GENERAL CHEMISTRY explained in 19 minutes

Summary

The video script offers an insightful exploration into the fundamental aspects of chemistry, starting with the atomic structure and the periodic table's organization. It delves into the behavior of valence electrons, which dictate chemical reactions, and explains the formation of molecules and compounds. The script further elucidates the different types of chemical bonds, such as ionic, covalent, and metallic, and how they influence a substance's properties. Intermolecular forces, including hydrogen bonding and Van der Waals forces, are also discussed, highlighting their role in solubility and molecular interactions. The concept of chemical equilibrium and the impact of temperature on reaction spontaneity are covered, along with a brief introduction to quantum mechanics and electron configurations. The summary concludes with the importance of understanding chemical reactions, emphasizing the relevance of stoichiometry and the conservation of mass.

Takeaways

• 🌟 Everything is composed of atoms, which consist of a core made of protons and neutrons, and electrons that orbit around the core.
• ⚛️ The number of protons in an atom's core determines the element, and atoms combine to form molecules, which can be either compounds or pure substances.
• 🔋 Atoms strive for a state of lower energy, often achieving this by sharing or transferring electrons to fill their outer electron shell, a process fundamental to chemical bonding.
• 🔗 There are three main types of chemical bonds: ionic, covalent, and metallic, each with distinct properties and forming under different conditions.
• 💧 Water is a universal solvent due to its polarity, which allows it to interact with and dissolve many substances with an uneven charge distribution.
• 📊 The periodic table organizes elements by their atomic number and provides valuable information about their properties, such as valence electrons and atomic mass.
• 🔑 Quantum mechanics describes the behavior of electrons in atoms using quantum numbers, which define the probability of finding an electron in a particular region around the nucleus.
• 🔥 Temperature and entropy are key factors in determining the state of matter and the spontaneity of reactions, with substances tending towards lower energy and higher disorder.
• 🔍 Isomers are molecules with the same molecular formula but different structural arrangements, leading to different physical and chemical properties.
• 🤝 Intermolecular forces, such as hydrogen bonds and Van der Waals forces, influence how molecules interact and the properties of substances like boiling points and solubility.
• ⚖️ Chemical reactions follow the law of conservation of mass and occur in specific stoichiometric ratios to achieve a more stable, lower-energy state.

Q & A

What are the basic components of an atom?

An atom consists of a core made of protons and neutrons, and electrons that orbit around the core. The core is the central part of the atom, while electrons are in the outer regions in what are known as electron shells.

How do the number of protons in an atom's core determine its properties?

The number of protons in an atom's core, also known as the atomic number, determines the element's identity. Different elements have different numbers of protons, which in turn affects their chemical properties and how they interact with other elements.

What is the significance of the periodic table in chemistry?

The periodic table is a tabular arrangement of the chemical elements, ordered by their atomic number, electron configuration, and recurring chemical properties. It helps predict the behavior of elements and their interactions, as elements in the same group or period tend to exhibit similar properties.

How does the number of valence electrons influence an element's chemical behavior?

Valence electrons, which are the electrons in the outermost shell of an atom, play a crucial role in chemical reactions. Elements with the same number of valence electrons often exhibit similar chemical behaviors, as they tend to follow similar patterns in forming chemical bonds.

What is the difference between isotopes of an element?

Isotopes of an element are variants of the same element that have the same number of protons but different numbers of neutrons in their atomic core. This difference in neutron count leads to variations in mass but does not change the chemical properties of the element.

How does the charge of an atom affect its classification as an ion?

An atom's charge determines whether it is classified as an ion. If an atom has the same number of electrons as protons, it is neutral. If it has more electrons, it becomes a negatively charged ion (anion), and if it has fewer electrons, it becomes a positively charged ion (cation).

What is a covalent bond, and how does it form?

A covalent bond is a type of chemical bond formed when two atoms share one or more pairs of electrons. This sharing allows each atom to achieve a stable electron configuration, typically a full outer shell, which is a state of lower energy.

How do ionic bonds differ from covalent bonds?

Ionic bonds form when there is a transfer of electrons from one atom to another, resulting in the formation of ions. This is common when there is a significant difference in electronegativity between the two atoms. In contrast, covalent bonds involve the sharing of electrons rather than their transfer.

What is the role of electronegativity in determining the type of chemical bond?

Electronegativity is a measure of an atom's ability to attract electrons in a chemical bond. The difference in electronegativity between two atoms influences the type of bond they form. A large difference (greater than about 1.7) typically results in an ionic bond, while a smaller difference (less than about 0.5) leads to a nonpolar covalent bond.

Why are noble gases generally unreactive?

Noble gases are unreactive because they already have a full outer shell of electrons, which is a state of lowest potential energy. This complete electron shell makes them stable and disinclined to form additional bonds with other atoms.

How do intermolecular forces (IMFs) affect the properties of substances?

Intermolecular forces, which include hydrogen bonding and Van der Waals forces, influence the physical properties of substances, such as their boiling and melting points, solubility, and viscosity. These forces act between molecules and affect how they interact with each other.

Outlines

00:00

🌟 The Basics of Atoms and the Periodic Table

This paragraph introduces the fundamental concept that everything is composed of atoms, which consist of a core made of protons and neutrons, and electrons. It explains the role of protons in determining the element's identity and how the periodic table categorizes elements by their valence electrons. The paragraph also touches on quantum mechanics, isotopes, and the concept of ions, anions, and cations. It concludes with a brief mention of the periodic table's division into metals, non-metals, and semimetals, and the formation of molecules and compounds.

05:04

🔬 Bonding and Molecular Structure

The second paragraph delves into the different types of chemical bonds, including covalent, ionic, and metallic bonds. It explains the driving force behind these bonds is the desire of atoms to achieve a lower energy state, often by filling their outer electron shell. The concept of electronegativity and how it influences bond polarity is discussed, along with the role of intermolecular forces, particularly hydrogen bonding and Van der Waals forces. The paragraph also covers the states of matter and the factors influencing their transitions, such as temperature and pressure.

10:08

🧪 Chemical Reactions and Equilibrium

This paragraph explores the types of chemical reactions, including synthesis, decomposition, single replacement, and double replacement, and emphasizes the importance of stoichiometry in balancing chemical equations. It discusses the concept of moles and molar ratios, activation energy, and the role of catalysts in facilitating reactions. The paragraph also explains enthalpy, entropy, and the Gibbs Free Energy equation, which together determine the spontaneity of a reaction. It concludes with an overview of chemical equilibrium, acid-base chemistry, and the principles of redox reactions.

15:11

⚛️ Quantum Mechanics and Electron Configuration

The final paragraph provides a deeper understanding of electron behavior within atoms, debunking the classical model of electron orbits with a quantum mechanical perspective. It introduces the concept of quantum numbers (N, l, ml, ms) that describe the properties and positions of electrons within an atom. The paragraph explains the structure of subshells and orbitals, the Pauli Exclusion Principle, and the Aufbau principle for filling electron subshells. It concludes with a method for determining valence electrons, particularly for transition metals, and a call to action for viewers to engage with the content.

Keywords

💡Atoms

Atoms are the basic units of matter, composed of a nucleus containing protons and neutrons, surrounded by electrons. They are the fundamental building blocks of all elements and compounds. In the video, atoms are introduced as the starting point for understanding the composition of everything from the human body to the elements listed in the periodic table.

💡Electron Shells

Electron shells, also known as energy levels, are regions around the nucleus of an atom where electrons are likely to be found. The outermost shell contains the valence electrons, which play a crucial role in chemical reactions. The concept is central to the video's discussion of chemical bonding and the behavior of elements.

💡Periodic Table

The periodic table is a tabular arrangement of the chemical elements, ordered by their atomic number, electron configuration, and recurring chemical properties. It is a fundamental tool in chemistry and the video uses it to explain how elements with the same number of valence electrons (group number) exhibit similar chemical properties.

💡Valence Electrons

Valence electrons are the outermost electrons in an atom and are responsible for an element's chemical bonding behavior. The video emphasizes that most of chemistry revolves around the interactions of these electrons, which dictate how elements react with each other to form compounds.

💡Isotopes

Isotopes are variants of a particular chemical element that have the same number of protons but different numbers of neutrons in their nuclei. This results in different atomic masses but does not change the chemical properties. The video mentions isotopes in the context of varying the number of neutrons and the resulting instability of many isotopes.

💡Ions

Ions are atoms or molecules that have a net electric charge due to the loss or gain of one or more electrons. The video explains that ions can be either anions, which have a negative charge, or cations, which have a positive charge. This concept is integral to understanding ionic bonding and the formation of compounds like table salt.

💡Covalent Bonds

Covalent bonds are a type of chemical bond where two atoms share one or more pairs of electrons. The video describes covalent bonds as a way for atoms to achieve a full outer shell of electrons, thus lowering their potential energy, which is why such bonds form spontaneously.

💡Electronegativity

Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. It is a key factor in determining the type of chemical bond formed between atoms. The video illustrates how the difference in electronegativity between atoms can lead to ionic or covalent bonds.

💡Metallic Bonds

Metallic bonds are a type of chemical bond that involves a regular arrangement of cations and free electrons, which are delocalized and can move throughout the metal lattice. The video explains that this type of bonding is responsible for the characteristic properties of metals, such as electrical conductivity and malleability.

💡Intermolecular Forces

Intermolecular forces (IMFs) are the forces that cause molecules to attract each other, which are weaker than chemical bonds. The video discusses various types of IMFs, including hydrogen bonds and Van der Waals forces, which play a significant role in determining the physical properties of substances.

💡States of Matter

The three traditional states of matter are solid, liquid, and gas. The video describes how the particles in each state are arranged and move, and how temperature and pressure affect the transition between these states. Understanding the states of matter is essential for discussing phase changes and the behavior of substances under different conditions.

Highlights

Atoms are the building blocks of all matter, consisting of a core made of protons and neutrons, and electrons orbiting the core.

The number of protons in an atom's core determines the element it forms.

Water is a compound made of hydrogen and oxygen, and mixing certain elements can lead to unexpected reactions.

Quantum mechanics provides a more accurate description of atomic structure than traditional models.

The periodic table organizes elements by their number of valence electrons and displays their properties systematically.

Elements in the same group of the periodic table tend to exhibit similar chemical behaviors.

The transition metals in the periodic table do not follow a simple pattern and are more complex.

Elements with the same number of valence electrons often react chemically in similar ways.

The periodic table categorizes elements into metals, non-metals, and semimetals.

Molecules are formed when two or more atoms bond together, and compounds are made from different elements.

Isomers are molecules with the same molecular formula but different structural arrangements.

Lewis-Dot-Structures illustrate valence electrons and bonds, helping to understand why atoms bond.

Atoms naturally seek to achieve a state of lower energy, often by sharing electrons to fill their outer shell.

Electronegativity is a measure of an atom's ability to attract electrons and varies across the periodic table.

Ionic bonds form when there is a significant difference in electronegativity between atoms, as seen in sodium chloride.

Metallic bonds are characterized by a 'sea' of delocalized electrons moving freely among positively charged nuclei.

Polar covalent bonds occur when there is a significant difference in electronegativity between bonded atoms, leading to an electric dipole.

Intermolecular forces, such as hydrogen bonds and Van der Waals forces, influence how molecules interact and dissolve in solvents.

The strength of different bonds and forces can be ranked from ionic to Van der Waals forces, with ionic bonds being the strongest.

States of matter (solid, liquid, gas) are defined by the energy and movement of particles, with temperature and pressure affecting their transitions.

Chemical reactions often aim to decrease energy and reach a more stable state, following the law of conservation of mass.

Activation energy is required for chemical reactions to occur, and catalysts can lower this energy threshold.

Enthalpy and Gibbs Free Energy are used to determine whether a reaction is exothermic, endothermic, spontaneous, or at equilibrium.

Acid-base chemistry is defined by the donation and acceptance of protons, with pH and pOH scales measuring their concentrations in solutions.

Redox reactions involve the transfer of electrons, changing the oxidation numbers of elements involved.

Quantum mechanics describes electron behavior in atoms using quantum numbers, which define electron shells, subshells, and orbitals.

The Aufbau principle dictates the order in which electron subshells and orbitals are filled in atoms.

Transcripts

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

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.

Quantum mechanics tells us that this is not what atoms actually look like, they look more

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.

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

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

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

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.

Depending on the number of neutrons in the core, you get different isotopes of the same

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

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”. The periodic table is also pretty much a dictionary,

as every cell tells you: The name and symbol of an element, the number

of protons in the core, which is also the total number of electrons and the atomic mass,

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.

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

example the Molecular formula, where you just count the number of each atom in a molecule,

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.

Showing this difference is probably kind of important: It’s the only thing that separates

graphite from diamonds, because they’re both just fancy versions of carbon, and I

don’t think anyone’s going to go “mmm, yes, this dusty black blob is indeed very

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.

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.

If two atoms don’t have a full outer shell, but can achieve one by sharing electrons,

they will naturally do so, the same way a ball will go downhill, as their combined energy

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

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

electron. If the difference in electronegativity is

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

loses an electron and turns into a cation, and chlorine gains an electron and turns into

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

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.

This kind of bond is responsible for the properties of metals, like conducting electricity and

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.

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

enough to completely steal an electron, but definitely hard enough to skew the electrons

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

is called a “electric dipole”. All permanent dipole molecules can interact

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 to something very electronegative, like Fluorine, Oxygen or Nitrogen, creating

strong dipoles that tug on each other. The polarity of water also explains why it’s

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.

This way, it can dissolve almost anything with an uneven distribution of charges.

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.

Just remember the ancient saying: “Similia Similibus Solventur”, or in a language that’s

actually spoken: similar things will dissolve similar things.

But even if molecules are not polar at all, there can be electrostatic forces acting between

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”.

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

ranked by strength: [Ionic Bonds, Covalent Bonds, Metallic Bonds, Hydrogen Bonds, Van

der Waals Forces]

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 is the average kinetic energy of particles in a system, or how much and

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, as they take a lot of energy and therefore a high temperature to break

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.

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.

Collosions of the ions with other particles makes their electrons move to a higher energy

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

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

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

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. This allows the particles to stay evenly distributed,

but not fully dissolved, placing them somewhere between solutions and suspensions.

 

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:

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,

for example, to produce water molecules, you need twice the amount of hydrogen compared

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.

As a rule of thumb, you should balance out the metals first, then the nonmetals, and

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

happen in a lab, how would we know that we have exactly twice the amount of hydrogen

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

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.

It’s important to differentiate between physical and chemical changes, as reactions

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

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

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

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.

To get the whole picture, we have to look at Gibbs Free Energy, which looks at the change

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

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

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

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

a base is called "amphoteric".

A strong acid will dissociate almost completely into its ionic form, giving off a lot of protons

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

log of the hydronium concentration, which means one step on the scale is a 10x change,

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

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

add up to 14, because they counteract each other, so by knowing one, you know both! Now,

if you have a strong base and a strong acid and you pour them together, no, they will

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

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.

Of course not, that would make sense, it’s other way around.

More accurately, redox reactions are reactions that change the oxidation numbers of elements,

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.

Just by looking at the oxidation numbers of reactants and products you can deduce the

flow of electrons, which gives you these equations. If redox reactions happen in acidic or basic

solutions, you can balance out the charges with the ions, and fix the stoichiometry with

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:

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

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

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.

Some guy named Pauli said two electrons can never have the exact same quantum numbers

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

total of 2 electrons, the second can hold 8, the third can hold 18, and generally, the

number of electrons a shell can hold follows the rule 2n2, with n being the principal quantum

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.

You can also look up the next smallest noble gas and shorten it by just referring to its

electron configuration as the base, because those shells are full, and don’t change

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

you one subscribe and a thumbs up, thank you very much, your comment is my delight, and

I shall now guide you, fine person, to the exit, where the next lesson is excitedly waiting

for you.