An Intuitive Understanding of Life at the Atomic Level


What is an Atom?

An atom is about 500,000 times smaller than the thickness of a piece of paper.

Atoms make up everything. 

The air you breathe, the water you drink, the wires that give electricity, the medicine you take, and even the DNA strands that make you who you are.

The reason the world is filled, with different colors, textures, tastes, temperatures, along with different properties that keep us living and/or make our lives easier is because of atoms. 

There are 118 foundational elements all representing a different kind of atom with a unique configuration.

And these atoms bond to make a much bigger variety of compounds and molecules. 

What is the periodic table?

The periodic table is a chart of all of the foundational elements in the universe. Each element represents atoms all having the same configuration that is unique to that element, which is why each element has unique physical and chemical properties.

Physical Properties Include

  • How hard an element is
  • How dense an element is
  • How much an element weighs
  • An element’s ability to conduct electricity
  • The color of an element
  • How much heat it takes to melt or boil an element

Chemical properties are how an element reacts with other elements over time.

Basic Components of an Atom

An atom is made of a nucleus that contains protons and neutrons. An atom also has high-energy electrons that move around in the surrounding space. 

The nucleus accounts for almost all of the weight of an atom, with protons and neutrons each weighing in at 1 atomic mass unit (AMU) while electrons are pretty much weightless in comparison.

In chemistry, there are positive charges and negative charges, just know that opposite charges attract and same charges repel.

Protons are positively charged, electrons are negatively charged, and neutrons are neutrally charged.

The neutrons stabilize the protons in the nucleus preventing the protons from colliding with the electrons, while the protons keep the high-energy electrons in the vicinity of the nucleus.

The number of electrons and the behavior of electrons in an atom are different for each element. This is referred to as the electron configuration of an atom and is what gives the elements their unique physical and chemical properties which is best demonstrated by how the elements are ordered on the perodic table.

Now let’s go over the first two trends on the periodic table, increasing nuclear force and the expansion of the electron cloud.

Nuclear Force (Nuclear Charge)

Nuclear force is the attraction between the protons in the nucleus and the electrons surrounding the nucleus.

There are neutrons and protons in the nucleus. The neutrons prevent the protons from colliding with the electrons, but the attraction pulls the electrons closer to the nucleus.

The more protons in the nucleus the more nuclear force that atom has and the more it can pull the electrons closer to the nucleus.

Take a look at the periodic table. Every element has an atomic number.

The atomic number represents the number of protons in an atom of that element. 

So a Hydrogen atom has 1 proton, Helium 2 protons, Lithium 3 protons, Beryllium 4 protons, and so on left to right throughout the periodic table.

Nuclear force increases from left to right.

Electron Cloud

An atom can have multiple “shells” or energy levels surrounding the nucleus. These shells can hold a set number of potential electrons.

Assume that these are neutrally charged atoms (protons = electrons)

In these 2-dimensional representations of an atom

The first shell can hold up to 2 electrons.

The second shell can hold up to 8 electrons.

The third shell can hold another 8 electrons.

This directly correlates with the first three rows of the periodic table.

The 1st row has 2 elements.  H, He

The 2nd row has 8 elements. Li, Be, B, C, N, O, F, Ne

The 3rd row has another 8 elements. Na, Mg, Al, Si, P, S, Cl, Ar

Let’s see what the atoms of these elements look like using the diagrams shown previously.

*Again assume that the protons in the nucleus are equivalent to the electrons*

Things to Know

  1. The electrons fill from the inside out. Electrons have to fill the first shell before filling the outer shells.
  1. All elements in the same group will have the same number of valence electrons (ve) or electrons in their outermost shell.
  1. Nobel Gases fulfill the maximum number of potential electrons in their outermost shell. Full Shell (FS).
  1. Elements seek to have a full stable configuration in their outermost shell, which we’ll touch on more when we go over the other periodic trends and the bonding between atoms.
  1. When you move to a new row, a new shell is added. This isn’t perfect logic especially as you move to the 4th row and beyond. The transition metals (groups 3-12) complicate this logical 2-dimensional progression. 
  1. Just know that as you move down the periodic table there are increasing levels of energy and we conceptualize that they hold a set number of electrons.
  1. We’ll address the complexity of this when we talk about electron configurations and quantum in future sections. But for now, this 2 dimensional understanding of an atom will do.periodic table.

Nuclear Force and the Expansion of The Electron Cloud set up the other trends of the periodic table. Click arrow on the right of the page to move on to the next article.

How Atomic Radius Affects Ionization Energy and Electronegativity

The trends of atomic radius, ionization energy, and electronegativity set up the circumstances for atomic bonding.

Note: The elements that are noble gases in the group on the far right of the periodic table, are stable elements and don’t want to bond with other elements. So noble gases will be excluded from these trends.

Atomic Radius

Atomic Radius is the distance between the nucleus and the outermost electron shell. 

It’s a combination of nuclear force and the number of electron shells.

Interior electron shells repel outer electrons away from the nucleus, while the nuclear force tries to pull the outermost electrons closer to the nucleus.

The greater amount of nuclear force relative to the number of electron shells the smaller the atomic radius.

I am going to show you three atoms: Hydrogen, Lithium, and Fluorine. 

*assume these are neutrally charged atoms (# of protons = # of electrons)*

Maybe you thought Hydrogen was going to have the smallest radius sin it’s the only element besides Helium with only one shell, but fluorine has the smallest atomic radius.

This is because hydrogen has a weak nuclear force with only one proton, while fluorine has 9  protons to pull in its two shells really close.

Lithium only has 3 protons which is why its two shells are further out.

The smallest elements are on the top right (including hydrogen) because the further up you go the less number of “shells” or energy levels and the further right you go the greater amount of nuclear force relative to the number of shells.

The biggest elements are on the bottom left because the further down you go the more energy levels there are and the further left you go the less amount of nuclear force relative to the number of energy levels.

                Ionization Energy (I.E)                                                  Electronegativity (E.N)

An atom’s ability to resist losing an electron.An atom’s ability to attract an electron.
Losing an electron is a + charge since electrons are – charges.Gaining an electron is a – charge.
An element with high I.E would say 
“I do not want a + charge” 
In other words “I don’t want to lose an electron”
An element with high E.N would say
“I want a negative charge”
In other words “I want to attract an electron”
I.E is used in the scenario of Ionic Bonding where one atom steals an electron from another atom. This is based on an individual element’s ability to resist losing an electron.
The element with lower I.E loses its electron to the element with higher I.E, this creates opposite charges which allows the elements to bond.
E.N is used in the scenario of Covalent Bonding where two atoms share electrons.
This is based on multiple elements’ ability to attract an electron.

Both elements need to have a high E.N to share the electrons, but one element can have a higher E.N and have more ownership over the shared electrons.
The smaller the atomic radius the better an element can protect its electrons. The smaller the atomic radius the better an element can attract an electron.
The nuclear force pulls the electrons closer to the nucleus.The nuclear force can attract additional electrons.
The fewer electron shells the easier its for an element to protect its outermost electrons.The fewer electron shells the less resistance there is to attract additional electrons. Since electrons repel electrons.

Let’s see these trends on the periodic table.

The trends of I.E and E.N go hand in hand because if an element can protect its electrons it can attract additional electrons.

These two trends are opposite to atomic radius because a smaller atomic radius means high I.E and E.N.

Let’s look at I.E and E.N  using valence electrons (ve).

Valence Electrons 

Valence Electrons are the electrons in an element’s outermost shell. These lewis dot structures don’t show the accumulation of shells just the dots resembling that element’s valence electrons.

The rule regarding an element’s valence shell is that all elements want to fill it to obtain a noble gas configuration.

So when you see that Flourine is one electron away from fulfilling a noble gas configuration like Neon you know that Flourine wants an electron or a -1 charge. This also means that Flourine is high in E.N and I.E as well because it doesn’t want to lose an electron and be two electrons away from ng configuration. 

Looking at Lithium you see it needs to lose an electron in its second shell to obtain the noble gas configuration of He, so it wants a +1 charge. 

This means it’s low in I.E, as well as E.N because if it were to attract an electron it would need to lose two to get ng configuration.

Carbon could want a +4 or -4 charge because it can lose its 4 electrons to get to He or gain 4 electrons to get to Ne.

Hydrogen has one electron, if it loses that electron it is no longer an atom so hydrogen wants an electron to fulfill the ng configuration of He, meaning it wants a -1 charge.

This makes hydrogen a nonmetal. 

Nonmetals have high I.E and E.N and want a negative charge. Nonmetals are on the top right of the periodic table but hydrogen is an exception.

Metals make up the first two groups of the periodic table. They form positive charges and have low I.E and E.N.

Now that we understand periodic trends we can go over atomic bonding. The first type of atomic bonding is called ionic bonding. 

Ionic Bonding, Covalent Bonding, and Polyatomic Ions

Ionic Bonding

This bond is created after a transfer of electrons, and involves an element with low I.E and high I.E.

In other words, it’s a bond between a metal and a nonmetal. The metal loses an electron(s) to the non-metal.

The metal forms a + charge or a cation, and the non-metal receives a negative charge or an anion.

This process is called ionization.

Since the two elements would have opposite charges they would then bond.

Let’s see some examples.

Since we know the valence electrons of both the metals and the nonmetals we know that they need to either give away or receive a certain number of electrons to obtain a stable ng configuration.

This makes figuring out the charges for the metals and nonmetals easy.

Ionic bonding with transition metals however is not that straightforward.

Ionic Bonding with Transition Metals 

Like the metals in groups 1 and 2, transition metals (groups 3-12) can form cations to bond with nonmetals.

Unlike the metals that need to lose a certain number of electrons, to form a specific cationic charge, the transition metals take into account the varying levels of energy in an atom and can have configurations that satisfy different orbitals of energy. 

This also means they don’t consistently follow the trends for ionization energy and electronegativity, unlike metals and nonmetals.

I will go over this more when I explain electron configurations and the 3-dimensionality of the atom in future sections, but for now, all you need to know is that transition metals can form multiple cations or lose different numbers of electrons because they can have multiple stable configurations besides ng configuration.

Here are some transition metals and all their different cationic charges.

Chromium (Cr)  +2, +3, +6

Vanadium (V) +2, +3, +4, +5

Iron (Fe) +2, +3, +4, +6

Gold (Au) +1, +3

Manganese (Mn) +2, +3, +4, +6, +7

Since the nonmetals and metals have predictable charges you can use those to figure out which cation the transition metal is.

Examples

FeCl2

Looking at the periodic table you can see that Chlorine is in group 17 like Fluorine, which means it can form a -1 charge.

Since there are two Chlorine atoms there a two negative charges.

This means that the cation for Iron (Fe)  is +2.

KMnO4

We know oxygen forms a -2 charge since it wants two electrons.

Since there are 4 oxygen atoms we know there is a charge of -8.

Potassium (K) is a metal in the first group so we know it for a consistent +1 charge.

This means that Manganese (Mn) must have a +7 charge.

Covalent Bonding

Now that we got ionic bonding out of the way let’s go over a different type of atomic bonding, covalent bonding.

Covalent bonding is a bond created when two elements share electrons and is between two elements with high electronegativity, which means it is between two nonmetals.

There are two types of covalent bonds nonpolar and polar.

Non-polar is when two elements of equal electronegativities share electrons, and polar is when two elements of different electronegativities share electrons, which creates dipoles but let’s go over nonpolar covalent bonding first.

Nonpolar

A good example of when two atoms have equal electronegativities is when they are atoms of the same element.

A hydrogen atom has one electron and needs one more to fulfill both electrons in that first shell.

Two hydrogen atoms can each donate their only electron so they’re sharing two in the middle, satisfying both of their needs to have two electrons in their only shell.

Since they have equal electronegativities they both have equal ownership of the electrons.


An oxygen atom has 6 electrons in its outermost shell and needs 2 to fulfill ng configuration.

2 oxygen atoms can each donate 2 electrons to the point where they both have 4 to themselves and sharing 4 equally in the middle.

Fulfilling 8 electrons in their second shells.

Polar

Let’s look at some examples of when nonmetals of different electronegativities create polar covalent bonds.

Both Hydrogen and Fluorine want one electron, hydrogen to fill 2 in the first, and Fluorine to fill 8 in the second.

Their both going to donate one electron so there both sharing 2, giving hydrogen and fluorine stable configurations, Fluorine still having 6 electrons to itself. 

Since fluorine is more electronegative than hydrogen it will have more ownership of the 2 shared electrons.

This creates dipoles (δ) which shows that Fluorine has a more negative pull, while Hydrogen has a weaker, more positive pull.


Oxygen wants 2 electrons to fill all 8 in the second shell, while two hydrogens want 1 electron to fill 2 in their first shells.

Oxygen will donate one electron to each hydrogen, while the hydrogens donate their electrons back to oxygen.

The hydrogens are both sharing 2 electrons with oxygen fulfilling their shells, and oxygen is sharing a total of 4 electrons, while still having 4 to itself, filling 8 in its second shell.

Oxygen has a stronger electronegativity than the hydrogens so it has more ownership over all 4 shared electrons.

Polyatomic Ions

Polyatomic ions are atoms that are covalently bonded but have an ionic charge.

Common Polyatomic Ions

[OH] -1 charge (Hydroxide)[NO2] -1 charge (Nitirite)
[NH4] +1 charge (Ammonium)[CO3] -2 charge (Carbonate]

When covalently bonded atoms don’t have an ideal number of electrons to fulfill a stable configuration, the compound as a whole receives an ion.

Ion- When # of electrons doesn’t equal # of protons.

Cation- When there are more protons than electrons.

Anion- When there are more electrons than protons.

If the polyatomic ion is a cation then it gives an electron(s) to a nonmetal.

If the polyatomic ion is an anion then it takes an electron(s) from a metal.

How Polyatomic Ions Form

With hydroxide, hydrogen and oxygen are in a polar covalent bond. Oxygen has 6 electrons in its outermost shell so ideally, it would give 2 electrons to get back 4 for 8 ve, but hydrogen only has one to give. Hydrogen gets 2 electrons back which satisfies it, but this leaves oxygen with only 7 ve, so it needs to acquire an electron from a metal in this case Lithium.


With ammonium, there is 1 nitrogen atom and 4 hydrogen atoms. Each hydrogen can donate 1 electron to get back 2, nitrogen giving 4 to get back 8, but nitrogen still has 1 electron more than it needs since it starts with 5 ve. So the ammonium compound must lose an electron to a nonmetal, in this case, Fluorine.


With nitrite, there is a nitrogen atom and two oxygen atoms. Each oxygen can donate 2 electrons to get back 4 which means nitrogen donates a total of 4 electrons to get back 8. But like the last one, it has an extra electron so it loses it to chlorine.


With carbonate, there are 3 oxygen atoms and one carbon atom. The carbon atom has 4 electrons to donate so it gives 2 to one oxygen, and 1 electron each to the other 2 oxygens. This gives carbon the 8 electrons it needs as well as one of the oxygens, while 2 of the oxygens need an electron since they’re only forming one bond with carbon. So in this case magnesium gives each of the two oxygens an electron.

That concludes this section. In the next section were going to delve into electron configuration and the 3-dimensional model of an atom.

An Intuitive Understanding of Electron Configurations

In this article, we will be making a connection between the 2-dimensional valence electron diagrams we were using in the last section and the new concept of the 3-dimensional atom. 

Let’s review some things about the atom.

In an atom, there is a positively charged nucleus, and negatively charged high-energy electrons surrounding the nucleus. 

The nucleus is made of positively charged protons and neutrally charged neutrons. The neutrons keep the protons bound to the nucleus.

The positive charge from the protons keeps the highly mobile electrons in the vicinity of the nucleus.

If the electrons get too close to each other they will start to repel and occupy more space further away from the nucleus. That’s why the “shells” or energy levels only hold a set number of electrons. When electrons fill a lower energy level, additional electrons will have to fill the next energy level.

The electrons on the interior energy levels shield the exterior electrons from the attraction of the nucleus, which is why electrons further away from the nucleus are more mobile and have more energy.

2-dimensional valence electron model

In a typical 2-dimensional atomic model, electrons orbit around the nucleus, in this model, there are 3 orbits or shells.

These three shells expand outward from the nucleus, each shell can hold a set number of electrons.

*Assume these are neutrally charged atoms (# of protons = # of electrons)*

Things to know

  1. Red dots are valence electrons.
  1. The elements in the same group have the same number of valence electrons.
  1. When you look at the next element to the right it will have 1 more proton and 1 more electron than the previous element.
  1. The electrons build from the inside out. So an element on the third row will fill the first 2 shells before filling the third.
  1. When you move to a new row you add a new shell, this gets a bit more complicated when you go to the 4th row which I’ll explain later in this article.

Now we’re going to try to develop an intuitive understanding of the 3-Dimensionality of the atom. We’ll refer back to the previous model to make a clear connection between what you already know and what is much harder to conceptualize. 

The Behavior of Electrons in the 3-Dimensional Atom

Let’s say, as opposed to having a set location on a shell or orbit, electrons have more room to move and are more unpredictable.

The same rules apply, the nuclear force keeps the electrons in the general vicinity, while the electrons repel other electrons to higher levels of energy and mobility.

But we’re going to account for the mobility and unpredictability of the electron by using  orbitals. 

Orbitals are areas in space that predict where 2 electrons can be over a period of time with about 90% accuracy.

There are different types of orbitals representing different geometric shapes where 2 electrons can be. 

These different geometric shapes are categorized by sublevels. Each sublevel contains a specific number of orbital(s). 

The letters s,p,d, and f are the 4 sublevels that will be used for the electron configurations of the elements on the periodic table. 

Remember each orbital represents the probable area of where 2 electrons can be, no matter the shape, so we know how many electrons each sublevel will hold. 

s-sublevel
1 orbital
2 electrons
p-sublevel
3 orbitals
6 electrons
d-sublevel
5 orbitals
10 electrons
f-sublevel
7 orbitals
14 electrons

Now that we know what the sublevels are we can use them to see how they fit with our traditional atomic model and follow our logic all the way down the periodic table.

Instead of using shells, I’m going to use rectangles to represent the rows on the periodic table.

*represents 1 atom of any element depending on how many electrons are filled*

Looking at the first shell, we know it holds 2 electrons. So we know that there’s an s-sublevel on the first shell.

Looking at the second shell, we know it holds 8 electrons. So we know there’s an s-sublevel and a p-sublevel on the second shell.

Looking at the third shell, we know it holds 8 electrons. So we know there’s an s-sublevel and a p-sublevel on the third shell.

Following our logic down the periodic table, looking at the 4th row, we see there are 18 elements because you add 10 transition metals. If we had a 4th shell we could split that up into an s-sublevel, d- sublevel, and p-sublevel.

The same thing if there were a 5th shell, you would add a d-sublevel to account for the 10 transition metals.

For the 6th row on the periodic table, you would insert the lanthanides, before the transition metals which would mean 32 elements. This means you would add the f-sublevel to account for 14* elements if there were a 6th shell.

For the 7th row on the periodic table, you would insert the actinoids, before the transition metals which would mean 32 elements. This means you would add the f-sublevel to account for 14* elements if there were a 7th shell.

Now that we know what the sublevels are and how they are ordered on the periodic table we can introduce different levels of energy.

Energy Levels

Energy levels progress outwards from the nucleus, the further away the electrons are from the nucleus, the more energy they have.

So the higher the energy level the more mobility and vibration the electrons have.

As the energy level increases the more freedom the electrons have to move, which means the size of the area where the electrons can be increases.

This means the sizes of the orbitals increase by energy level.

This seems pretty straightforward, the only problem is that the rows/shells don’t represent the energy levels. The energy levels vary within the rows once you get to the 4th row.

Energy LevelsRows on Periodic Table
1st EL | 1s2 | 2 electrons1st Row | 1s2 | 2 electrons
2nd EL | 2s2, 2p6 | 8 electrons2nd Row | 2s2, 2p6 | 8 electrons
3rd EL | 3s2, 3p6, 3d10 | 18 electrons3rd Row | 3s2, 3p6 | 8 electrons
4th EL | 4s2, 4p6, 4d10, 4f14 | 32 electrons4th Row | 4s2, 3d10, 4p6 | 18 electrons
5th EL | 5s2,5p6,5d10,5f14 | 50 electrons*5th Row | 5s2, 4d10, 5p6 | 18 electrons
6th EL | 6s2,6p6,6d10 | 72 electrons* 6th Row | 6s2, 4f14,5d10,6p6 | 32 electrons
7th EL | 7s2,7p6 |  88 electrons*7th Row | 7s2, 5f14, 6d10, 7p6 | 32 electrons

Understanding Energy in an Atom

Why is this? Why does 4s fill before 3d if it’s on a higher energy level or further away from the nucleus?  

One thing to understand is that energy levels are not based on atoms with multiple electrons. 

There based on the Neils Bohr model which showed that hydrogen’s one electron can jump to specific levels of energy and take the shape of different orbitals over time. 

Atoms with multiple electrons are best represented by how the sublevels are ordered on the periodic table because it best represents our probabilistic understanding of electron behavior when electron shielding is taken into consideration.

Electron Shielding and Nodes

Before we even discussed orbitals and sublevels, we knew that electrons repelled other electrons further away from the nucleus. 

We interpreted this by using a 3 shell model based on the first three rows on the periodic table. 

Once elements in the second row filled the first two electrons on the first shell, then other electrons would fill the second shell that was further away from the nucleus.

 But no electron could exist in between the first and second shells because the interior electronics were shielding that area. Same thing with the electrons on the second shell shielding the area below the third shell.

These areas are called nodes, which represent a 0% probability of electrons existing there.

But now we understand that the levels of energy in an atom or more complex containing different sublevels and orbitals with geometric shapes.

Each orbital gives off a node that affects the same orbital on the next energy level. These nodes are called radial nodes.

Looking at the s-sublevel you can see that 2s has one node because of shielding from the 1s. 

3s will have two radial nodes from both 1s and 2s. 

4s will have 3 radial nodes from 1s,2s, and 3s.


The thing with the orbitals in the p sublevel is that they each have two lobes which cause   electron shielding in between the lobes. 

This shielding creates an angular node which prevents any electron in a p orbital from being right by the nucleus. 

This one angular node created by the 2p orbitals affects all of the electrons in the p sublevels no matter what energy level they are in.

The orbitals in 2p also give off radial nodes which just affect the orbitals in 3p. 

Then the 4p orbitals will have radial nodes from both 2p and 3p.

But again every orbital in the p sublevels will have that one angular node created by the 2p orbitals.


The 3d orbitals, similar to the 2p orbitals, create 2 angular nodes. These angular nodes affect any electron in a d-sublevel no matter what energy level they are in.

The 3d orbitals also give off a radial node to the orbitals in 4d.


The 4f orbitals have 3 angular nodes and would give off a radial node to the 5f orbitals.


Now let’s answer the question of why 4s gains electrons before 3d, and why 6s and 5p gain electrons before 4f.

Remember an orbital is only a prediction of where the electrons are most likely to be over time, and we know they can’t exist at the nodes.

This means that they have to exist at other energy levels some of the time.

The way sublevels are ordered on the periodic table is by the probability of finding an electron close to the nucleus. 

So 4s is higher in energy than 3d, but a 4s electron is also more likely to be close to the nucleus.

How is this possible?

Penetration Effect/Penetrating Nodes

The penetration effect is an electron’s ability to penetrate a node and exist on lower energy levels close to the nucleus. 

We talked about radial nodes and angular nodes, to put it bluntly, angular nodes are harder to penetrate than radial nodes.

The fact that the s sublevels don’t have an angular node means they can access lower energy levels easier than the other sublevels.

You might be thinking the p sublevel has an angular node but you don’t see 3p before 2s on the periodic table. 

The logic that seems to stand is if there is a disparity of two angular nodes between different sublevels then that means the sublevel with fewer angular nodes will have orbitals one energy level higher with a greater probability of having electrons close to the nucleus.

Since the s sublevels have two fewer angular nodes than the d sublevels, 4s has a greater probability of being close to the nucleus than 3d.

This logic works for 5p and 4f as well, with the p sublevels having 1 angular node and the f sublevels having 3 angular nodes.

If there is a disparity of 3 angular nodes then the sublevels with lesser angular nodes will have orbitals two energy levels ahead, like 6s and 4f.

Radial Probability Distribution 

The radial probability distribution is a graphical representation of how an electron in a specified sublevel can exist on lower sublevels. 

These graphs show the sublevels (up to 4s) and their probability of having their electrons close to the nucleus, which is what determines the order for electron configurations. 

Although the periodic table is set up in such a way where the electrons in an atom are situated in specific orbitals, what goes on at the quantum level is much more abstract. I think it’s best to develop a solid understanding of calculus and classical physics before delving into quantum concepts. 

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