Short Notes on Types of Thermal reactions.
Exothermic Reaction –
- negative heat flow
- heat loss to the surroundings (get hot)
- produce heat or can be explosive
- can be spontaneous
Na + Cl  NaCl gives off 411 KJ of energy
Zn + HCl  ZnCl + H2 (corrosion)
How can you tell that a chemical change has taken place?
What problem could result if you put the cork on too tight?
You produced hydrogen gas (what are 4 properties)
How is hydrogen like oxygen?
How is hydrogen different?
Endothermic Reaction –
- must absorb energy to occur
- cannot be spontaneous – work must be done
- absorb heat so reaction “gets cold”
photosynthesis requires 15MJ of energy for 1 kg of glucose (what energy is put into the reaction)
Title: Reaction Demos
Purpose: recognize the parts and importance of chemical equation when describing how things
chemically react to form new substances. Identify exothermic and endothermic reaction.
8.9 substances have chemical physical properties
8.10 complex interactions occur between mater and energy
Materials: labeled chemicals, test tubes, test tube rack, goggles, apron
Procedure:
1. On your reaction chart (for each demonstration) place a red box around the reactants and a blue
circle around the products.
2. Follow the exact instructions which go along with the chemical reaction. Be sure to wear goggles
and aprons at all times. Dispose of wastes properly.
3. Fill in the rest of chart for the chemical reaction. Use your notes if needed to help identify the type
(if possible) be sure to provide at least two (2) pieces of evidence for the reaction.
Teacher demo and discussion:
2Zn + 2HCl  2ZnCl + H2
How many molecules of HCl are in the reactants?
How many atoms of Zn are on the products side of the equation?
4Fe + 3O2  2Fe2O3 (rust)
Identify one molecule in the equation.
How many atoms are in one compound of Fe2O3?
Equation
Thermal Type Reaction Type
single,
(exo or endo) (syn, double,
decomp)
Other Evidence
6CO2 + 6H2O C6H12O6 + 6O2
Photosynthesis is the conversion of light energy into chemical energy by living organisms. The
raw reactants are carbon dioxide and water; the energy source is sunlight; and the products
include oxygen and (energy rich) carbohydrates, for example sucrose and starch.
Photosynthesis occurs in the plastid, chloroplast. The chloroplast is the green colored
organelle found in plant cells. Some plants have different colored pigments in their chloroplast
such as red.
2NH4NO3  2N2 + O2 + 4H2O
1. Add 10 ml of H2O to the test tube. Add the
thermometer and observe the initial
temperature.
2. Place 1 scoop of NH4NO3 in the test tube.
3. Gently swirl the test tube.
4. Observe the reaction and fill in your data
chart.
5. When finished, pour products through the
strainer in the back sink. Place remaining
NH4NO3 in the proper container.
6. Clean up when finished and return all
materials to the bin.
C6H8O7 + 3NaHCO3 3CO2+ 3H2O + NaC6H5O7
1. Place 10 ml of C6H8O7 to the test tube. Add
the thermometer and observe the initial
temperature.
2. Add 1tsp of NaHCO3 to the test tube.
3. Stir or swirl the chemical together.
4. Observe the reaction, including the
temperature change. Fill in your data chart.
5. When finished, pour products into the sink.
6. Clean up your lab area and return all materials
to the bin.
CaCl2 + 2Na(HCO3) 2NaCl + CO2 + Ca(HCO3)2
1. Measure 2 eye droppers full of C19H1305S and place
into the test tube. The phenol red is a combination of
C19H1305S powder and water. Phenol Red is an acid
base/base indicator. It is generally red in the presence of
bases and yellow in presence of acids.
2. Add the thermometer and observe the initial
temperature.
3. Put 10 ml of NaHCO3 into the test tube. Add 1 scoop of
CaCl2.
3. Gently swirl the contents together.
4. Observe the reaction and write your observations in
the data chart.
6. Pour contents into sink.
7. Clean up your lab station and return all materials to
the bin. You may need to scrub out the test tube.
2H2O2 ----> 2H2O + O2
In this lab, you will study an enzyme that is found in the cells of many living
tissues. The name of the enzyme is catalase (KAT-uh-LAYSS); it speeds up a
reaction which breaks down hydrogen peroxide, a toxic chemical, into 2
harmless substances--water and oxygen.
This reaction is important to cells because hydrogen peroxide (H2O2) is
produced as a waste product (byproduct) of many normal cellular reactions.
If the cells did not break down the hydrogen peroxide, they would be
poisoned and die. In this lab, you will study the catalase found in liver cells.
It might seem strange to use dead cells to study the function of enzymes.
This is possible because when a cell dies, the enzymes remain intact and
active for several weeks, as long as the tissue is kept refrigerated.
1. Place 5 ml of H2O2 to a test tube.
2. Place a thermometer in the test tube. Observe the
initial temperature.
3. Add a small piece of liver to the test tube.
4. Observe the reaction and write your observations in
the data chart.
6. Pour contents into sink.
7. Clean up your lab station and return all materials to
the bin. You may need to scrub out the test tube.
Na(OH) + HCl  NaCl +H2O
1. Add 5 ml of HCl to the test tube. Add the thermometer
and observe the initial temperature.
2. Put 1 scoop of Na(OH) into the test tube.
3. Gently swirl the contents together.
4. Observe the reaction and write your observations in
the data chart.
5. Pour contents into sink.
6. Clean up your lab station and return all materials to
the bin. You may need to scrub out the test tube.
NaHCO3 + HOOCCH3  NaOOCCH3 + H2O + CO2
1. Add 5 ml of HOOCCH3 to the test tube. Add the
thermometer and observe the initial temperature.
2. Put 1 scoop NaHCO3 of into the test tube.
3. Gently swirl the contents together.
4. Observe the reaction and write your observations in
the data chart.
5. Pour contents into sink.
6. Clean up your lab station and return all materials to
the bin. You may need to scrub out the test tube.
Al2(SO4)3 + 6 NaHCO3 → 2 Al(OH)3 + 3 Na2SO4 + 6 CO2
1. Add 5 ml of NaHCO3 to the test tube. Add the
thermometer and observe the initial temperature.
2. Put 1 scoop of
Al2(SO4)3
into the test tube.
3. Gently swirl the contents together.
4. Observe the reaction and write your observations in
the data chart.
5. Pour contents into sink.
6. Clean up your lab station and return all materials to
the bin. You may need to scrub out the test tube.
CaCO3 + HCl  CO2 + CaCl2 + H20
1. Put 5 ml of CaCO3 into the test tube.
2. Add the thermometer and observe the initial
temperature.
3. Put 10 ml of HCl into the test tube.
3. Gently swirl the contents together.
4. Observe the reaction and write your observations in
the data chart.
6. Pour contents into sink.
7. Clean up your lab station and return all materials to
the bin. You may need to scrub out the test tube.
CO2 + Ca(OH)2  CaCO3 + H2O
1. Place 10 ml of Ca(OH)2 to the test tube. Add
the thermometer and observe the initial
temperature.
2. Add 5 ml of CO2 to the test tube.
3. Stir or swirl the chemical together.
4. Observe the reaction, including the
temperature change. Fill in your data chart.
5. When finished, pour products into the sink.
6. Clean up your lab area and return all materials
to the bin.
Key for Reactions
Zn + HCl
Exothermic, single replacement
Fe + O2
Exothermic, synthesis
Photosynthesis
Endothermic, synthesis
NaH4NO3 (ammonium nitrate)
Endothermic, decomposition
H3C6H5O7 (citric acid) + NaHCO3 (sodium bicarbonate) Endothermic, Decomposition
CaCl2 (calcium chloride) + NaHCO3 (baking soda)
Exothermic, Double replacement
CO2 (from sparkling water)+ Ca(OH)2 (lime water) CaCO3 + H2O
white precipitant
Liver and H2O2 peroxide
single replacement,
Exothermic, decomposition (enzymatic action)
Na(OH) + HCl > NaCl +H2O
double replacement
Lithium chloride + water exothermic (40g LiCl 50ml H2O)
Potassium Chloride + water endothermic
Ammonium chloride + water endothermic
C3H5O(COOH)3 (fruit juice) with ammonia makes a salt?
Al2(SO4)3 + 6 NaHCO3 → 2 Al(OH)3 + 3 Na2SO4 + 6 CO2
Aluminium Sulfate + baking soda endothermic = used for fire extinguishers
Calcium hydroxide in water Ca(OH)2 (lime water) (see above)
CaCO3 + H2O + 2HCl (any acid)  CO2 + CaCl2 + H20 (exothermic?) turns from milky
white to clear
Barium chloride and sodium sulfate
Epsom salt and alum
Aluminum and copper chloride
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KCl(aq) + Pb(NO3)2(aq) --> KNO3(aq) + PbCl2 precipitate
The only way you can form a precipitate is if the cation and anion
combination have a very low solubility product. Unfortunately, most
nitrates and carbonates are soluble, and of the chlorides only silver
chloride and lead chloride are insoluble. You may have to get access to
some silver nitrate or lead(II) nitrate - which when combined with the
chlorides should form a precipitate. Alternatively, barium hydroxide and
calcium hydroxide are only sparingly soluble - so if you can get a hold of
sodium or potassium hydroxide - that would work too.
There are two issues here. The first is safety. Strontium, and Barium can be
toxic if ingested so cautious handling and supervision are necessary.
Nitrates can also be toxic but to a lesser extent. An interesting
precipitation reaction is to take a solution of freshly prepared calcium
chloride and have the students exhale, that is, blow bubbles through a soda
straw immersed in the calcium chloride solution. Carbon dioxide exhaled from
the lungs forms calcium carbonate which causes the solution to turn cloudy.
Keep the stock solution of CaCl2 protected from the atmosphere because it
will absorb CO2 from the atmosphere. Strontium and barium chloride will
undergo the same precipitation reaction, but I would be cautious about
handling these salts by 6-8th graders without careful supervision. You could
compare the results of exhaling through a solution of Ca(Cl)2 and Mg(SO4)
[Epsom's salt] which is available at any pharmacy and most grocery stores.
The solubility product constant of Mg(CO3)and Ca(CO3) is 6x10^-6 and 3x10^-9,
so the difference in the amount of breath necessary to precipitate the two
solutions should be observable, although
I have not done the experiment.
Another demonstration could be done with Ca(CO3). While it is only very
slightly soluble in water, you could start with a slurry of the solid and add
HCl. This will cause the precipitate to dissolve, and the solution to become
clear. Then, raise the pH with KOH (or NaOH). As the solution becomes more
alkaline, Ca(OH)2 will precipitate because its solubility product constant is
about 5x10^-6.
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Endothermic Reaction
reaction of barium hydroxide octahydrate crystals with dry ammonium chloride
dissolving ammonium chloride in water
reaction of thionyl chloride (SOCl2) with cobalt(II) sulfate heptahydrate
mixing water and ammonium nitrate
mixing water with potassium chloride
reacting ethanoic acid with sodium carbonate
photosynthesis (chlorophyll is used to react carbon dioxide plus water plus energy to make
glucose and oxygen)
Endothermic and Exothermic Reactions
Introduction:
This demonstration illustrates how chemical reaction can either give off heat (exothermic) or
absorb heat (endothermic). The crystallization of a supersaturated sodium acetate solution is an
exothermic process that is available commercially is the form of hand warmers. This
demonstration can also be carried out in a large flask by seeding the supersaturated solution with
a small sodium acetate crystal.
To demonstrate an endothermic process, Barium hydroxide (octahydrate) and Ammonium nitrate
are mixed (in a 2:1 ratio) in a small beaker. This reaction displays the endothermic process and
illustrated the interaction between changes in enthalpy and entropy is spontaneous chemical
reactions. For a process to take place spontaneously at constant temperature and pressure, the
change in free energy must be negative. An endothermic reaction may thus be spontaneous at
constant pressure if the positive value of the heat absorbed is offset by a sufficient increase in
entropy (randomness). In the reaction between barium hydroxide octahydrate and ammonium
nitrate, the large increase in entropy is related to the increase in the number of particles present
and their states (remember that two solids are combining to form a solid product and some
liquid). As the reaction below shows, there are three molecules that combine to form 13 product
molecules.
With this reaction, temperatures of -20o C can be achieved. Both of these demos can be done
hands-on by volunteers from the audience.
Materials:
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Flask of supersaturated sodium acetate
Sodium acetate seed crystals
Handwarmers (one for each or every other student)
Approximately 32g of barium hydroxide octahydrate
Approximately 17g of ammonium nitrate
125 mL Erlenmeyer flask
Glass stirring rod
Small piece of cardboard
Squeeze bottle of water
Exothermic Reactions
Procedure:
Handwarmers can be passed out among the audience (for smaller groups there should be enough
for each or at least every other student). Have the student click the metal disk inside the
handwarmers to activate the crystallization. They should notice that the crystallization will begin
at the metal disk and spread outward though the whole package. As an alternative, the large flask
containing the sodium acetate solution can be used. To activate the crystallization, add a single
seed crystal of sodium acetate. If the solution has been regenerated properly you should observe
a long crystalline spike run out from the seed crystal and eventually spread through the whole
solution. As the solution crystallizes, the system will give off heat.
Endothermic Reactions
Procedure:
Place a pre-weighed amount of solid barium hydroxide octahydrate and ammonium nitrate into a
Erlenmeyer flask and stir the two components together. Within about 30 seconds the odor of
ammonia will become evident and the liquid product will begin to form.
Within a few minutes the temperature of the flask will drop to about -20o or -30o C. If a small
amount of water is placed on a small piece of cardboard, the flask will freeze the water and the
cardboard will stick to the flask after a minute or two of contact. Alternatively, this may be son
with a small wooden block. The temperature of the flask may be measured directly if a
thermometer is available.
Helpful Hints:
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It should be noted that the hand warmers can be recharged by boiling them in water for
six to ten minutes.
Be sure to collect all of the hand warmers when you are done, the have a tendency to
disappear.
Ask the students see if they can come up with any practical applications for both
endothermic and exothermic reactions. They may suggest something like a car battery
heater or a chemical ice pack for injured muscles.
Safety:
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Be aware that the temperatures achieved in the endothermic reaction are well below the
normal body temperature and the reaction flask should not be handled for prolonged
periods of time.
Inhalation of concentrated ammonia vapor can cause edema of the respiratory tract,
spasm of the glottis, and asphyxia. Show the students how to properly smell a reaction
flask.
Soluble barium salts are poisonous if ingested. Upon contact with the skin, barium and
ammonium salts may produce minor irritations or cause allergic reactions. If the flask is
spilled or broken, its contents should be flushed down the drain with copious amounts of
water.
Handwarmers do tend to get very warm. It is possible that someone who is sensitive to
heat or cold could be burned.
Due to the ammonia vapor given off, the endothermic demonstration should not be done
in a room with poor ventilation.
Always were safety goggles when working with these chemicals. Be sure to provide
safety goggles for any of the audience members who may be helping you or who are
handling these demonstrations.
Exo
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Digestion of food releases energy
All combustion reactions (fires)
C + O2  CO2 + energy
Adding an alkali metal to water
2 Na + 2 H2O  2 NaOH + H2 + energy
Condensation of water
Explosion of bombs
Hand warmers
Hand blasters
Flameless ration heaters
Exothermic or endothermic?
This is a useful class practical to introduce energy changes in chemical reactions. The students
measure the temperature changes in four reactions, and classify the reactions as exothermic or
endothermic. The experiments can also be used to revise different types of chemical reaction
and, with some classes, chemical formulae and equations.
Read our standard health & safety guidance
Lesson organisation
There are five solutions and three solids involved. Careful consideration will need to be given as
to the most appropriate way to dispense these to the class. Special care should be taken with the
magnesium ribbon and magnesium powder and, with some classes, teachers may prefer to
dispense these materials directly.
The length of time required for carrying out the actual reactions is around 30 minutes, but this
will depend on the nature of the class and how the practical is organised.
Apparatus and chemicals
Eye protection
Each group of students will need:
Polystyrene cup (expanded polystyrene)
Beaker (250 cm3) in which to stand the polystyrene cup for support (see note 1)
Thermometer (–10°C to 110°C)
Measuring cylinder (10 cm3), 2
Spatula
Absorbent paper
Access to the following solutions:
(all at approx 0.4 mol dm–3 concentration); (see note 2)
Copper(II) sulfate (Low hazard)
Hydrochloric acid (Low hazard)
Sodium hydrogencarbonate (Low hazard)
Sodium hydroxide (Irritant)
Sulfuric acid (Low hazard)
Access to the following solids (see note 3):
Magnesium ribbon (Highly flammable), cut into 3 cm lengths.
Magnesium powder (Highly flammable).
Citric acid (Irritant).
Technical notes
Copper(II) sulfate (Low hazard) Refer to CLEAPSS Hazcard 27C and CLEAPSS Recipe Card
19
Hydrochloric acid (Low hazard) Refer to CLEAPSS Hazcard 47A and CLEAPSS Recipe Card
31
Sodium hydrogencarbonate (Low hazard) Refer to CLEAPSS Hazcard 95C and CLEAPSS
Recipe Card 64
Sodium hydroxide (Irritant) Refer to CLEAPSS Hazcard 91 and CLEAPSS Recipe Card 65
Sulfuric acid (Low hazard) Refer to CLEAPSS Hazcard 98A and CLEAPSS Recipe Card 69
Magnesium ribbon (Highly flammable) Refer to CLEAPSS Hazcard 59A
Magnesium powder (Highly flammable) Refer to CLEAPSS Hazcard 59A
Citric acid (Irritant) Refer to CLEAPSS Hazcard 36C
1 Typical expanded polystyrene cups fit snugly into 250 cm3 squat form beakers. This provides a
more stable reaction vessel and also prevents spillage if the polystyrene cup splits.
2 At the suggested concentrations, the solutions (except for sodium hydroxide) represent
minimal hazards, although it is probably advisable to label them as Harmful. If the
concentrations are increased then the solutions must be labelled with the correct hazard warning.
The solutions could be provided in small (100 cm3) labelled conical flasks or beakers.
3 Small amounts of the solids can be provided in plastic weighing boats or similar. The teacher
may prefer to keep the magnesium ribbon and powder under their immediate control and to
dispense on an individual basis.
Procedure
SAFETY: Wear eye protection throughout.
Reaction of sodium hydroxide solution and dilute hydrochloric acid
a Stand the polystyrene cup in the beaker.
b Use the measuring cylinder to measure out 10 cm3 of sodium hydroxide solution and pour it
into the polystyrene cup.
c Measure the initial temperature of the sodium hydroxide solution and record it in a suitable
table.
d Measure out 10 cm3 of hydrochloric acid and carefully add this to the sodium hydroxide
solution in the polystyrene cup. Stir with the thermometer and record the maximum or minimum
temperature reached.
e Work out the temperature change and decide if the reaction is exothermic or endothermic.
f Discard the mixture (in the sink with plenty of water). Rinse out and dry the polystyrene cup.
Reaction of sodium hydrogencarbonate solution and citric acid
a Repeat steps a – c of the previous experiment, using sodium hydrogencarbonate solution in
place of sodium hydroxide solution.
b Add 4 small (not heaped) spatula measures of citric acid. Stir with the thermometer and record
the maximum or minimum temperature reached.
c Work out the temperature change and decide if the reaction is exothermic or endothermic.
d Discard the mixture (in the sink with plenty of water). Rinse out and dry the polystyrene cup.
Reaction of copper(II) sulfate solution and magnesium powder
a Repeat steps a – c of the first experiment, using copper(II) sulfate solution in place of sodium
hydroxide solution.
b Add 1 small (not heaped) spatula measure of magnesium powder. Stir with the thermometer
and record the maximum or minimum temperature reached.
c Work out the temperature change and decide if the reaction is exothermic or endothermic.
d Discard the mixture (in the sink with plenty of water). Rinse out and dry the polystyrene cup.
Reaction of sulfuric acid and magnesium ribbon
a Repeat steps a – c of the first experiment, using sulfuric acid in place of sodium hydroxide
solution.
b Add one 3 cm piece of magnesium ribbon. Stir with the thermometer and record the maximum
or minimum temperature reached.
c Work out the temperature change and decide if the reaction is exothermic or endothermic.
d Once all the magnesium ribbon has reacted, discard the mixture (in the sink with plenty of
water). Rinse out and dry the polystyrene cup.
Teaching notes
The reactions and types of reaction involved are:
Sodium hydroxide + hydrochloric acid → sodium chloride + water (Neutralisation)
NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)
Copper(II) sulfate + magnesium → magnesium sulfate + copper (Displacement, Redox)
CuSO4(aq) + Mg(s) → MgSO4(aq) + Cu(s)
Sulfuric acid + magnesium → magnesium sulfate + hydrogen (Displacement, Redox)
H2SO4(aq) + Mg(s) → MgSO4(aq) + H2(g)
At this level the neutralisation reaction between sodium hydrogen carbonate and citric acid may
be a bit complicated – it may be better to just use the word equation. More able students could
use H+(aq) to represent the acid.
Sodium hydrogencarbonate + citric acid → sodium citrate + water + carbon dioxide
NaHCO3(aq) + H+(aq) → Na+(aq) + H2O(l) + CO2(g)
Health & Safety checked, June 2007