Crash Course Chemistry

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Atoms, Protons, Neutrons & Electrons: Crash Course Chemistry #1

Difference Between An Atom & A Molecule/Compound

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Unit Conversion & Significant Figures: Crash Course Chemistry #2

IU: The International System of Units is a system of measurement based on 7 base units. 

Every other unit is derived from these 7 base units. 

• Speed = length/time
• Acceleration = m/s2 (speed/time)
• Force = N = kg x m/s2 (F=M x A)
• Work (Joules) = J = kg x m2/s2 (F x distance)
• Power = W = J/s = kg x m2/s3 (work/time)

Exact vs Measured Numbers

When you’re doing experimental calculations, there’s two kinds of numbers. There’s exact and measured.

Exact numbers are like the number of seconds in a minute or the number of eggs in a dozen. They’re defined that way and thus we know them in effect all the way out to an infinite number of decimal places. If I say that there are a dozen eggs you know that that’s 12. It’s not 12.0000000001 or 11.9999999. It’s 12.

Measured Numbers:  Measured numbers you never ever know the exact number.  No matter how well I measure a car’s speed, I will never know it at the same level of precision that I know the number of eggs in a dozen. So the car could have been going 59.87390039 mph or 60.49321289 mph; the speedometer would still say 60. But with a measured number you just have to remember that the actual number goes out to infinite decimal places, you just never know all of them. You can’t. It’s impossible, so when my scale says 175 lbs, that doesn’t mean 175.000000 lbs. It means 175.something lbs.

Practical Implication of Measured Numbers

A measured number can be pretty unhelpful if you don’t have knowledge of the precision of the measurement. So you have to conserve the precision through your calculations or else you might end up killing someone with an imprecise dose of insulin or something. So we have a set of rules for what are called significant figures:

Significant Figures (Sig Figs)

These are the digits in your number that you actually know.

So when a number ends in a zero, or two or three zeroes, it’s hard to tell if those zeroes are significant.

Scientific Notation

This all gets so much simpler when you use scientific notation.

So 60 mph would instead be 6.0 x 10¹.

1 = ‘To the 1st power.’

We get that 0 is significant because we wrote it. Otherwise, it would just be 6 x 10¹. We keep that zero around because we actually know it.

The Exponent

The exponent of a number says how many times to use that number in a multiplication. It is written as a small number to the right and above the base number. In this example: 8² = 8 × 8 = 64 (= 8×8)

The number of the exponent just tells you how many places to move the decimal point. E.G: 

6.0 x 10¹

So to the 1st power you move it the decimal one to the right and you get 60.

To the negative 1st power you move the decimal point one place to the left and you get 0.60.

6.0 x 10^{5  =} To the fifth power you move the decimal place 5 zeroes to the right = 600,000

Your significant figures get preserved so 2.4590 x 10^-4 is 0.00024590 and you still get the same five sig figs.

How many sig figs should your answer have?

There are two simple rules for this.

Significant Figures (Addition)

If it’s addition or subtraction it’s only the number of figures after the decimal point that matters. The number with the fewest figures after the decimal point decides how many figures you can have after the decimal in your answer.

So 1,495.2 + 1.9903 = 1,497.1903

and then you round to the first decimal, because that first number only had one figure after the decimal.

So you get 1,497.2.

Significant Figures (Multiplication)

And for multiplication just make sure the answer has the same sig figs as your least precise measurement.

60 x 5.0839 = 305.034

but we only know 2 sig figs (60) so everything after those first two numbers (30) in 305 becomes zeroes making it round to 300.

Of course then we’d have to point out to everyone that the second zero but not the third is significant so we’d write it out with scientific notation: 3.0 x 10².

I know it feels counterintuitive not to show all of the numbers that you have at your fingertips, but you’ve got to realize: all of those numbers beyond the number of sig figs you have? They’re lies. They’re big lying numbers. You don’t know those numbers. And if you write them down people will assume that you do know those numbers. And you will have lied to them.

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The Creation of Chemistry – The Fundamental Laws: Crash Course Chemistry #3

Avogadro’s Law:

Equal volumes at the same temperature and pressure contain the same number of molecules (in reference to gaseous substances).

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The Periodic Table: Crash Course Chemistry #4

The periodicity of elements is a physical phenomenon. It’s a function of electrons. The periodic table is a representation of reality; a way of understanding and sorting the universe as it exists.

• The periodic table starts from the lowest atomic weight (H) to the highest row by row.
• Row #/Period # = how many shells there are
• Column #/Group # = how many valence electrons there are on the outermost shell
  •  E.G. Period 1 = H and He only have 1 shell with 2 or fewer electrons in their outer shell.

Alkali Metals (Very Reactive)

• Starting at the left, we have the soft, shiny, extremely reactive alkali metals, so reactive, in fact, that they have to be stored in inert gases or oil, to prevent them from reacting with the atmosphere.
• Alkali metals want nothing more than to dump off an electron and form a positive ion (cation).
• seeing as they’re so reactive, you don’t find hunks of them lying around in nature; instead, chemists must extract them from compounds containing them.

Alkali Earth Metals (Not As Reactive)

• Are reactive metals, but not as reactive as the alkali metals, for cations with two positive charges instead of just one. 

Transition Metals & Pre-Transition Metals (Fairly Unreactive) 

• The middle body area of the table is made up of a nice, solid rectangle of transition metals – these are the metals you think of as metal, with iron, and nickel, and gold, and platinum.
• The majority of elements are metals – they’re fairly unreactive, great conductors of heat, but more importantly for us, good conductors of electricity, they’re malleable, and can be bent and formed and hammered into sheets, and they’re extremely important in chemistry but overall surprisingly similar to each other.

Reactive Non Metals (Halogens) (Extremely Reactive)

• Extremely reactive gases that form negative ions (anions), with one negative charge, and love to react with the alkali and alkaline earth metals.

Metals, metalloids, gases, and nonmetals;

• The rectangle between the halogens and the transition metals contain a peculiar scatter shot of metals, metalloids, gases, and nonmetals; these guys don’t end up as ions unless you take extreme action and start shooting other ions at them,
  • Metal: solid at room temp (except for mercury)
  • Metalloids: possess characteristics of metals and non-metals
  • Nonmetals: are brittle solids and gain electricity easily

Lanthanides & Aactinides

• Down below, in their own little island, are the lanthanides and actinides, metals that were largely undiscovered in Mendeleev’s day because they’re so similar that it’s next to impossible to separate them from each other.

Noble Gases (Unreactive/Inert) 

• Also undiscovered when Mendeleev built his chart, the completely unreactive noble gases.

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The Electron: Crash Course Chemistry #5

Because building the 3d orbital requires a lot of energy, electrons actually go into the s-orbital of the fourth shell, 4s, before going into the third shell’s d-orbital.

These are the orbitals we know and all the shells that we’ve seen exist:

To figure out what order to fill them in, you just draw a diagonal line from the top right to the bottom left as you go.  So 1s first, then 2s, then 2p 3s, then 3p 4s, 3d 4p 5s, 4d 5p 6s, and so on.

• An interesting thing about d-orbitals, and the even bigger, more electron rich f-orbitals, is they don’t really need to be filled quite as much as the s and p’s, because they’re literally shielded beneath the s-orbitals of the next shell.
  • The s and p-orbitals I think of kinda like the trumpets and violins: it’s really terrible when they sound bad, but the base notes, deep and rich, hide a bit underneath the rest of the orchestra, just like the d-orbitals literally hide underneath the s-orbitals that have already filled above them.
  • Yes these incomplete orbitals affect them, but because their shielded, these middle of the chart elements are generally less reactive and happier to bump electrons along from atom to atom making them conductive, or just hanging out together in big masses of electron-sharing lumps of metal.

Ionization & Electron Affinities

So why are orbitals useful when it comes to understanding how an atom is likely to react?

• Well first, it really matters how much energy is required to remove an electron from an atom to form a positively charged ion. This energy is called “ionization energy.” [analogous to ionising radiation?]
• If there are several electrons being removed this is a step-wise process, starting with the electron at the highest energy level, the outermost one. Since the outermost electron has the highest energy, there’s the least energy necessary to remove it. More energy is needed to remove the second furthest one out and so on.
  • And of course when all the electrons in the outermost shell are removed there’s a really large energy jump necessary to remove an electron from the next shell down because that shell will be isoelectrically analogous to a noble gas.
• Just like how atoms are isotopically the same when they have the same number of protons and neutrons, atoms are isoelectrically the same when they have the same number of electrons.
• And just like there’s energy associated with removing an electron to form cations (positively charged ions), there’s energy associated with adding electrons, usually to fill an orbital to achieve a stable two or eight electron shell configuration.
• Just like with the ionization energy there’s a discrete energy jump involved with the adding of an electron. That energy is called “electron affinity.” 

 Big Picture

There are, a number of everywhere-permeating fields in our universe. One of those, is the electron field.

The Electron Field:

• In order for an electron to exist, there has to be an excitation of the electron field and we can describe those excitations as waves, just as a wave in the ocean is an excitation of the water. At any given moment, the electron can be anywhere within the function of the wave.
• But waves are defined not by harsh boundaries, instead they’re strong in some areas and weak in others. The strength of the wave at one certain point in space determines how likely it is that you will find an electron there at any given time if you measure.

And so if we’re trying to understand reality we should not think of electrons as circling around the nucleus of an atom like planets around a star,

but instead as an excitation around the nucleus and the shape of that excitation is the orbital.

Orbitals are precisely the reason that everything exists. They are the root and the key and the nexus and the crux and the keystone and every other metaphor of not just chemistry but existence.

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Stoichiometry – Chemistry for Massive Creatures: Crash Course Chemistry #6

Why this is important?

• So mass is how we massive beings tend to understand the world; in our day-to-day dealings with substances, we need to have some sense of how much of it there is before we can use it or predict how it’s going to act.
• E.G. Chemistry will be happy to tell me that the atomic structure of the sugar in this packet is 12 carbon atoms, 22 hydrogen atoms, and 11 oxygen atoms in every molecule. But I don’t have any idea how many molecules of sugar I want to put in my tea! Or how that one molecule will react with other chemicals in my body.
• To understand that kind of stuff, I need to know the mass of the sugar that I’m dealing with. In other words, I need to measure it. And that, is why there’s stoichiometry – the science of measuring chemicals that go into and come out of any given reaction.
  • In Greek, it literally means measuring elements, and, in essence, it allows us to count up atoms and molecules by weighing them.
• Stoichiometry, yes, contains a fair bit of math, but it’s one of the most important decoders that we have as chemists. It’s what we use to translate from the very small to the very big, to parley the stuff that we can’t see into the stuff that we can. And because of that, chemists use it all the time.

Relative Atomic Mass

• We already have a way to measure elements and that is by using relative atomic mass = The average atomic mass of all of the naturally occurring isotopes of a given element.
• So for example, all of the natural carbon on earth occurs as one of three and only three isotopes: C-12, C-13, and C-14. They all have six protons, but the number of neutrons varies. And these isotopes show up on our planet in totally different proportions. So the relative atomic mass of carbon is a weighted average of these three masses, which comes out to 12.01 amu. But 12.01 what?

This is how we weigh atoms:

Atomic mass units (amu) = 1/12th of the mass of an atom of ¹²C (‘carbon 12’)

But I don’t know how many amu’s of these molecules together are going to make this tea good to me, or how many other molecules of sugar I can consume while maintaining my slim yet robust physique. In order to make these calculations and predict reactions, I first need to be able to convert the atomic mass of this sugar, into a standard amount of substance. Not weight, not volume, just purely, objective amount of stuff.  That, my friends, is what moles are for.

Moles

A mole is arguably the most important unit in all of chemistry, because it allows us to express a chemical’s atomic mass in terms of grams. 

Defining A Mole: Avogadros Number

• To define what a mole is, no matter what it’s a mole of, we use,¹²C.
• There are 6.022 x 10²³ atoms in 12 grams of ¹²C and by definition, that number of anything is a mole of that thing.

So there are this many carbon atoms in a mole of carbon-12 and there are the same number of anything in a mole of anything else.

• Like a dozen roses is twelve roses, but a mole of roses is 6.022 x 10²³ roses
• A mole of sand would be 6.022 x 10²³ grains of sand

But in chemistry the thing to remember is this:

A mole of any element contains 6.022 x 10²³ atoms of that element no matter what. This is what lets us translate number of atoms into grams. It lets us weigh elements.

Examples:

• One mole of carbon-12 contains 6.022 x 10²³ atoms and weighs 12 grams, right? So one mole of oxygen also contains 6.022 x 10²³ atoms but because oxygen atoms are more massive it weighs 16 grams and you’ll recall that oxygen’s relative atomic mass is 16 amus.

The number of atoms per mole remains the same, but the mass of a mole depends on the average mass of the element. This simply means that one mole of any element equals its relative atomic mass in grams. 

• So now you’ve got it, 1 mole of hydrogen weighs 1.008 g, a mole of iron is 55.85 g, and a mole of natural carbon is 12.01 grams. This is known as an element’s molar mass.
• And now that we know the molar mass of elements we can calculate the molar mass of any compound. All we have to do is add up the molar masses of its component elements.

1. For instance, the formula for this sugar or sucrose is C₁₂H₂₂O₁₁. One mole of sucrose, by definition contains 6.022 x 10²³ molecules, and since each molecule contains 12 carbon atoms and 22 hydrogen atoms and 11 oxygen atoms, then one mole of sucrose contains 12 moles of carbon, 22 moles of hydrogen, and 11 moles of oxygen.
2.  Multiply the number of moles of each element by its molar mass and add them all up, that’s the molar mass of the whole compound.

How did they figre out 12 mol, 22 mol and 11 mol? 

See, the mole is like our chemical Rosetta Stone, with it, we can translate anything from the level of atoms and molecules to the level of grams and kilograms.

Let’s say, y’know, hypothetically, that I’m watching my weight, so I want to know what it’ll take for me to burn a certain amount of sugar that I consume. That’s a reaction!–and it’s a pretty simple one, my body uses sucrose by combining it with oxygen to create energy plus CO₂ and H₂0 as waste.

You can write this out as an equation, in which the reactants combine on the left to yield the products on the right. But there’s a problem here: this equation doesn’t reflect chemical reality.

Equation Balancing

• During a reaction, bonds are broken and new ones are formed but the number of atoms of each element remains the same.
• The sugar and oxygen molecules may be busted apart and mixed up but the number of each kind of atom that you start with ends up being exactly the same after the reaction = conservation of mass.
• So when writing a reaction out as an equation the number of atoms of each element has to be exactly the same on both sides. Reconciling the reactants with the products is called equation balancing and it’s a good bit of what stoichiometry is all about, because from a chemical perspective an unbalanced equation is pretty useless – it doesn’t tell you how much is going in and how much is coming out.

How do balance an equation?

Dealing with the C’s:

The best way is to start with the most complicated molecule, which in this case is, of course, the sucrose. For every molecule of sucrose that goes into the reaction, you know that you’re gonna have twelve carbon atoms so right off the bat you know that you’re gonna have to end up with at least 12 molecules of CO₂ as a product, because that’s the only molecule where those carbon atoms end up.

Dealing with the H’s:

• Now let’s deal with the hydrogen, because that also shows up in only one molecule on both sides of the equation so that’s easier.
• You know that at least 22 atoms of hydrogen go into the reaction and the product contains some multiple of 2 hydrogen atoms (that’s the H₂ in the water molecule), so if there were eleven water molecules produced that would balance the hydrogen with 22 hydrogen atoms on each side.

Dealing with the O:

Finally, the oxygen. Since we know we have 12 CO₂ molecules and 11 water molecules as products so far, we also know that we’re gonna end up with thirty-five oxygen atoms.

If you look at your reactants, on the left, you see that you have 11 oxygen atoms in the sucrose molecule and 2 in the molecular oxygen, O₂.

The carbon and hydrogen are balancing nicely with only one molecule of sucrose so let’s leave that alone but there could be any number of paired oxygen atoms involved.

Since you need 35 and you know you have 11 to start with in the sucrose you just need 24 more, which would equal 12 molecules of O₂ 

And now, the equation is balanced! You know exactly what my body is producing. For every molecule of sucrose I’m metabolizing I have to inhale 12 molecules of oxygen and in return, in addition to a little sugar buzz, I’ll produce 12 molecules of carbon dioxide and 11 molecules of water. This is incredibly useful in helping us to understand the proportions of chemicals as they react at the molecular level.

Molar Ratios

• But in a lab, or in life, you have to work with measurable amounts of stuff, so the last stoichiometric trick you need up your sleeve is to calculate specific masses of the reactants and products.
• So for instance, how much oxygen will I need to inhale in order to burn 5 grams of sugar? To figure that out, we just need to focus on the left part of the equation, because we only need to quantify the reactants. First, convert your balanced equation into molar masses; in order to get from molecules to grams, you need to go through moles first. (NOTE: In the video, it says 348 g of O₂, it should be 384 g. This is just a typo, all calculations from there on are correct.) 

When you figure out the molar masses, you see that the ratio of sucrose to oxygen is actually pretty close: 384 grams of oxygen for every 342.3 grams of sucrose.

Then you simply compare this ratio to the masses of reactants in your experiment, 5 grams of sugar to X grams of oxygen, and hopefully you know how to solve for X.

For every 5 grams of sugar I ingest I’ll need to inhale 5.6 grams of oxygen, which I happen to know is about 35 breaths’ worth.

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