The Periodic Table | Calm Bedtime Reading For Sleep

Relax with this calm bedtime reading about the periodic table, perfect for sleep and easing insomnia. Drift off as you learn about the history, structure, and significance of this fundamental tool in chemistry.

In this gentle episode, we explore how the periodic table organizes elements, the contributions of scientists like Mendeleev, and the properties that define each group. Benjamin’s soothing voice guides you through the facts in a peaceful, thoughtful tone—no whispering, no hypnosis, just calm education to help you relax and unwind. This episode is ideal for relieving stress, anxiety, and sleepless nights. Press play, get cozy, and let your mind settle.

Happy sleeping!

Transcript

Welcome to the I Can't Sleep podcast,

Where I help you drift off one fact at a time.

I'm your host,

Benjamin Boster,

And today's episode is about the periodic table.

The periodic table,

Also known as the periodic table of the elements,

Is an ordered arrangement of the chemical elements into rows,

Or periods,

And columns,

Or groups.

An icon of chemistry,

The periodic table is widely used in physics and other sciences.

It is a depiction of the periodic law,

Which states that when the elements are arranged in order of their atomic numbers,

An approximate recurrence of their properties is evident.

The table is divided into four roughly rectangular areas called blocks.

Elements in the same group tend to show similar chemical characteristics.

Vertical,

Horizontal,

And diagonal trends characterize the periodic table.

Metallic character increases going down a group and from right to left across a period.

Non-metallic character increases going from the bottom left of the periodic table to the top right.

The first periodic table to become generally accepted was that of the Russian chemist Dmitry Mendeleev in 1869.

He formulated the periodic law as a dependence of chemical properties on atomic mass.

As not all elements were then known,

There were gaps in his periodic table,

And Mendeleev successfully used the periodic law to predict some properties of some of the missing elements.

The periodic law was recognized as a fundamental discovery in the late 19th century.

It was explained early in the 20th century with the discovery of the atomic numbers and associated pioneering work in quantum mechanics,

Both ideas serving to illuminate the internal structure of the atom.

A recognizably modern form of the table was reached in 1945 with Glenn T.

Seaborg's discovery that the actinides were in fact F-block rather than D-block elements.

The periodic table and law have become a central and indispensable part of modern chemistry.

The periodic table continues to evolve with the progress of science.

In nature,

Only elements up to atomic number 94 exist.

To go further,

It was necessary to synthesize new elements in the laboratory.

By 2010,

The first 118 elements were known,

Thereby completing the first seven rows of the table.

However,

Chemical characterization is still needed for the heaviest elements to confirm that their properties match their positions.

New discoveries will extend the table beyond these seven rows,

Though it is not yet known how many more elements are possible.

Moreover,

Theoretical calculations suggest that this unknown region will not follow the patterns of the known part of the table.

Some scientific discussion also continues regarding whether some elements are correctly positioned in the table.

Many alternative representations of the periodic law exist,

And there is some discussion as to whether there is an optimal form of the periodic table.

Each chemical element has a unique atomic number,

Z for Zoll,

German for number,

Representing the number of protons in its nucleus.

Each distinct atomic number therefore corresponds to a class of atom.

These classes are called the chemical elements.

The chemical elements are what the periodic table classifies and organizes.

Hydrogen is the element with atomic number 1,

Helium,

Atomic number 2,

Lithium,

Atomic number 3,

And so on.

Each of these names can be further abbreviated by a one- or two-letter chemical symbol.

Those for Hydrogen,

Helium,

And Lithium are respectively H,

He,

And Li.

Neutrons do not affect the atom's chemical identity,

But do affect its weight.

Atoms with the same number of protons but different numbers of neutrons are called isotopes of the same chemical element.

Naturally occurring elements usually occur as mixes of different isotopes.

Since each isotope usually occurs with a characteristic abundance,

Naturally occurring elements have well-defined atomic weights,

Defined as the average mass of a naturally occurring atom of that element.

All elements have multiple isotopes,

Variants,

With the same number of protons but different numbers of neutrons.

For example,

Carbon has three naturally occurring isotopes.

All of its atoms have 6 protons,

And most have 6 neutrons as well,

But about 1% have 7 neutrons,

And a very small fraction have 8 neutrons.

Isotopes are never separated in the periodic table.

They are always grouped together under a single element.

When atomic mass is shown,

It is usually the weighted average of naturally occurring isotopes.

But if no isotopes occur naturally in significant quantities,

The mass of the most stable isotope usually appears often in parentheses.

In the standard periodic table,

The elements are listed in order of increasing atomic number.

A new row,

Period,

Is started when a new electron shell has its first electron.

Columns,

Groups,

Are determined by the electron configuration of the atom.

Elements of the same number of electrons in a particular subshell fall into the same columns.

E.

G.

Oxygen,

Sulfur,

And Selenium are in the same column because they all have 4 electrons in the outermost P subshell.

Elements with similar chemical properties generally fall into the same group in the periodic table,

Although in the F block,

And to some respect in the D block,

The elements in the same period tend to have similar properties as well.

Thus,

It is relatively easy to predict the chemical properties of an element if one knows the properties of the elements around it.

Today,

118 elements are known,

The first 94 of which are known to occur naturally on Earth.

The remaining 24,

Americium to Organescent,

95 to 118,

Occur only when synthesized in laboratories.

Of the 94 naturally occurring elements,

83 are primordial,

And 11 occur only in decay chains of primordial elements.

A few of the latter are so rare that they were not discovered in nature,

But were synthesized in the laboratory before it was determined that they exist in nature.

Technetium,

Element 43,

Promethium,

Element 61,

Astatine,

Element 85,

Neptunium,

Element 93,

And Plutonium,

Element 94.

No element heavier than Einsteinium,

Element 99,

Has ever been observed in macroscopic quantities in its pure form,

Nor has Astatine.

Francium,

Element 87,

Has been only photographed in the form of light emitted from microscopic quantities.

Of the 94 natural elements,

80 have a stable isotope,

And one more,

Bismuth,

Has an almost stable isotope,

With a half-life of 2.

01 x 10 to the 19th years,

Over a billion times the age of the universe.

Two more,

Thorium and Uranium,

Have isotopes undergoing radioactive decay,

With a half-life comparable to the age of the Earth.

The stable elements,

Plus Bismuth,

Thorium,

And Uranium,

Make up the 83 primordial elements that survive from the Earth's formation.

The remaining 11 natural elements decay quickly enough that their continued trace occurrence rests primarily on being constantly regenerated as intermediate products of the decay of Thorium and Uranium.

All 24 known artificial elements are radioactive.

Under an international naming convention,

The groups are numbered numerically from 1 to 18 from the left-most column,

The alkali metals,

To the right-most column,

The noble gases.

The F-block groups are ignored in this numbering.

Groups can also be named by their first element,

E.

G.

The Scandium group for group 3.

Previously,

Groups were known by Roman numerals.

In the United States,

The Roman numerals were followed by either an A if the group was in the S or P-block,

Or a B if the group was in the D-block.

The Roman numerals used correspond to the last digit of today's naming convention,

E.

G.

The group 4 elements were group Roman numeral 4B,

And the group 14 elements were group Roman numeral 4A.

In Europe,

A was used for groups 1 through 7,

And B was used for groups 11 through 17.

In addition,

Groups 8,

9,

And 10 used to be treated as one triple-sized group,

Known collectively in both notations as group Roman numeral 8.

In 1988,

The new IUPAC,

International Union of Pure and Applied Chemistry Naming System,

1 through 18,

Was put into use,

And the old group names,

Roman numerals 1 through 8,

Were deprecated.

For reasons of space,

The periodic table is commonly presented with the F-block elements cut out,

And positioned as a distinct part below the main body.

This reduces the number of element columns from 32 to 18.

Both forms represent the same periodic table.

The form with the F-block included in the main body is sometimes called the 32-column or long form.

The form with the F-block cut out,

The 18-column or medium-long form.

The 32-column form has the advantage of showing all elements in their correct sequence,

But it has the disadvantage of requiring more space.

The form chosen is an editorial choice,

And does not imply any change of scientific claim or statement.

For example,

When discussing the composition of group 3,

The options can be shown equally,

Unprejudiced,

In both forms.

Periodic tables usually at least show the element symbols.

Many also provide supplementary information about the elements,

Either via color coding or as data in the cells.

Tables may include extra information such as the names and atomic numbers of the elements,

Their blocks,

Natural occurrences,

Standard atomic weight,

States of matter,

Melting and boiling points,

Densities,

As well as provide different classifications of the elements.

The periodic table is a graphic description of the periodic law,

Which states that the properties and atomic structures of the chemical elements are a periodic function of their atomic number.

Elements are placed in the periodic table according to their electron configurations,

The periodic reoccurrences of which explain the trends in properties across the periodic table.

An electron can be thought of as inhabiting an atomic orbital,

Which characterizes the probability it can be found in any particular region around the atom.

Their energies are quantized,

Which is to say that they can only take discrete values.

Furthermore,

Electrons obey the Pauli exclusion principle.

Different electrons must always be in different states.

This allows classification of the possible states an electron can take in various energy levels known as shells,

Divided into individual subshells,

Which each contain one or more orbitals.

Each orbital can contain up to two electrons.

They are distinguished by a quantity known as spin,

Conventionally labeled up or down.

In a cold atom,

One in its ground state,

Electrons arrange themselves in such a way that the total energy they have is minimized by occupying the lowest energy orbitals available.

Only the outermost electrons,

Valence electrons,

Have enough energy to break free of the nucleus and participate in chemical reactions with other atoms.

The others are called core electrons.

Elements are known with up to the first seven shells occupied.

The first shell contains only one orbital,

A spherical s-orbital.

As it is in the first shell,

This is called the 1s orbital.

This can hold up to two electrons.

The second shell similarly contains a 2s orbital,

And it also contains three dumbbell-shaped 2p orbitals,

And can thus fill up to eight electrons.

2 times 1 plus 2 times 3 equals 8.

The third shell contains one 3s orbital,

Three 3p orbitals,

And five 3d orbitals,

And thus has a capacity of 2 times 1 plus 2 times 3 plus 2 times 5,

Which equals 18.

The fourth shell contains one 4s orbital,

Three 4p orbitals,

Five 4d orbitals,

And seven 4f orbitals,

Thus leading to a capacity of 2 times 1 plus 2 times 3 plus 2 times 5 plus 2 times 7,

Which equals 32.

Higher shells contain more types of orbitals that continue the pattern,

But such types of orbitals are not filled in the ground states of known elements.

The subshell types are characterized by the quantum numbers.

Four numbers describe an orbital in an atom completely.

The principal quantum number n,

The azimuthal quantum number l,

The orbital type,

The orbital magnetic quantum number m l,

And the spin magnetic quantum number m s.

The sequence in which the subshells are filled is given in most cases by the Aufbau principle,

Also known as the Madelung or Klitschkofsky rule,

After Erwin Madelung and Veslovalod Klitschkofsky respectively.

This rule was first observed empirically by Madelung,

And Klitschkofsky and later authors gave it theoretical justification.

The shells overlap in energies,

And the Madelung rule specifies the sequence of filling according to 1s,

Then 2s,

Then 2p,

Followed by 3s,

Then 3p,

4s,

3d,

4p,

5s,

4d,

5p,

6s,

4f,

5d,

6p,

7s,

5f,

6d,

And 7p.

Phrased differently,

Electrons enter orbitals in order of increasing n plus l,

And if two orbitals are available with the same value of n plus l,

The one with a lower n is occupied first.

In general,

Orbitals with the same value of n plus l are similar in energy,

But in the case of the s-orbital,

With l equals zero,

Quantum effects raise their energy to approach that of the next n plus l group.

Hence,

The periodic table is usually drawn to begin each row,

Often called a period,

With the filling of a new s-orbital,

Which corresponds to the beginning of a new shell.

Thus,

With the exception of the first row,

Each period length appears twice,

2,

8,

8,

18,

18,

32,

32,

And so on.

The overlaps get quite close at the point where the d-orbitals enter the picture,

And the order can shift slightly with atomic number and atomic charge.

Starting from the simplest atom,

This lets us build up the periodic table one at a time in order of atomic number,

By considering the cases of single atoms.

In hydrogen,

There is only one electron,

Which must go in the lowest energy orbital 1s.

This electron configuration is written as 1s1,

Where the superscript indicates the number of electrons in the subshell.

Helium adds a second electron,

Which also goes into 1s,

Completely filling the first shell and giving the configuration 1s2.

Starting from the third element,

Lithium,

The first shell is full,

So its third electron occupies a 2s orbital,

Giving a 1s2,

2s1 configuration.

The 2s electron is lithium's only valence electron,

As the 1s subshell is now too tightly bound to the nucleus to participate in chemical bonding to other atoms.

Such a shell is called a core shell.

The 1s subshell is a core shell for all elements from lithium onward.

The 2s subshell is completed by the next element beryllium,

1s2,

2s2.

The following elements then proceed to follow the 2p subshell.

Boron,

1s2,

2s2,

2p1,

Puts its new electron in a 2p orbital.

Carbon,

1s2,

2s2,

2p2,

Fills a second 2p orbital.

And with nitrogen,

1s2,

2s2,

2p3,

All three 2p orbitals become singly occupied.

This is consistent with Hund's rule,

Which states that atoms usually prefer to singly occupy each orbital of the same type before filling them with the second electron.

Oxygen,

1s2,

2s2,

2p4,

Fluorine,

1s2,

2s2,

2p5,

And Neon,

1s2,

2s2,

2p6,

Then complete the already singly filled 2p orbitals.

The last of these fills the second shell completely.

Starting from element 11,

Sodium,

The second shell is full,

Making the second shell a core shell for this and all heavier elements.

The 11th electron begins the filling of the third shell by occupying a 3s orbital,

Giving a configuration of 1s2,

2s2,

2p6,

3s1 for sodium.

This configuration is abbreviated Ne3s1,

Where Ne represents Neon's configuration.

Magnesium,

Ne3s2,

Finishes this 3s orbital,

And the following six elements,

Aluminum,

Silicon,

Phosphorus,

Sulfur,

Chlorine,

And argon,

Fill the three 3p orbitals.

Ne3s2,

3p1,

Through Ne3s2,

3p6.

This creates an analogous series in which the outer shell structures of sodium through argon are analogous to those of lithium through neon,

And is the basis for the periodicity of chemical properties that the periodic table illustrates.

At regular but changing intervals of atomic numbers,

The properties of the chemical elements approximately repeat.

The first 18 elements can thus be arranged as the start of a periodic table.

Elements in the same column have the same number of valence electrons,

And have analogous valence electron configurations.

These columns are called groups.

The single exception is helium,

Which has two valence electrons,

Like beryllium and magnesium,

But is typically placed in the column of neon and argon,

To emphasize that its outer shell is full.

Some contemporary authors question even this single exception,

Preferring to consistently follow the valence configurations and place helium over beryllium.

There are eight columns in this periodic table fragment,

Corresponding to at most eight outer shell electrons.

A period begins when a new shell starts filling.

Finally,

The coloring illustrates the blocks.

The elements in the s-block,

Colored red,

Are filling s-orbitals,

While those in the p-block,

Colored yellow,

Are filling p-orbitals.

Starting the next row for potassium and calcium,

The 4s subshell is the lowest in energy,

And therefore they fill it.

Potassium adds one electron to the 4s shell,

Ar4s1,

And calcium then completes it,

Ar4s2.

However,

Starting with scandium,

Ar3d1 4s2,

The 3d subshell becomes the next highest in energy.

The 4s and 3d subshells have approximately the same energy,

And they compete for filling the electrons,

And so the occupation is not quite consistently filling the 3d orbitals one at a time.

The precise energy ordering of 3d and 4s changes along the row,

And also changes depending on how many electrons are removed from the atom.

For example,

Due to the repulsion between the 3d electrons and the 4s ones,

At chromium the 4s energy level becomes slightly higher than 3d,

And so it becomes more profitable for a chromium atom to have an Ar3d5 4s1 configuration than an Ar3d4 4s2 one.

A similar anomaly occurs at copper,

Whose atom has an Ar3d10 4s1 configuration rather than the expected Ar3d9 4s2.

These are violations of the Modulung Rule.

Such anomalies,

However,

Do not have any chemical significance.

Most chemistry is not about isolated gaseous atoms,

And the various configurations are so close in energy to each other that the presence of a nearby atom can shift the balance.

Therefore,

The periodic table ignores them,

And considers only idealized configurations.

At zinc,

Ar3d10 4s2,

The 3d orbitals are completely filled with a total of 10 electrons.

Next comes the 4p orbitals,

Completing the row,

Which are filled progressively by gallium Ar3d2.

Ar3d10 4s2 4p1 through krypton Ar3d10 4s2 4p6 in a manner analogous to the previous p-block elements.

From gallium onwards,

The 3d orbitals form part of the electronic core,

And no longer participate in chemistry.

The s- and p-block elements,

Which fill their outer shells,

Are called main group elements.

The d-block elements,

Which fill an inner shell,

Are called transition elements,

Or transition metals,

Since they are all metals.

The next 18 elements fill the 5s orbitals,

Rubidium and strontium,

Then 4d,

Yttrium through cadmium,

Again with a few anomalies along the way,

And then 5p,

Indium through xenon.

Again,

From indium onward,

The 4d orbitals are in the core.

Hence,

The fifth row has the same structure as the fourth.

The sixth row of the table likewise starts with two s-block elements,

Cesium and barium.

After this,

The first f-block elements begin to appear,

Starting with lanthanum.

These are sometimes termed inner transition elements,

As there are now not only 4f,

But also 5d and 6s subshells at similar energies.

Competition occurs once again with many irregular configurations.

This resulted in some dispute about where exactly the f-block is supposed to begin,

But most who study the matter agree that it starts at lanthanum in accordance with the Aufbau principle.

Even though lanthanum does not itself fill the 4f subshell as a single atom because of repulsion between electrons,

Its 4f orbitals are low enough in energy to participate in chemistry.

At yttrium,

The 7 4f orbitals are completely filled with 14 electrons.

Thereafter,

A series of 10 transition elements,

Lutetium through mercury follows,

And finally 6 main group elements,

Thallium through radon,

Complete the period.

From lutetium onwards,

The 4f orbitals are in the core,

And from thallium onwards,

So are the 5d orbitals.

The 7th row is analogous to the 6th row.

7s fills,

Francium and radium,

Then 5f,

Actinium to nobelium,

Then 6d,

Lorencium to copernicium,

And finally 7p,

Nionium to organesson.

Starting from the lorencium,

The 5f orbitals are in the core,

And probably the 6d orbitals join the core starting from nionium.

Again,

There are a few anomalies along the way.

For example,

As single atoms,

Neither actinium or thorium actually fills the 5f subshell.

The lorencium does not fill the 6d shell.

But all these subshells can still become filled in chemical environments.

For a very long time,

The 7th row was incomplete as most of its elements do not occur in nature.

The missing elements beyond uranium started to be synthesized in the laboratory in 1940,

When neptunium was made.

However,

The first element to be discovered by synthesis rather than in nature was technetium in 1937.

The row was completed with the synthesis of tennessine in 2010.

The last element,

Organesson,

Had already been made in 2002.

And the last elements in this 7th row were given names in 2016.

Although the modern periodic table is standard today,

The placement of the period 1 elements hydrogen and helium remains an open issue under discussion,

And some variation can be found.

Following their respective S1 and S2 electron configurations,

Hydrogen would be placed in group 1,

And helium would be placed in group 2.

The group 1 placement of hydrogen is common,

But helium is almost always placed in group 18 with the other noble gases.

The debate has to do with conflicting understandings of the extent to which chemical or electronic properties should decide periodic table placement.

Like the group 1 metals,

Hydrogen has one electron in its outermost shell,

And typically loses its only electron in chemical reactions.

Hydrogen has some metal-like chemical properties,

Being able to displace some metals from their salts,

But it forms a diatomic,

Non-metallic gas at standard conditions,

Unlike the alkali metals which are reactive solid metals.

This,

And hydrogen's formation of hydrides,

In which it gains an electron,

Brings it close to the properties of the halogens,

Which do the same,

Though it is rarer for hydrogen to form H- than H+.

Moreover,

The lightest two halogens,

Fluorine and chlorine,

Are gaseous like hydrogen at standard conditions.

Some properties of hydrogen are not a good fit for either group.

Hydrogen is neither highly oxidizing nor highly reducing,

And is not reactive with water.

Hydrogen thus has properties corresponding to both those of the alkali metals and the halogens,

But matches neither group perfectly,

And is thus difficult to place by its chemistry.

Therefore,

While the electronic placement of hydrogen in group 1 predominates,

Some rarer arrangements show either hydrogen in group 17,

Duplicate hydrogen in both groups 1 and 17,

Or float it separately from all groups.

This last option has nonetheless been criticized by the chemist and philosopher of science,

Eric Skerry,

On the grounds that it appears to imply that hydrogen is above the periodic law altogether,

Unlike all the other elements.

Helium is the only element that routinely occupies a position in the periodic table that is not consistent with its electronic structure.

It has two electrons in its outermost shell,

Whereas the other noble gases have eight,

And it is an s-block element,

Whereas all other noble gases are p-block elements.

However,

It is unreactive at standard conditions,

And has a full outer shell.

These properties are like the noble gases in group 18,

But not at all like the reactive alkaline earth metals of group 2.

For these reasons,

Helium is nearly universally placed in group 18,

Which its properties best match.

A proposal to move helium to group 2 was rejected by IUPAC in 1988 for these reasons.

Nonetheless,

Helium is still occasionally placed in group 2 today,

And some of its physical and chemical properties are closer to the group 2 elements and support the electronic placement.

Solid helium crystallizes in a hexagonal close-packed structure,

Which matches beryllium and magnesium in group 2,

But not the other noble gases in group 18.

Recent theoretical developments in noble gas chemistry,

In which helium is expected to show slightly less inertness than neon and to form HeO-LiF2 with a structure similar to the analogous beryllium compound,

But with no expected neon analog,

Have resulted in more chemists advocating a placement of helium in group 2.

This relates to the electronic argument,

As the reason for greater inertness is repulsion from its filled p-shell that helium lacks,

Though realistically it is unlikely that helium-containing molecules will be stable outside extreme low temperature conditions,

Around 10 Kelvin.

The first row anomaly in the periodic table has additionally been cited to support moving helium to group 2.

It arises because the first orbital of any type is usually small,

Since unlike its higher analogs,

It does not experience inter-electronic repulsion from a smaller orbital of the same type.

This makes the first row of elements in each block usually small,

And such elements tend to exhibit characteristic kinds of anomalies for their group.

Some of the chemists arguing for the repositioning of helium have pointed out that helium exhibits these anomalies if it is placed in group 2,

But not if it is placed in group 18.

On the other hand,

Neon,

Which would be the first group 18 element if helium was removed from that spot,

Does exhibit those anomalies.

The relationship between helium and beryllium is then argued to resemble that between hydrogen and lithium,

A placement which is much more commonly accepted.

For example,

Because of this trend in the sizes of orbitals,

A large difference in atomic radii between the first and second members of each main group is seen in groups 1 and 13-17.

It exists between neon and argon,

And between helium and beryllium,

But not between helium and neon.

This similarly affects the noble gases,

Boiling points,

And solubilities in water,

Where helium is too close to neon,

And the large difference characteristic between the first two elements of a group appears only between neon and argon.

Moving helium to group 2 makes this trend consistent in groups 2 and 18 as well,

By making helium the first group 2 element and neon the first group 18 element.

Both exhibit the characteristic properties of kinosymmetric first element of a group.

The group 18 placement of helium nonetheless remains near-universal due to its extreme inertness.

Additionally,

Tables that float both hydrogen and helium outside all groups may rarely be encountered.

That concludes this episode about the periodic table of elements.

Reviews

Benjamin, how did you even stay awake to record this one? 😂😂😂
Thank you, this put me right out! 😻

Yep, learnt a couple of facts and fell asleep quickly. This track delivered as advertised ! 😁Perfect! Thank you 🙏

I loved this. I almost didn’t want to fall asleep. Thank you.