I Can’t Sleep - Atoms | Calm Bedtime Reading for Sleep

Episode Date: July 4, 2024

Drift off with this calm bedtime reading on atoms, created to ease insomnia and bring peace to restless nights. In this gentle episode, Benjamin explores the fundamental building blocks of matter, from early philosophical ideas to modern scientific understanding. His soothing cadence blends education with relaxation—no whispers, no hypnosis, just calm, fact-filled narration to help quiet your mind. Whether you’re struggling with sleeplessness, stress, or anxiety, this peaceful reading invites you to press play, unwind, and drift into restful sleep while learning about the fascinating world of atoms. Want More? Request a Topic: https://www.icantsleeppodcast.com/request-a-topic Ad-Free Episodes: https://icantsleep.supportingcast.fm/ Shop Sleep-Friendly Products: https://www.icantsleeppodcast.com/sponsors Join the Discussion on Discord: https://discord.gg/myhGhVUhn7 This content is derived from the Wikipedia article on Atoms, available under the Creative Commons Attribution-ShareAlike (CC BY-SA) license. Read the full article: Wikipedia – Atoms. Happy sleeping! Learn more about your ad choices. Visit megaphone.fm/adchoices

You're listening to a Glassbox media podcast. What if I told you that most of the modern day self-help advice you've been hearing could actually make you worse? The key to a better life isn't about feel-good gimmicks that sound catchy. The Mentally Stronger Podcast gives you access to a licensed therapist who shares science-backed tools that will actually change your life. Hi, I'm Amy Morin, psychotherapist, mental strength trainer, and international best-selling author. In each episode, we cover research-back strategies, like how to stop relying on willpower and start creating habits for lasting change. And the five mental strength-building exercises you can do from your couch.

I also speak to world-class experts like Dr. Nicole Kane, who shares how to permanently heal anxiety by addressing the root cause. With over 200 episodes in our catalog, this podcast is for you if you're ready to crush self-doubt, conquer challenges, and become stronger than ever with therapist-approved strategies that can change your life. Listen to Mentally Stronger with Therapist Amy Morin, wherever you get your podcasts. Welcome to the I Can't Sleep podcast, where I read random articles from across the web to bore you to sleep with my soothing voice. I'm your host, Benjamin Boster. Today's episode is from a Wikipedia article titled, Adam.

Atoms are the basic particles of the chemical elements. An atom consists of a nucleus of protons and generally neutrons, surrounded by an electromagnetically bound swarm of electrons. The chemical elements are distinguished from each other by the number of protons that are in their atoms. For example, any atom that contains 11 protons is sodium, and any atom that contains 29 protons is copper. Atoms with the same number of protons, but a different number of neutrons, are called isotopes of the same element.

atoms are extremely small, typically around 100 pichromedars across. A human air is about a million carbon atoms wide. This is smaller than the shortest wavelength of visible light, which means humans cannot see atoms with conventional microscopes. atoms are so small that accurately predicting their behavior, using classical physics, is not possible due to quantum effects. More than 99.94% of an atom's mass is in the nucleus. Protons have a positive electric charge,

and neutrons have no charge, so the nucleus is positively charged. The electrons are negatively charged, and this supposing charge is what binds them to the nucleus. If the numbers of protons and electrons are equal, as they normally are, then the atom is electrically neutral as a whole. If an atom has more electrons than protons,

then it has an overall negative charge and is called a negative ion, or anion. Conversely, if it has more protons than electrons, it has a positive charge, and is called a positive ion, or a cation. The electrons of an atom are attracted to the protons in an atomic nucleus by the electromagnetic force. The protons and neutrons in the nucleus are attracted to each other by the nuclear force. This force is usually stronger than the electromagnetic force that repels the positively charged protons from one another. Under certain circumstances, the repelling electromagnetic force becomes stronger than the nuclear force. In this case, the nucleus splits and leaves behind different elements.

This is a form of nuclear decay. atoms can attach to one or more other atoms by chemical bonds to form chemical compounds to form chemical consequences. compounds, such as molecules or crystals. The ability of atoms to attach and attach from each other is responsible for most of the physical changes observed in nature. Chemistry is the science that studies these changes. The basic idea that matter is made up of tiny indivisible particles is an old idea that appeared in many ancient cultures. The word atom is derived from the ancient Greek word atomous. which means uncutable.

This ancient idea was based in philosophical reasoning rather than scientific reasoning. Modern atomic theory is not based on these old concepts. In the early 19th century, the scientist John Dalton noticed that chemical substances seemed to combine with each other by a basic unit of weight, and he decided to use the word atom to refer to these units as he thought they were indivisible in essence. In the early 1800s, the English chemist John Dalton compiled experimental data gathered by him and other scientists

and discovered a pattern now known as the Law of Multiple Proportions. He noticed that in any group of chemical compounds, which all contain two particular chemical elements, the amount of element A per measure of element B will differ across these compounds by ratios of small whole numbers. This pattern suggested that the elements combine with each other in multiples of a basic unit of weight, with each element having a unit of unique weight.

Dalton decided to call these units atoms. For example, there are two types of tin oxide. One is a gray powder that is 88.1% tin and 11.9% oxygen, and the other is a white powder that is 78.7% tin and 21.3% oxygen. Adjusting these figures in the gray powder, there is about 13.5 grams of oxygen for every 100 grams of 10. And in the white powder, there is about 27 grams of oxygen for every 100 grams of 10. 13.5 and 27 form a ratio of 1 to 2. Dalton concluded that in the gray oxide, there is one atom of oxygen for every atom of 10.

And in the white oxide, there are two atoms of oxygen for every atom of 10, SNO and SNO2. Dalton also analyzed iron oxides. There is one type of iron oxide that is a black powder, which is 78.1% iron and 21.9% oxygen. and there is another iron oxide that is a red powder, which is 70.4% iron and 29.6% oxygen. Adjusting these figures in the black powder, there is about 28 grams of oxygen for every 100 grams of iron, and in the red powder, there is about 42 grams of oxygen for every 100 grams of iron. 28 and 42 form a ratio of 2 to 3. Dalton concluded that in these oxides, for every two atoms of iron, there are two or three atoms of oxygen respectively, F.E.202 and F.E.203.

As a final example, nitrous oxide is 63.3% nitrogen and 36.7% oxygen. Nitric oxide is 44.05% nitrogen. and 55.95% oxygen, and nitrogen dioxide is 29.5% nitrogen and 70.5% oxygen. Adjusting these figures in nitrous oxide, there is 80 grams of oxygen for every 140 grams of nitrogen. In nitric acid, there is about 160 grams of oxygen for every 140 grams of nitrogen. and in nitrogen dioxide, there is 320 grams of oxygen for every 140 grams of nitrogen. 80, 160, and 320 form a ratio of 1 to 2 to 4. The respective formulas for these oxides are N2, N2O, N2. In 1897, J.J. Thompson discovered that cathode rays are not a form of light,

but made of negatively charged particles because they can be deflected by electric and magnetic fields. He measured these particles to be at least a thousand times lighter than hydrogen, the lightest atom. He called these new particles core puzzles, but they were later renamed electrons, since these are the particles that carry electricity. Thompson also showed that electrons were identical to particles given off by photoelectric

and radioactive materials. Thompson explained that an electric current is the passing of electrons from one atom to the next, and when there was no current, the electrons embedded themselves in the atoms. This, in turn, meant that atoms were not indivisible as scientists thought. The atom was composed of electrons whose negative charge was balanced out by some source of positive charge to create an electrically neutral atom. Ions, Thompson explained, must be atoms which have an excess or shortage of electrons. The electrons in the atom logically had to be balanced out by a commensurate amount of positive charge,

but Thompson had no idea where this positive church came from, so he tentatively proposed that this positive charge was everywhere in the atom, the atom being in the shape of a sphere. Following from this, he imagined the balance of electrostatic forces would distribute the electrons throughout the sphere in a more or less even manner. Thompson's model is popularly known as the Plum Pudding model, though neither Thompson nor his colleagues used this analogy. Thompson's model was incomplete. It was unable to predict any of the other known properties of atoms, such as emission, spectra, and valencies. It was soon rendered obsolete by the discovery of the atomic nucleus.

Between 1908 and 1913, Ernest Rutherford and his colleagues, Hans Geiger and Ernest Marston, performed a series of experiments in which they bombarded thin foils of metal with a beam of alpha particles. They did this to measure the scattering patterns of the alpha particles. They spotted a small number of alpha particles being deflected by angles greater than 90 degrees. This shouldn't have been possible, according to the Thompson model of the atom, whose charges were too diffused to produce a sufficiently strong electric field. The deflections should have all been negligible. Rutherford proposed that the positive charge of the atom, along with most of the atom's mass,

is concentrated in a tiny nucleus at the center of the atom. Only such an intense concentration of positive charge, anchored by its high mass, and separated from the negative charge, could produce an electric field that could deflect the alpha particles so strongly. In 1913, the physicist Niels Bohr proposed a model in which the electrons of an atom were assumed to orbit the nucleus, but could only do so in a finite set of orbits, and could jump between these orbits only in discrete changes of energy, corresponding to absorption or radiation of a photon.

This quantization was used to explain why the electron orbits are stable, given that in classical physics, charges in acceleration, including circular motion, lose kinetic energy, which is emitted as electromagnetic radiation, and why elements absorb and emit electromagnetic radiation in discrete spectra. Back in 1815, William Prout noticed that the atomic weights of the elements were all multiples of hydrogen's atomic weight, which is true if one takes isotopes into account. In 1911, based on his alpha particles scattering experiments, Rutherford estimated that an atom's nuclear charge expressed in units of hydrogen's nuclear charge

was about half its atomic weight. In 1913, Henry Mosley discovered that the frequency of x-ray emissions from an excited atom was a function of the element's atomic number and the charge of a hydrogen nucleus. In 1917, Rutherford bombarded nitrogen gas with alpha particles and observed hydrogen nuclei being emitted from the gas and concluded that the hydrogen nuclei emerged from the nuclei of the nitrogen atoms. In effect, he had split the atom. These observations led Rutherford to conclude that the hydrogen nucleus is a singular particle

with a positive charge equal to the electron's negative charge. He named this particle proton, the element's atomic number, which up to that point had been defined as the element's position on the periodic table, was evidently the number of protons it had in the nucleus. The atomic weight of each element was higher than its atomic number, so Rutherford hypothesized that the surplus mass was carried by unknown particles with no charge, with a mass equal to that of the proton. In 1928, Walter Boat observed that beryllium emitted a highly penetrating electrically neutral radiation when bombarded with alpha particles. It was later discovered that this radiation could knock hydrogen atoms out of paraffin wax. Initially, it was thought to be high-energy gamma radiation, since gamma radiation had a similar effect on electrons and metals. But James Chadwick found that the ionization effect was too strong for it to be due to electrical. electromagnetic radiation, so long as energy and momentum were conserved in the interaction.

In 1932, Chadwick exposed various elements, such as hydrogen and nitrogen to the mysterious beryllium radiation, and by measuring the energies of the recoiling charged particles, he deduced that the radiation was actually composed of electrically neutral particles, which could not be massless like the gamma ray, but instead were required. to have a mass similar to that of a proton. Chadwick now claimed these particles as Rutherford's neutrons. In 1925, Werner Heisenberg published the first consistent mathematical formulation of quantum mechanics, matrix mechanics.

One year earlier, Lewis de Broly had proposed that all particles behave like waves to some extent, and in 1926 Erwin Schrodinger used his idea to develop a Schrodinger equation, which describes electrons as three-dimensional waveforms rather than points in space. A consequence of using waveforms to describe particles is that it is mathematically impossible to obtain precise values for both the position and momentum of a particle at a given point in time. This became known as the uncertainty principle, formulated by Werner-Heisenberg in 1927. In this concept, for a given accuracy in measuring a position, one could only obtain a range of probable values for momentum and vice versa.

Thus, the planetary model of the atom was discarded in favor of one that described atomic orbital zones around the nucleus, where a given electron is most likely to be found. This model was able to explain observations of atomic behavior that previous models could not, such as certain structural and spectral patterns of atoms larger than hydrogen. Though the word atom originally denoted a particle that cannot be cut into smaller particles, in modern scientific usage, the atom is composed of various subatomic particles. The constituent particles of an atom are the electron, the electron, the proton, and the neutron. The electron is the least massive of these particles by four orders of magnitude,

at 9.11 times 10 to the negative 31 kilograms, with a negative electrical charge in a size that is too small to be measured using available techniques. It was the lightest particle with a positive rest mass measured until the discovery of neutrino mass. Under ordinary conditions, electrons are bound to the positively charged nucleus by the attraction created from opposite electric charges. If an atom has more or fewer electrons than its atomic number, then it becomes respectively negatively or positively charged as a whole.

A charged atom is called an ion. Electrons have been known since the late 19th century, mostly thanks to J.J. Thompson. Protons have a positive charge and a mass of 1.6726 times 10 to the negative kilograms. The number of protons in an atom is called its atomic number. Ernest Rutherford 1919 observed that nitrogen under alpha particle bombardment ejects what appeared to be a hydrogen nuclei. By 1920, he had accepted that the hydrogen nucleus is a distinct particle within the atom and named it proton. Neutrons have no electrical charge and have a mass of 1.6749

times 10 to the negative 27 kilograms. Neutrons are the heaviest of the three constituent particles, but their mass can be reduced by the nuclear binding energy. Neutrons and protons, collectively known as nucleons, have comparable dimensions, although the surface of these particles is not sharply defined. The neutron was discovered in 1932 by the English physicist James Chadwick. In the standard model of physics, electrons are truly elementary particles with no internal structure,

whereas protons and neutrons are composite particles, composed of elementary particles called quarks. There are two types of quarks and atoms, each having a fractional electric charge. Protons are composed of two upcorks, each with charge plus 23, and one down quark with a charge of negative, of 13. Neutrons consist of one up cork and two down quarks. This distinction accounts for the difference in mass and charge between the two particles. The corks are held together by the strong interaction or strong force, which is mediated by gluons. The protons and neutrons in turn are held to each other in the nucleus by the nuclear force, which is a residuum of the strong force that is somewhat different range properties.

The gluon is a member of the family of gauge bosons, which are elementary particles that mediate physical forces. All the bound protons and neutrons in an atom make up a tiny atomic nucleus and are collectively called nucleons. Atoms of the same element have the same number of protons called the atomic number. Within a single element, the number of neutrons may vary, determining the isotope of that element. The total number of protons and neutrons determine the nucleide.

The number of neutrons relative to the protons determines the stability of the nucleus, with certain isotopes undergoing radioactive decay. The proton, the electron, and the neutron are classified as fermions. Fermions obey the poly-exclusion principle, which prohibits identical fermions, such as multiple protons from occupying the same quantum state at the same time. Thus, every proton in the nucleus must occupy a quantum state different from all other protons, and the same applies to all neutrons of the nucleus,

and to all electrons of the electron cloud. A nucleus that has a different number of protons and neutrons can potentially drop to a lower energy state, through a radioactive decay that causes the number of protons and neutrons to more closely match. As a result, atoms with matching numbers of protons and neutrons are more stable against decay, but with increasing atomic number, the mutual repulsion of the protons requires an increasing proportion of neutrons to maintain the stability of a nucleus. The number of protons and neutrons in the atomic nucleus can be modified, although this can require very high energies because of the strong force. Nuclear fusion occurs when multiple atomic particles join to form a heavier nucleus, such as through the energetic collision of two nuclei.

For example, at the core of the sun, protons require energies of 3 to 10 killa electron volts to overcome their mutual repulsion. the coulum barrier, infused together into a single nucleus. Nuclear fission is the opposite process, causing a nucleus to split into two smaller nuclei, usually through radioactive decay. The nucleus can also be modified through bombardment by high-energy subatomic particles or photons. If this modifies a number of protons in a nucleus, the atom changes to a different chemical element. If the mass of the nucleus following a fusion reaction is less than the sum of the masses of the separate particles, then the difference between these two values can be emitted as a type of usable energy, such as a gamma ray or the kinetic energy of a beta particle, as described by

Albert Einstein's mass energy equivalence formula, E equals mc squared, where M is the mass loss, and C is the speed of light. This deficit is part of the binding energy of the new nucleus, and it is the non-recoverable loss of the energy that causes the fused particles to remain together in a state that requires this energy to separate. The fusion of two nuclei that create larger nuclei with lower atomic numbers in iron and nickel, a total nucleon number of about 60, is usually an exothermic process that releases more energy than is consumed to bring them together. It is this energy-releasing process that makes nuclear fusion in stars a self-sustaining reaction. For heavier nuclei, the binding energy per nucleon begins to decrease.

That means that a fusion process producing a nucleus that has an atomic number higher than about 26, and a mass number higher than about 60 is an endothermic process. Thus, more massive nuclei cannot undergo an energy-producing fusion reaction that can sustain the hydrostatic equilibrium of a star. The electrons in an atom are attracted to the protons in the nucleus by the electromagnetic force. This force binds the electrons inside an electrostatic potential well surrounding the smaller nucleus, which means that an external source of energy is needed for the electron to escape. The closer an electron is to the nucleus, the greater the attractive force.

Hence, electrons bound near the center of the potential well require more energy to escape than those at greater separations. Electrons, like other particles, have properties of both the particle and a wave. The electron cloud is a region inside the potential well, where each electron forms a type of three-dimensional standing wave, a waveform that does not move relative to the nucleus. This behavior is defined by an atomic orbital, a mathematical function that characterizes the probability that an electron appears to be at a particular location when its position is measured. Only a discrete or quantized set of these orbitals exist, around the nucleus, as other possible wave patterns rapidly decay into a more stable form.

Orbitals can have one or more ring or node structures and differ from each other in size, shape, and orientation. Each atomic orbital corresponds to a particular energy level of the electron. The electron can change its state to a higher energy level by absorbing a photon with sufficient energy to boost, to boost it into the new quantum state. Likewise, through spontaneous emission, an electron in a higher energy state can drop to a lower energy state while radiating the excess energy as a photon. These characteristic energy values defined by the differences in the energies of the quantum states are responsible for atomic spectral lines. The amount of energy needed to remove or add an electron, the electron binding energy,

is far less than the binding energy of nucleons. For example, it requires only 13.6 electron volts to strip a ground-state electron from a hydrogen atom, compared to 2.23 million electron volts for splitting a deuterium nucleus. Atoms are electrically neutral if they have an equal number of protons and electrons. atoms that have either a deficit or a surplus of electrons are called ions. Electrons that are farthest from the nucleus may be transferred to other nearby atoms or shared between atoms. By this mechanism, atoms are able to bond into molecules and other types of chemical compounds like ionic and covalent network crystals. By definition, any two atoms with an identical number of protons in their nuclei belong to the same chemical.

element. Atoms with equal numbers of protons, but a different number of neutrons are different isotopes of the same element. For example, all hydrogen atoms admit exactly one proton, but isotopes exist with no neutrons, one neutron, two neutrons, and more than two neutrons. The known elements form a set of atomic numbers from the single proton element hydrogen, up to the 118 proton element organesan. All known isotopes of elements with atomic numbers greater than 82 or radioactive. Although the radioactivity of element 83 Ebersmith is so slight as to be practically negligible, about 339 nuclides occur naturally on Earth, of which 251, about 74%,

have not been observed to decay and are referred to as stable isotopes. Only 90 nuclides are stable theoretically, while another 161, bringing the total to 251, have not been observed to decay even though, in theory, it is energetically possible. These are also formally classified as stable. An additional 35 radioactive nuclides have half-lives long. longer than 100 million years, and are long-lived enough to have been present since the birth of the solar system. This collection of 286 nuclides are known as primordial nuclides.

Finally, an additional 53 short-lived nuclides are known to occur naturally as daughter products or primordial nucleide decay, such as radium or uranium, or as products of natural energetic processes, energetic processes on Earth, such as cosmic ray bombardment, for example, carbon 14. For 80 of the chemical elements, at least one stable isotope exists. As a rule, there is only a handful of stable isotopes for each of these elements, the average being 3.1 stable isotopes per element. 26 mono-isotopic elements have only a single stable isotope, while the largest number of stable isotopes observed for any element is 10, for the element 10. Element 43, 61, and all elements numbered 83 or higher have no stable isotopes.

Stability of isotopes is affected by the ratio of protons to neutrons, and also by the presence of certain magic numbers of neutrons or protons that represent closed and filled quantum shells. These quantum shells correspond to a set of energy levels within the shell model of the nucleus. Filled shells, such as the filled shells of 50 protons for 10, confers unusual stability on the nucleide. Of the 251 known stable nuclides, only four have both an odd number of protons and odd number of neutrons. hydrogen 2, deuterium, lithium 6, boron 10, and nitrogen 14. Tantalum 180M is odd-od-od and observational stable, but is predicted to decay with a very long half-life. Also, only four naturally occurring radioactive odd-odd-nuclides have a half-life over a billion years.

Potassium 40, vanadium-50, L anthonym 138 and Lutetium 176. Most odd-odd nuclei are highly unstable with respect to beta decay because the decay products are even even and are therefore more strongly bound due to nuclear pairing effects. The large majority of an atom's mass comes from the protons and neutrons that make it up. The total number of these particles called nucleons in a given atom is, called the mass number. It is a positive integer and dimensionless instead of having dimension

of mass because it expresses a count. An example of use of a mass number is carbon 12, which has 12 nucleons, six protons and six neutrons. The actual mass of an atom at rest is often expressed in Dalton's DA, also called the Unified Atomic Mass Unit, U. This unit is different defined as a 12th of the mass of a free neutral atom of carbon 12, which is approximately 1.66 times 10 to the negative 27th kilograms. Hydrogen 1, the lightest isotope of hydrogen, which is also the nuclide with the lowest mass, has an atomic weight of 1.007825 DA. The value of this number is called the atomic mass. A given atom has an atomic mass approximately equal within 1%, to its mass number times the atomic mass unit. For example, the mass of nitrogen 14 is roughly 14 DA.

But this number will not be exactly an integer, except by definition, in the case of carbon 12. The heaviest stable atom is lead-208, with a mass of 207.976651 DA. As even the most massive atoms are far too light to work with directly, chemists instead use the unit of moles. One mole of atoms of any element always has the same number of atoms, about 6.022 times 10 to the 23. This number was chosen so that if an element has an atomic mass of one unit, a mole of atoms of that element has a mass close to one gram. Because of the definition of the unified atomic mass unit, each carbon 12 atom has an atomic mass of exactly 12 DA,

and so a mole of carbon 12 atoms weighs about 0.012 kilograms.