Chem - Blog Posts

Things that happened in chemistry today:

1. Teacher said this weekend was better than any weekend he'd ever had

Why? He took his baby to this gymnastics thingy (valid)

This was rated better than:

1. When he beat an American record in pole vaulting

2. When he played in Germany for American pole vaulting

3. His wedding

4. When he took said baby home for the first time

2. A mole went 'missing' for st.patricks day

3. Said mole is apparently named Moledo (Like waldo)

4. Teacher apparently lets Moledo go out at night to catch worms n shit

5. Moledo is dead. He was murdered.

Teacher said his goodbyes to Moledo

We couldn't hear him cause he was in the courtyard and we were in the halls

6. We were sent to investigate Moledos death

7. We were the only class to sprint (If we collected the evidence cards in under 7 min we got extra credit)

RIP Moledo

He will be missed 😔

The Chemical Structure of Redstone

So I was curious about what the chemical structure of Redstone looks like, and Minecraft Education Edition, albeit unintentionally, gives us a canon look into what Redstone is made of:

In Minecraft Education Edition, putting a Redstone Block into a Material Reducer shows that it's composed of 31 Carbon, 31 Uranium, and 38 Unobtanium, which we can assume to be measured in grams

Dividing the Redstone Block into Redstone Dust, each Redstone Dust is then composed of approximately 3.4 Carbon, 3.4 Uranium, and 4.2 Unobtanium

Again assuming that's measured in grams, that's 0.17 cm³ of Uranium, 1.496 cm³ of Carbon, and ???³ of Unobtanium per Redstone Dust

So what does this tell us about the chemical structure of Redstone? Basing this on Redstone Dust's composition, we can estimate that each Redstone molecule is composed of 3 Carbon atoms, 3 Uranium atoms, 4 Unobtanium atoms, a little under half of the time it binds to an extra Uranium and/or Carbon, and 20% of the time it binds to an extra Unobtanium

This also has some horrifying implications for how Redstone works:

Redstone would be extremely volatile as the radioactive decay from Unobtanium and Uranium would occasionally release Helium ions through alpha radiation, sometimes breaking apart Carbon into two Beryllium atoms (as it absorbs the extra proton and neutron from the Uranium) or merging into Oxygen

So Redstone should, in theory, be extremely flammable and potentially explosive, which implies that cave static, or the player mining Redstone with an Iron Pickaxe, could lead to a spark that causes an explosive cave-in

As Unobtanium is just a placeholder for unobtainable elements (hence the name), I'm going to estimate Unobtanium in this case as Unbinilium, the placeholder name for element 120

Why?

I'm estimating the Unobtanium as Redstone as being larger than the largest man-made element, Oganesson, which holds an impressive 118 protons

Each valence electron shell, from innermost to outermost, can bind with 2, 8, 18, 32, 32, 18, and 8 shells respectively, so I'd like Unobtanium to be an element we haven't discovered yet, and consequently I'd like to jump up to the next shell

While I could estimate with element 119's placeholder, Ununennium, it would have one electron in the next shell, so Unbinilium allows for easier chemical binding

So what does this molecule look like then? Well, horrifyingly...

It looks like this. As Redstone forms in crystal lattices, and only two Carbon atoms are free to bind, I can absolutely see why it's so brittle that it breaks into powder.

This makes the structure of Redstone:

C3U3Uno4 (55% of molecules) C4U3Uno4 (13% of molecules) C3U4Uno4 (13% of molecules) C4U4Uno4 (7% of molecules) C3U3Uno5 (5% of molecules) C4U3Uno5 (3% of molecules) C3U4Uno5 (3% of molecules) C4U4Uno5 (1% of molecules)

An extremely radioactive, flammable, and explosive compound.

Cesium-133, let it be. Cesium-134, let it be even more.

Periodic Table Regions [Explained]

Transcript

[A periodic table with regions labeled.]

[Hydrogen:] Slightly fancy protons [Lithium and Beryllium:] Weird dirt [Group 1 & 2 metals, Periods 3-4:] Regular dirt [Group 1 & 2 metals, Periods 5-7:] Ends in a number, let it slumber ends in a letter, not much better [Left side of the transition metals group:] Boring alloy metals Probably critical to the spark plug industry or something (but one of them is radioactive so stay on your toes) [Most of the top row of the transition metals + aluminum:] Regular metals [Below the rightmost "regular metals" - the "ordinary metals" and some transition metals:] Weird metals [The platinum group:] $$$$ [Boron:] Boron (fool's carbon) [Carbon, Nitrogen, Oxygen, and Phosphorus:] You are here [The Halogens:] Safety goggles required [Noble Gases:] Lawful neutral [Iodine and Radon:] Very specific health problems [Ordinary metals and metalloids - Arsenic, Antimony, Tellurium, Thallium, Lead, Bismuth, Polonium] Murder weapons [Astatine and Period 7 from Rutherfordium onwards:] Don't bother learning their names - they're not staying long [Lanthanides and Actinides:] Whoever figures out a better way to fit these up there gets the next Nobel Prize

Covalent and Dative Bonds

Covalent and dative (sometimes called co-ordinate) bonds occur between two or more non-metals,  e.g. carbon dioxide, water, methane and even diamond. But what actually are they?

A covalent bond is a chemical bond that involves the sharing of electron pairs between atoms. They are found in molecular elements or compounds such as chlorine or sulfur, but also in macromolecular elements and compounds like SiO2 and graphite. Covalent bonds are also found in molecular ions such as NH4+ and HCO3-.

Single covalent bonds have just one shared pair of electrons. Regularly, each atom provides one unpaired electron (the amount of unpaired electrons is usually equal to the number of covalent bonds which can be made) in the bond. Double covalent bonds have two shared pairs of electrons, represented by a double line between atoms, for example, O=C=O (CO2). Triple covalent bonds can also occur such as those in N ≡ N.

Dot and cross diagrams represent the arrangement of electrons in covalently bonded molecules. A shared pair of electrons is represented by a dot and a cross to show that the electrons come from different atoms.

Unpaired electrons are used to form covalent bonds as previously mentioned. The unpaired electrons in orbitals of one atom can be shared with another unpaired electron in an orbital but sometimes atoms can promote electrons into  unoccupied orbitals in the same energy level to form more bonds. This does not always occur, however, meaning different compounds can be formed - PCl3 and PCl4 are examples of this.

An example where promotion is used is in sulfur hexafluoride (SF6). The regular configuration of sulfur atoms is 1s2 2s2 2p6 3s2 3p4. It promotes, as shown in the diagram (see excited state), two electrons: one from the 3s electrons to the 3d orbital and one from the 3p to the 3d. Therefore there are 6 unpaired electrons for fluorine atoms to join. It has an octahedral structure.

An atom which has a lone pair (a pair of electrons uninvolved in bonding) of electrons can form a coordinate bond with the empty orbital of another atom. It essentially donates an electron into this orbital which when formed, acts the same as a normal covalent bond. A coordinate bond therefore contains a shared pair of electrons that have come from one atom.

When ammonia reacts with a H+ ion, a coordinate bond is formed between the lone pair on the ammonia molecule and the empty 1s sub-shell in the H+ ion. An arrow represents the dative covalent bond (coordinate bond). Charges on the final ion must be showed.

Summary

A covalent bond is a chemical bond that involves the sharing of electron pairs between atoms. They are found in molecular elements or compounds as well as in macromolecular elements and compounds. Also found in molecular ions.

Single covalent bonds have just one shared pair of electrons.  Double covalent bonds have two shared pairs of electrons, represented by a double line between atoms. Triple covalent bonds can also occur.

Dot and cross diagrams represent the arrangement of electrons in covalently bonded molecules. A shared pair of electrons is represented by a dot and a cross to show that the electrons come from different atoms.

Unpaired electrons are used to form covalent bonds - they can be shared with another unpaired electron in an orbital but sometimes atoms can promote electrons into unoccupied orbitals in the same energy level to form more bonds. This does not always occur, however, meaning different compounds can be formed.

An example where promotion is used is in sulfur hexafluoride (SF6). 

An atom which has a lone pair (a pair of electrons uninvolved in bonding) of electrons can form a coordinate bond with the empty orbital of another atom.

 It donates an electron into this orbital which when formed, acts the same as a normal covalent bond. A coordinate bond therefore contains a shared pair of electrons that have come from one atom.

When ammonia reacts with a H+ ion, a coordinate bond is formed between the lone pair on the ammonia molecule and the empty 1s sub-shell in the H+ ion. An arrow represents the dative covalent bond (coordinate bond). Charges on the final ion must be showed.

Breaking Down Alkanes - isn’t it cracking?

Unfortunately, if you’re sitting your A Level chemistry exam, you need to know a little more than the basic properties of alkanes outlined in my last post. Luckily though, this post takes you through fractional distillation and the two types of cracking - isn’t that convenient?

Crude oil contains carbon compounds formed by the effects of pressure and high temperature on plant and animal remnants. It is viscious, black and found in rocks beneath the earth’s surface. It is a mixture of mainly alkane hydrocarbons which are separated by a process called fractional distillation. Crude oil is essential because it is burned as a fuel and each fraction has different properties e.g. diesel, petrol, jet fuel.

Fractional distillation is the continual evaporation and condensation of a mixture which causes fractions to split due to a difference in boiling point. It is important to note that fractional distillation does not separate crude oil into pure compounds but rather less complex mixtures. Fractions are groups of compounds that have similar boiling points and are removed at the same level of a fractionating column.

The first step in this process is to heat crude oil in a furnace until some changes state from a liquid to a vapour. This mixture goes up a fractionating tower or column which is hotter at the bottom than the top and reaches a layer which is cool enough to condense and be collected. Shorter chain molecules are collected at the top where it is cooler since they have lower boiling points.

As you go down the fractionating column, bear in mind that: the column temperature increases, the boiling point increases, the number of carbon atoms increases and the strength of the Van der Waals’ between molecules increases.

Different fractions have different usefulnesses and often, it is the fractions with lower boiling points and shorter chains which are much more purposeful. Therefore there needs to be a process to getting shorter chains because they are the least abundant in crude oil samples. To meet demand, long chain molecules that are less useful are broken down into shorter chain molecules. This is done by cracking.

Cracking is a process where long chain hydrocarbon molecules are broken down into shorter chain molecules which are in high demand. This can be done one of two ways - thermal or catalytic.

Thermal cracking involves heating long chain alkanes to high temperatures - usually between 1000 - 1200K. It also uses high pressures up to 70atm and takes just one second. It only needs a second because the conditions could decompose the molecule completely to produce carbon and hydrogen instead. The conditions produce shorter chain alkanes and mostly alkenes.

A typical equation for this:

Decane -> octane + ethene

C10H22 -> C8H18 + C2H4

Catalytic cracking also breaks down long alkanes by heat under pressure using the presence of a zeolite catalyst. Temperature used is approx. 800-1000K and the pressure is often between 1-2 atm. Zeolite is an acidic mineral with a honeycomb structure, made from aluminium oxide and silicion dioxide. The honeycomb structure gives the catalyst a larger surface area which increases ROR. Factories which catalytically crack are often operated continuously for around 3 years at a time and produce branched alkanes, cycloalkanes and aromatic compounds.

You need to be able to compare the conditions of catalytic and thermal cracking for the A Level exam. Know that thermal cracking has a high temperature and pressure, a short duration, no catalyst and produces a high percentage of alkenes and some short chain alkanes. Catalytic uses a catalyst, a high temperature, a low pressure and produces aromatic hydrocarbons and motor fuels.

SUMMARY

Crude oil contains carbon compounds formed by the effects of pressure and high temperature on plant and animal remnants. I It is a mixture of mainly alkane hydrocarbons which are separated by a process called fractional distillation.

Fractional distillation is the continual evaporation and condensation of a mixture which causes fractions to split due to a difference in boiling point. 

It is important to note that fractional distillation does not separate crude oil into pure compounds but rather less complex mixtures.

Fractions are groups of compounds that have similar boiling points and are removed at the same level of a fractionating column.

The first step in this process is to heat crude oil in a furnace until some changes state from a liquid to a vapour. This mixture goes up a fractionating tower or column which is hotter at the bottom than the top and reaches a layer which is cool enough to condense and be collected. Shorter chain molecules are collected at the top where it is cooler since they have lower boiling points.

As you go down the fractionating column, bear in mind that: the column temperature increases, the boiling point increases, the number of carbon atoms increases and the strength of the Van der Waals’ between molecules increases.

Fractions with lower boiling points and shorter chains are much more purposeful but are the least abundant in crude oil samples. To meet demand, long chain molecules that are less useful are broken down into shorter chain molecules. 

Cracking is a process where long chain hydrocarbon molecules are broken down into shorter chain molecules which are in high demand. 

Thermal cracking involves heating long chain alkanes to high temperatures - usually between 1000 - 1200K. It also uses high pressures up to 70atm and takes just one second. It only needs a second because the conditions could decompose the molecule completely to produce carbon and hydrogen instead. The conditions produce shorter chain alkanes and mostly alkenes.

Catalytic cracking also breaks down long alkanes by heat under pressure using the presence of a zeolite catalyst. Temperature used is approx. 800-1000K and the pressure is often between 1-2 atm. Zeolite is an acidic mineral with a honeycomb structure, made from aluminium oxide and silicion dioxide. The honeycomb structure gives the catalyst a larger surface area which increases ROR. 

You need to be able to compare the conditions of catalytic and thermal cracking for the A Level exam. Know that thermal cracking has a high temperature and pressure, a short duration, no catalyst and produces a high percentage of alkenes and some short chain alkanes. Catalytic uses a catalyst, a high temperature, a low pressure and produces aromatic hydrocarbons and motor fuels.

Happy studying!

Alkanes: Saturated Hydrocarbons

So you want to be an organic chemist? Well, learning about hydrocarbons such as alkanes is a good place to start…

Alkanes are a homologous series of hydrocarbons, meaning that each of the series differs by -CH2 and that the compounds contain carbon and hydrogen atoms only. Carbon atoms in alkanes have four bonds which is the maximum a carbon atom can have - this is why the molecule is described to be saturated. Saturated hydrocarbons have only single bonds between the carbon atoms.

The general formula of an alkane is CnH2n+2 where n is the number of carbons. For example, if n = 3, the hydrocarbon formula would be C3H8 or propane. Naming alkanes comes from the number of carbons in the chain structure.

Here are the first three alkanes. Each one differs by -CH2.

Shorter chain alkanes are gases at room temperature, medium ones are liquids and the longer chain alkanes are waxy solids.

Alkanes have these physical properties:

1. They are non-polar due to the tiny difference in electronegativity between the carbon and hydrogen atoms.

2. Only Van der Waals intermolecular forces exist between alkane molecules. The strength of these increase as relative molecular mass increases therefore so does the melting/boiling point.

3. Branched chain alkanes have lower melting and boiling points than straight chain isomers with the same number of carbons. Since atoms are further apart due to a smaller surface area in contact with each other, the strength of the VDWs is decreased.

4. Alkanes are insoluble in water but can dissolve in non-polar liquids like hexane and cyclopentane. Mixtures are separated by fractional distillation or a separating funnel.

The fractional distillation of crude oil, cracking and the combustion equations of the alkanes will be in the next post.

SUMMARY

Alkanes are a homologous series of hydrocarbons. Carbon atoms in alkanes have four bonds which is the maximum a carbon atom can have - this is why the molecule is described to be saturated. Saturated hydrocarbons have only single bonds between the carbon atoms.

The general formula of an alkane is CnH2n+2 where n is the number of carbons.

Shorter chain alkanes are gases at room temperature, medium ones are liquids and the longer chain alkanes are waxy solids.

They are non-polar.

Only Van der Waals intermolecular forces exist between alkane molecules. The strength of these increase as relative molecular mass increases therefore so does the melting/boiling point.

Branched chain alkanes have lower melting and boiling points than straight chain isomers with the same number of carbons.

Alkanes are insoluble in water but can dissolve in non-polar liquids like hexane. Mixtures are separated by fractional distillation or a separating funnel.

Nomenclature - what in the organic chemistry is it?

Organic chemistry is so widely studied it requires a standard system for naming compounds, developed by IUPAC. Nomenclature is simply naming these organic compounds.

So, you want to be an organic chemist? Well, it starts here. Are you ready?

(psst… once you’ve learnt this theory, try a quiz here!)

1. Count your longest continuous chain of carbons.

Bear in mind that some chains may be bent. You’re looking for the longest chain of subsequent carbon atoms. This number correlates to root names that indicate the carbon chain length, listed below:

The second part of naming your base comes from the bonding in the chain. Is it purely single bonds or are there double bonds in there? If you are familiar with carbon chemistry, you’ll know that saturated hydrocarbons are called alkanes and unsaturated hydrocarbons are called alkenes. Therefore, the syllable -ane is used when it has only single bonds and the syllable -ene is used when it has some double bonds. For example:

Sometimes carbon chains exist in rings rather than chains. These have the prefix of -cyclo.

2. Identify your side chains attached to this main carbon and name them.

Side chains are added as prefixes to the root names. Sometimes called substituents, these are basically anything that comes off the carbon chain. Examples of the prefixes are listed below:

There are other prefixes such as fluoro (-F) and chloro (-Cl) which can describe what is coming off the chain.

3. Identify where each side chain is attached and indicate the position by adding a number to the name. 

We aim to have numbers as small as possible. For example, if bromine is on the second carbon of a 5-carbon saturated chain, we number it as 2-bromopentane instead of 4-bromopentane, since it would essentially be 2-bromopentane if it was flipped. Locant is the term used for the number which describes the position of the substitute group, e.g. the ‘2′ in 2-chlorobutane is the locant.

Sometimes there are two or more side chains e.g. a methyl group and a chlorine attached to a pentane. In these cases, these rules apply:

1. Names are written alphabetically.

2. A separate number is needed for each side chain or group.

3. Hyphens are used to separate numbers and letters.

This would be named 2-chloro-3-methyl-pentane. This is because the longest chain of carbons is 5 (pentane), the chlorine is on the second carbon (2-chloro) and the methyl group is on the third carbon (3-methyl). It is 2-chloro rather than 4-chloro as we aim to have as small as numbers as possible.

Another variation of this step to be aware of is how many of the same side chains or groups there are, for example, having two methyl groups would be dimethyl rather than solely methyl. Each group must also be given numbers separated by commas to show where each one is located. 

The list of these prefixes is found here:

Convention does not usually require mono- to go before a single group or side chain.

4. Number the positions of double bonds if applicable.

Alkenes and other compounds have double bonds. These must be indicated with numbers. For example, pent-2-ene shows that the double bond is between carbon 2 and carbon 3. The number goes in the middle of the original root name e.g. butene, pentene.

(!) Below is a list of functional groups that you may need to study for the AS and A Level chemistry exams. “R” represents misc. carbons. It is important to know that some groups are more prioritised than naming. From the most to least priority: carboxylic acid, ester, acyl chloride, nitrile, aldehyde, ketone, alcohol, amine, alkene, halogenalkane. It is worthwhile learning these.

bigger version here (I suggest downloading and printing it)

But wait, there’s more:

Here are some things to bear in mind when naming organic compounds:

1. The letter ‘e’ is removed when there are two vowels together e.g. propanone rather than propaneone. The ‘e’ isn’t removed when it is next to consonant, e.g. propanenitrile isn’t propannitrile.

2. When compounds contain two different, one is named as part of the unbranched chain and the other is named as a substituent. Which way round this goes depends on the priority. 

SUMMARY

Count your longest continuous chain of carbons.

Chains may be bent. You’re looking for the longest chain of subsequent carbon atoms. This number correlates to root names that indicate the carbon chain length, e.g. pentane.

The second part of naming your base comes from the bonding in the chain. Is it purely single bonds or are there double bonds in there? The syllable -ane is used when it has only single bonds and the syllable -ene is used when it has some double bonds.

Rings have the prefix of -cyclo.

Identify your side chains attached to this main carbon and name them.

Side chains are added as prefixes to the root names. Sometimes called substituents, these are basically anything that comes off the carbon chain. 

There are other prefixes such as fluoro (-F) and chloro (-Cl) which can describe what is coming off the chain.

Identify where each side chain is attached and indicate the position by adding a number to the name.

We aim to have numbers as small as possible. Locant is the term used for the number which describes the position of the substitute group, e.g. the ‘2′ in 2-chlorobutane is the locant.

Sometimes there are two or more side chains e.g. a methyl group and a chlorine attached to a pentane. In these cases, names are written alphabetically, a separate number is needed for each side chain or group and hyphens are used to separate numbers and letters.

When there are two or more of the same side chains or substituent groups, these must also be given numbers separated by commas to show where each one is located.

Number the positions of double bonds if applicable.

Alkenes and other compounds have double bonds. These must be indicated with numbers. The number goes in the middle of the original root name e.g. butene, pentene.

It is worthwhile learning the other functional groups that can be added on.They have varying priorities.

The letter ‘e’ is removed when there are two vowels together e.g. propanone rather than propaneone. The ‘e’ isn’t removed when it is next to consonant, e.g. propanenitrile isn’t propannitrile.

When compounds contain two different, one is named as part of the unbranched chain and the other is named as a substituent. Which way round this goes depends on the priority.

Happy studying guys!

Haloalkanes and Their Angelic Reactions: Part One

Haloalkanes are more commonly referred to as halogenoalkanes. Obviously you’ve already read my post on halogenoalkanes and their properties so there’s no surprise that you’re itching to read what I’ve got to say about these beauties and their reactions! Should we delve in?

There are a few different kinds of reactions you must learn for the A Level exam that involve halogenoalkanes. 

The first is the synthesis of chloroalkanes via the photochemical chlorination of the alkanes. I know it looks scary, but don’t worry, it is simpler than it sounds. It essentially means “forming chloroalkanes through chlorinating an alkane in the presence of sunlight”.

Chlorine will react with methane when UV light is present and will form several kinds of chloroalkanes and fumes of hydrogen chloride gas. Chloromethane was once commonly used as a refridgerant. Depending on how many chlorine molecules there are, there will be different compounds formed:

methane + chlorine -> chloromethane + hydrogen chloride

CH4 + Cl2 -> CH3Cl + HCl

or

methane + chlorine -> trichloromethane + hydrogen chloride

CH4 + 3Cl2 -> CHCl3 + 3HCl

When undergone in real life, mixtures of halogenoalkanes are produced with some long chain alkanes which can be separated out with fractional distillation. 

To understand what happens in an overall chemical reaction, chemists use mechanisms. These basically show the step-by-step process that is usually shown by a simple symbol equation that summarises everything.

The chlorination of methane is something you must learn the mechanism for. It’s pretty easy but involves a lot of steps and must be revised periodically to remember them.

The actual reaction is a substitution reaction because one atom or group is replaced by another. Since the chlorine involved is a free radical, it can also be called a free-radical substitution reaction.

1. Initiation

UV light is essential for the first step in the mechanism. This breaks the Cl-Cl covalent bond so that each chlorine leaves with one electron from the shared pair. Chlorine free radicals, with one unpaired electron in the outer shell, are formed. Free radicals are only formed if a bond splits evenly - each atom getting one of the two electrons. The name given to this is homolytic fission.

2. Propagation

This has two sub-steps

(a) Chlorine free radicals (highly reactive) react with methane to form hydrogen chloride and leave a methyl free radical.

Cl• + CH4 -> HCl + •CH3

(b) This free radical then reacts with another chlorine to form chloromethane and another chlorine free radical. Producing free radicals is a chain reaction which is why it is such a problem in ozone depletion - a little amount can cause a lot of destruction.

•CH3 + Cl2 -> CH3Cl +  •Cl

3. Termination

This step stops the chain reaction. It only happens when two free radicals collide to form a molecule in several ways:

Cl• + Cl• -> Cl2

UV light would just break down the chlorine molecule again, so although this is technically a termination reaction it is not the most efficient.

Cl• +  •CH3 -> CH3Cl

Forming one molecule of methane uses one chlorine and one methyl free radical.

•CH3 +  •CH3 -> C2H6

Ethane can be formed from two methyl free radicals - this is why there are longer chain alkanes in the mixture. 

This whole process is how organic halogenoalkanes are the product of photochemical reactions of halogens with alkanes in UV light - made via free radical substitution mechanisms in chain reaction.

Another reaction you need to know is a nucleophilic substitution reactions. A nucleophile is an electron pair donor or proton acceptor - the name comes from Greek origins (”loves nucleus”) - such as hydroxide ions, cyanide ions or ammonia molecules. Hydroxide and cyanide ions are negative but ammonia is neutral.

Halogenoalkanes have a polar bond because of the difference between the highly electronegative halogen and the carbon atom. The 𝛿+ carbon can go under nucleophilic attack. The mechanism for negatively charged nucleophiles these in general is:

Nu represents the nucleophile. This example is with a bromoalkane. Make sure to include curly arrows that begin at a lone pair or the centre of a bond and end at an atom or centre of bond, and delta (slight) charges.

Lets look at a more specific example:

One nucleophile that can be used is a hydroxide ion, found in either water or sodium hydroxide. In this case, you need to know about aqueous sodium hydroxide or potassium hydroxide and a halogenoalkane. This takes place at room temperature but is slow so is often refluxed (continuously boiled and condensed back into the reaction flask). Reflux apparatus is shown below:

The halogenoalkane is dissolved into ethanol since it is insoluable in water and this solution along with the aqueous hydroxide can mix. The product produced is an alcohol, which is organic.

The general reaction is:

R-CH2X + NaOH -> CH3CH2OH + NaX

Where X represents a halogen.

You must learn the mechanism for this reaction. The lone pair on the hydroxide attacks the carbon atom attached to the halogen and this causes both carbon electrons to move to the halogen which becomes a halide ion.

The reaction of a hydroxide ion can also be classed as a hydrolysis reaction as it breaks down chemical bonds with water or hydroxide ions. The speed of reaction depends on the strength of the bond - a stronger carbon-halogen bond, a slower reaction.

C-I is the most reactive (reactivity increases down group 7) and C-F is therefore the least reactive and strongest.

Part two of this post will cover nucleophilic substitution of cyanide ions and ammonia molecules, as well as elimination reactions.

SUMMARY

You need to know about the synthesis of chloroalkanes via the photochemical chlorination of the alkanes. - “forming chloroalkanes through chlorinating an alkane in the presence of sunlight”.

Chlorine will react with methane when UV light is present and will form several kinds of chloroalkanes and fumes of hydrogen chloride gas. Depending on how many chlorine molecules there are, there will be different compounds formed.

When undergone in real life, mixtures of halogenoalkanes are produced with some long chain alkanes which can be separated out with fractional distillation. 

To understand what happens in an overall chemical reaction, chemists use mechanisms. These basically show the step-by-step process.

The chlorination of methane is something you must learn the mechanism for. The actual reaction is a substitution reaction because one atom or group is replaced by another. 

The first step is initiation - UV light is essential for the first step in the mechanism. This breaks the Cl-Cl covalent bond so that each chlorine leaves with one electron from the shared pair. Chlorine free radicals, with one unpaired electron in the outer shell, are formed. Free radicals are only formed if a bond splits evenly - each atom getting one of the two electrons.

Step two is propagation: (a) Chlorine free radicals (highly reactive) react with methane to form hydrogen chloride and leave a methyl free radical (b) this free radical then reacts with another chlorine to form chloromethane and another chlorine free radical. Producing free radicals is a chain reaction which is why it is such a problem in ozone depletion - a little amount can cause a lot of destruction.

To stop the chain reaction, the final step is termination. It only happens when two free radicals collide to form a molecule in several ways: two chlorine free radicals forming a chlorine molecule, two methyl FRs forming ethane or a chlorine FR and a methyl FR forming chloromethane.

Ethane contributes to the longer chain alkanes in the mixture. 

Another reaction you need to know is a nucleophilic substitution reactions. A nucleophile is an electron pair donor or proton acceptor, such as hydroxide ions, cyanide ions or ammonia molecules. Hydroxide and cyanide ions are negative but ammonia is neutral.

Halogenoalkanes have a polar bond because of the difference between the highly electronegative halogen and the carbon atom. The 𝛿+ carbon can go under nucleophilic attack. 

Nu represents the nucleophile. Make sure to include curly arrows that begin at a lone pair or the centre of a bond and end at an atom or centre of bond, and delta (slight) charges.

One nucleophile that can be used is a hydroxide ion, found in either water or sodium hydroxide. In this case, you need to know about aqueous sodium hydroxide or potassium hydroxide and a halogenoalkane. This takes place at room temperature but is slow so is often refluxed (continuously boiled and condensed back into the reaction flask). The halogenoalkane is dissolved into ethanol since it is insoluable in water and this solution along with the aqueous hydroxide can mix. The product produced is an alcohol, which is organic.

The general reaction is :R-CH2X + NaOH -> CH3CH2OH + NaX where X represents a halogen

The lone pair on the hydroxide attacks the carbon atom attached to the halogen and this causes both carbon electrons to move to the halogen which becomes a halide ion.

The reaction of a hydroxide ion can also be classed as a hydrolysis reaction as it breaks down chemical bonds with water or hydroxide ions. 

The speed of reaction depends on the strength of the bond - a stronger carbon-halogen bond, a slower reaction. C-I is the most reactive (reactivity increases down group 7) and C-F is therefore the least reactive and strongest.

Metallic Bonding

A short one to finish off my first ever mini-series on bonding – ionic, covalent and finally metallic. There are metallic and metallic compounds and elements but for the A Level exam, we must look at the bonding within metals themselves. Don’t worry – I saved the easiest to last!

Metals are most usually solid so have particles packed close together. These are in layers which mean that the outer electrons can move between them rather than being bound to particular atoms. These are referred to as delocalised electrons because of this.

It’s pretty common knowledge that metals are good conductors of heat and electricity and it’s these delocalised electrons that give them this property.

Metals are therefore without their electrons so become positive ions. The metallic bond is actually the attraction between delocalised electrons and positive metal ions in the lattice. And that’s pretty much metallic bonding, you just need to know the properties of metals which are touched upon at lower levels of education.

These are the properties of metals:

1.       High melting points

Metals have large regular structures with strong forces between the oppositely charged positive ions and negative electrons, meaning these must be overcome to melt the metal – this requires a large amount of heat energy. Transition metals tend to have higher melting points than the main group metals because they have large numbers of d-shell electrons which can become delocalised creating a stronger metallic bond. Melting points across a period increase because they can have progressively more delocalised electrons: Na+, Mg 2+ and Al 3+ for example.

2.       Heat conductivity

Heat is conducted if particles can move and knock against each other to pass it on. Delocalised electrons allow this to happen.  Silver is a particularly good conductor of heat.

3.       Electrical conductivity

Delocalised electrons can carry charge and move, the two requirements of electrical conductivity. Current can flow because of these delocalised electrons.

4.       Ductile and malleable

Metals can be stretched and hammered into shape, making them ideal for things such as wires. Layered lattices mean that layers can slide over each other without disrupting the bonding – it is all still held together by the delocalised electrons and their strong attraction to the positive metal ions.

5.       High densities

Being a solid, metal ions are packed closely together so they have a high density, which makes them ideal for musical instrument strings. These can withstand the frequency of vibration whilst also being thinner.

 SUMMARY

Metals are  solid so have particles packed close together. These are in layers which mean that the outer electrons can move between them rather than being bound to particular atoms. These are referred to as delocalised electrons because of this.

Metals are therefore without their electrons so become positive ions. The metallic bond is actually the attraction between delocalised electrons and positive metal ions in the lattice. 

Metals have high melting points.

Metals have large regular structures with strong forces between the oppositely charged positive ions and negative electrons, meaning these must be overcome to melt the metal – this requires a large amount of heat energy. Transition metals tend to have higher melting points than the main group metals because they have large numbers of d-shell electrons which can become delocalised creating a stronger metallic bond. 

Metals conduct heat.

Heat is conducted if particles can move and knock against each other to pass it on. Delocalised electrons allow this to happen.

Metals have good electrical conductivity

Delocalised electrons can carry charge and move, the two requirements of electrical conductivity. Current can flow because of these delocalised electrons.

Metals are ductile and malleable.

Metals can be stretched and hammered into shape, making them ideal for things such as wires. Layered lattices mean that layers can slide over each other without disrupting the bonding – it is all still held together by the delocalised electrons and their strong attraction to the positive metal ions.

Being a solid, metal ions are packed closely together so they have a high density. 

 Happy studying!

The Name’s Bond ... Ionic Bond.

This is the first in my short series of the three main types of bond - ionic, metallic and covalent. In this, you’ll learn about the properties of the compounds, which atoms they’re found between and how the bonds are formed. Enjoy!

When electrons are transferred from a metal to a non-metal, an ionic compound is formed. Metals usually lose electrons and non-metals usually gain them to get to a noble gas configuration. Transition metals do not always achieve this.

Charged particles that have either lost or gained electrons are called ions and are no longer neutral - metal atoms lose electrons to become positive ions (cations) whereas non-metals gain electrons to become negative ions (anions).

The formation of these ions is usually shown using electron configurations. Make sure you know that the transfer of electrons is not the bond but how the ions are formed. 

An ionic bond is the electrostatic attraction between oppositely charged ions.

You need to know how to explain how atoms react with other atoms and for this the electron configurations are needed. You can use dot and cross diagrams for this. 

Ionic solids hold ions in 3D structures called ionic lattices. A lattice is a repeating 3D pattern in a crystalline solid. For example, NaCl has a 6:6 arrangement - each Na+ ion is surrounded by 6 Cl- and vice versa. 

Ionic solids have many strong electrostatic attractions between their ions. The crystalline shape can be decrepitated (cracked) on heating. Ionic Lattices have high melting and boiling points since they need more energy to break because atoms are held together by lots of strong electrostatic attractions between positive and negative ions. The boiling point of an ionic compound depends on the size of the atomic radius and the charge of the ion. The smaller the ion and the higher the charge, the stronger attraction.

Look at this diagram. It shows how atomic radius decreases across a period regularly. This is not the case with the ions. Positive ions are usually smaller than the atoms they came from because metal atoms lose electrons meaning the nuclear charge increases which draws the electrons closer to the nucleus. For negative ions, they become larger because repulsion between electrons moves them further away - nuclear charge also decreases as more electrons to the same number of protons.

Ionic substances can conduct electricity through the movement of charged particles when molten or dissolved (aqueous). This is because when they are like this, electrons are free to move and carry a charge. Ionic solids cannot conduct electricity.

Ionic compounds are usually soluble in water. This is because the polar water molecules cluster around ions which have broken off the lattice and so separate them from each other. Some substances like aluminium oxide have too strong electrostatic attractions so water cannot break up the lattice - it is insoluble in water.

Molecular ions such as sulfate, nitrate, ammonium or carbonate can exist within ionic compounds. These compounds may have covalent bonds within the ions but overall they are ionic and exhibit thee properties described above.

SUMMARY

When electrons are transferred from a metal to a non-metal, an ionic compound is formed.

Charged particles that have either lost or gained electrons are called ions and are no longer neutral - metal atoms lose electrons to become positive ions (cations) whereas non-metals gain electrons to become negative ions (anions).

The formation of these ions is usually shown using electron configurations. The transfer of electrons is not the bond but how the ions are formed.

An ionic bond is the electrostatic attraction between oppositely charged ions.

Ionic solids hold ions in 3D structures called ionic lattices. A lattice is a repeating 3D pattern in a crystalline solid.

Ionic solids have many strong electrostatic attractions between their ions. The crystalline shape can be decrepitated (cracked) on heating. 

Ionic Lattices have high melting and boiling points since they need more energy to break because atoms are held together by lots of strong electrostatic attractions between positive and negative ions.

The boiling point of an ionic compound depends on the size of the atomic radius and the charge of the ion. The smaller the ion and the higher the charge, the stronger attraction.

 Positive ions are usually smaller than the atoms they came from because metal atoms lose electrons meaning the nuclear charge increases which draws the electrons closer to the nucleus. Negative ions become larger because repulsion between electrons moves them further away - nuclear charge also decreases as more electrons to the same number of protons.

Ionic substances can conduct electricity through the movement of charged particles when molten or dissolved (aqueous). This is because when they are like this, electrons are free to move and carry a charge. Ionic solids cannot conduct electricity.

Ionic compounds are usually soluble in water because the polar water molecules cluster around ions which have broken off the lattice and so separate them from each other.

 Some substances like aluminium oxide have too strong electrostatic attractions so water cannot break up the lattice - it is insoluble in water.

Molecular ions such as sulfate, nitrate, ammonium or carbonate can exist within ionic compounds. These compounds may have covalent bonds within the ions but overall they are ionic and exhibit thee properties described above.

Shapes of Molecules

his post is more information than trying to explain something - the truth is, you just need to learn shapes of molecules like you do with anything. I’ve got a physical chemistry mock tomorrow that I’m dreading since I’ve done zero revision. The fact that I run a study blog yet don’t revise myself is odd, but what else can I do? Oh, wait … revise. So here it is, my last minute revision for myself and you too - I present, shapes of molecules!

VSEPR stands for valence shell electron pair repulsion theory. If you’ve ever seen a moly-mod or a diagram of a molecule in 3D space, you may wonder how they decided it was that shape. Well, VSEPR answers all.

The theory essentially states that electron pairs are arranged to minimise repulsions between themselves - which makes sense, since electrons carry the same charge and therefore try to repel each other. Of course, there are different types of electron pairs, lone and bonding. The strongest repulsions happen between lone pair - lone pair followed by lone pair - bonding pair and finally, bonding pair - bonding pair have the least repulsion. 

Since the repulsion governs the shape of the molecule, to work out a molecule’s shape you must look at dot and cross diagrams or electron configurations to see how a molecule is bonded. There are many methods to do this, but the bottom line is that you must work out how many bonding pairs of electrons and how many lone pairs are involved.

The easiest shape to learn is linear. This has two bonding pairs and no lone pairs at an angle of 180 degrees, since that is the furthest the two can get away from each other. Examples of linear molecules include carbon dioxide and beryllium chloride.

Next up is trigonal planar. This has three bonding pairs and no lone pairs, each at the angle of 120 degrees. Trigonal means three and planar means on one plane, this should help you in identifying the molecules since after a fourth pair of electrons, the shape becomes 3D. Examples of trigonal planar molecules include boron trifluoride and sulfur trioxide.

What if you were to have two bonding pairs and two lone pairs? Well, then you’d have a bent molecule. Water is a good example of a bent molecule. Since it has two lone pairs that repel the other two bonding pairs more than they repel each other, the bond angle is 104.5. I’d be careful though, as in many textbooks it shows a bent molecule to have one lone pair and a different bond angle.

Another variation of the bent molecule I’ve seen is the one with two bonding pairs and one lone pair. It is deemed as bent with a bond angle of 109 or sometimes less than 120 degrees.

Tetrahedral molecules have four bonding pairs and no lone pairs. The bond angle is 109.5 degrees. Examples of this include the ammonium ion, methane and the phosphate ion. A good thing to note here is how these molecules are drawn. To demonstrate the 3D shape, where the molecule moves onto a plane, it is represented with a dashed line and triangular line along with a regular straight line. 

Trigonal pyramidal, sometimes just called pyramidal, is where there are three bonding pairs and a lone pair. Bond angles are roughly 107 degrees due to the repulsion from the lone pairs. An example of a trigonal pyramidal molecule is ammonia, which has a lone pair on the nitrogen.

Having five bonding pairs gives a trigonal bipyramidal structure. I guess the three bonding pairs on the trigonal plane accounts for that part of the name, where the rest comes from the position of the remaining two. These molecules have no lone pairs and have a bond angle of 90 degrees between the vertical elements and 120 degrees around the plane. Diagrams below are much clearer than my description! Examples of this include phosphorus pentachloride.

Six bonding pairs is an octahedral structure. I know this is confusing because octahedral should mean 8 but it’s one of those things we get over, like the fact sulfur isn’t spelt with a ph anymore. It’s actually to do with connecting the planes to form an octahedral shape.There are no lone pairs and each bond angle is a nice 90 degrees. Common examples include sulfur hexafluoride.

Square planar shapes occur when there are six bonding pairs and two lone pairs. All bond angles are 90 degrees! They take up this shape to minimise repulsions between electrons - examples include xenon tetrafluoride.

The final one to know is T-shape. This has three bonding pairs and two lone pairs. These molecules have bond angles of (less than) 90 degrees, usually a halogen trifluoride like chlorine trifluoride.

There are plenty more variations and things you could know about molecular geometry, but the truth is, there won’t be an extensive section on it. It’s a small part of a big topic!

I’m not going to do a summary today since I’d just be repeating the same information (I tried to keep it concise for you guys) so instead I’ll just leave you with, 

Happy studying!

bernadette banner just referred to someone as her "resident chemist" and i am quite sure that this is the profession i aspire to.

there are some days where i'm like "hell yeah! how many moles of photons do i need to raise a 400 grams of water 5 degrees using a CO2 laser at a specific wavelength? i don't know but i'll figure it out!!"

and then there are some days where I forget how to balance the reaction of sodium bicarb and vinegar...