## Intro [00:00] [00:00] Everything is made of atoms. Yes, even you. Atoms consist of a core and some electrons. The core is made of protons and neutrons. And depending on the number of protons, you [00:07] get different elements. Water is made of Hydrogen and Oxygen. This is some Sodium. Hm, I wonder what happens when you mix them…oh, whoopsie. ## Valence Electrons [00:16] [00:20] Quantum mechanics tells us that this is not what atoms actually look like, they look more [00:24] like this, but we’ll get to that later. For now, just think of atoms as having multiple electron “shells”. The electrons in the outermost shell are called “valence electrons”. Most of chemistry is really just the behaviour of these electrons. ## Periodic Table [00:34] [00:35] Every element is listed in the periodic table. All elements in the same column or “group” have the same number of valence electrons. For the main groups, the number of valence electrons is just the group number from 1 to 8. Except for helium. It’s too small [00:46] to have 8 electrons, it can only have 2. But still, it acts like a noble gas, so it’s kind of just grouped in with those. Luckily, the transition metals also follow a nice pattern! That was a lie, it’s kind of a mess. But that’s not so important for now, so we’ll [00:58] get to that later. Elements with the same number of valence electrons tend to show similar behaviour in chemical reactions. For example, the first group, without hydrogen, is called the “alkali metals”. Here’s some things they have in common: They have one valence electron. They’re shiny metals. They’re kind of soft. And [01:13] they do this sometimes. All elements in the same row or “period” have the same number of shells. This number increases from top to bottom. Also, the mass gets bigger from left to right, as each element gains a proton, an electron and some neutrons. ## Isotopes [01:24] [01:24] Depending on the number of neutrons in the core, you get different isotopes of the same [01:27] element, most of which are pretty unstable, and fall apart, releasing ionizing radiation. Fun fact! That stuff will kill you. If an atom has the same amount of electrons ## Ions [01:34] [01:36] as protons, it has no charge. If it has more, it has a negative charge, and if it has less, it has a positive charge. Charged atoms are called “ions”, negative ions are “anions” and positive ions are “cations”. ## How to read the Periodic Table [01:47] [01:47] The periodic table is also  pretty much a dictionary, [01:50] as every cell tells you: The name and symbol of an element, the number [01:52] of protons in the core, which is also the total number of electrons and the atomic mass, [01:56] which is the mass of neutrons and protons combined. The periodic table is roughly divided into three categories: Left of this line are the metals. Right of it are the non-metals, which are mostly gases, and the line is called the “semimetals”, which have properties that fall somewhere inbetween. ## Molecules & Compounds [02:09] [02:09] Two or more atoms bonded together form a molecule. If you have at least two different elements, you get a “compound”. Oh yeah, this is probably a good time to mention that compounds often behave completely differently than the elements they’re made of. Like, put together an explosive metal and a toxic gas, and you get, of course, an even more explo- table salt. You get tablesalt. There’s many ways to write molecules, for ## Molecular Formula & Isomers [02:27] [02:28] example the Molecular formula, where you just count the number of each atom in a molecule, [02:32] and write them as a subscript number next to the element symbol. But that has some problems. Look at these two molecules: They have the same molecular formula, but obviously, they’re not the same. They’re isomers. [02:41] Showing this difference is probably kind of important: It’s the only thing that separates [02:45] graphite from diamonds, because they’re both just fancy versions of carbon, and I [02:48] don’t think anyone’s going to go “mmm, yes, this dusty black blob is indeed very ## Lewis-Dot-Structures [02:53] [02:53] expensive”. One way to show the structure of an atom is a Lewis-Dot-Structure, which represents the valence electrons and bonds as dots and lines. That is also going to help us understand why atoms bond in the first place. ## Why atoms bond [03:03] [03:03] You see, everything in the universe wants to get to a state of lower energy. That’s why a ball on a hill will roll down by itself, because that decreases its potential energy. This trend also applies to atoms: The state of lowest potential energy is having a full outer shell of electrons, which is most often eight, or in the case of hydrogen and helium, two. If you think back to the periodic table, you’ll see that all noble gases already have a full outer shell, which is why they don’t really want to react with anything. ## Covalent Bonds [03:26] [03:26] If two atoms don’t have a full outer shell, but can achieve one by sharing electrons, [03:29] they will naturally do so, the same way a ball will go downhill, as their combined energy [03:33] would be lower than if they were separate. The sharing of electrons is called a “covalent bond”. These bonds are also caused by the positively charged nucleus of an atom tugging ## Electronegativity [03:37] [03:40] on electrons of another atom. The strength of this pull is called “electronegativity”. In the periodic table, the electronegativity increases from bottom left to the top right. Therefore, Fluorine has the strongest pull. It’s just, unbelievably desperate for an [03:53] electron. If the difference in electronegativity is ## Ionic Bonds & Salts [03:54] [03:55] bigger than about 1.7, you get an ionic bond. A good example is Sodium Chloride. Chlorine would do anything for an electron, while Sodium has one too many and just kind of wants to get rid of it anyway. “Perfect!” they both say, forming an ionic bond, where sodium [04:08] loses an electron and turns into a cation, and chlorine gains an electron and turns into [04:12] an anion. That seems pretty important…you might wanna remember it. The most common place you see Ionic bonds is in salt, yes, also table salt, but more generally, when metals and nonmetals bond, you get “a” salt, which is just a grid of ions. Speaking of metals, a pure metal forms “metallic ## Metallic Bonds [04:25] [04:27] bonds”. You can imagine this as a huge grid of the positively charged nuclei, which are surrounded by freely moving electrons. You see, in a metal grid, the valence electrons are kind of promiscuous, or as nerds call it, “delocalized”. They can move freely in a giant playground of nuclei, instead of being loyal to just one. [04:43] This kind of bond is  responsible for the properties [04:45] of metals, like conducting electricity and [04:47] heat, and also, being malleable, as in being kind of bendy. Like, you can hammer on this stuff until it’s the most deformed, unelegant and ugly looking piece of material ever known, and it will just limp along as if nothing ever happened. ## Polarity [04:59] [04:59] If the difference in electronegativity is lower than about 0.5, the electrons are shared pretty equally and you get a nonpolar covalent bond. If it’s bigger than 0.5 but smaller than around 1.7, one of the elements is pulling on the electrons pretty hard. Not quite hard [05:14] enough to completely steal an electron, but definitely hard enough to skew the electrons [05:18] a bit, making it a polar covalent bond. An example is water. Oxygen has a very high electronegativity compared to hydrogen. As a result, it pulls the electrons of hydrogen so hard, that they kind of belong to oxygen, giving it a partial negative charge, and leaving hydrogen with a partial positive charge. The presence of two poles with opposite charge [05:36] is called a “electric dipole”. All permanent dipole molecules can interact ## Intermolecular Forces [05:37] [05:39] with each other, and really, with anything that has a charge. As a result, the molecules will tug on each other and arrange themselves in a way that oppositely charged ends are next to each other. The forces acting between them are called “intermolecular forces” or IMFs. A specific example is hydrogen bonds, where ## Hydrogen Bonds [05:51] [05:53] hydrogen bonds to something very electronegative, like Fluorine, Oxygen or Nitrogen, creating [05:58] strong dipoles that tug on each other. But even if molecules are not polar at all, ## Van der Waals Forces [06:00] [06:02] there can be electrostatic forces acting between [06:04] them. How? Electrons move around inside atoms, and by pure chance they can end up on one side of the atom, creating a momentary dipole, which influences other particles next to it to become a dipole as well. At least for a very short time, as the electrons keep moving and the dipole disappears. This is called “Van der Waals forces”. ## Solubility [06:20] [06:20] The polarity of water also explains why it’s [06:22] one of the most versatile solvents to exist. It can pull apart molecules by tugging on charges, and it keeps them apart by surrounding a particle with its oppositely charged end. Water cannot dissolve nonpolar molecules though. It’s the reason why water and oil don’t mix, since fat molecules are nonpolar, while water is polar. [06:37] Just remember the ancient saying: “Similia Similibus Solventur”, [06:42] or in a language that’s actually spoken: similar [06:44] things will dissolve similar things. ## Surfactants [06:46] [06:46] Fun fact! Soap works because the molecules that it’s made from, which  are called “surfactants” have a polar “head”, and a nonpolar “tail”. This way, when in water, they can surround for example nonpolar fat molecules and form “micelles”, which, along with the water, transport the dirt particles away. These are the most important bonds and forces ## Forces ranked by Strength [07:00] [07:02] ranked by strength: [Ionic Bonds, Covalent Bonds, Metallic Bonds, Hydrogen Bonds, Van [07:06] der Waals Forces] ## States of Matter [07:07] [07:07] There are three main states of matter: Solid, liquid and gas. Solids are tightly packed in a fixed structures, where the particles can only wiggle. Unless, you know, you smash them. In liquids, the particles can move freely but are still confined to a fixed volume, as the forces between them are still strong enough to keep them together, and the particles in gases have enough energy to just do whatever they want and fill up all the volume you give them. Knowing this we can define two important words: ## Temperature & Entropy [07:28] [07:30] Temperature is the average kinetic energy of particles in a system, or how much and [07:33] how fast they move and entropy is the amount of disorder. Substances tend to be solid at low temperature and/or high pressure, which is a state of low entropy, as they’re neatly organized and don’t move that much, and gas at high temperature and/or low pressure, where they move around like crazy, so it’s a state of high entropy. Strong bonds, like ionic bonds, lead to high ## Melting Points [07:49] [07:51] melting points, as they take a lot of energy and therefore a high temperature to break [07:54] apart. That’s why most salts are solid at room temperature, whereas water, which is only being held together by hydrogen bonds, is a liquid. ## Plasma & Emission Spectrum [08:01] [08:01] Well, actually (!), there’s another state called “plasma” which is ionized gas and can exist at very high temperatures, such as in stars, or very high electric potential. The latter is used for neon lights. Gas is ionized in a tube with a very high voltage. [08:14] Collosions of the ions with other particles makes their electrons move to a higher energy [08:18] state. Once they falls back down, the difference in energy is released as light. The colour of the light depends on the element that’s used in the tube, as each element has different, but fixed energy levels, and the difference between those determines the energy and therefore the frequency of the released light, which is what changes the colour. All possible frequencies, that an element can emit, are called the “emission ## Mixtures [08:35] [08:35] spectrum”. All matter can be divided into two categories: Pure substances, which can consist of one element or one compound, and mixtures. Mixtures consist of at least two pure substances and can be homogeneous or  heterogeneous. Homogeneous means the substances will mix evenly and the mixture looks the same everywhere, like salt in water, which is a “solution”. Heterogeneous mixtures look different depending [08:54] on where you look. They have distinct regions made of separate substances. One example is sand in water, which is called a “suspension”. Okay, well what about milk? That looks the [09:03] same everywhere, so it must be homogeneous! Uhhh, no. Milk is something we call a “colloid”, or more precisely an “emulsion”. The difference between salt water and milk is that the solute doesn’t fully dissolve in the solvent, meaning there are bigger particles than in a solution, but smaller particles than in a suspension. [09:18] This allows the particles  to stay evenly distributed, [09:20] but not fully dissolved, placing them somewhere between solutions and suspensions. ## Types of Chemical Reactions [09:24] [09:24] Hey remember sodium and water? What’s going on here? Explosions are really just chemical reactions that release a lot of energy in a very short amount of time. Also, they expand, like, a lot. There’s a couple types of chemical reactions: [09:36] synthesis, decomposition, single replacement, and double replacement. Here’s an example for each one. They all happen mainly for one reason: To decrease energy and get to a more stable state. Chemical reactions happen in certain ratios, ## Stoichiometry & Balancing Equations [09:45] [09:47] for example, to produce water molecules, you need twice the amount of hydrogen compared [09:51] to oxygen. This is called “Stoichiometry”. These ratios are based on the conservation of mass, which says that mass cannot be created or destroyed, only converted. Practically, when dealing with reaction equations, you have to make sure that there’s the same amount of atoms on each side of the equation, and if not, balance it out element by element. [10:08] As a rule of thumb, you should balance out the metals first, then the nonmetals, and [10:11] hydrogen and oxygen at the end. But, it’s really just trial and error until everything is balanced. Okay, but if we wanted to make this reaction ## The Mole [10:16] [10:18] happen in a lab, how would we know that we have exactly twice the amount of hydrogen [10:22] compared to oxygen? You can’t just take 20 grams of this and mix it with 10 grams of that, because the atoms don’t weigh the same, so 10 grams of both contain a different amount of particles. What to do? Just look up the atomic mass of the reactants and take that amount in grams. You’ll get [10:35] exactly this amount of particles. That is 1 mole, which is just an amount of something, kind of like “a dozen”. In other words, we can interpret the reaction as 2 moles of this react with 1 mole of that, which we can easily measure. ## Physical vs Chemical Change [10:46] [10:46] It’s important to differentiate between physical and chemical changes, as reactions [10:50] only take place in the latter. Physical change happens when the appearance changes but the substance does not, for example hammering metal. A chemical change happens when the substances themselves change and this is often accompanied by bubbles, a funky smell, or, you guessed it, explosions. All chemical reactions need activation energy ## Activation Energy & Catalysts [11:05] [11:07] to take place. Wood won’t just spontaneously react with oxygen and start burning, or else, you know, the planet would be on fire, but if you give it enough energy, it will. Catalysts reduce the activation energy needed for a reaction, which makes it happen easier and faster. And as a neat bonus, they don’t even get used up during the reaction, so you can just reuse them! Because chemical reactions are changes in ## Reaction Energy & Enthalpy [11:24] [11:26] energy, it’s quite useful to keep track of it. “Enthalpy” is, simply put, the internal energy or heat content of a system. If the total enthalpy of a reaction is lower at the end than at the beginning, heat was given off to the surroundings, which is an “exothermic” reaction. If it’s the other way around a reaction is “endothermic”. It’s easy to see how exothermic reactions ## Gibbs Free Energy [11:42] [11:44] can be spontaneous. It’s kind of like a ball on a hill. It will only start rolling if you push it a little bit, but then it will keep rolling on its own, just like wood keeps burning on its own. But in endothermic reactions, you have to keep putting in energy, like pushing a ball uphill. That doesn’t just spontaneously happen, right? Well, yes, actually, it does. [12:00] To get the whole picture, we have to look at Gibbs Free Energy, which looks at the change [12:03] of enthalpy but also entropy of a system which is dependent on temperature. If this whole thing is less than zero, the reaction is “exergonic”, or spontaneous, because free energy was released. If it’s bigger than zero, it’s “endergonic”, or not spontaneous, because free energy was needed and absorbed. Here’s where temperature and entropy come into play: Even if delta H is positive, so the reaction is endothermic, if the change in entropy is big enough, it can offset this and make the total free energy negative, which means a reaction is spontaneous. But this [12:31] is strongly dependent on the temperature. For example, melting an ice cube is endothermic, because it absorbs heat, but also, it increases the entropy a lot, as the neatly organized ice turns to water, which is just kind of a mess. This can happen spontaneously, but only if the temperature is above 0. If it’s below 0, the water will spontaneously freeze, which is exothermic. If it’s exactly 0, then no reaction will ## Chemical Equilibriums [12:50] [12:53] take place spontaneously. In other words, if delta G is 0, we’re at equilibrium. Chemical equilibriums exist when reversible reactions take place at the same speed in both directions, which means that even if reactions are taking place, the concentrations of both sides stay the same, and to someone watching from the outside, nothing seems to be happening. We often find chemical equilibriums in phase changes, but also acid base chemistry. According to Brondsted-Lowry, an acid is a ## Acid-Base Chemistry [13:15] [13:16] molecule that donates protons, while bases accept protons. A proton in this case is just a hydrogen ion. So, with this definition, a molecule with at least one hydrogen that it can throw away can be an acid, and anything that can pick it up can be a base. This also means that once they react, they turn into the conjugate opposite, as an acid that gave away a proton can now accept one back, which is what bases do. A molecule that can act as both an acid and [13:39] a base is called "amphoteric". ## Acidity, Basicity, pH & pOH [13:41] [13:41] A strong acid will dissociate almost completely into its ionic form, giving off a lot of protons [13:46] to the water and therefore creating lots of hydronium ions. A weak acid just won’t dissociate nearly as much, giving us a lower concentration of hydronium ions. So, to measure the strength of an acid we can measure the concentration of Hydronium ions. This is called the “pH”. Mathematically, it’s defined as the negative [14:01] log of the hydronium concentration, which means one step on the scale is a 10x change, [14:05] and also, since it’s a negative log, the higher the concentration, the lower the pH. For example. Pure water is  in a chemical equilibrium. There’s exactly one hydronium ion for every 10 million water molecules. In other words, the concentration of hydronium is 1 over 10 [14:19] million, or 1 times 10^-7. Taking the negative log of this gives us a pH  of 7, which is considered neutral. Anything lower than 7 is acidic, and anything above is “basic”, unlike you. You can do the same thing with hydroxide ions and you will get the pOH, which keep track of basicity. Fun Fact! The pH and pOH always [14:38] add up to 14, because they counteract each other, so by knowing one, you know both! Now, ## Neutralisation Reactions [14:43] [14:43] if you have a strong base and a strong acid and you pour them together, no, they will [14:47] not explode, they will neutralize by forming water along with a salt, which is neutral. For example, Hydrochloride and Sodium Hydroxide will form water and table salt. Oh yeah, speaking of table salt, remember how it consists of ionic bonds, because sodium ## Redox Reactions [14:56] [14:59] transfers an electron to chlorine? Well that is called a Reduction-Oxidation reaction or “redox”. If sodium chloride forms out of it’s pure elements, the sodium gets oxidized as it loses an electron, and the chlorine gets reduced, as it gains an electron. Logically, Sodium is the oxidant, and chlorine is the reductant. [15:15] Of course not, that would make sense, it’s other way around. ## Oxidation Numbers [15:18] [15:18] More accurately, redox reactions are reactions that change the oxidation numbers of elements, [15:22] which are kind of like imaginary charges. There’s just a few rules you have to know to figure those out: Hydrogen is mostly +1, Oxygen is mostly -2, halogens are mostly -1, single elements are always 0, and the numbers of all atoms in a molecule always have to add up to the molecule’s charge. So this would total 0, while single ions just have their charge as the oxidation number. For example, in sulfuric acid, we have 4 oxygens, which totals -8, we have two hydrogens, which brings the total to -6, and since the whole molecule is neutral, sulfur must have an oxidation number of +6. [15:53] Just by looking at the oxidation numbers of reactants and products you can deduce the [15:56] flow of electrons, which  gives you these equations. If redox reactions happen in acidic or basic ## Quantum Chemistry [16:01] [16:01] solutions, you can balance out the charges with the ions, and fix the stoichiometry with [16:05] water. Okay, now to this weird looking thing. I spared you from it because for describing electrons, this is very simple, and this not. But, this is actually like, pretty wrong, electrons don’t orbit in circles. Here’s how it actually works: [16:17] All electrons inside an atom are described by four quantum numbers. N, l, ml, and ms. N corresponds to the shells, so all electrons with the same n are in the same shell. Within the shells we have subshells, with multiple orbitals, which are three dimensional regions in space where electrons could be. We know these exist thanks to schrödinger’s [16:35] equation, which gives a probabilistic wave function. You can imagine it as cloud, and the denser it is, the more likely an electron is to be there if we were to look for it. L describes the shape and ml the orientation of orbitals in a subshell. There are four subshells called s, p, d and f. If electrons have the same l, they’re [16:52] in the same subshell. If electrons have the same n, l, and ml, they are in the same orbital. Also, the number of orbitals increases by two for every bigger subshell, starting at just one for s. The last quantum number describes an intrinsic property of electrons called “spin”, which can have two values. [17:07] Some guy named Pauli said two electrons can never have the exact same quantum numbers [17:11] inside one atom. Since ms can only have two values, every orbital defined by n l and ml, can hold a maximum of 2 electrons with opposite spin. Therefore the s subshell can hold 2 electrons, the p subshell can hold 6, d can hold 10, and f can hold 14. Now, the quantum numbers restrain each other like this, which means that the first shell can only have an s subshell, the second can have an s and a p subshell, and so on. This means that the first shell can hold a [17:37] total of 2 electrons, the second can hold 8, the third can hold 18, and generally, the [17:42] number of electrons a shell can hold follows the rule 2n2, with n being the principal quantum [17:46] number. The principal quantum number, and therefore total number of shells increases from top to bottom in the periodic table, from 1 to 7. Every element has a different number of electrons that fill up these orbitals, and the different subshells and orbitals are filled in a specific order, called the “Aufbauprinciple”: just write down the subshells like this and draw diagonal lines from top right to bottom left. To get an electron configuration, just look up the number of electrons of the element in the periodic table, and fill up the subshells in this order, until there are no electrons left. This would be the electron configuration of Sodium. [18:16] You can also look up the next smallest noble gas and shorten it by just referring to its [18:19] electron configuration as the base, because those shells are full, and don’t change [18:23] for any bigger elements. This is also how you can figure out the valence electrons for transition metals. Just look up their electron configuration, ignore the full shells of the next smallest noble gas, and the remaining electrons are the valence electrons! Easy peasy. Anyways! All this knowledge going to cost [18:37] you one subscribe and a thumbs up, thank you very much, your comment is my delight, and [18:41] I shall now guide you, fine person, to the exit, where the next lesson is excitedly waiting [18:46] for you.