without being consumed by the reaction itself. In the reaction of potassium chlorate, which breaks down into potassium chloride and oxygen, a catalyst is available that allows this reaction to proceed much faster than it would by itself under room conditions. The reaction is: \[2 \ce{KClO_3} \left( s \right) \overset{\ce{MnO_2} \left( s \right)}{\longrightarrow} 2 \ce{KCl} \left( s \ right) + 3 \ce{O_2} \left( g \right)

Number \] The catalyst is manganese dioxide, and its presence causes the reaction shown above to proceed several times faster than it would without the catalyst. When the reaction is complete, the \(\ce{MnO_2}\) can be removed from the reaction vessel and its state is exactly the same as before the reaction. This is part of the definition of a catalyst – that it is not consumed by the reaction. Note that the catalyst is not included in the equation as a reactant or product, but is listed above the yield arrow. This is the standard notation for using a catalyst. Some reactions proceed very slowly in the absence of a catalyst. In other words, the activation energy for these reactions is very high. When the catalyst is added, the activation energy is lowered because the catalyst provides a new reaction pathway with lower activation energy. In the figure on the right, the endothermic reaction shows the catalyst reaction in red with the lower activation energy, labeled \(\text{E}’_a\). The new reaction pathway has a lower activation energy but does not affect the energy of the reactants, the products, or the value of \(\Delta H\). The same applies to the exothermic reaction. The activation energy of the catalyzed reaction is lower than that of the uncatalyzed reaction. The new reaction pathway provided by the catalyst affects the energy required for reactant bonds to break and product bonds to form. While many reactions in the laboratory can be enhanced by increasing temperature, not all reactions that occur in our bodies throughout our lives can do so. In fact, the body has to be kept at a very specific temperature: \(98.6^\text{o} \text{F}\) or \(37^\text{o} \text{C}\). Of course, there are times when the body is fighting infections, for example, when body temperature may be elevated. But in general, in a healthy person, the temperature is fairly constant. However, many of the reactions that a healthy body depends on could never occur at body temperature. The answer to this dilemma are catalysts – also called enzymes. Many of these enzymes are made in human cells because human DNA contains the instructions for making them. However, there are some enzymes required by the body that are not made by human cells. These catalysts must be supplied to our body with food and are called vitamins. Reversible Reactions When we think of a chemical reaction, we usually think of the reactants being completely consumed, so there aren’t any left and we end up with only products. Also, we generally consider chemical reactions to be one-way events. You may have learned in previous science classes that this is one way to distinguish chemical changes from physical changes—physical changes (like ice melting and freezing) are easily reversed, but chemical changes are irreversible (quite hard to eliminate). fry an egg). In this chapter we will see that this is not always the case. We will see that many chemical reactions are actually reversible under the right conditions. And because many reactions can be reversed, our notion of a reaction ending with no remaining reactants, only products, needs to be modified. Here are some examples of reversible reactions: 1. Nitrogen dioxide, \(\ce{NO_2}\), a red-brown gas, reacts to colorless dinitrogen tetroxide, \(\ce{N_2O_4}\) : \ (\ce{2NO_2(g) \ rightarrow N_2O_4(g)}\) But the reaction can also go the other way around—nitrous oxide also easily decomposes to form nitrogen dioxide: \(\ce{N_2O_4(g) \rightarrow 2NO_2 (g)}\) Usually we write a reaction that splits into both Directions can go by using a double arrow (which sometimes appears as ↔ in these online notes): \(\ce{2NO_2(g) \leftrightarrow N_2O_4(g) }\) Since the reaction proceeds in both directions simultaneously, go us neither \(\ce{NO_2}\) nor \(\ce{N_2O_4}\). \(\ce{NO_2}\) is continuously consumed to form \(\ce{N_2O_4}\), but at the same time \(\ce{N_2O_4}\) forms more \(\ce{NO_2} \) 2. When hydrogen gas is passed over heated iron oxide, iron and steam are formed: (1) \(\ce{Fe_3O_4(s) + 4H_2 (g) \rightarrow 3Fe(s) + 4H_2O(g)}\ ) The reverse reaction can occur, when steam is passed over red-hot iron: (2) \(\ce{3Fe(s) + 4H_2O(g) \rightarrow Fe_3O_4(s) + 4H_2(g)}\) We can write these two equations together as: (3 ) \(\ce{Fe_3O_4(s) + 4H_2(g) \leftrightarrow 3Fe(s) + 4H_2O(g) }\) If we have written a reversible reaction in this way, we need to be able to distinguish in which direction the reaction goes. As written in reaction (3) above, we would say that in the forward reaction, iron oxide and hydrogen gas, the reactants, produce the products iron and steam. In the reverse reaction, iron reacts with steam to produce iron oxide and hydrogen gas. It is important to understand the terminology and use the terms correctly. Does it matter how we write our reversible reaction? It can also be written as \(\ce{3Fe(s) + 4H_2O(g) \leftrightarrow Fe_3O_4(s) + 4H_2(g)}\) Now iron and steam are the forward direction reactants and iron oxide and hydrogen gas would be the reactants the opposite direction.

Summary The collision theory explains why reactions between atoms, ions and molecules take place.

For a reaction to be effective, particles must collide with enough energy and be in the right orientation.

With increasing temperature, the energy that can be converted into activation energy in a collision increases and increases the reaction speed. Lowering the temperature would have the opposite effect.

As the temperature increases, the number of collisions increases.

Increasing the concentration of a reactant increases the frequency of collisions between reactants and will therefore increase the reaction rate.

Increasing the surface area of ​​a reactant (by breaking up a solid reactant into smaller particles) increases the number of particles available for collision and increases the number of collisions between reactants per unit time.

A catalyst is a substance that accelerates the rate of a reaction without being consumed by the reaction itself. When a catalyst is added, the activation energy is lowered because the catalyst provides a new reaction pathway with lower activation energy.

Vocabulary Catalyst – A substance that accelerates the rate of a reaction without being consumed by the reaction itself.

– A substance that accelerates the reaction rate without being consumed by the reaction itself. Surface Area to Volume Ratio – The comparison of the volume within a solid to the area exposed on the surface.

teacher overview

summary

In this investigation, students classify chemical reactions as exothermic or endothermic. Next, students examine the relationship between an observed change in temperature and the classification of a change as chemical or physical.

objective

Students will study energy changes during chemical reactions, heat of reaction (ΔH), and the connection between energy changes and chemical changes.

security

Make sure you and the students wear well-fitting safety goggles.

Acetic acid (vinegar) fumes can be irritating. Work in a well-ventilated area. In case of eye contact, rinse with water. The concentration of acetic acid in this experiment poses no significant hazards.

Calcium chloride can be irritating to body tissues. In case of contact, flush affected areas with water. Dispose of calcium chloride solutions in accordance with local regulations.

materials for each group

Vinegar

baking powder

calcium chloride

water

thermometer

4 small clear plastic cups

1 cup measuring cup

Measuring spoon (1 tablespoon, ½ teaspoon)

Required time

One lesson, approximately 45-50 minutes.

lab tips

After students have studied an example of an endothermic change and an example of an exothermic change, they are asked to explore the relationship between energy changes and chemical reactions. To do this, students may need guidance to get the idea that temperature changes can also accompany dissolution.

Students will find it easier to design a fair test when they are familiar with the definitions of physical and chemical changes. Students should suggest an experiment to you before testing their hypothesis. To observe a change in temperature during a physical change, students should develop a procedure like the following:

Put 10 ml of water in a small plastic cup and put a thermometer in the water. Record the initial temperature (T i ).

). Add ½ teaspoon of calcium chloride to the water and swirl the cup. After it has stopped changing, record the final temperature (T f ).

Discussion in front of the lab

This study introduces the concepts of enthalpy (heat) of ΔH in the context of exothermic and endothermic reactions. To give students a deeper foundation in fundamentals and to reinforce the basic concepts covered previously, you may want to review the mechanics of chemical changes, how to write balanced chemical equations, and the law of conservation of energy.

inclusion in the curriculum

This study could be integrated into a unit on chemical changes or thermochemistry.

Chris

Chemical Reactions: They are fundamental to chemistry; They make new things by rearranging other things. They can blow things up… or freeze things quickly. In short, they are great.

Brittny

But why do some chemical reactions release enormous amounts of energy while others absorb energy? In a chemical reaction, the main change relates to the way atoms are joined (or bonded) together. In order to change these connections, bonds must be broken and new bonds formed. Let’s break down how energy is transferred in these reactions.

Chris

To understand the energetic effects of chemical reactions, it is important to keep two key ideas in mind:

It takes energy to break bonds. When bonds are formed, energy is released.

To understand this, consider the chemical reaction between vinegar and baking soda. That’s right – the classic soda volcano experiment. The chemical reaction behind this science fair favorite involves baking soda — also known to chemists as sodium bicarbonate — and vinegar, also known as acetic acid.

These compounds react to form the molecules sodium acetate, water and carbon dioxide. The baking soda and vinegar are called the reactants. The formed sodium acetate, water and carbon dioxide are referred to as products.

Before the atoms in acetic acid and sodium bicarbonate can rearrange to form the products, some of the bonds between the atoms in these molecules must be broken, and because the atoms are attracted to each other, it takes energy to pull them apart.

Then, as the products (sodium acetate, water, and carbon dioxide) are formed, energy is released because atoms that have an attraction for one another are brought back together.

By comparing the energy absorbed in breaking bonds in the reactants with the energy released when bonds are formed in the products, you can determine whether a chemical reaction releases energy or absorbs energy overall.

Brittny

Chemical reactions that release energy are called exothermic. In exothermic reactions, more energy is liberated in forming bonds in the products than is expended in breaking bonds in the reactants. Exothermic reactions are accompanied by an increase in the temperature of the reaction mixture.

Chemical reactions that absorb (or expend) energy overall are called endothermic. In endothermic reactions, more energy is absorbed when bonds are broken in the reactants than is released when new bonds are formed in the products. Endothermic reactions are accompanied by a decrease in the temperature of the reaction mixture.

Chris

You can use energy level diagrams to visualize the change in energy during a chemical reaction. To understand these charts, compare the energy levels of the reactants on one side and the products on the other side.

For example, consider a graph showing the change in energy as a candle burns. Wax (C 34 H 70 ) burns in the presence of oxygen (O 2 ) to form carbon dioxide (CO 2 ) and water (H 2 O). Since more energy is released in the formation of the products than is expended in breaking down the starting materials, this reaction is exothermic.

Brittny

All of these things relate to thermodynamics – the study of heat and its relationship to energy and work. Using thermodynamics, you will learn how to calculate the exact amount of energy consumed or released by chemical reactions. Classifying a chemical reaction as exothermic or endothermic is straightforward. It comes down to balancing the energy required to break bonds in the reactants with the energy released when the products are formed.

It’s a simple idea, but one with a lot of power.

Potential energy is the energy due to position, composition, or arrangement. It is also the energy associated with forces of attraction and repulsion between objects. Any object that is lifted from its resting position has stored energy, hence it is called potential energy because it has a potential to do work when released.

Introduction For example, when a ball is released from a certain height, gravity pulls it and the potential energy is converted to kinetic energy as it falls. Since this energy converts from potential to kinetic energy, it is important to remember that energy can neither be created nor destroyed (law of conservation of energy). This potential energy becomes kinetic energy as the ball accelerates towards the ground. The total energy of the object can be found by summing these two energies. In an exothermic chemical reaction, potential energy is the source of energy. During an exothermic reaction, bonds break and new bonds form, and protons and electrons transition from a higher potential energy structure to a lower potential energy structure. During this change, potential energy is converted to kinetic energy, which is the heat released in reactions. The opposite occurs in an endothermic reaction. The protons and electrons move from an area of ​​low potential energy to an area of ​​high potential energy. That costs energy. Potential energy at the molecular level: Energy stored in bonds and static interactions are: Covalent bonds

Electrostatic Forces

Nuclear Forces

Gravitational potential energy \[PE= Fx\] where \(F\) is the opposing force and \(x\) is the distance travelled. To calculate the potential energy of an object on Earth or in another force field, the formula \[PE=mgh \label{pe1}\] is \(m\) the mass of the object in kilograms

\(g\) is the acceleration due to gravity. On Earth, that’s 9.8 meters/second 2

\(h\) is the height of the object. The height should be given in meters. If the above units for \(m\), \(g\), and \(h\) are used, the final answer should be in joules. Example \(\PageIndex{1}\) A 15 gram ball lies on top of a 2 m high refrigerator. What is the potential energy of the ball at the top of the fridge? Solution Use equation \ref{pe1} with \(m =15\, grams\). However, this mass must be specified in kilograms. The conversion from grams to kilograms is: 1,000 grams per 1 kg \(\text{height}=2\, m\)

\(g=9.8 \, m/s^2\) \[PE=(0.015 \, kg)(9.8 \, m/s^2)(2\,m)=0.294\, J

onumber\] Example \(\PageIndex{2}\) What is the mass of a shopping cart full of groceries standing on a 2 m high hill if its gravitational potential energy is 0.3 J? Solution Use equation \ref{pe1} \[0.3\,J=(m)(9.8\,m/s^2)(2\,m)

onumber\] and solve for mass \[m=0.015 \,kg=15\, g.

onumber\] Example \(\PageIndex{3}\) A 200 gram weight is placed on a shelf with a potential energy of 5 J. What is the weight? Solution \[5\,J=\left(\dfrac{200\,g}{1000\,g/kg}\right)(9.8 m/s^2)(h)

onumber\] and solve for height \[h=2.55\, m

Number\]

Coulomb potential energy The potential energy of two distant charged particles can be found by the following equation: \[E= \dfrac{q_1 q_2}{4π \epsilon_o r} \label{Coulomb}\] where \(r\) is distance

\(q_1\) and \(q_2\) are the charges

\(ε_0= 8.85 \times 10^{-12} C^2/J\,m\) For charges with the same sign, \(E\) has a + sign and tends to be smaller than \(r\) increases. This may explain why like charges repel each other. Systems prefer low potential energy and therefore repel each other, increasing the distance between them and decreasing potential energy. For charges with different charges, the opposite of what was said above applies. E has a – sign, which becomes even more negative as the oppositely charged particles attract or get closer together. Example \(\PageIndex{4}\) Calculate the potential energy of two particles with the charges \(3 \times 10^{-6}\, C\) and \(3.9 \times 10^{-6}\ , C\) are separated by a distance of \(1\, m\) Solution with equation \ref{Coulomb} \[\begin{align*} E &=\dfrac{(3\times 10^{-6} \ ,C)(3.9 \times 10^{-6}\,C)}{4π \,8.85 \times 10^{-12} \,C^2/Jm} \\[4pt] &=0.105 \, J \end{align*}\] Example \(\PageIndex{5}\) Find the distance between two particles that have a potential energy of \(0.2\, J\) and charges of \(2.5 \times 10^{ have -6}\, C\) and \(3.1 \times 10^{-6}\, C\). Solution \[\begin{align*} 0.2 &=\dfrac{(2.5 \times 10^{-6}\,C)(3.1 \times 10^{-6} \,C)}{4\pi (8.85 \times 10^{-12} \,C^2/Jm) r} \\[4pt] &=\dfrac{(8.99 \times 10^9)(7.75 \times 10^{-11})}{r } \\[4pt] &=\dfrac{0.6967}{r} \end{align*}\] Cross multiply and solve for \(r\) \[r=3.5\, m

onumber\] Includes all interactions in the system such as: in the nucleus of atoms; in atoms; between atoms in a molecule (intramolecular forces); and between different molecules (intermolecular forces).

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A catalyst is a substance that accelerates a chemical reaction, or lowers the temperature or pressure needed to start it, without itself being consumed during the reaction. Catalysis is the process of adding a catalyst to facilitate a reaction.

During a chemical reaction, the bonds between atoms in molecules are broken, rearranged, and rebuilt, causing the atoms to recombine into new molecules. Catalysts make this process more efficient by lowering the activation energy, which is the energy barrier that must be overcome for a chemical reaction to take place. As a result, catalysts facilitate the breaking of atoms and the formation of chemical bonds to make new combinations and new substances.

The use of catalysts leads to faster and more energy-efficient chemical reactions. Catalysts also have a key property called selectivity, which allows them to steer a reaction to increase the amount of the desired product and decrease the amount of unwanted by-products. You can create entirely new materials with entirely new uses.

In recent decades, scientists have increasingly developed specialized catalysts for essential real-world applications. High-performance catalysts in particular have changed the chemical industry. These advances have led to biodegradable plastics, new medicines, and greener fuels and fertilizers.

DOE Office of Science: Contributions to Catalyst Research

The Department of Energy (DOE) Office of Science Basic Energy Sciences program actively supports basic research on catalysts. DOE focuses on the design of new catalysts and the use of catalysts to control chemical transformations at the molecular and submolecular levels. DOE research emphasizes understanding these reactions and how to make them more efficient and targeted. The overarching goal of DOE is the development of new catalysis concepts and new catalysts to support the industry in producing fuels and chemicals from fossil and renewable raw materials more efficiently and sustainably. This research is helping advance solar fuels, which companies are making from the sun and common chemicals like carbon dioxide and nitrogen. This research is also creating advanced methods of converting discarded plastic into new products.

brief info

Humans have been using catalysts for thousands of years. For example, the yeast we use to bake bread contains enzymes, which are natural catalysts that help turn flour into bread.

The 2005 Nobel Prize in Chemistry was awarded to three researchers (Yves Chauvin, Robert H. Grubbs, and Richard R. Schrock) for their work on metathesis catalysts. dr Grubbs and Schrock were funded in part by the DOE for their Nobel Prize research. dr Schrock continues to be funded by the DOE.

The 2018 Nobel Prize in Chemistry was awarded to Frances H. Arnold for her pioneering work directing the evolution of enzymes for environmentally benign applications such as renewable fuels. It is partially funded by DOE.

Visit the Argonne National Lab for seven more things you might not know about catalysis

resources

Scientific terms can be confusing. DOE Explains provides simple explanations of key words and concepts in basic science. It also describes how these concepts apply to the work of the Department of Energy’s Office of Science as it helps the United States excel in research across the scientific spectrum.

What would not increase reaction speed?

a) Increasing the volume of the container for a gaseous reaction.

b) addition of a catalyst.

d) increasing the concentration of the reactants.

d) increasing the surface area of ​​a solid reactant.

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Set up the equipment with 30cc^3 of diluted muriatic acid in a styrofoam cup standing in a cup. With a lid with a hole and a thermometer in the hole in the lid.