Why is mgo a base

Stone, E. Garrone and A. Zecchina, Mater. Coluccia and A. Tench, in: Proc. Congress on Catalysis , Tokyo Elsevier, Amsterdam, p. Wu, J. Corneille, C. Estrada, J. He and D. Goodman, Chem. Wu, C. Estrada and D. Goodman, Phys. Corneille and D. Goodman, J. Petrie and J. Vols, Surf. Thiel and T. Madey, Surf. Andersson and J. Davenport, Solid State Commun. Chen, J. Crowell and J. Yaates, J. 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:.

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

Chlorine VII oxide is also known as dichlorine heptoxide, and chlorine I oxide as dichlorine monoxide. It continues the trend of the highest oxides of the Period 3 elements towards being stronger acids. 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. Chlorine VII oxide itself also reacts directly with sodium hydroxide solution to give the same product:. 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. Jim Clark Chemguide. Sodium Oxide Sodium oxide is a simple strongly basic oxide.

Magnesium oxide Magnesium oxide is another simple basic oxide, which also contains oxide ions. Aluminum Oxide Describing the properties of aluminum oxide can be confusing because it exists in a number of different forms. Silicon dioxide silicon IV oxide Silicon is too similar in electronegativity to oxygen to form ionic bonds.

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.

Argon is obviously omitted because it doesn't form an oxide. Note: If you haven't already been there, you might be interested in looking at the page about the structures and physical properties of the Period 3 oxides as a useful introduction before you go any further.

Use the BACK button on your browser to return quickly to this page later if you choose to follow this link. The trend is from strongly basic oxides on the left-hand side to strongly acidic ones on the right, via an amphoteric oxide aluminium oxide in the middle.

An amphoteric oxide is one which shows both acidic and basic properties. For this simple trend, you have to be looking only at the highest oxides of the individual elements.

Those are the ones on the top row above, and are where the element is in its highest possible oxidation state. The pattern isn't so simple if you include the other oxides as well. For the non-metal oxides, their acidity is usually thought of in terms of the acidic solutions formed when they react with water - for example, sulphur trioxide reacting to give sulphuric acid.

They will, however, all react with bases such as sodium hydroxide to form salts such as sodium sulphate. Warning: The rest of this page contains quite a lot of detail about the various oxides. Don't lose sight of the overall trend in the period with respect to the highest oxides when you are looking at all this detail. It is essential to know what your syllabus says about this topic, and to explore past papers and mark schemes - otherwise you are going to end up bogged down in a mass of detail that you don't actually need to know about.

If you are working towards a UK-based exam A level or its equivalent and haven't got any of these things follow this link before you go any further to find out how to get them. 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. Sodium oxide reacts exothermically with cold water to produce sodium hydroxide solution.

Depending on its concentration, this will have a pH around As a strong base, sodium oxide also reacts with acids. For example, it would react with dilute hydrochloric acid to produce sodium chloride solution.

Magnesium oxide is again a simple basic oxide, because it also contains oxide ions. However, it isn't as strongly basic as sodium oxide because the oxide ions aren't so free. It takes more energy to break these. Even allowing for other factors like the energy released when the positive ions form attractions with water in the solution formed , the net effect of this is that reactions involving magnesium oxide will always be less exothermic than those of sodium oxide. If you shake some white magnesium oxide powder with water, nothing seems to happen - it doesn't look as if it reacts.

However, if you test the pH of the liquid, you find that it is somewhere around pH 9 - showing that it is slightly alkaline. There must have been some slight reaction with the water to produce hydroxide ions in solution. Some magnesium hydroxide is formed in the reaction, but this is almost insoluble - and so not many hydroxide ions actually get into solution.

Magnesium oxide reacts with acids as you would expect any simple metal oxide to react. For example, it reacts with warm dilute hydrochloric acid to give magnesium chloride solution. Describing the properties of aluminium oxide can be confusing because it exists in a number of different forms. One of those forms is very unreactive.

It is known chemically as alpha-Al 2 O 3 and is produced at high temperatures. Aluminium oxide doesn't react in a simple way with water in the sense that sodium oxide and magnesium oxide do, and doesn't dissolve in it. Although it still contains oxide ions, they are held too strongly in the solid lattice to react with the water.

Note: Some forms of aluminium oxide do, however, absorb water very effectively. I haven't been able to establish whether this absorption just involves things like hydrogen bonds or whether an actual chemical reaction to produce some sort of hydroxide occurs. If you have any firm information on this, could you contact me via the address on the about this site page.

Aluminium oxide contains oxide ions and so reacts with acids in the same way as sodium or magnesium oxides. That means, for example, that aluminium oxide will react with hot dilute hydrochloric acid to give aluminium chloride solution. In this and similar reactions with other acids , aluminium oxide is showing the basic side of its amphoteric nature. Aluminium oxide has also got an acidic side to its nature, and it shows this by reacting with bases such as sodium hydroxide solution.

Various aluminates are formed - compounds where the aluminium is found in the negative ion. This is possible because aluminium has the ability to form covalent bonds with oxygen. In the case of sodium, there is too much electronegativity difference between sodium and oxygen to form anything other than an ionic bond. But electronegativity increases as you go across the period - and the electronegativity difference between aluminium and oxygen is smaller. That allows the formation of covalent bonds between the two.

Magnesium oxide is used extensively in the soil and groundwater remediation, wastewater treatment, drinking water treatment, air emissions treatment, and waste treatment industries for its acid buffering capacity and related effectiveness in stabilizing dissolved heavy metal species. Many heavy metals species, such as lead and cadmium are most soluble in water at acidic pH below 6 as well as high pH above Granular MgO is often blended into metals-contaminated soil or waste material, which is also commonly of a low acidic pH, in order to drive the pH into the 8—10 range where most metals are at their lowest solubilities.

Metal-hydroxide complexes have a tendency to precipitate out of aqueous solution in the pH range of 8— In effect, magnesia and other alkaline earths act as alkalies in some reactions displacement of other metal hydroxides but not in others failing to form concentrated strongly basic solutions. They represent an intermediate stage between full-fledged alkalis and the majority of metal oxides and hydroxides which are inert in the absence of added acid or strong base.

The word alkali is reserved to hydroxides of elements of the first column of the periodic table Li, Na, K, Rb, Cs, Fr. So oxides are not alkalies. Furthermore Magnesium is an element of the 2nd column. It is an alkaline earth metal. It does not produce an alkali when oxidized to MgO. Nevertheless, the "Concise Encyclopedia Chemistry" states that "in a broad sense, the alkalies also include the alkaline earth hydroxides and aqueous ammonia, because of their weaker basicity".

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