aluminum and hydrochloric acid experiment

Foil + Hydrochloric Acid = Awesome Chemistry Experiment

Introduction: Foil + Hydrochloric Acid = Awesome Chemistry Experiment

WARNING: This experiment involves hazardous materials and fire. DANGER: Possibility of injury!

Please don't try this at home. I recommend suggesting the activity to a chemistry teacher or someone who will have the proper safety tools at their disposal. I am not responsible for any injury you may sustain during the recreation of this experiment. You assume all risk and liability. The chemical reaction is exothermic and generates a significant amount of heat along with an aqueous aluminum chloride solution.

The reaction of Aluminum and HCl: 2Al + 6HCl ----> 2AlCl3 + 3H2

Now let's have some fun!

Safety Goggles

Aluminum Foil (aka Tin Foil)

250 mL Flask

9 inch balloon

Binder Clip

Matches (now you know this is going to be awesome!)

You may want to have a fire extinguisher nearby just in case.

Step 1: Foil + Hydrochloric Acid Chemistry Experiment

This experiment is perfect for the mad scientists out there because of the intense reaction. It is also perfect for science teachers because the materials are relatively cheap and easy to come by (always a plus on an educators' budget), and the reaction that takes place is quite simple for students to understand.

Educational objectives can be modified for different ages and ability levels. NOT recommended for students younger than 8th grade.

Student will demonstrate ability to:

• create a replacement reaction
• describe what happens during a chemical change and four possible clues that it has taken place
• apply the law of conservation of mass to chemical reactions
• relate pure chemistry to applied chemistry
• describe the relationships among the temperature, pressure and volume of a gas
• write a word equation
• write a balanced chemical equation
• Extension: identify areas affected by chemistry research

Step 2: Procedure

• Read all steps of the procedure before performing the experiment.
• Place safety goggles on your face.
• Add 100 mL of Hydrochloric Acid to the flask
• Measure and cut a 6 inch by 6 inch square of aluminum foil. (If you use more than 6x6, you risk the balloon bursting and spraying droplets of very hot aqueous aluminum chloride on everything within a 10 foot radius :(
• Get ready for your adrenaline to start pumping.
• Before proceeding with the experiment, read the chemical formula below to help understand the chemistry behind the reaction.The reaction of Aluminum and HCl: 2Al + 6HCl --> 2AlCl3 + 3H2
• Ball up your foil
• Drop the foil into the flask. You will have approximately 30 seconds before the reaction gets going.
• Quickly stretch the balloon over the mouth of the flask. Don't waste a single second when fixing the balloon over the mouth of the flask. The balloon must be placed over the mouth of the flask immediately after the aluminum and HCl are combined.
• Observe the reaction from a safe distance. When the reaction really gets going, bubbles will form. These bubbles will climb higher and higher inside the flask. It is not uncommon for some of the bubbles to enter into the balloon. If you use more than a 6x6 sheet of foil, chances are very good the experiment will end with less than desirable results. These bubbles are very hot and the flask will become too hot to touch.
• When the bubbles settle down and the reaction is finished, twist the balloon 3 or 4 times to trap the hydrogen gas inside.
• Pinch the area you twisted and remove the mouth of the balloon from the flask.
• Tie off the end of the balloon.

Step 3: Experiment Part 2: Flammability of Hydrogen Gas

Now we want to double check to make sure we captured hydrogen. Here's how I tested it. You can use your brains to figure out the best way with what you've got.

I tied string to a beam running along the ceiling of my classroom. I had the string hang down near the middle of the room where I made sure there was no danger of anything catching fire.

Next, I tied a large binder clip to the end of the string (more mass).

I attached the hydrogen-filled balloon to the binder clip.

I used duct tape to combine two meter sticks. A small binder clip was attached to the end of one of the meter sticks with a small strip of duct tape. Two matches were placed into the clamp of the small binder clip. A third match was used to ignite the matches in the grasp of the small binder clip.

Hold the matches near the bottom of the balloon, but not directly at the bottom. I have found that water vapor (or some other liquid) ends up in the balloon with the heat of the reaction. If you hold the match at the lowest point of the balloon, there is a good chance some of the liquid has pooled there. You won't get the big BOOM the students are expecting. Instead, the hydrogen escapes in a narrow stream and takes off like a rocket on a string Although it's still pretty cool to watch, the BOOM is better. Students will feel the blast in the form of a mini shock wave (another teachable moment).

Finally, make sure the balloon isn't on fire. I have a sink near the test site where I can place the balloon into a water bath. It might also be a good idea to place a bucket of water beneath the test site so that if the balloon does catch you could simply cut the string and allow it to fall into the water.

If you turn out the lights and get a slow motion video it looks pretty awesome.

Participated in the Tinfoil Speed Challenge

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How hydrochloric acid reacts with aluminum. Formulas and description of the process

Features of hydrochloric acid and aluminum interaction

Alu­minum is a mal­leable, light, sil­very-white met­al. It is a good elec­tri­cal con­duc­tor. It can re­act with both acids and bases. Com­bin­ing alu­minum with an acid re­sults in a typ­i­cal sin­gle dis­place­ment re­ac­tion, form­ing alu­minum salt and gaseous hy­dro­gen. This can be seen from a sim­ple ex­am­ple – how hy­drochlo­ric acid re­acts with alu­minum.

With al­ka­lis, the re­ac­tion pro­ceeds dif­fer­ent­ly: in ad­di­tion to a hy­dro­gen re­lease, the re­ac­tion forms MeAlO₂ alu­mi­nate (where Me is the cation of a met­al from the al­ka­li) and a com­plex com­pound with the for­mu­la Me[Al(OH)₄] in the so­lu­tion.

How alu­minum re­acts with hy­drochlo­ric acid

Alu­minum re­acts with di­lut­ed hy­drochlo­ric acid at room tem­per­a­ture. The met­al dis­solves in hy­drochlo­ric acid, yield­ing alu­minum chlo­ride and col­or­less hy­dro­gen gas. This re­ac­tion is ir­re­versible, as the fi­nal prod­ucts will not re­act with each oth­er. The re­ac­tion be­tween metal­lic alu­minum and hy­drochlo­ric acid is what is known as an ox­i­da­tion-re­duc­tion re­ac­tion. Alu­minum acts as the re­duc­ing agent, giv­ing up elec­trons:

Al⁰ - 3e = Al³⁺

Cations of hy­drochlo­ric acid take these elec­trons and are re­duced to molec­u­lar hy­dro­gen:

2H⁺ + 2e = H₂

The com­plete ion­ic re­ac­tion equa­tion reads:

2Al⁰ + 6H⁺ + 6Cl⁻ = 2Al³⁺ + 6Cl⁻ + 3H₂↑

Net-ion­ic form:

2Al⁰ + 6H⁺ = 2Al³⁺ + 3H₂↑

In molec­u­lar form, the re­ac­tion looks as fol­lows:

2Al + 6HCl = 2Al­Cl₃ + 3H₂↑

Metal­lic alu­minum is not the only sub­stance ca­pa­ble of re­act­ing with hy­drochlo­ric acid – many met­al com­pounds pos­sess this prop­er­ty. An ex­change re­ac­tion oc­curs with the salts, when ions or re­ac­tive groups of both reagents “change places.” In or­der for the re­ac­tion with alu­minum or its com­pounds to be ir­re­versible, the reagents must form a gas, a pre­cip­i­tate, or a poor­ly-sol­u­ble sub­stance. The nec­es­sary reagent quan­ti­ties must be cal­cu­lat­ed with pre­ci­sion.

Re­ac­tions of alu­minum hy­drox­ides and ox­ides with hy­drochlo­ric acid

Al(OH)₃ is an am­pho­ter­ic base, which is a white jel­ly-like pre­cip­i­tate that dis­solves poor­ly in wa­ter.

Alu­minum hy­drox­ide en­ters into a neu­tral­iza­tion re­ac­tion with hy­drochlo­ric acid (the hy­drox­ide must be fresh­ly-pre­cip­i­tat­ed for the re­ac­tion to pro­ceed re­li­ably):

Al(OH)₃ + 3HCl = Al­Cl₃ + 3H₂O

One can ob­serve the dis­so­lu­tion of the white pre­cip­i­tate of alu­minum hy­drox­ide (alu­minum chlo­ride Al­Cl₃ dis­solves well in wa­ter). With alu­minum ox­ide, the re­ac­tion yields salt and wa­ter ac­cord­ing to the fol­low­ing equa­tion:

Al₂O₃ + 6HCl = 2Al­Cl₃ + 3H₂O

Re­ac­tions of salts, hy­drides, and alu­minum com­plex­es with hy­drochlo­ric acid

Hy­drochlo­ric acid also re­acts with many oth­er alu­minum com­pounds.

With alu­minum car­bide

Al₄C₃ + 12H­Cl = 4Al­Cl₃ + 3CH₄↑

(alu­minum car­bide dis­solves when treat­ed with an ex­cess of hy­drochlo­ric acid)

With alu­minum ac­etate

(CH₃­COO)₃Al + 3HCl = Al­Cl₃ + 3CH₃­COOH

With alu­minum ni­tride

AlN + 4HCl = Al­Cl₃ + NH₄­Cl

(hot con­cen­trat­ed acid is used; the re­ac­tion takes place slow­ly)

With alu­minum sul­fide

Al₂S₃ + 6HCl = 2Al­Cl₃ + 3H₂S↑

With alu­minum phos­phide

AlP + 3HCl = Al­Cl₃ + PH₃↑

(The re­ac­tion pre­sup­pos­es the treat­ment of phos­phide with hot con­cen­trat­ed acid)

With alu­minum phos­phate

AlPO₄ + 3HCl = Al­Cl₃ + H₃PO₄

With lithi­um alanate (tetrahy­droa­lu­mi­nate)

Li[AlH₄] + 4HCl = Al­Cl₃ + LiCl + 4H₂↑

(The re­ac­tion is car­ried out at a low tem­per­a­ture)

With sodi­um alu­mi­nate

NaAlO₂ + 4HCl = NaCl + Al­Cl₃ + 2H₂O

With sodi­um tetrahy­drox­oa­lu­mi­nate

Na[Al(OH)₄] + 4HCl = Al­Cl₃ + NaCl + 4H₂O

Alu­minum sul­fates and ni­trates do not re­act with hy­drochlo­ric acid, as all the com­pounds in the mix­ture are sol­u­ble – no pre­cip­i­tate forms, no poor­ly-sol­u­ble sub­stances form, and gas is not re­leased.

Click here to learn more about alu­minum and its prop­er­ties.

How mix­tures of met­als re­act with hy­drochlo­ric acid

If you take a mix­ture of sev­er­al met­als and treat them with hy­drochlo­ric acid, each met­al will re­act sep­a­rate­ly. For ex­am­ple, if you add HCl to a mix­ture of alu­minum and iron shav­ings, the re­ac­tion will pro­ceed as fol­lows:

Fe + 2HCl = Fe­Cl₂ + H₂↑

As di­lut­ed hy­drochlo­ric acid is a weak ox­i­diz­er, iron is only re­duced to a +2 ox­i­da­tion state.

Prod­ucts of the re­ac­tion be­tween alu­minum and hy­drochlo­ric acid, and their ap­pli­ca­tion

Al­most all re­ac­tions of hy­drochlo­ric acid and alu­minum (or its com­pounds) re­sult in the for­ma­tion of alu­minum chlo­ride (Al­Cl₃). The salt dis­solves well in or­gan­ic sol­vents (ni­troben­zene, dichloroethane, ace­tone) and wa­ter. In aque­ous so­lu­tions, one can ob­serve the hy­drol­y­sis of Al­Cl₃, as this salt is formed by the strong acid HCl and the weak base Al(OH)₃.

Al­Cl₃ is used as a cat­a­lyst in or­gan­ic syn­the­sis. For in­stance, it is uti­lized in the iso­mer­iza­tion of paraf­fins, ini­ti­a­tion of alky­la­tion re­ac­tions, acy­la­tion, and the break­down of oil into frac­tions. Alu­mini­um chlo­ride hex­ahy­drate Al­Cl₃・6H₂O is used to treat tim­ber ma­te­ri­als, pu­ri­fy waste­water, and man­u­fac­ture an­tiper­spi­rants.

The re­ac­tion of alu­minum with a so­lu­tion of hy­drochlo­ric acid can be used as a lab­o­ra­to­ry method for ob­tain­ing hy­dro­gen (but metal­lic zinc is more com­mon­ly used for these pur­pos­es).

Dozens of experiments you can do at home

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Positive Reactions

• Apr 16, 2022

Single Displacement Reaction with HCL and Aluminum

Description:

In this experiment, we will show case another very fast reaction where Hydrochloric acid reacts with Aluminum foils creating Aluminum Chloride and releasing hydrogen gas.

Difficulty: Do it outdoors with adult supervision, handle the acid caution. Gloves and goggles are mandatory like all experiments. Easy experiment to conduct

Hydrochloric acid(HCL) diluted at 31% (Muriatic acid) : 50 ml

Aluminum foil: 1 gram

An Erlenmeyer flask

Add 50 ml of Hydrochloric acid to the Erlenmeyer flask very carefully

Tearing the Aluminum foils into smaller pieces and make small balls

Drop the Aluminum foil balls into the flask

Stand back and watch the reaction unfold. Don't try to touch the flask, it will be very hot from the exothermic reaction

How It Works:

Hydrochloric acid reacts with Aluminum producing Aluminum Chloride and releasing hydrogen gas. The reaction is as follows

This is also an example of a displacement reaction. Aluminum displaces the hydrogen in hydrochloric acid creating Aluminum chloride.

Aluminum chloride is a widely used commercial product, it is a key ingredient in anti perspirants. The reaction described above is one of the commercial ways of making Aluminum Chloride

This reaction is highly exothermic and the water is converted into steam that is released along with hydrogen.

The reaction takes a few seconds to start because Aluminum foil has a small coating of Aluminum oxide on it and it takes a few seconds for the Aluminum oxide to dissolve in the acid to expose the Aluminum metal

Here is a video of this reaction

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What Happens when Aluminium Reacts with Hydrochloric acid

Have you ever wondered what happens when aluminum comes into contact with hydrochloric acid? In this blog post, we will explore the chemical reaction between aluminum and hydrochloric acid in detail. We’ll look at why this reaction occurs, what it looks like, and how it can be used in various applications

At normal temperatures, aluminum reacts with dilute hydrochloric acid. Aluminium chloride and colourless hydrogen gas are produced when the metal dissolves in hydrochloric acid . Because the final products do not react with each other, this reaction is irreversible.

When aluminum comes into contact with hydrochloric acid, a single displacement reaction takes place. This means that the more reactive element—in this case, hydrogen—is displaced from the less reactive element—hydrochloric acid. This happens because hydrogen is more electronegative than aluminum so it will take its place in the compound.

As a result of this reaction, aluminum chloride (AlCl3) and hydrogen gas (H2) are formed.

The Appearance of the Reaction

When aluminum is exposed to hydrochloric acid, a fizzing noise can be heard as bubbles of hydrogen gas are released. This signifies that the reaction has taken place and is happening at a rapid rate. Additionally, there may be some white-colored smoke or fumes coming from the mixture due to the formation of AlCl3 particles.

Uses of the Reaction

The reaction between aluminum and hydrochloric acid can be used to produce various products, such as alumina (Al2O3), which is used as an abrasive material for polishing surfaces. Additionally, it can also be used as an ingredient in certain commercial cleaning products such as oven cleaners and bathroom cleaners due to its ability to dissolve grease quickly. Furthermore, AlCl3 can also be used in water treatment systems as a coagulant agent for removing suspended solids from water supplies.

The reaction between aluminum and hydrochloric acid is an important one that results in several important products being produced. It is not only useful for producing alumina or AlCl3 but also for producing commercial cleaning products such as oven cleaners and bathroom cleaners. Finally, it can also be used in water treatment systems as a coagulant agent for removing suspended solids from water supplies. Understanding this reaction is essential for anyone looking to use these materials or products effectively in their day-to-day life or business operations!

Abhishek is a seasoned blogger and industry expert, sharing his insights and knowledge on various topics. With his research, Abhishek offers valuable insights and tips for professionals and enthusiasts. Follow him for expert advice on the latest trends and developments in the metal industry.

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Reactivity Of Aluminum And Hydrochloric Acid

Aluminum is a reactive metal and different grades of aluminum react differently with chemicals especially acids. Certain acid types are harmless to some aluminum grades, while others are harmful.

Acidic solutions can sometimes remove other things from aluminum machine parts without harming the metal, depending on the aluminum grade and acid type.

What Happens When Hydrochloric Acid Is Poured On Aluminum Foil?

Aluminum is a light silvery white metal that effectively reacts with hydrochloric acid at room temperature. The reaction occurs at normal temperature and produces chlorides.

When aluminum reacts with hydrochloric acid, aluminum chloride and hydrogen gas are produced. The reaction is irreversible as the products cannot be converted into reactants again. The balanced equation of this reaction is as follows:

2Al + 6HCl → 2AlCl 3 + 3H 2 ↑

What Type Of Reaction Occurs Between Hydrochloric Acid And Aluminum?

A typical single displacement reaction takes place when hydrochloric acid combines with aluminum foil at room temperature. The metal dissolves in hydrochloric acid, yielding aluminum chloride and colorless hydrogen gas.

Dilute acids react with relatively reactive metals such as magnesium, aluminum, and iron. In general, the more reactive the metal, the faster the reaction. In case of aluminum, it reacts slowly with acids because it has a protective oxide layer.

What Is The Mechanism Of Reaction Between Aluminum And Hydrochloric Acid?

An oxidation-reduction process, also known as a redox reaction, occurs when metallic aluminum reacts with hydrochloric acid. Both oxidation and reduction occur concurrently.

Reactions Mechanism

The reaction requires a ratio of two aluminum molecules to six hydrochloric acid molecules because aluminum has three electrons in its outer shell. The reaction starts when each chlorine atom in hydrochloric acid gains an electron from aluminum while losing a hydrogen atom.

Because hydrogen does not like to exist as a single atom, it takes six molecules of hydrochloric acid to make three molecules of the diatomic hydrogen molecule, therefore two molecules of aluminum are needed to keep things balanced.

Listed below are the stepwise reaction taking place in between aluminum and hydrochloric acid

Step 1: Aluminum acts as the reducing agent, giving up electrons:

Al⁰ – 3e → Al³⁺

Step 2: Cations of hydrochloric acid take these electrons and are reduced to molecular hydrogen:

2H⁺ + 2e → H₂↑

The complete ionic reaction equation is as follows:

2Al⁰ + 6H⁺ + 6Cl⁻ → 2Al³⁺ + 6Cl⁻ + 3H₂↑

Net-ionic form:

2Al⁰ + 6H⁺ → 2Al³⁺ + 3H₂↑

In molecular form, the reaction looks as follows:

2Al + 6HCl = 2Al­Cl₃ + 3H₂↑

Metallic aluminum isn’t the only substance that can react with hydrochloric acid. Many metal compounds have this property. When ions or reactive groups of both reagents “change places,” this is called an exchange reaction.

The reagents must generate a gas, a precipitate, or a poorly soluble substance in order for the reaction with aluminum or its compounds to be irreversible. Precision is required for calculating the required reagent quantities.

What Are The Applications Of Products Formed By The Reaction Of Aluminum And Hydrochloric Acid?

Almost all the products formed as a result of chemical reactions are useful, same is the case with the products formed when aluminum and hydrochloric acid react with each other.

• Aluminum chloride (AlCl 3 ) or its compounds are formed in almost all reactions between hydrochloric acid and aluminum. Organic solvents (nitrobenzene, dichloroethane, acetone) and water readily dissolve the salt. The hydrolysis of AlCl 3 can be observed in aqueous solutions, as this salt is produced by the strong acid HCl and the weak base Al(OH) 3 .
• In organic synthesis, AlCl 3 is utilized as a catalyst. It is used in the isomerization of paraffins, the initiation of alkylation reactions, acylation, and the fraction of oil.
• The aluminum chloride hexahydrate AlCl 3 6H 2 O is used to cure timber materials, purify wastewater, and manufacture antiperspirants.
• The reaction of aluminum with a solution of hydrochloric acid can be utilized as a laboratory method for obtaining hydrogen (although metallic zinc is more commonly used for these purposes).

How To Protect Aluminum From Acid Damage?

Aluminum parts in motors, drives, and gears have been reported to be damaged by hydrochloric and sulfuric acids. Dilution can help to lessen the effects of hydrochloric acid and prevent these parts from acidic damage.

If kept at ambient temperature, very dilute sulfuric acid solutions will not harm aluminum parts. Boric, carbonic, lactic, and nitric acids do not commonly cause aluminum to corrode. Depending on the concentration of the acidic solution and the temperature, chromic acids inflict significant damage.

Related Questions

Why does aluminum not react with hydrochloric acid?

Aluminum does react with hydrochloric acid. When aluminum is poured into an acid, it may appear that it does not react at first. This is because, as a result of the preceding interaction with the air, a coating of aluminum oxide forms on the surface of the aluminum, which acts as a protective barrier. As aluminum chloride forms, the hydrochloric acid quickly turns a dismal grey color.

Does acid affect aluminum?

Aluminum is a flexible, light, silvery-white metal with a high melting point. It’s a good conductor of electricity. It has the ability to react with both acids and bases. When aluminum is combined with an acid, a classic single displacement reaction occurs, resulting in aluminum salt and gaseous hydrogen.

What is the aluminum foil’s coating?

Aluminum foil is usually covered with lacquer or water-based latex or bonded to a polymer film with adhesive or extrusion coating or lamination. Many lidding applications use vinyl-based lacquer-coated aluminum foils , particularly for dairy goods like yogurt.

Aluminum being a metal is reactive towards acids, but due to the presence of protective oxide layer, it is somewhat resistant to acids. However, the above-mentioned reactions show the formation of aluminum chloride which is a useful chemical substance.

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Acid-base Behavior of the Oxides

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• Page ID 3665

• Truro School in Cornwall

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This page discusses the reactions of the oxides of Period 3 elements (sodium to chlorine) with water, and with acids or bases where relevant (as before, argon is omitted because it does not form an oxide).

A quick summary of the trend

The oxides: The oxides of interest are given below:

The trend in acid-base behavior can be summarized as follows:

Acidity increases from left to right, ranging from strongly basic oxides on the left to strongly acidic ones on the right, with an amphoteric oxide (aluminum oxide) in the middle. An amphoteric oxide is one which shows both acidic and basic properties.

This trend applies only to the highest oxides of the individual elements (see the top row of the table), in the highest oxidation states for those elements. The pattern is less clear for other oxides. Non-metal oxide acidity is defined in terms of the acidic solutions formed in reactions with water—for example, sulfur trioxide reacts with water to forms sulfuric acid. They will all, however, react with bases such as sodium hydroxide to form salts such as sodium sulfate as explored in detail below.

Sodium Oxide

Sodium oxide is a simple strongly basic oxide. It is basic because it contains the oxide ion, O 2- , which is a very strong base with a high tendency to combine with hydrogen ions.

Reaction with water : Sodium oxide reacts exothermically with cold water to produce sodium hydroxide solution. A concentrated solution of sodium oxide in water will have pH 14.

\[ Na_2O + H_2O \rightarrow 2NaOH \nonumber \]

Reaction with acids: As a strong base, sodium oxide also reacts with acids. For example, it reacts with dilute hydrochloric acid to produce sodium chloride solution.

\[Na_2O + 2HCl \rightarrow 2NaCl + H_2O \nonumber \]

Magnesium oxide

Magnesium oxide is another simple basic oxide, which also contains oxide ions. However, it is not as strongly basic as sodium oxide because the oxide ions are not as weakly-bound. In the sodium oxide, the solid is held together by attractions between 1+ and 2- ions. In magnesium oxide, the attractions are between 2+ and 2- ions. Because of the higher charge on the metal, more energy is required to break this association. Even considering other factors (such as the energy released from ion-dipole interactions between the cations and water), the net effect is that reactions involving magnesium oxide will always be less exothermic than those of sodium oxide.

Reaction with water: At first glance, magnesium oxide powder does not appear to react with water. However, the pH of the resulting solution is about 9, indicating that hydroxide ions have been produced. In fact, s ome magnesium hydroxide is formed in the reaction, but as the species is almost insoluble, few hydroxide ions actually dissolve. The reaction is shown below:

\[MgO + H_2O \rightarrow Mg(OH)_2 \nonumber \]

Reaction with acids: Magnesium oxide reacts with acids as predicted for a simple metal oxide. For example, it reacts with warm dilute hydrochloric acid to give magnesium chloride solution.

\[MgO + 2HCl \rightarrow MgCl_2+H_2O \nonumber \]

Aluminum Oxide

Describing the properties of aluminum oxide can be confusing because it exists in a number of different forms. One of those forms is very unreactive (known chemically as alpha-Al 2 O 3 ) and is produced at high temperatures. The following reactions concern the more reactive forms of the molecule. Aluminium oxide is amphoteric . It has reactions as both a base and an acid.

Reaction with water: Aluminum oxide is insoluble in water and does not react like sodium oxide and magnesium oxide. The oxide ions are held too strongly in the solid lattice to react with the water.

Reaction with acids: Aluminum oxide contains oxide ions, and thus reacts with acids in the same way sodium or magnesium oxides do. Aluminum oxide reacts with hot dilute hydrochloric acid to give aluminum chloride solution.

\[Al_2O_3 + 6HCl \rightarrow 2AlCl_3 + 3H_2O \nonumber \]

This reaction and others display the amphoteric nature of aluminum oxide.

Reaction with bases: Aluminum oxide also displays acidic properties, as shown in its reactions with bases such as sodium hydroxide. Various aluminates (compounds in which the aluminum is a component in a negative ion) exist, which is possible because aluminum can form covalent bonds with oxygen. This is possible because the electronegativity difference between aluminum and oxygen is small, unlike the difference between sodium and oxygen, for example (electronegativity increases across a period)

Aluminum oxide reacts with hot, concentrated sodium hydroxide solution to produce a colorless solution of sodium tetrahydroxoaluminate:

\[Al_2O_3 + 2NaOH +3H_2O \rightarrow 2NaAl(OH)_4 \nonumber \]

Silicon dioxide (silicon(IV) oxide)

Silicon is too similar in electronegativity to oxygen to form ionic bonds. Therefore, because silicon dioxide does not contain oxide ions, it has no basic properties. In fact, it is very weakly acidic, reacting with strong bases.

Reaction with water: Silicon dioxide does not react with water, due to the thermodynamic difficulty of breaking up its network covalent structure.

Reaction with bases : Silicon dioxide reacts with hot, concentrated sodium hydroxide solution, forming a colorless solution of sodium silicate:

\[SiO_2 + 2NaOH \rightarrow Na_2SiO_3 + H2O \nonumber \]

In another example of acidic silicon dioxide reacting with a base, the Blast Furnace extraction of iron, calcium oxide from limestone reacts with silicon dioxide to produce a liquid slag, calcium silicate:

\[SiO_2 + CaO \rightarrow CaSiO_3 \nonumber \]

Phosphorus Oxides

Two phosphorus oxides, phosphorus(III) oxide, P 4 O 6 , and phosphorus(V) oxide, P 4 O 10 , are considered here.

Phosphorus(III) oxide: Phosphorus(III) oxide reacts with cold water to produce a solution of the weak acid, H 3 PO 3 —known as phosphorous acid, orthophosphorous acid or phosphonic acid:

\[P_4O_6 + 6H_2O \rightarrow 4H_3PO_3 \nonumber \]

The fully-protonated acid structure is shown below:

The protons remain associated until water is added; even then, because phosphorous acid is a weak acid, few acid molecules are deprotonated. Phosphorous acid has a pK a of 2.00, which is more acidic than common organic acids like ethanoic acid (pK a = 4.76).

Phosphorus(III) oxide is unlikely to be reacted directly with a base. In phosphorous acid, the two hydrogen atoms in the -OH groups are acidic, but the third hydrogen atom is not. Therefore, there are two possible reactions with a base like sodium hydroxide, depending on the amount of base added:

\[ NaOH + H_3PO_3 \rightarrow NaH_2PO_3 + H_2O \nonumber \]

\[ 2NaOH + H_3PO_3 \rightarrow Na_2HPO_3 + 2H_2O \nonumber \]

In the first reaction, only one of the protons reacts with the hydroxide ions from the base. In the second case (using twice as much sodium hydroxide), both protons react.

If instead phosphorus(III) oxide is reacted directly with sodium hydroxide solution, the same salts are possible:

\[4NaOH + P_4O_6 + 2H_2O \rightarrow 4NaH_2PO_3 \nonumber \]

\[9NaOH + P_4O_6 \rightarrow 4Na_2HPO_3 + 2H_2O \nonumber \]

Phosphorus(V) oxide: Phosphorus(V) oxide reacts violently with water to give a solution containing a mixture of acids, the nature of which depends on the reaction conditions. Only one acid is commonly considered, phosphoric(V) acid, H 3 PO 4 (also known as phosphoric acid or as orthophosphoric acid).

\[P_4O_{10} + 6H_2O \rightarrow 4H_3PO_4 \nonumber \]

This time the fully protonated acid has the following structure:

Phosphoric(V) acid is another weak acid with a pK a of 2.15, marginally weaker than phosphorous acid. Solutions of each of these acids with concentrations around 1 mol dm -3 have a pH of about 1.

Phosphoric (V) oxide is also unlikely to be reacted directly with a base, but the hypothetical reactions are considered. In its acid form, molecule has three acidic -OH groups, which can cause a three-stage reaction with sodium hydroxide:

\[ NaOH + H_3PO_4 \rightarrow NaH_2PO_4 + H_2O \nonumber \]

\[ 2NaOH + H_3PO_4 \rightarrow Na_2HPO_4 + 2H_2O \nonumber \]

\[ 3NaOH + H_3PO_4 \rightarrow Na_3PO_4 + 3H_2O \nonumber \]

Similar to phosphorus (III) oxide, if phosphorus(V) oxide reacts directly with sodium hydroxide solution, the same possible salt as in the third step (and only this salt) is formed:

\[12NaOH + P_4O_{10} \rightarrow 4Na_3PO_4 + 6H_2O \nonumber \]

Sulfur Oxides

Two oxides are considered: sulfur dioxide, SO 2 , and sulfur trioxide, SO 3 .

Sulfur dioxide: Sulfur dioxide is fairly soluble in water, reacting to give a solution of sulfurous acid (also known as sulfuric(IV) acid), H 2 SO 3 , as shown in the reaction below. This species only exists in solution, and any attempt to isolate it gives off sulfur dioxide.

\[ SO_2 + H_2O \rightarrow H_2SO_3 \nonumber \]

The protonated acid has the following structure:

Sulfurous acid is also a relatively weak acid, with a pK a of around 1.8, but slightly stronger than the two phosphorus-containing acids above. A reasonably concentrated solution of sulfurous acid has a pH of about 1.

Sulfur dioxide also reacts directly with bases such as sodium hydroxide solution. Bubbling sulfur dioxide through sodium hydroxide solution first forms sodium sulfite solution, followed by sodium hydrogen sulfite solution if the sulfur dioxide is in excess.

\[ SO_2 + 2NaOH \rightarrow Na_2SO_3 + H_2O \nonumber \]

\[Na_2SO_3 + H_2O \rightarrow 2NaHSO_3 \nonumber \]

Another important reaction of sulfur dioxide is with the base calcium oxide to form calcium sulfite (also known as calcium sulfate(IV)). This is of the important methods of removing sulfur dioxide from flue gases in power stations.

\[CaO + SO_2 \rightarrow CaSO_3 \nonumber \]

Sulfur trioxide: Sulfur trioxide reacts violently with water to produce a fog of concentrated sulfuric acid droplets.

\[ SO_3 + H_2O \rightarrow H_2SO_4 \nonumber \]

Pure, fully-protonated sulfuric acid has the structure:

Sulfuric acid is a strong acid, and solutions will typically have a pH around 0. The acid reacts with water to give a hydronium ion (a hydrogen ion in solution) and a hydrogen sulfate ion. This reaction runs essentially to completion:

\[ H_2SO_4 (aq) + H_2O (l) \rightarrow H_3P^+ + HSO_4^- (aq) \nonumber \]

The second proton is more difficult to remove. In fact, the hydrogen sulfate ion is a relatively weak acid, similar in strength to the acids discussed above. This reaction is more appropriately described as an equilibrium:

\[ HSO_4^- (aq) + H_2O \rightleftharpoons H_3O^+ (aq) + SO_4^{2-} (aq) \nonumber \]

It is useful if you understand the reason that sulfuric acid is a stronger acid than sulfurous acid. You can apply the same reasoning to other acids that you find on this page as well.

Sulfuric acid is stronger than sulfurous acid because when a hydrogen ion is lost from one of the -OH groups on sulfuric acid, the negative charge left on the oxygen is spread out (delocalized) over the ion by interacting with the doubly-bonded oxygen atoms. It follows that more double bonded oxygen atoms in the ion make more delocalization possible; more delocalization leads to greater stability, making the ion less likely to recombine with a hydrogen ion and revert to the non-ionized acid.

Sulfurous acid only has one double bonded oxygen, whereas sulfuric acid has two; the extra double bond provides much more effective delocalization, a much more stable ion, and a stronger acid. Sulfuric acid displays all the reactions characteristic of a strong acid. For example, a reaction with sodium hydroxide forms sodium sulfate; in this reaction, both of the acidic protons react with hydroxide ions as shown:

\[2NaOH +H_2SO_4 \rightarrow Na_2SO_4 + 2H_2O \nonumber \]

In principle, sodium hydrogen sulfate can be formed by using half as much sodium hydroxide; in this case, only one of the acidic hydrogen atoms is removed.

Sulfur trioxide itself also reacts directly with bases such as calcium oxide, forming calcium sulfate:

\[ CaO + SO_3 \rightarrow CaSO_4 \nonumber \]

This reaction is similar to the reaction with sulfur dioxide discussed above.

Chlorine Oxides

Chlorine forms several oxides, but only two (chlorine(VII) oxide, Cl 2 O 7 , and chlorine(I)oxide, Cl 2 O) are considered here. Chlorine(VII) oxide is also known as dichlorine heptoxide, and chlorine(I) oxide as dichlorine monoxide.

Chlorine(VII) oxide: Chlorine(VII) oxide is the highest oxide of chlorine—​the chlorine atom is in its maximum oxidation state of +7. It continues the trend of the highest oxides of the Period 3 elements towards being stronger acids. Chlorine(VII) oxide reacts with water to give the very strong acid, chloric(VII) acid, also known as perchloric acid.

\[ Cl_2O_7 + H_2O \rightarrow 2HClO_4 \nonumber \]

As in sulfuric acid, the pH of typical solutions of perchloric acid are around 0. Neutral chloric(VII) acid has the following structure:

When the chlorate(VII) ion (perchlorate ion) forms by loss of a proton (in a reaction with water, for example), the charge is delocalized over every oxygen atom in the ion. That makes the ion very stable, making chloric(VII) acid very strong.

Chloric(VII) acid reacts with sodium hydroxide solution to form a solution of sodium chlorate(VII):

\[ NaOH + HClO_4 \rightarrow NaClO_4 + H2O \nonumber \]

Chlorine(VII) oxide itself also reacts directly with sodium hydroxide solution to give the same product:

\[ 2NaOH + Cl_2O_7 \rightarrow 2NaClO_4 + H_2O \nonumber \]

Chlorine(I) oxide: Chlorine(I) oxide is far less acidic than chlorine(VII) oxide. It reacts with water to some extent to give chloric(I) acid, \(HOCl^-\) also known as hypochlorous acid.

\[ Cl_2O + H_2O \rightleftharpoons 2HOCl \nonumber \]

The structure of chloric(I) acid is exactly as shown by its formula, HOCl. It has no doubly-bonded oxygens, and no way of delocalizing the charge over the negative ion formed by loss of the hydrogen. Therefore, the negative ion formed not very stable, and readily reclaims its proton to revert to the acid. Chloric(I) acid is very weak (pK a = 7.43) and reacts with sodium hydroxide solution to give a solution of sodium chlorate(I) (sodium hypochlorite):

\[ NaOH + HOCl \rightarrow NaOCl + H_2O \nonumber \]

Chlorine(I) oxide also reacts directly with sodium hydroxide to give the same product:

\[2NaOH + Cl_2O \rightarrow 2NaOCl + H_2O \nonumber \]

Contributor

Jim Clark ( Chemguide.co.uk )

Practical: Investigate Metals Reacting with Acids ( Edexcel IGCSE Chemistry )

Revision note.

Chemistry Lead

Practical: Investigate Metals Reacting with Acids

To investigate the reactions between dilute hydrochloric and sulfuric acids with the metals magnesium, iron and zinc

Investigating the reactions of dilute acids with metals

• Wear some safety glasses before handling acids
• Using a small measuring cylinder, add 5 cm 3 of dilute hydrochloric acid to each of three test tubes
• Add about 1 cm length of magnesium ribbon to the first tube, observe and note down what you see
• Use a lighted splint to test for any gases given off
• To the second test tube add a few pieces of iron filings and to the third some zinc turnings
• Observe what happens, test for any gases and note down your observations
• Repeat the experiment with dilute sulfuric acid

Metals with Acids Observations Table

Equations for the Reactions

• The metals can be ranked in reactivity order Mg > Zn > Fe
• The three metals react in the same with both acids
• Hydrogen and a metal salt solution is produced

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What happens when aluminum reacts with hydrochloric acid?

Reaction of metal with acid when a metal reacts with an acid, a salt is formed along with the liberation of hydrogen gas. when aluminum al reacts with hydrochloric acid hcl , aluminum chloride alcl 3 is formed along with colorless hydrogen gas h 2 . this reaction is irreversible in nature. chemical equation the chemical equation is depicted below. 2 al ( s ) + 6 hcl ( aq ) → 2 alcl 3 ( aq ) + 3 h 2 ( g ).

What happens when magnesium metal reacts with the hydrochloric acid?

what happens when :

sodium reacts with oxygen

sulphur reacts with oxygen

aluminium reacts with dilute hydrochloric acid

1. What happens when sodium hydroxide reacts with hydrochloric acid ?

2. What happens when Potassium hydroxide reacts with hydrochloric acid ?

what happens when magnesium reacts with hydrochloric acid?

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How can I balance this chemical equations? Aluminum and hydrochloric acid react to form aluminum chloride and hydrogen gas.

Explanation:

This reaction is between a metal and an acid which typically results in a salt and the release of hydrogen gas. The unbalanced reaction is

#Al + HCl -> H_2 + AlCl_3# .

This is a redox reaction , whose half-reactions are and become:

#2("Al"(s) -> "Al"^(3+)(aq) + cancel(3e^(-)))# #3(2"H"^(+)(aq) + cancel(2e^(-)) -> "H"_2(g))# #"-----------------------------------------------"# #2"Al"(s) + 6"H"^(+)(aq) -> 3"H"_2(g) + 2"Al"^(3+)(aq)#

Aluminum oxidizes as #"Al" -> "Al"^(3+)# , while hydrogen reduces as #2"H"^(+) -> "H"_2^0# .

If we add back the spectator #"Cl"^(-)# , we get:

I hope this was helpful.

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The reaction of aluminium and copper(II) sulfate

In association with Nuffield Foundation

Try this class practical or demonstration to illustrate the displacement of copper from copper(II) sulfate using aluminium foil

In this experiment, students add aluminium cooking foil to copper(II) sulfate solution and observe no reaction. They then add and dissolve sodium chloride, producing a vigorous displacement reaction which illustrates the reactivity of aluminium. The solution gets very hot, the aluminium dissolves and red copper becomes visible.

The class practical can take about 30 minutes to complete. A flexicam would work well if this is to be done as a demonstration and allow students a clearer view of what is going on.

• Eye protection (goggles)
• Conical flask, 100 cm 3
• Aluminium foil, 2 cm x 2 cm
• Copper(II) sulfate solution, 0.8 M (HARMFUL), 20 cm 3
• Sodium chloride, 2–3 g

Health, safety and technical notes

• Read our standard health and safety guidance.
• Wear eye protection (goggles) throughout and disposable nitrile gloves.
• Aluminium foil, Al(s) – see CLEAPSS Hazcard HC001A .
• Copper(II) sulfate solution, CuSO 4 (aq), 0.8 M (HARMFUL, DANGEROUS FOR THE ENVIRONMENT) – see CLEAPSS Hazcard  HC027c  and CLEAPSS Recipe Book RB031.
• Sodium chloride, NaCl(s), (table salt) – see CLEAPSS Hazcard HC047b .
• Ensure the aluminium foil is completely consumed by the reaction before disposal to prevent a continued exothermic reaction in the rubbish bin. Use plenty of copper(II) sulfate solution and sodium chloride to ensure a complete reaction.

Source: Royal Society of Chemistry

The equipment required for illustrating the reaction between copper(II) sulfate and aluminium, before sodium chloride is added to disrupt the oxide layer on the aluminium foil

• Measure approximately 20 cm 3 of copper(II) sulfate solution into the conical flask.
• Add a square of aluminium foil.
• Look for signs of a reaction.
• Add a spatula of sodium chloride and stir to dissolve.
• Observe any changes. If nothing happens, add more sodium chloride. Has displacement of copper from copper(II) sulfate occurred?

Student questions and example table

• Before the sodium chloride is added, does any reaction occur?
• After adding sodium chloride, does the aluminium appear more or less reactive?
• How does the addition of sodium chloride affect this change?
• Write ‘yes’ or ‘no’ to fill in the table below.

Observations	Before the sodium chloride is added	After the sodium chloride is added
Bubbles observed
Colour changes
Temperature change
Copper observed

Teaching notes

Aluminium does not show its true reactivity until the oxide layer is disturbed. Sodum chloride disturbs this oxide layer. Scratches on the surface of the oxide layer allow chloride ions to react with aluminium, this effects the cohesiveness of the oxide layer. This allows reaction with the copper(II) sulfate. Remind students what copper looks like, so that they know what they are looking for.

Answers to student questions

• Aluminium appears less reactive than copper. The aluminium foil appears unable to displace copper from copper(II) sulfate solution.
• Now aluminium is more reactive because it displaces copper. Aluminium + copper(II) sulfate → copper + aluminium sulfate
• Scratches on the surface of the oxide layer allow chloride ions to react with aluminium, this effects the cohesiveness of the oxide layer. This allows a simple exchange reaction with the copper(II) sulfate. The protective oxide layer forms  instantly the aluminium is exposed to the air.

Additional information

This is a resource from the  Practical Chemistry project , developed by the Nuffield Foundation and the Royal Society of Chemistry.

Practical Chemistry activities accompany  Practical Physics  and  Practical Biology .

© Nuffield Foundation and the Royal Society of Chemistry

• 11-14 years
• 14-16 years
• 16-18 years
• Practical experiments
• Demonstrations
• Elements and the periodic table
• Reactions and synthesis

Specification

• A more reactive metal can displace a less reactive metal from a compound.
• 4.2 Explain displacement reactions as redox reactions, in terms of gain or loss of electrons
• C3.2.1 deduce an order of reactivity of metals based on experimental results including reactions with water, dilute acid and displacement reactions with other metals
• (c) the relative reactivities of metals as demonstrated by displacement (e.g. iron nail in copper(II) chloride solution) and competition reactions (e.g. thermit reaction)
• (i) the properties and uses of iron (steel), aluminium, copper and titanium
• 2.1.4 explain and describe the displacement reactions of metals with other metal ions in solution;
• Corrosion of metals.
• Protective layers on Al, Cr.

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