Acids and Bases...
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So, the question goes like this:
Which of the following elements burns in air to give a compound which, in water, gives a solution of pH 2?
1) Aluminium
2)Calcium
3)Hydrogen
4)Sulphur
An oxide (burns in air), which gives H+ ions when dissociated in water (hydrogen?) Am I on the correct train of thought? Is an acid and base reaction reversible in the sense that the compound formed can be broken back into their original forms?
Secondly, whats the difference between acid and base? And whats the main point behind memorising the stupid basic/acidic/neutral/amphoteric oxide thing?
How are we supposed to differentiate hydrochloric acid from sulphuric acid?
sorry for the multiple questions >.>
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Originally posted by donkhead333:
So, the question goes like this:
Which of the following elements burns in air to give a compound which, in water, gives a solution of pH 2?
1) Aluminium
2)Calcium
3)Hydrogen
4)Sulphur
An oxide (burns in air), which gives H+ ions when dissociated in water (hydrogen?) Am I on the correct train of thought? Is an acid and base reaction reversible in the sense that the compound formed can be broken back into their original forms?
Secondly, whats the difference between acid and base? And whats the main point behind memorising the stupid basic/acidic/neutral/amphoteric oxide thing?
How are we supposed to differentiate hydrochloric acid from sulphuric acid?
sorry for the multiple questions >.>
>>> Which of the following elements burns in air to give a compound which, in water, gives a solution of pH 2?
1) Aluminium
2)Calcium
3)Hydrogen
4)Sulphur <<<
pH 2 indicates an acidic solution, which is produced by dissolving non-metal oxides in water, which must hence be sulfur dioxide or sulfur trioxide (which forms sulfurous acid and sulfuric acid respectively).
>>> An oxide (burns in air), which gives H+ ions when dissociated in water (hydrogen?) Am I on the correct train of thought? <<<
Yes, for instance, SO3 + H2O --> H2SO4 --> 2H+ + SO4 2-
>>> Is an acid and base reaction reversible in the sense that the compound formed can be broken back into their original forms? <<<Yes of course, you'll explore more about this at 'A' levels. Here's a sneak preview :
Consider the proton transfer (ie. acid-base reaction) between ammonia and water.
NH3 + H2O <<<---> NH4+ + OH-
NH4+ + OH- <--->>> NH3 + H2O
When ammonia (base) is protonated (ie. proton transferred from the acid, water, to ammonia), it forms ammonium cation (conjugate acid).
When water (acid) is deprotonated (ie. proton transferred from the acid, water, to ammonia), it forms hydroxide anion (conjugate base).
The NH3 (base) has become NH4+ which is now a conjugate acid, because it now has a proton that it can give (which is the function of any acid, to give protons).
The H2O (acid) has become OH- which is now a conjugate base, because it now has the capacity to accept (back) a proton (which is the function of any base, to accept protons).
>>> Secondly, whats the difference between acid and base? <<<
'O' level definition uses Arrhenius definition :
Arrhenius: According to this definition developed by the Swedish chemist Svante Arrhenius, an acid is a substance that increases the concentration of hydrogen ions (H+), which are carried as hydronium ions (H3O+) when dissolved in water, while bases are substances that increase the concentration of hydroxide ions (OH-). This definition limits acids and bases to substances that can dissolve in water. Around 1800, many French chemists, including Antoine Lavoisier, incorrectly believed that all acids contained oxygen. Indeed the modern German word for oxygen is Sauerstoff (lit. sour substance), as is the Afrikaans word for oxygen suurstof, with the same meaning. English chemists, including Sir Humphry Davy, at the same time believed all acids contained hydrogen. Arrhenius used this belief to develop this definition of acid.
'A' level definition uses Bronsted-Lowry definition and Lewis definition :
Brønsted-Lowry: According to this definition, an acid is a proton (hydrogen nucleus) donor and a base is a proton acceptor. The acid is said to be dissociated after the proton is donated. An acid and the corresponding base are referred to as conjugate acid-base pairs. Brønsted and Lowry independently formulated this definition, which includes water-insoluble substances not in the Arrhenius definition. Acids according to this definition are variously referred to as Brønsted acids, Brønsted-Lowry acids, proton acids, protic acids, or protonic acids.
Lewis: According to this definition developed by Gilbert N. Lewis, an acid is an electron-pair acceptor and a base is an electron-pair donor. (These are frequently referred to as "Lewis acids" and "Lewis bases," and are electrophiles and nucleophiles, respectively, in organic chemistry; Lewis bases are also ligands in coordination chemistry.) Lewis acids include substances with no transferable protons (ie H+ hydrogen ions), such as iron(III) chloride, and hence the Lewis definition of an acid has wider application than the Brønsted-Lowry definition. In fact, the term Lewis acid is often used to exclude protic (Brønsted-Lowry) acids. The Lewis definition can also be explained with molecular orbital theory. In general, an acid can receive an electron pair in its lowest unoccupied orbital (LUMO) from the highest occupied orbital (HOMO) of a base. That is, the HOMO from the base and the LUMO from the acid combine to a bonding molecular orbital.
>>> And whats the main point behind memorising the stupid basic/acidic/neutral/amphoteric oxide thing? <<<
So that you know which substances can react with which, and how they react, and what are the products.
For 'O' levels, you are only required to be familiar with 3 amphoteric oxides : aluminium oxide, zinc oxide, and lead(II) oxide.
Being metals, they react normally as basic oxides. But because they are amphoteric oxides, they can react as acidic oxides too. At 'O' levels, you need only know the name (and not the formula, that's 'A' levels) of the salts that are produced when these amphoteric oxides act as acidic oxides.
Eg. aluminium oxide + sodium hydroxide --> sodium aluminate.
Eg. zinc oxide + sodium hydroxide --> sodium zincate.
Eg. lead(II) oxide + sodium hydroxide --> sodium plumbate*
(*plumbing pipes in the toilets were traditionally made of lead metal)
>>> How are we supposed to differentiate hydrochloric acid from sulphuric acid? <<<
HCl is a strong monoprotic (monobasic) acid, while H2SO4 is a strong diprotic (dibasic) acid.
HCl contains Cl- ions, while H2SO4 contains SO4 2- ions.
Hence, you can distinguish the two acids on the basis of the above points. The easiest way involves the latter point, ie. by verifying the presence of the SO4 2- sulfate ions, by use of aqueous barium nitrate, which produces a ppt (which is barium sulfate solid) when sulfate ions are present.
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umm...regarding non-metal oxides dissolving in water to form an acid, im not sure how it is derived because all i know is that non-metal oxides include neutral and acidic oxides which has got totally nothing to do with water..
and i was referring to the difference between base and alkali, typo above sry.
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alkali is a soluble base. Is calcium hydroxide soluble?
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Originally posted by donkhead333:
umm...regarding non-metal oxides dissolving in water to form an acid, im not sure how it is derived because all i know is that non-metal oxides include neutral and acidic oxides which has got totally nothing to do with water..
and i was referring to the difference between base and alkali, typo above sry.
>>> umm...regarding non-metal oxides dissolving in water to form an acid, im not sure how it is derived because all i know is that non-metal oxides include neutral and acidic oxides which has got totally nothing to do with water.. <<<The reason why you have this problem (regarding this concept), is (not your fault at all, the blame lies with, or is) due to the oversimplified nature of the 'O' levels.
Perhaps you can understand better after reading this (if not, then no worries, you'll eventually understand when you explore Chemistry at 'A' levels).
Carbon monoxide is a neutral oxide, which means it does not react with bases/alkalis.
(Btw, an alkali is a base in aqueous state. Not all bases can exist as alkalis. But all alkalis are themselves bases. Eg. Na2O(s) and NaOH(aq) are bases, but only NaOH(aq) is an alkali, Na2O(s) is not an alkali.)
Carbon dioxide is an acidic oxide, which means it will react with bases/alkalis.
And how does carbon dioxide function as an acidic oxide? (This is the 1st part which donkhead333 is confused over, so here goes...) Carbon dioxide is acidic because it dissolves in water to form carbonic acid, which has the capacity to donate protons (ie. H+ ions), which (the protons) are accepted by the base/alkali, eg. sodium hydroxide.
CO2 + H2O --> H2CO3
H2CO3 + NaOH --> H2O + NaHCO3
or H2CO3 + 2NaOH --> 2H2O + Na2CO3
At 'O' levels, you may also write it summarily (ie. you summarize the above equations) as :
CO2 + NaOH --> NaHCO3
CO2 + 2NaOH --> H2O + Na2CO3
(This is the 2nd part which donkhead333 is confused over, so here goes...)
Since CO2 has the capacity (or potential) to form carbonic acid when dissolved in water, we consider CO2 an acidic oxide, even if we don't actually dissolve it in water. (In other words, you don't have to first dissolve CO2 in water before you call it an acidic oxide.)
Note that CO2 (or SO2 or SO3) has the capacity to react with aqueous bases/alkalis, because CO2 (or SO2 or SO3) first dissolves to form carbonic acid (or sulfurous acid or sulfuric acid). However, many 'O' level students are not aware of this because they are taught to write the summary equation directly, without actually understanding the mechanism properly.
However, at this point (if you're not confused and want to learn more), it may be relevant to point out that certain acidic oxides, may indeed react directly with certain bases, under certain conditions, without water present.
Eg. Silicon dioxide, SiO2, is an acidic oxide; but it is inert (ie. does not react, does not behave as an acid) with aqueous alkalis or solid bases. It will only react with molten bases (ie. under conditions of high temperature), such as molten sodium hydroxide, NaOH, or molten sodium oxide, Na2O.
SiO2 + Na2O (molten) --> Na2SiO3
SiO2 + NaOH (molten) --> Na2SiO3 + H2O
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Originally posted by davidche:
alkali is a soluble base. Is calcium hydroxide soluble?
Partially. In other words, you will obtain both SOLID calcium hydroxide (known as slaked lime) as well as AQUEOUS calcium hydroxide (known as limewater). -
Cool, small question: zinc hydroxide which is insoluble should not be considered an alkali even though it is called a hydroxide right?
Also, only ''SPA'' hydroxides are totally soluble right?
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Originally posted by davidche:
Cool, small question: zinc hydroxide which is insoluble should not be considered an alkali even though it is called a hydroxide right?
Also, only ''SPA'' hydroxides are totally soluble right?
>>> Cool, small question: zinc hydroxide which is insoluble should not be considered an alkali even though it is called a hydroxide right? <<<
Yes. Further info :
Zinc hydroxide, is not an "alkali" (because as you say, it's not soluble), but it is more properly called an "amphoteric hydroxide" (so it's both a basic hydroxide, as well as an acidic hydroxide).
See http://en.wikipedia.org/wiki/Zinc_hydroxide
>>> Also, only ''SPA'' hydroxides are totally soluble right? <<<
It's dangerous to over-generalize to say "ONLY" such and such are fully soluble. Yes, these 3 hydroxides are certainly fully soluble, but there are also others.
Don't worry too much about this at 'O' levels; you have your Qualitative Analysis data sheet to help you figure out which hydroxides are soluble and which are not.
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Eg. Silicon dioxide, SiO2, is an acidic oxide; but it is inert (ie. does not react, does not behave as an acid) with aqueous alkalis or solid bases. It will only react with molten bases (ie. under conditions of high temperature), such as molten sodium hydroxide, NaOH, or molten sodium oxide, Na2O.
SiO2 + Na2O (molten) --> Na2SiO3
SiO2 + NaOH (molten) --> Na2SiO3 + H2O
I would prefer that you use the one that is in the "O" level syllabus.
The blast furnace, where silica is reacted with limestone to produce slag which is then tapped off. Alright O level pals, remember this
CaCO3 (thermal decomposition of carbonates in blast furnace due to high heat)--> CaO(main agent to react silica to form slag) + CO2(escapes as waste gases)
CaO + SiO2 --> CaSiO3
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Originally posted by Garrick_3658:
I would prefer that you use the one that is in the "O" level syllabus.
The blast furnace, where silica is reacted with limestone to produce slag which is then tapped off. Alright O level pals, remember this
CaCO3 (thermal decomposition of carbonates in blast furnace due to high heat)--> CaO(main agent to react silica to form slag) + CO2(escapes as waste gases)
CaO + SiO2 --> CaSiO3
Yes, that's a good example.An additional example (similarly perhaps more relevant to your 'O' level syllabus), would be "Flue gas desulfurization".
From : http://en.wikipedia.org/wiki/Flue_gas_desulfurization
Scrubbing with a basic solid or solution
SO2 is an acidic gas and thus the typical sorbent slurries or other materials used to remove the SO2 from the flue gases are alkaline. The reaction taking place in wet scrubbing using a CaCO3 (limestone) slurry produces
CaSO3 (calcium sulfite) and can be expressed as:
- CaCO3 (solid) + SO2 (gas) → CaSO3 (solid) + CO2 (gas)
When wet scrubbing with a Ca(OH)2 (lime) slurry, the reaction also produces CaSO3 (calcium sulfite) and can be expressed as:
- Ca(OH)2 (solid) + SO2 (gas) → CaSO3 (solid) + H2O (liquid)
When wet scrubbing with a Mg(OH)2 (magnesium hydroxide) slurry, the reaction produces MgSO3 (magnesium sulfite) and can be expressed as:
- Mg(OH)2 (solid) + SO2 (gas) → MgSO3 (solid) + H2O (liquid)
Some FGD systems go a step further and oxidize the CaSO3 (calcium sulfite) to produce marketable CaSO4 · 2H2O (gypsum):
- CaSO3 (solid) + ½O2 (gas) + 2H2O (liquid) → CaSO4 · 2H2O (solid)
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wow, thanks mod for such quick and detailed replies!
this is regarding electrical conductivity in molten states. i understand that aqueous solutions are electrical conductors because of the free moving ions separated by the electronegativity of water(is this even correct?). however, i am unable to capture why a molten state enables compounds such as sodium chloride to become an electrical conductor.
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CaCO3 (solid) + SO2 (gas) → CaSO3 (solid) + CO2 (gas)
That's all we need to know.
To add on, maybe you can talk about reacting CaSO3 with oxygen to produce unreactive CaSO4, that can be disposed easily with minimal environmental damage.
i understand that aqueous solutions are electrical conductors because of the free moving ions separated by the electronegativity of water(is this even correct?)
You do not have to state that at O level (what's electro-negativity anyway?). It is because water is a polar solvent. For example, sodium chloride in water conducts electricity because the Na+ cations are attracted to the OH- ions in water, while the Cl- ions are attracted to the H+ ions in water. The sodium cations and chloride anions have dispersed into water as free mobile ions. Hence these mobile ions are able to act as charge carriers hence allow aqueous NaCl solution to conduct electricity.
In short, your theory about free moving ions is CORRECT.
however, i am unable to capture why a molten state enables compounds such as sodium chloride to become an electrical conductor.
Please be specific. I believe you are asking about IONIC compounds.
In molten state, the sodium cations and chloride anions are not longer held in fixed positions in their crystal lattice structure. If you need an example, compare the arrangement and movement of SOLID and LIQUID. Therefore, the mobile ions are able to act as charge carriers thus molten NaCl can conduct electricity.
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Originally posted by Garrick_3658:
CaCO3 (solid) + SO2 (gas) → CaSO3 (solid) + CO2 (gas)
That's all we need to know.
To add on, maybe you can talk about reacting CaSO3 with oxygen to produce unreactive CaSO4, that can be disposed easily with minimal environmental damage.
i understand that aqueous solutions are electrical conductors because of the free moving ions separated by the electronegativity of water(is this even correct?)
You do not have to state that at O level (what's electro-negativity anyway?). It is because water is a polar solvent. For example, sodium chloride in water conducts electricity because the Na+ cations are attracted to the OH- ions in water, while the Cl- ions are attracted to the H+ ions in water. The sodium cations and chloride anions have dispersed into water as free mobile ions. Hence these mobile ions are able to act as charge carriers hence allow aqueous NaCl solution to conduct electricity.
In short, your theory about free moving ions is CORRECT.
however, i am unable to capture why a molten state enables compounds such as sodium chloride to become an electrical conductor.
Please be specific. I believe you are asking about IONIC compounds.
In molten state, the sodium cations and chloride anions are not longer held in fixed positions in their crystal lattice structure. If you need an example, compare the arrangement and movement of SOLID and LIQUID. Therefore, the mobile ions are able to act as charge carriers thus molten NaCl can conduct electricity.
Garrick_3658 wrote :
>>> You do not have to state that at O level (what's electro-negativity anyway?). It is because water is a polar solvent. For example, sodium chloride in water conducts electricity because the Na+ cations are attracted to the OH- ions in water, while the Cl- ions are attracted to the H+ ions in water. The sodium cations and chloride anions have dispersed into water as free mobile ions. Hence these mobile ions are able to act as charge carriers hence allow aqueous NaCl solution to conduct electricity. <<<
The idea is more or less as Garrick described. But let's briefly run through it (this is for your understanding, not memorization).
The reason why in the AQUEOUS (which means "dissolved in water") state, ionic compounds (specifically those that are soluble in water, eg. nitrates... as to why some compounds are soluble while others are not, that's for you to enjoy at 'A' levels... don't unwrap all your gifts at once!) are able to apparently 'separate' from each other, and move freely in solution, is because :
Instead of the cations and anions being 'stuck' to each other, they are now 'stuck' or 'bonded' to water molecules instead (though it is not always 'covalent bonds', depending on the ions involved, the interactions may simply be electrostatic attraction with water).
Garrick mentioned the ions are attracted to the "H+ ions and OH- ions of water", to be somewhat more precise, the cations are attracted (or "bonded"; to be precise, "ion-dipole interactions") with the partially negatively charged oxygen atoms of water molecules, while the anions are attracted (or "bonded"; to be precise, "ion-dipole interactions") with the partially positively charged hydrogen atoms of the water molecules.
As to why the hydrogen atoms of water are partially positively charged, and why the oxygen atoms of water are partially negatively charged (hence, as Garrick mentioned, water is a POLAR molecule), it is because (and donkhead333 mentioned this term) oxygen is more ELECTRONEGATIVE than hydrogen (a more electronegative atom will draw/shift/polarize the electrons in its bonds to become slightly closer to itself, and slightly away from the less electronegative atom it is bonded with. For instance, in the H-O-H water molecule, the electrons in the two bonds are closer to oxygen and further from hydrogen, hence resulting in partial charges on the atoms).
So in effect, the ions (in aqueous state, ie. dissolved in water) are now "bonded" or "attracted" to water molecules, instead of being "stuck" to each other. This is the reason why soluble ionic compounds in the aqueous state can conduct electricity.
Finally, as to why ionic compounds in the molten state conducts electricity, it is because the high heat energy causes the ions to become so kinetically excited that they are able to overcome the electrostatic attraction (not completely, but just enough) so that they (all the ions, both cations and anions) are able to move relatively freely, just like a liquid (eg. in liquid water, water molecules slide over each other freely). Hence, with mobile ions, electricity can be conducted.
Note that in metal (eg. wires), electricity is the flow of electrons.
In ionic solutions or molten state, electricity is the flow of ions.
It's fine whether it's electrons or ions, because both are CHARGED PARTICLES. Which is what electricity is all about.