1. Thermal changes in chemical reactions
Experimental Principle:
The change in energy in a chemical reaction, usually manifested as a change in heat. If there is heat released in the reaction, it is an exothermic reaction; On the contrary, if there is heat absorption in the reaction, it is an endothermic reaction.
Experimental Operation:
Add a small amount of hydrochloric acid to a test tube, add a small amount of NaOH solution, touch the outer wall of the test tube with your hands, measure the temperature with a thermometer, and note the temperature change.
Put a few zinc granules into a test tube, add 5ml of 2mol l hydrochloric acid, and touch the outer wall of the test tube with your hands, how does it feel?
Add a certain amount of Ba(OH)2 crystals to a small beaker, and then add a certain amount of NH4Cl crystals to the beaker, and stir quickly with a glass rod to make a full reaction, touch the outer wall of the beaker with your hands, what is the feeling?
Experimental Phenomena and Conclusions:
The outer wall of the test tube is heated and the temperature rises, indicating that heat is released during the reaction.
NaOH + HCl NACL + H2O (exothermic).
The reaction produces a large number of bubbles, and at the same time, the outer wall of the test tube heats up and the temperature rises, indicating that there is heat released during the reaction. Zn+2HCl, ZNCL2+H2 (Exothermic).
The outer wall of the beaker is cold and the temperature decreases, indicating that the reaction absorbs a large amount of heat.
Ba(OH)2+2NH4Cl BACl2+2NH3 +2H2O (endothermy).
Question:
What are the common exothermic reactions? What are the common endothermic reactions?
Everyone has felt the endothermic reaction and exothermic reaction, what are the applications in life and production?
Why does a chemical reaction have a change in energy? How to explain it from two different perspectives: macro and micro.
2. The influence of temperature, concentration and catalyst on the rate of chemical reaction
Experimental Principle:
The effect of conditions such as concentration, temperature, pressure, and catalyst on the rate of chemical reactions.
Experimental Operation:
Add the same volume of 0. to two test tubes containing magnesium strips and iron sheets with approximately the same surface area at the same time5mol l hydrochloric acid, observe the phenomenon that occurs.
Add the same volume of 0. to two tubes containing iron sheets with the same surface area at the same time5mol l and 3mol l of hydrochloric acid, observe the phenomenon that occurs.
Iron sheets with the same surface area were added to test tubes filled with hydrochloric acid at different temperatures at the same time, and the phenomenon was observed.
Manganese dioxide powder was added to two test tubes containing the same volume and concentration of hydrogen peroxide and the phenomenon was observed.
The same volume and concentration of hydrochloric acid were added dropwise to two test tubes containing lumpy calcium carbonate and calcium carbonate powder of the same mass at the same time, and the phenomenon was observed.
Experimental Phenomena and Conclusions:
Air bubbles are formed on the magnesium strips and quickly, and the magnesium strips disappear quickly. Air bubbles are produced on the iron sheet but slowly, and the iron sheet does not disappear.
mg+2hcl＝mgcl2+h2↑
fe+2hcl＝fecl2+h2↑
Mg is more active than Fe and reacts violently and quickly with acid.
The results show that when other conditions are equal, the more reactive the metal is, the faster the reaction with the acid. The rate of chemical reaction is determined by the properties of the reactants themselves (intrinsic causes).
Add 0The bubbles produced in the test tube with 5 mol l of hydrochloric acid are slow, and the air bubbles produced in the test tube with 3 mol l of hydrochloric acid are faster.
It shows that the greater the concentration of reactants, the faster the chemical reaction when other conditions are equal.
Bubbles are produced quickly at high temperatures, and bubbles produced slowly at low temperatures.
It shows that the higher the temperature of the reactants, the faster the chemical reaction when other conditions are equal.
The bubbles produced in the test tube with manganese dioxide and manganese dioxide were fast, and the bubbles produced in the test tube without manganese dioxide were slow.
In the hydrogen peroxide splitting reaction, manganese dioxide, ferric chloride, etc. can be used as catalysts, 2H2O2 (manganese dioxide) 2H2O+O2
It shows that when other conditions are equal, the chemical reaction rate is accelerated by using a catalyst.
The test tube containing calcium carbonate powder produces bubbles quickly, and the test tube containing lumpy calcium carbonate produces bubbles slowly.
caco3+2hcl＝cacl2+h2o+co2↑
Other things being equal, the larger the surface area of the solid reactants, the faster the chemical reaction.
In general, the rate of chemical reactions can be increased by increasing the concentration of reactants, increasing the temperature, increasing the contact area of reactants, and using catalysts.
Question:
A certain amount of hydrochloric acid reacts with an excess of iron powder, in order to slow down the reaction rate and not affect the total amount of H2 generated, an appropriate amount of ( ) can be added to the hydrochloric acid
a．naoh(s)
b．h2oc．na2co3(s)
d．ch3coona(s)
e NaCl solution.
A certain amount of hydrochloric acid reacts with an excess of iron powder, in order to speed up the reaction rate without affecting the total amount of H2 generated, the measures that can be taken are ( ).
a Appropriate temperature increase
b Add a small amount of CuSO4(S).
c Add a small amount of concentrated H2SO4
d Add a certain amount of magnesium powder
e．nano3(s)
Please use what you have learned to answer the questions:
Firefighters are careful when opening the door when fighting a fire, because it can happen as soon as the door is opened**. Please explain why.
From the perspective of chemical reactions, what is the main purpose of using a refrigerator?
Why is pressure cooker cooking fast?
Why does enzyme laundry detergent remove stains quickly?
Why do you have to chop small firewood when cooking rice?
Why should chlorinated water, concentrated Hno3, Agno3 solution, KMno4 solution, AGBR(S), etc. be stored in brown bottles and kept in a cold and dark place? Why is it easy to irradiate the mixture of H2 and Cl2 and methane and Cl2 when exposed to strong light**?
Take a certain quality of sandpaper sandpaper and mix it with a sufficient amount of concentrated hydrochloric acid, why does the reaction speed increase first, then decrease, and finally stop?
3. The influence of concentration and temperature on chemical equilibrium
Experimental Principle:
Effect of concentration, temperature, pressure, etc., on chemical equilibrium.
Experimental Operation:
Xiang Sheng has 5ml0Add 5 ml of 01 mol of LFECl3 solution to the tube03mol lkscn solution, the solution is red. The solution was evenly divided into three test tubes, and a small amount of 1mol LFECl3 solution was added to the first test tube, and the color change of the solution was observed by full shaking. A small amount of 1mol lkscn solution was added dropwise to the second test tube, and the color change of the solution was observed. Add a small amount of NaOH solution dropwise to the third tube and observe the observed phenomenon.
The gas mixture of NO2 and N2O4 is placed in two connected flasks and the rubber tube is clamped with clamps. Place one flask in hot water and the other in ice water (or cold water) and observe the phenomenon.
Experimental Phenomena and Conclusions:
The solution in the first tube darkens in color; The solution in the second tube darkens in color; There was a reddish-brown precipitate in the third tube, and the solution became lighter in color. Since there is the following equilibrium in this reaction system:
fe3++3scn－＝fe(scn)3
The dropwise concentrated FeCl3 solution and the dropwise concentrated KSCN solution both increased the concentration of reactants, shifted the equilibrium to the right, and deepened the color. Dropwise NaOH solution reacts with Fe3+ to form a reddish-brown Fe(OH)3 precipitate, which reduces the concentration of reactants, shifts the equilibrium to the left, and lightens the color.
It shows that when other conditions remain unchanged, the concentration of reactants decreases or the concentration of products increases, and the equilibrium shifts to the direction of consuming products, that is, the equilibrium shifts to the left. When the concentration of reactants increases or the concentration of products decreases, the equilibrium shifts in the direction of consuming reactants, that is, the equilibrium shifts to the right.
In a beaker of hot water, the gas darkens in color; In a beaker of cold water, the gas becomes lighter in color. Due to the following equilibrium of the gases in the flask:
2no2(g)＝n2o4(g)，h=-56.9kj/mol
This reversible reaction is exothermic in the forward reaction and endothermic in the reverse reaction. NO2(G) is reddish-brown and N2O4(G) is colorless, so the color of the gas mixture is determined by the concentration of NO2. The color of the heating gas deepened, indicating that the concentration of NO2 increased, that is, the equilibrium shifted to the direction of heat absorption. The color of the cooling gas becomes lighter, indicating that the concentration of NO2 decreases, that is, the equilibrium shifts in the direction of exothermy.
It shows that when other conditions remain constant, the temperature increases and the equilibrium moves to the endothermic direction. The temperature is lowered and the equilibrium is shifted in the direction of exothermy.
Question:
How do the rates of forward and reverse reactions change when the concentration of reactants increases or decreases?
How do the rates of forward and reverse reactions change during heating or cooling? How do these changes in reaction rates affect the chemical equilibrium shift?
What is the relationship between the rate of chemical reaction and chemical equilibrium?
If a reversible reaction reaches equilibrium under certain conditions, if the rate of chemical reaction changes, does the equilibrium necessarily move? If the equilibrium shifts, does the rate of chemical reaction necessarily change?
Can the amount of oxygen be increased by adding manganese dioxide when decomposing oxygen from potassium chlorate?
The following equilibrium is known to exist in the solution of K2Cr2O7.
cr2o72－+h2o2＝cro42－+2h+
Orange. Yellow).
What is the phenomenon of adding NaOH solution (6mol L) or dilute H2SO4 solution?
In a constant-capacity vessel, after 2NO2(G) N2O4(G) is equilibrated, the reactant NO2 is reintroduced into the vessel, or the product N2O4 is reintroduced into the vessel, how does the equilibrium move in both cases? In both cases, how does the percentage of NO2 change after the equilibrium is re-equilibrated?