Guide Inorganic Chemistry Quick Review: Acid and Base Properties (Quick Review Notes)

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The overall electrolytic decomposition is Ions are charged particles e. What does the complete electrical circuit consist of?

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There are two ion currents in the electrolyte flowing in opposite directions: positive cations e. O 2— attracted to the positive anode electrode, BUT remember no electrons flow in the electrolyte, only in the graphite or metal wiring! The circuit of 'charge flow' is completed by the electrons moving around the external circuit e. Electron balancing : In the above process it takes the removal of four electrons from two oxide ions to form one oxygen molecule and the gain of three electrons by each aluminium ion to form one aluminium atom.

Therefore for every 12 electrons you get 3 oxygen molecules and 4 aluminium atoms formed. The properties and uses of aluminium. Aluminium can be made more resistant to corrosion by a process called anodising. Iron can be made more useful by mixing it with other substances to make various types of steel. Many metals can be given a coating of a different metal to protect them or to improve their appearance. Aluminium is a reactive metal but it is resistant to corrosion.

This is because aluminium reacts in air to form a layer of aluminium oxide which then protects the aluminium from further attack. This is why it appears to be less reactive than its position in the reactivity series of metals would predict. For some uses of aluminium it is desirable to increase artificially the thickness of the protective oxide layer in a process is called anodising. This involves removing the oxide layer by treating the aluminium sheet with sodium hydroxide solution.

The aluminium is then placed in dilute sulphuric acid and is made the positive electrode anode used in the electrolysis of the acid. Oxygen forms on the surface of the aluminium and reacts with the aluminium metal to form a thicker protective oxide layer. There is a note about structure of metal alloys on the metallic bonding page. More on the reactions of aluminium Reaction with aluminium with chlorine If dry chlorine gas Cl 2 is passed over heated iron or aluminium, the chloride is produced.

The experiment shown above should be done very carefully by the teacher in a fume cupboard. This is because on evaporation the compounds contain 'water of crystallisation'. Reaction of chloride with water : With a little water it rapidly, and exothermically hydrolyses to form aluminium hydroxide and nasty fumes of hydrogen chloride gas.

Theoretically its quite a reactive metal but an oxide layer is readily formed even at room temperature and this has quite an inhibiting effect on its reactivity. Although this again illustrates the 'under—reactivity' of aluminium, the Thermit Reaction shows its rightful place in the reactivity series of metals. The Thermit reaction : However the true reactivity of aluminium can be spectacularly seen when its grey powder is mixed with brown iron III oxide powder.

When the mixture is ignited with a magnesium fuse needed because of the very high activation energy!


Note the high temperature of the magnesium fuse flame is so high, the oxide layer to the delight of all pupils fails to inhibit the displacement reaction! Slow reaction with dilute hydrochloric acid to form the colourless soluble salt aluminium chloride and hydrogen gas. A mphoteric nature of aluminium hydroxide and acidity of the hexaaquaaluminium ion.

The addition of limited amounts of the bases sodium hydroxide or ammonia solution to an aluminium salt solution.

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A white gelatinous precipitate of aluminium hydroxide is formed. The further addition of excess sodium hydroxide or ammonia solution. With excess ammonia there is no effect, but with excess sodium hydroxide the aluminium hydroxide dissolves to form a soluble aluminate complex anion — therefore exhibiting amphoteric behaviour. You could write the equation in terms of forming these species too and any of the three possibilities should get you the marks.

To complete the 'amphoteric' picture of aluminium hydroxide we consider it dissolving in mineral acids to form typical salts e.

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Bubbles of carbon dioxide and a white gelatinous precipitate of aluminium hydroxide are formed. There several equation 'permutations' to represent this quite complicated reaction, so I've just composed one that shows the formation of both observed products. Since sodium carbonate solution is alkaline you can legitimately write a hydroxide ppt.

You can write an equation to show the formation of carbon dioxide leaving a soluble cationic complex of aluminium in solution and this equation fits in well with the acid—base nature of this reaction. This equation shows the hexaaquaaluminium ion acting as a Bronsted—Lowry acid donating two protons to the carbonate ion B—L base to form carbon dioxide and water.

This reaction shows why 'aluminium carbonate' 'Al 2 CO 3 3 ' cannot exist. The hydrated highly charged central metal ion is too acidic to co—exist with a carbonate ion. However with a lesser charged, lesser acidic ion, carbonates can exist, so there is an iron II carbonate FeCO 3. Aluminium salt solutions are slightly acidic for the same reasons as the carbonate reaction — namely the acidity of the hexaaquaaluminium ion i. The addition of excess sodium carbonate solution has no further effect.

Sodium carbonate is too weak a base to effect the amphoteric nature of aluminium hydroxide and dissolve the aluminium hydroxide precipitate. For strong alkalis like sodium hydroxide the whole sequence of each theoretical step of aluminium hydroxide precipitation and its subsequent dissolving in strong base—alkali is shown the series of diagrams below.

All are, for simplicity, treated as octahedral complexes of 6 ligands — either water H2O or hydroxide ion OH—. Aluminium compound reducing agents in organic chemistry. L ithium tetrahydridoaluminate III , LiAlH 4 lithium tetrahydride reduces aldehydes to primary alcohols and ketones to secondary alcohols. LiAlH 4 is a more powerful reducing agent than NaBH 4 and reacts violently with water, so the reaction must be carried out in an inert solvent like ethoxyethane ' ether '. The initial product is hydrolysed by dil.

LiAlH 4 is a more powerful reducing agent than NaBH 4 , and in ether solvent, readily reduces carboxylic acids to primary alcohols. What does that mean, exactly? How do these principles relate to alcohols? Cyclohexanol has the pKa of a typical alcohol about The pKa of phenol, however, is about That means the charge can be spread out throughout the molecule, which is stabilizing.

Any factor which stabilizes the conjugate base will increase acidity.

Acidity and Basicity of Alcohols – Master Organic Chemistry

Compare ethanol pKa 16 to 2,2,2-trifluoroethanol pKa about Why do you think trifluoroethanol is more acidic? Compare their conjugate bases. What is fluorine doing here to make the conjugate base more stable? This is an example of an inductive effect. Fluorine, being highly electronegative, pulls electron density away from the neighbouring carbon. That carbon, now being electron poor, pulls electron density away from the carbon next door. And that carbon, being slightly electron poor, can pull some electron density away from the oxygen.

The net result is that the oxygen has lower electron density, which is stabilizing. This also works if we compare alcohol variations where we change the distance between the OH and the fluorine atom. We can also use electronegativity trends to determine the order of acidity in these molecules. Since fluorine is more electronegative than chlorine which is more electronegative than bromine which is more electronegative than iodine, the inductive effect will be highest for CF3 and lowest for CI3.

Thank you for sharing the article. Regarding the phenol vs cyclohexanol example, since most resonance forms break the aromaticity of the ring, is the charge delocalization still significant enough as to make it a more stable conjugate base? Reopen my mind of organic chem. I have a question. Specifically for protonated amino acids. OK, so I am curious here..

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I posted a question regarding electronegativity vs polarizability. I see it said that polarizability takes presidence yet as you have stated he shows opposing acidity trends in the example as you have shown. Or maybe the characteristics oh the double bonded oxygen contributing?? The person you replied to is correct on one point and wrong on everything else. Furthermore, aside from 2,2,2-trichloroethanol, the pattern is in fact correct. It should not be wholly reversed like the person you replied to stated.

The polarizability of larger molecules is relevant for you in explaining intermolecular forces induced dipole attractions. The stability of halogen anions the conjugate base of a halogenic acid has to do with the charge density. Fluorine is very small, so carrying a negative charge by itself concentrates the negative charge in a very small space alongside other electrons, which leads to a repulsive and destabilizing interaction. In iodine, the single extra electron is spread out over a much larger volume which minimizes destabilizing interactions.

This along with orbital overlap HSAB theory — traditionally covered in your first inorganic chemistry course should more or less account for the differences in halogenic acid pka. The difference of 0. I am confused here. Yet starting a new sentence, the higher electronegativity due higher attraction generated through closness of opposing charges of the smaller halides mainly in question here can b justified to be electron withdrawing to carbon and oxygen there fore stabilizing the anion from both points being electronegativity, and polarizability..

Any meaningful way to spread out the charge density ex. For example, phenol the right-most molecule in question before the conclusion above , has a pka of 10 due to resonance structures. CF3 is a lot of inductive power, but resonance is still more important. Another example for polarizability: primary alcohols OH group have a pka around 16 while primary thiols SH group have a pka of around 10 — a million times more acidic.