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Polarity, Resonance, and Electron Pushing: Crash Course Organic Chemistry #10:

We’ve all heard the phrase “opposites attract.” It may or may not be true for people, but it’s definitely true in organic chemistry. In this episode of Crash Course Organic Chemistry, we’re learning about electronegativity, polarity, resonance structures, and resonance hybrids. We’ll practice a very important skill for this course that will help us avoid a lot of memorization in the future: electron pushing. It’ll be a lot of trial and error at first, but we all start somewhere!

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.

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!

Covalent Bonds: Sharing Is Caring!

Welcome to my second out of three posts on bonding - ionic, covalent and metallic. This post also covers the coordinate/ dative bond which I can’t remember if I’ve covered before. Only one more of this series left! Find the others here.

Covalent bonding involves one or more shared pairs of electrons between two atoms. These can be found in simple molecular elements and compounds like CO2 , macromolecular structures like diamond and molecular ions such as ammonium. Covalent bonds mostly occur between non-metals but sometimes metals can form covalent bonds.

Single covalent bonds share just one pair of electrons. Double covalent bonds share two. Triple covalent bonds share three.

Each atom usually provides one electron – unpaired in the orbital – in the bond. The number of unpaired electrons in an atom usually shows how many bonds it can make but sometimes atoms promote electrons to fit in more. Covalent bonds are represented with lines between the atoms – double and triple bonds represented with two and three lines respectively.

Dot and cross diagrams show the arrangement of electrons in covalent bonds. They use dots and crosses to demonstrate that the electrons come from different places and often only the outer shell is shown.

The simple explanation as to how atoms form covalent bonds is that one unpaired electron in the orbital of one atom overlaps with one in another atom. Sometimes atoms promote electrons in the same energy level to form more covalent bonds. For example, if an atom wants to make three covalent bonds but has a full 3s2 shell and a 3p1 shell, it can promote one of its 3s2 electrons so that an electron from the other atoms can fill the 3s shell and pair with the new 3p2 shell.

Sometimes promotion does not occur and that means different compounds can be made such as PCl3 or PCl5.

A lone pair of electrons is a pair of electrons from the same energy sub-level uninvolved in bonding. Sometimes these can form something called a coordinate bond, which contains a shared pair of electrons where both come from one atom. The lone pair of electrons is “donated” into the empty orbital of another atom to form a coordinate bond.

This is an example of a coordinate (sometimes called dative) bond between ammonia and a H+ ion which has an empty orbital. The lone pair on the ammonia overlaps with this H+ ion and donates its electrons. Both electrons come from the ammonia’s lone pair so it is a coordinate bond. This is demonstrated with an arrow. The diagram is missing an overall charge of + on the ammonium ion it produces. Coordinate bonds act the same as covalent bonds.

Once you have your covalent bonds, you need to know about covalent substances and their properties. There are two types of covalent substance: simple covalent (molecular) and macromolecular (giant covalent).

Molecular simply means that the formula for the compound or element describes exactly how many atoms are in one molecule, e.g. H2O. Molecular covalent crystalline substances usually exist as single molecules such as iodine or oxygen. They are usually gases or liquids at room temperature but can be low melting point solids.

Solid molecular covalent solids are crystalline so can be called molecular covalent crystals. Iodine and ice are examples of these. Iodine (shown below) has a regular arrangement which makes it a crystalline substance and water, as ice, has a crystalline structure as well.

The properties of these crystals are that they have low melting points, are very brittle due to the lack of strong bonds holding them together and also do not conduct electricity since no ions are present.

The other kind of covalent substance you need to know is macromolecular. This includes giant covalent structures such as diamond or graphite, which are allotropes of carbon. Non-metallic elements and compounds usually form these crystalline structures with a regular arrangement of atoms.

Allotropes are different forms of the same element in the same physical state.

Diamond is the hardest naturally occurring substance on earth therefore is good for cutting glass and drilling and mining. It has a high melting point due to the many covalent bonds which require a lot of energy to break. Each carbon has four of these bonds joining it to four others in a tetrahedral arrangement with a bond angle of 109.5 degrees and it does not conduct electricity or heat because there are no ions free to move.

Graphite, on the other hand, can conduct electricity. This is because it has delocalised electrons between the layers which move and carry charge. Carbon atoms within the structure are only bonded to three others in a hexagonal arrangement with a bond angle of 120 degrees. Since only three of carbon’s unpaired electrons are used in bonding, the fourth becomes delocalised and moves between the layers of graphite causing weak attractions, explaining why it can conduct electricity.

Graphite’s layered structure and the weak forces of attractions between it make it a good lubricant and ideal for pencil lead because the layers can slide over each other. The attractions can be broken easily but the covalent bonds within the layers give graphite a high melting point due to the amount of energy needed to break them.

SUMMARY

Covalent bonding involves one or more shared pairs of electrons between two atoms. Covalent bonds mostly occur between non-metals but sometimes metals can form covalent bonds.

Single covalent bonds share just one pair of electrons. Double covalent bonds share two. Triple covalent bonds share three.

Each atom usually provides one electron – unpaired in the orbital – in the bond. The number of unpaired electrons in an atom usually shows how many bonds it can make but sometimes atoms promote electrons to fit in more. Covalent bonds are represented with lines between the atoms.

Dot and cross diagrams use dots and crosses to demonstrate that the electrons come from different places and often only the outer shell is shown.

The simple explanation as to how atoms form covalent bonds is that one unpaired electron in the orbital of one atom overlaps with one in another atom. Sometimes atoms promote electrons in the same energy level to form more covalent bonds. 

Sometimes promotion does not occur and that means different compounds can be made such as PCl3 or PCl5.

A lone pair of electrons is a pair of electrons from the same energy sub-level uninvolved in bonding. Sometimes these can form something called a coordinate bond, which contains a shared pair of electrons where both come from one atom. The lone pair of electrons is “donated” into the empty orbital of another atom to form a coordinate bond.

The formation of ammonium is an example of this.

There are two types of covalent substance: simple covalent (molecular) and macromolecular (giant covalent).

Molecular simply means that the formula for the compound or element describes exactly how many atoms are in one molecule, e.g. H2O. Molecular covalent crystalline substances usually exist as single molecules such as iodine or oxygen. They are usually gases or liquids at room temperature but can be low melting point solids.

Solid molecular covalent solids are crystalline so can be called molecular covalent crystals. Iodine and ice are examples of these. 

The properties of these crystals are that they have low melting points, are very brittle due to the lack of strong bonds holding them together and also do not conduct electricity since no ions are present.

Giant covalent structures such as diamond or graphite are allotropes of carbon. Allotropes are different forms of the same element in the same physical state.

Diamond has a high melting point due to the many covalent bonds which require a lot of energy to break. Each carbon has four of these bonds joining it to four others in a tetrahedral arrangement with a bond angle of 109.5 degrees and it does not conduct electricity or heat because there are no ions free to move.

Graphite can conduct electricity. This is because it has delocalised electrons between the layers which move and carry charge. Carbon atoms within the structure are only bonded to three others in a hexagonal arrangement with a bond angle of 120 degrees. Since only three of carbon’s unpaired electrons are used in bonding, the fourth becomes delocalised and moves between the layers of graphite causing weak attractions, explaining why it can conduct electricity.

Graphite’s layered structure and the weak forces of attractions between it make it a good lubricant and ideal for pencil lead because the layers can slide over each other. The attractions can be broken easily but the covalent bonds within the layers give graphite a high melting point due to the amount of energy needed to break them.

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!

14/100 days of productivity: Worked all day long. I did 2 DM ( long and more difficult exercice than usual ) today. My boyfriend helped me for the second and we are quite a nice team together.

im gonna kill myself

i love describing acid base buffers as "they turn the extra pH-ey things into the less pH-ey things."

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...

*waters plants, chugs three cups of oversweetened black tea, pulls up six half-read, poorly annotated papers, knits three rows of a frog stuffie, re-reads prompt and guidelines, tries to command f keywords in the paper, finds nothing, re-reads the abstract, realizes all the papers are useless, goes onto googles scholar and searches for the same keyword, finds new papers, reads abstract and intro, finds something super interesting, reads discussion and realizes the methods for this study were slightly different than what you were expecting, has an existential crisis, wonders why tf they're doing genomics when they're interested in hydrocarbons, sheds one tear, shrugs, types furiously and writes it up anyway with a quick sentence explaining that it's slightly different but still relevant, cries, does citations and slams computer shut, stares at the clock which has somehow gone from 11AM to 6PM in the blink of an eye, and proceeds to not sleep for another ten hours*

"yEaH, i'M a ScIenTiSt."

cool facts

- AgNO3 (silver nitrate) clots blood! it's actually got several medical applications! (and is often in military first aid kits!)

- when oil spills happen, the bacteria which can metabolize polycyclic aromatic hydrocarbons (PAHs) tend to rise in population, and those that cannot tend to decrease!

intro chem tips (high school)

#intro

welcome to ramblings on how to survive your first chem class, from someone who’s taken 3 3/4* chem classes and is still kicking.

*1/2 of an outside course and 1/4 of comparing notes with friends taking a companion chem class)

reminder before we start

- not everything works for everyone, so you may have to do some finagling, and that's okay.

- it's important to ask your teacher for help too. asking for help doesn't mean you're dumb, it means you take responsibility for your own understanding, and that is impressive.

- chemistry classes are hard, but they are not a judgment on your soul. you will mess up sometimes, and it will be okay.

#basics

calculations

- remember when your elementary school teacher told you to put units on your work? yeah? okay. it's coming back to haunt you, except now, with algebra. aka, 2000 times worse.

- show. your. work. i don't care how smart you are, or how well you can do mental math. this isn't math, it's chemistry, and we have an added layer of real world context. building on the units bit; if you don't write out what you're doing, you won't know what your calculations are doing, and you'll have to do it all over again. (note: once you're onto your next year or two of chem, you won't always need to write out simple molecular mass calculations, or going from grams to moles. but otherwise, label them well, and ensure that future you knows what you're looking at.)

- get a good calculator. one that you like. one that you will take good care of. one that you can spend the rest of your high school (and maybe college) career with. one that you'll marry. okay, maybe not that. but the point still stands. get a calculator you like, and become well acquainted with it. in the words of my ChemE teacher, "many tears have been shed over mistyped calculations." get a good one, practice using it, and optionally, put nice stickers on it.

- brush up on your basic math. i'm talking like, the elementary school stuff. addition, subtraction, multiplication, division. and also some of the other stuff, like logs and exponents, and general algebra. also, a key concept that has saved my ass many times doing equilibria—if the numerator of a fraction is large, the final number it divides out to will be large, and if the denominator is large, the final number will be small. or, knowing that log of 1 is 0. simple algebra/basic function stuff will help you gut-check your answers! (this is actually part of why i love chem; because of the real world context, you can almost always tell if your answer makes sense, and therefore, the algebra can actually be *gasp* fun!)

labs

- use metric. for those americans out there, this might feel like it sucks. coming from a fellow american—it doesn't. it actually makes a hell of a lot more sense. and all of most science is done this way. so suck it up and get used to it.

- get a watch. it doesn't have to be expensive, but something that has a second hand is useful. for a few reasons. a) measuring seconds is important for some labs, and you don't wanna use your phone because reason b) you don't wanna spill chemicals on and/or contaminate timing tools like stopwatches and phones with whatever you're working with, yet you don't want to take your gloves off.

- when you measure liquids, measure the meniscus. (the "vertex" of whatever parabola shows up where the liquid meets air.) generally, it's concave (upside down rainbow shaped) so you wanna make sure that the lowest point is what you're measuring. in order to see this, you will have to get down to eye level of whatever you're measuring. this... can get interesting. in the words of my lab partner while doing titrations "sweet jesus i feel like my back has aged sixty years in the course of twenty minutes." so. that may kill your back. but it will give you an accurate measurement :D

- the difference between precision and accuracy. when you measure things in chemistry, generally, we take measurements in triplicate (3x), just to make sure we are actually getting the right value, so that if we f*** up once, it doesn't matter as much. but what is the right value? well, if the value we measured is close to the general accepted value of the number (say, if the solution is hypothetically 0.5M, and you got an average of 0.48M from your trials), then we were accurate. if we got 0.30M three times, that means we were very precise, or that our measurements were very close together. generally, we want to be both precise and accurate when doing labs.

- wear your safety goggles and gloves, wear close toed shoes, and tie your hair back—follow the safety procedures. especially the gloves. i can't stress how important this is. for a few reasons. first of all, the obvious. if you are working with highly concentrated hydrochloric acid (a very strong acid), you do not want to accidentally kill your hand/your eyes. i do not recommend being blinded by chemicals. second. contamination. this applies more to biology, but we're gonna talk about it anyway. let's say you're dealing with a bacterial culture, and you get some on your hands. your hands are going to touch a bunch of stuff afterwards, and you might accidentally turn not only your lab, but that bus you took home, and your house into a huge culture surface. unrecommended. third, just straight up peace of mind about the first two. (i spilled low concentration sodium hydroxide on my arm, and even though we took care of it, i was paranoid about that spot of skin itching for the next four days. science classes can be difficult, and none of use need any extra anxiety.)

- label your beakers. i cannot stress how frustrating it can be to be halfway through a lab with ten different beakers half filled with clear liquids, and to then suddenly go "wait, which one is this?" and have to restart the whole thing, because it'd be worse to pour an unknown substance into something else, lest it be the wrong one. also, if you're doing serial dilutions (where you repeatedly dilute something to get specific concentrations of a substance) all the tubes will look the same. and they suck to repeat. so please. grab a sharpie and some tape, and just label it. (pro tip: fold over the edge of the tape on itself, so it's easier to get off when you clean up)

- this one should be pretty obvious, but don't use something if you don't know what it is. avoid unlabeled objects like the plague. (see: potassium metal and water are highly reactive. you think you see some copper metal strips for your galvanic cell? cool, let's just take that and stick it in this beaker of water. oh. oh no. shit, now there's a huge smoke cloud and i think it exploded.) in short, avoid this by making sure you know what you're using, and in what concentration. there is a very large difference between 0.1M hydrochloric acid (very concentrated) and 0.00001M hydrochloric acid (not so concentrated, mostly water.) namely, that 1M hydrochloric acid has a pH of 1 (very very acidic), and that 0.00001M hydrochloric acid has a pH of 5 (like coffee.)

habits/best practices

- take notes. preferably by hand. in the words of my intro chem teacher "this is a hill i will die on." i'm inclined to agree with him. chemistry has a lot of equations that cannot be easily typed out, and on top of that, it is often useful to annotate your calculations later, either with the logic behind it, or with mistakes you've made. some people *cough that one kid in my chemE class* manage to use latex software in order to do their calculations, but from my perspective, it just doesn't make sense; your teacher will be writing on the board, most of your classmates will be doing things on paper, you probably don't want to use your computer during the labs for fear of spilling chemicals on it, and your tests will probably be on paper. all in all, i would recommend taking notes by hand (if you're an ipad kid, that's also fine). for those of you who can do it on a computer, you have my respect, y'all are built different.

- pick a lab partner you can work well with. if you can. come early on the first day of class, and choose carefully who you sit next to. some teachers will say that that person is your lab partner for the rest of the year. general rules of thumb for picking a lab partner boil down to having ARC with them. A) you have an amicable relationship with them—you may not want to pick your bestie though. as i have learned the hard way, not all of your friends are good academic partners. R) they're reliable. they turn things in on time, and you can be sure they'll have your back in whatever shenanigans you get into. C) you can communicate with them. arguably the most important point, 'cuz if you can always work on your friendship and boundaries/expectations using communication, but not always the other way around. good luck, and pick a partner you have ARC with :) (we'll expand on this in the lab partner section)

#lab partners

intro

earlier we discussed having ARC relationship with your lab partner, which is an amicable, reliable, and communicative relationship. see best practices of #basics.

misc tips

- get their phone number for easy communication

- have an established way of sharing files, be it email, text, or a shared folder.

- ask them about their work style- do they like to do things early, or closer to the deadline. this can be helpful to know if you have different work styles, so there's no frustration or anxiety stemming from miscommunication.

flag system - an expansion of the ARC system

green flags and their reasons

1) they're nice - you don't want to work with someone you dislike

2) they pay attention in class - you want to have a lab partner who understands the activity and doesn't goof off.

3) they respect the teacher - hopefully the teacher is very neutral and doesn't treat students differently based on how respectful they are. unfortunately, this isn't always true. while i'm not saying you should work with the teacher's favorite, it can be awkward to work with someone who disrespects the teacher, and as a result, causes the teacher to look at your work unfavorably.

4) they're a reliable student (they do work, they turn things in on time, they will communicate and advocate when they can't.) - it's nice to have a partner you can depend on, but most importantly, to know when you can't expect them to fulfill their part. sure, it sucks when someone else doesn't do their work, but it's even worse when they don't tell you, and you're left hanging the day of.

5) they'll communicate with you - see earlier point about being left hanging. in addition, clear communication can help set expectations, and help both of you feel better about making constructive criticism about each other's work, and make both of you better.

6) they show up to class - it sucks to be stranded without your partner, cuz then you do 2x the work, so good to know your partner will be there when you need them.

red flags and their reasons

1) they don't communicate with you, and won't tell if you if they need to bail on something - see green flag 4.

2) they wait till the last minute for assignments - if you're an anxious bean, you may do everything early. having your partner do their work last minute may add to your anxiety and cause an "should i do it, cuz it's like 11:00pm?!?" dynamic, which is awkward.

3) they don't take feedback from you/teacher (assuming it's politely and respectfully given) - it can be difficult to feel confident about your work if your partner isn't okay with having their work modified. however, some tact is required, as no one wants to feel criticized. i recommend things like saying "hey, i was wondering if we could review x section, i had some ideas for *insert thing in x section,* but since i know we agreed to split up the work, and we decided that was your section, i wanted to get your input."

dividing work

- often, it can be difficult to work efficiently in pairs if you don't do some level of "divide and conquer," especially if you're working asynchronously. in order to make sure everyone participates a little bit in each section, my partner and i often agree on a solution where, say, i'll take the intro and the data section, along with crunching it into graphs, and they'll take the methods and results section with the calculations. then, we'll meet up later, and go over each other's sections, and make sure everything is in order. then, we'll both do a final readthrough to ensure everything seems seamless. if you have an ARC relationship with your lab partner, this should work pretty well.

troubleshooting

lab partner issues, are, unfortunately, all too common. i have yet to find a perfect partnership, only better ones. here are some trouble shooting tips for when things go awry.

make sure you divide work. it makes it easier to negotiate in situations when someone doesn't hold up their half of the bargain.

if you're having trouble communicating with them, finding a time to chat with them, say, after class, can be useful. general tenets of good communication apply here, and are listed below.

a) not accusing them, but rather explaining which actions, or things you have percieved, have frustrated/angered/saddened you.

b) proposing solutions. having a lot of problems without solutions can be frustrating. your solution can be "we shouldn't be lab partners anymore."

c) listening to them, too.

if they aren't doing work, talk to them about it, ideally early, and in a gentle manner. try "hey, i noticed that you haven't done the section we agreed to have you do last week. do you need any help, or are you planning on doing it later?" this give them an opportunity to explain they've got a different work style, or that they don't know what's going on.

alternatively, if this is a repeated problem, go to the teacher and let them know. try "hello [teacher name], i wanted to talk about some of the lab reports we've been doing lately. i just wanted to let you know that i've been doing most of the work, even though i've asked my partner about finding a more equal distribution of work. given that this work is designed for two people, i've been finding it to be a bit much, and was wondering if i might be able to split the sections with my partner, and label the ones i've done to show that i've done my half. if i do this, would you consider not grading it incomplete for me, if i've done my half of the work?"

if they don't show up, i'd also recommend talking to your teacher and asking if you can be added to another partnership or be given a modified version of the work. especially with this type of problem, your teacher will likely understand and agree that it's difficult for one person to keep up with a workload designed for two people.

#tests

preparation

- practice problems, practice problems, practice problems. these are the key to doing well on tests. after all, most questions aren't going to be just theory, they're going to ask an applied question. practice problems will help you learn to turn words into equations, and to be familiar with any strange formats your teacher may throw at you. my personal recommendation would be to do the studyguide (if you are fortunate enough to have one), redo some homework problems, and then turn to google. literally anything you can find on google (as long as it's actually correct.) the weirder and stranger and more difficult the problems are, the better the chances of you actually understanding the core of the content, and being prepared for any weirdness your teacher gives you. just grind them for a few hours (take breaks though!) and you'll find that at the very least, the basic calculations are easier, so even if it takes you a while to figure out the individual steps in a larger question, the component steps will be easier. (also, write up generalized steps of how to solve problems. eg: first, divide by this, then convert to xyz, and then take the negative log, etc. these generalized steps can help if you ever get stuck in the thick of a problem)

- bang your head against the wall memorization. sometimes you will have to do this, especially with equations. my recommendation? chant like you are summoning a demon. especially if there are mneumonics. there are amazing. just chant until you think the demon is intentionally ghosting you because of how many times you've repeated the summoning. (to this, quizlet is also very helpful for non-formula things, like solubility rules.)

- memorize, don't derive. going back to equations here, you will not have time to rederive the whole equation during a test. memorize it, and get if out of your system. efficiency is key.

- fix your sleep schedule. i know, it sounds ridiculous, but make sure you get good sleep the entire week before for best retention and performance.

the day of

- go through your useful formula, chant anything you were supposed to memorize, and look at some practice problems. you may not have time to solve them, but look at the problems, and walk yourself through how you would solve them.

- bring a small fidget, like a hairband. you may not be allowed to have a real fidget like a spinner or cube, but having a small outlet for nervous energy can be helpful.

- spend time with people who are calming. if this is nobody, that's okay. find a nice spot by yourself, put in your headphones, and breathe. (try box breathing.) if you have friends in the class, try not to let yourself listen to them talk about the test, as their nervousness can impact your ability. (honestly, for more on this, just search up sports psychology stuff, alternatively, https://bulletproofmusician.com/ is a great site with lots of helpful tips for performing under pressure.)

- eat well the day of, and don't up your caffeine intake, especially if you're anxious. too much caffeine can make you extra jittery, which can detract from focus.

- bring a watch to measure time.

- if you feel yourself panicking during the test, breathe a little bit, close your eyes, re-read the question, gut check your answer, and check your algebra.

#resources

- your teacher. your teacher should always be your primary resource. ask them for help, extra problems, or even tangentially related knowledge. a good relationship with your teacher can make your chem experience exponentially better.

what to do if your teacher sucks/can't provide the resources you need/is busy/literally any other reason they're not a good resource for you. listed in the order i would recommend them. (but not between online resources and human resources; both are great, preference depends on the person.

online resources

- khan academy. khan academy has great resources, and a semi-structure curriculum, which is great if you're stuck. they've got practice problems, video explanations, and articles.

- chemistry libre texts. literally my savior. unfortunately, tends to be rather surface level, but great for clarifying questions.

- ap chem/gcse chem/a-level chem practice tests/textbooks/websites. these only cover specific sections of chemistry material, but given that it's an intro class, you should be able to find all you need here.

- youtube. further into void territory, and not always correct, but good if you're on something somewhat obscure/have run out of khan academy stuff.

- googling the topic you need help with. the most sus.

human resources

- TAs. not every class has them, but they're also great at explaining things, and if they've taken the class, can give you helpful tips and tricks. they're also usually less intimidating then the teachers themselves.

- tutors. obviously, these cost money, and are not available to everyone, but if you are able to get one, they are an amazing resource.

- upperclassmen. while they usually can't reteach you the curriculum, they can usually give you quick tips and tricks for solving quickly, or tips on the teacher's style. take their advice with a pinch of salt though, because there's quite a possibility they remember something wrong.

- discords, specifically those geared towards homework help. obviously, the advice is being given by other people, so it can be a little sus, but it is much better than nothing. take their advice with a pinch of salt though, because there's quite a possibility the are still students themselves, or are just synthesizing information they've learned after a quick google.

#closing

don’t stress out! (you will inevitably, but try to mitigate it.) you’ll be fine. have fun! chemistry is fun! channel your inner mad scientist. good luck. remember that this class is not a judgement on your soul. take breaks. hydrate. and lastly, remember your units.

really trying my best to get into organic chemistry 2 studying but i’m so over it i need to start pchem and continue with my other math courses like im tired

okay so for anyone that got an A in physical chemistry … HOW? i was originally pre health but i want to go into research now, so my physics is only algebra based bleh but so far have gotten A in my calculus classes (will be taking calculus 3 spring)

techinally they are my last major class excluding physics 2 and an elective, i have seen the math of physical chemistry it looks interesting but very daunting . SOS send help please with any advice !!!

Did I skip my lecture to work on my lab report that I didn't do this weekend...? Yes. But at least I'm getting it done. Been feeling incredibly tired lately, however.

Today's goals:

● Finish Lab Report

● Take notes on lecture slides

● Sleep early

Sometimes you just have to lower your expectations a little and take care of yourself I guess.

🎼 : Blue World - Mac Miller

📖 : To Kill A Mockingbird - Harper Lee

woaaahhhh sig figs

AAAAAAHHHH I HAVE A MATH TEST TMRW

Normally, I’m not so stressed, but I missed class this week (I was skipping shame on me 😭) and I have no clue what’s going on… Me and the math teacher are besties because he’s the coach of the golf team, so I’m scared of disappointing him with my bad test scores.

I also had a long lab in chem this week. It’s really sad because we got bad results even though we ran our trial for over and hour lol. I think it’s my lab partner’s fault because he keeps on putting his crusty fingers all over the cuvettes which is messing up the spectrophotometer. Hopefully he actually writes his part so it’s not another repeat of last year.

- Practice midterms for matrices

- Grade practice midterms

- Write lab report

- Debug binary system simulation

- Work on astro research paper

- Work on program applications

- Finish new deal research

- Finish college assessment

- Email people back 😬

Don’t ask me why we have a midterm in January… I have no clue.