If you were searching for methods to get rid of the moles that are ruining your garden, then unfortunately this is the wrong place. Nor is it the mole that you may see on someone’s chin. But the mole we will be talking about is also very important and infests throughout chemistry. It is the international unit for the amount of substance. And if you came exactly for this mole, then “Welcome”.

Which came before? The mole or Avogadro’s constant? This is a question harder than “Which came first, the egg or the chicken?” Let’s look a bit at the history. Oh, if you are still wondering about the egg and the chicken, the answer is “The egg”.

Origin of Avogadro’s Law and The Mole

The Avogadro’s constant and the mole have an intertwined history.  

In the year 1803, English chemist John Dalton revitalized the idea of Democritus that every substance is made up of tiny atoms. And this got Amedeo Avogadro, an Italian scientist, thinking about the relation between the amount of substance and the number of atoms.

He observed many things, like the electrolysis of water produced hydrogen gas and oxygen gas in a 2:1 volume ratio. From his observations, in the year 1811, he proposed that an equal volume of gasses contains an equal number of atoms or molecules under the same condition regardless of the nature of the gas. Which is known as Avogadro’s law or hypothesis. But he could not calculate the number of molecules in an amount of gas yet.

Not much was done about this number of atoms or molecules after that for a while. But the concept of the mole was developing. Because, by that time, the term relative mass was already familiar to scientists and they calculated in terms of mass. Relative mass is in the end, the ratio of masses of atoms that has no unit. So, even if you expressed these relative masses in other units their ratio would hold. 

And this relative mass expressed in gram was called a mole of that substance. It is not as complex as it sounds. Let’s look at an example. The relative mass of the carbon atom is 12. So 12g carbon means 1mole carbon. Similarly, 

1 mole of hydrogen= 1g hydrogen

1 mole of aluminum= 27g aluminum

I know what you’re thinking: “How did it get the name mole?” German chemist Ostwald coined the unit Mol from the word Molekül in the year 1894. Mole is the translated form of mol.

Note that the number of atoms in a mole is the same for every element, the same as Avogadro’s law.  But the value of this number could not be obtained for a long time. No, its approximate value was actually determined by Joseph Loschmidt in 1865 albeit indirectly and the term, Avogadro’s constant, was still not there. And so, people were still unaware of this discovery.

This term was first introduced by Jean Perrin in 1909 who defined it as the number of molecules in 32 grams of oxygen as it was the standard of the relative mass at that time. But as said earlier, they had yet to know the value. 

Perrin finally had a breakthrough when Robert Millikan determined the charge of an electron through his oil-drop experiment. And the value of the total charge of 1mole of electrons, the Faraday constant, was already known to scientists since 1834. Although you read a bit earlier that the term mole was introduced in 1894, there was already a term known as the equivalent mass which is similar to the mole but depends on valency. For electrons the mole and the equivalent mass have the same value and so is the total charge.

You might now wonder how did the scientists manage to determine the total charge of a mole of electrons without knowing their number. The answer is electrolysis. Faraday is the pioneer in this case. He determined the Faraday constant by measuring the amount of electricity needed to obtain one mole of mono-valent metal through electrolysis.

The value of Faraday constant= 96485 C = Total charge of one mole of electrons.

Charge of an electron= 1.602×10−19 C

So, Avogadro’s constant NA= 96485 C1.602×10-19 C = 6.022×1023

Jean Perrin later calculated the value of Avogadro’s constant in many different ways and was awarded the nobel prize for his work in 1926.

Knowing the past helps us learn new things from the people that came before us. We can learn both from their success and failure. Although you will not need history to do calculations involving moles or Avogadro’s constant, this story teaches us how the things we need might be in front of us and be never noticed.

This should be enough about history. Let’s look at the exact information you need to know.

Avogadro’s Constant or Avogadro’s Number

You have seen that the Avogadro’s constant is defined as the number of atoms in 16g of oxygen or the number of molecules in 32g of oxygen. We have to clarify because oxygen is found as O2 in nature which contains 2 atoms per molecule.

But this definition changed over time. Carbon-12 isotope replaced oxygen as the new standard. And so, the definition changed to- the number of atoms in 12.00g of carbon-12.

The definition changed yet again and in the year 2017, the BIPM (Bureau International des Poids et Mesures) defined Avogadro’s constant as the exact value of 6.02214076×1023. It is expressed with N or NA. The value can be rounded up to 6.0226×1023.

This value is insanely larger than it looks. Let’s write it without using the scientific form.

602,214,076,000,000,000,000,000- That’s 602 sextillion or two times billion.

There is no analogy that you can easily imagine and understand. Suppose, the earth was completely made up of softballs. The seas, mountains, buildings, everything. The number of softballs you would need is Avogadro’s number.

Or suppose, you stacked 1 mole of papers. Paper sheets are very thin. So how high will the stack be? Any guess? If your answer is up to space then you are not right. The stack will reach space and beyond. It is so high that it will go up to the moon and back 80 billion times. How big is that!!

Now imagine 6.02214076×1023 water molecules. How much would that be? Hold tight, it’s about the amount you drink in a sip.  That is 18 mL of water. No need to worry, atoms are that small. It’s a perfect contrast.

It would take 20 drops of water to form a milliliter. Just think, for each drop of water people waste, 1.67 x 10^21 molecules of water are wasted.

Never waste even a drop of water.

The Mole

As said earlier, it is not the furry earth animal, it is a unit. Specifically, a counting unit.

It is defined as the amount of substance that contains 6.02214076×1023 number of atoms, molecules, ions, particles, or any entities is called a mole of that substance. It is expressed with n. It can also be abbreviated to mol.

But 6.02214076×1023 is Avogadro’s constant. So, do you get it now? A mole is any amount that contains Avogadro’s number of particles.

The mole is similar to the dozen. Let’s see,

1 dozen apples= 12 apples

1 dozen eggs= 12 eggs

So, 

1 mole apples= 6.02214076×1023 apples

1 mole eggs= 6.02214076×1023 eggs.

It is as simple as it looks.

Now let’s see how the mole relates to the mass of a substance. As you’ve already seen, the mole is defined in a way so that the mass of a mole of any substance is the same as the relative mass expressed in gram. So, 

1 mole of carbon= 12g carbon 

1 mole of aluminum= 27g aluminum

And this mass of 1 mole of any substance is called the molar mass of that substance which has the unit gmol-1.

Thus, you can say,

number of moles (mol), n=mass of substance in grams gmolar mass g mol–1

or n= mM

You must have thought many times by now, “Why is mole so important?” It’s because most of the measurements in chemistry are done in terms of moles. The mole can at the same time relate both the mass of the substance and the number of atoms or molecules. Also, chemical substances react in terms of numbers and the mole represents a whole number. So, the mole is also used in expressing reactions and plays a great role in quantitative chemistry. 

It’s common knowledge that 2 atoms of hydrogen react with one atom of oxygen to produce 1 molecule of water. Now, this can be also said in terms of mole.

2 mol hydrogen + 1 mol oxygen= 1 mol water.

H2 + 12 O2 = H2O  [Hydrogen and oxygen are diatomic gases]

You don’t get much information using only numbers of atoms. But from this mole equation, we can immediately know,

2g hydrogen+ 16g oxygen= 18g water

The benefits of using the mole unit are way more than this.

Let’s do some mathematical problems involving the mole now.

How to calculate the mole numbers of a given amount of substance

The equation that relates the mole and mass of the substance is, 

number of moles (mol), n=mass of substance in grams g, mmolar mass g mol–1, M

Note that the mass of the substance has to be in grams. And you already know that the molar mass is basically the relative mass expressed in grams or in this case, gmol-1.

Problem 1: Use these Ar (relative atomic mass) values (Fe = 55.8, N = 14.0, O = 16.0, S = 32.1) to calculate the amount of substance in moles of 10.7 g of sulfur atoms

Solution: 

Molar mass of sulfur atoms= 32.1 gmol-1

Mass of given sulfur atoms= 10.7 g

So, number of moles= 10.7 g32.1 gmol-1

   = 0.33 mol

Suppose, what’s given is the number of moles and you are told to determine the amount of substance, how would you do that? Have you already figured it out? Yes, you just need to rearrange the above equation. After rearranging we have,

Mass of substance in grams (g)= number of moles (mol)×molar mass (gmol-1)

Problem 2: Calculate the mass of 0.050 moles of sodium carbonate, Na2CO3 (Ar values: C = 12.0, O = 16.0, Na = 23.0).

Solution: Molar mass of Na2CO3 = (23×2) + 12 + (16×3) gmol-1

 = 106 gmol-1

Number of moles= 0.050 mol

Thus, mass of 0.050 moles of Na2CO3= 106 gmol-1 × 0.050 mol

         =5.3 g

Pretty easy, right?

With this, our discussion ends today. If you liked the content, please visit our website for learning more about chemistry.

Have a good day!!

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Before talking about subatomic particles, let’s know what an atom is. Atoms are the building blocks of matters. It is the smallest unit of an element retaining its properties.

Democritus was the first to propose that all matters were made up of indivisible particles called atoms. The word, “atom” comes from the Greek word, “Atomos” which is divided into, “a” and, “Tomos”. “A” means, “not” and, “Tomos” means, “to cut”. Thus combined, “atom” means “uncuttable” or “indivisible”.

 Suppose, if you were to break your phone (do not do it actually), you will see different parts that made up the phone. Now, grind it, and it will turn into very small particles. You can tell by looking at or touching what those originally were- plastic, metal, etc. and you can isolate these particles individually. Let’s not stop here and grind this even more and more. How much far can you go? By normal means, this process can’t go past a point. But, hypothetically, if you keep dividing those particles, you will reach a point that dividing it will make the particles lose their properties and you will be left with elementary particles and that is the smallest possible unit. In this case, carbon, hydrogen from plastic, copper, other metals, and silicon, oxygen from glass, etc. We are down to a scale that can’t be seen even with the strongest of microscopes. If you were to put 10 million hydrogen atoms in a line, it will only be 1 meter long. 

And this is what you can call an atom.

Particles in an Atom

But things don’t end here. People started to wonder if this is unbreakable or is it made of even smaller particles. And the answer is, “Yes”.

An atom is made up of about 200 different types of particles such as proton, neutron, electron, muon, positron, tau, and blah blah blah. These are called subatomic particles. Among these only protons, neutrons, and electrons are permanent and contribute to the structure of the atom while the others act as force carriers and not permanent. You can compare the protons, neutrons, and electrons to bricks, rods, and cement. And other particles can be compared to paints, tiles, or such.

Many scientists have worked hard to discover the structure of the atom. Particularly, Rutherford’s alpha-particle scattering experiment’s contribution is noteworthy.

Primarily, an atom can be divided into two parts: 1. The nucleus and 2. Electron orbits. Protons and neutrons make up the nucleus of the atom while electrons move around the nucleus in certain orbits. Protons and neutrons are also called nucleons as they make up the nucleus. Let’s elaborately talk about the atomic structure some other say. Today is for subatomic particles only. 

If are very much shocked about the fact that atoms can be divided then hold tight. There is an even bigger surprise for you. Even the particles proton and neutron can be divided. They are composite particles meaning they are made up of even smaller particles.

A proton is made up of two up quarks and one down quark while a neutron is made up of two down quarks and one up quark. The quarks themselves are elementary particles and so is an electron. There are mainly six types of quarks in the universe.

Electrons

One of the blessings of the modern world is electricity and electrons are the particle that carry it through wires. It is one of the twelve elementary particles meaning it can’t be divided anymore. Electrons were first discovered by J. J. Thompson in the year 1897 through the cathode tube ray experiment. He passed high voltage electricity through a vacuum tube. When electricity flows, a ray is produced known as the cathode ray. This ray was attracted to a positively charged plate placed on the tube. Later it was discovered that this ray was, in fact, a flow of negatively charged particles named, “electron” by J. J. Thompson. 

Electrons are denoted by, “e”. It has a mass of 0.000549 a.m.u. or 9.11×10^-31 kg and it has a charge of – 4.8×10^-10 e.s.u. or – 1.6×10^-19 C. An electron is almost 1836 times lighter than a proton.

Protons 

Protons are positively charged particles that can be found in the nucleus of an atom. Protons were discovered by Rutherford in 1917. In the year 1886, German physicist Eugen Goldstein discovered a positively charged ray in gas discharge which travels in a direction opposite to the cathode ray. He named them, “Canal ray”. This beam consists of positively charged ions. Using this the charge-to-mass ratio of various positive ions was calculated. It was then proved that the hydrogen ion had the smallest size among all ionized gases of elements.

In 1911 Rutherford conducted his alpha-particle-scattering experiment and concluded that all the positive charge of an atom was concentrated in the center of an atom and named it nucleus. He also observed that hydrogen nuclei were produced when the alpha beam was shot through the air. After investigation, he proceeded to bombard nitrogen gas with alpha particles which produced a greater number of hydrogen nuclei. He concluded that the hydrogen nucleus was a part of all the atoms and named it proton.

Protons are denoted by, “p”. As a hydrogen nucleus without an electron is but a proton, it is also expressed as H⁺.

The mass of a proton is 1.007 a.m.u. or 1.6725×10^-27 kg. It has the same charge as an electron but opposite i.e., 4.8×10^-10 e.s.u. or 1.6×10^-19 C.

Neutrons

Neutrons are as the name suggests, electrically neutral particles. They make up the nucleus along with protons. British physicist Sir James Chadwick was the one to discover neutrons in the year 1932. After the alpha-particle scattering experiment, Rutherford had already predicted the existence of a neutral particle. It was then known that the atomic number was the number of protons in the nucleus. But atomic numbers and relative mass were different. So, where did this extra mass come from? In 1913 isotopes were discovered which raised more questions about the difference in mass. It was thought that there were excess protons in the nucleus, with an equal number of electrons to cancel out the additional charge. But there was no evidence.

In 1928, German physicist Walter Bothe and Herbert Becker found out that if alpha particles emitted from polonium is incident on beryllium, it gives off a penetrating and electrically neutral radiation. This radiation was interpreted as high-energy photons.

But then something interesting happened. In 1932, Irene Joliot-Curie, one of Madame Curie’s daughters, and her husband, Frederic Joliot-Curie, studied the then-unidentified radiation from beryllium for further investigations. They found that this radiation ejected protons of high velocity from a paraffin target. They tried to associate with the Compton Effect. Here Compton Effect is a phenomenon, where, if photons with high enough energy are incident on a metal surface, they knock out protons from the metal. The Compton effect was observed by Arthur Holly Compton in 1923 and was named after him.

Now, the problem is that protons are about 1836 times heavier than electrons. Yes, an ant might be able to carry weights six times their own, but imagine an ant trying to knock away an elephant. So, the neutral beam being a high energy photon is quite unlikely. 

This is where Sir Chadwick comes in. He could not accept Compton Effect as the conclusion and tried similar experiments with other elements in addition to the paraffin wax, including helium, nitrogen, and lithium as targets. By measuring the kinetic energies of ejected protons, he found out that the beryllium emissions contained a neutral component with a mass approximately equal to that of the proton and named this particle neutron.

He won the Nobel Prize in the year 1935 from Physics for this discovery. Though the Juliot-Curie pair missed their chance of getting a Nobel prize, in this case, they got a Nobel prize for discovering artificial radioactivity. So, there is no reason to feel sorry.

But this is not the end. For all we know, there might be other particles playing hide and seek only to be found by you. By the way, the Nobel prize is worth about 1.5 million US dollars .

The post Subatomic Particles In An Atom: Their Discovery and Properties appeared first on Chemniverse.

The Main Parts of an Atom

Atoms contain three primary particles: protons, neutrons, and electrons. The protons and neutrons exist in the core or nucleus of the atom. The electrons orbit around this nucleus in shells.

Protons – These are positively charged particles in the nucleus. The number of protons defines what element the atom is. For example, all carbon atoms contain 6 protons.

Neutrons – Neutrons have a neutral charge. They are also found in the nucleus and contribute to the atom’s mass.

Electrons – Negatively charged electrons orbit outside the nucleus. Atoms want to have a neutral charge, so they contain equal protons and electrons.

The number of protons and neutrons determines the atom’s atomic mass number. Isotopes are variations of an element with the same proton number but different neutron numbers.

Atomic Structure and Subatomic Particles

The smallest unit of an element that still retains the properties of that element is an atom. Atoms are extremely small, around 0.1 nanometers in diameter. But they contain three distinct subatomic particles:

Protons – Positively charged particles found within the nucleus of an atom. The number of protons defines an element’s atomic number.

Neutrons – Electrically neutral particles found within the nucleus. Neutrons help stabilize the nucleus.

Electrons – Negatively charged particles that orbit the nucleus. The number of electrons is equal to the number of protons in an electrically stable atom.

These fundamental particles interact to form the basic atomic structure that makes up all matter. The properties of different elements come from the unique number of protons, neutrons, and electrons within their atoms.

What are Isotopes?

Isotopes are variations of the same element with different numbers of neutrons. Isotopes of an element have atoms with the same number of protons but differing numbers of neutrons. Because they have the same number of protons and electrons, isotopes exhibit the same chemical behaviors. But differences in neutrinos change their mass and some physical properties.

Some elements, like carbon, hydrogen, and oxygen, have naturally occurring stable isotopes. Other isotopes may be unstable or radioactive, like uranium-235. The various isotopes of an element are denoted by their mass number. This is the sum of the number of protons and neutrons in the nucleus.

For example, the three main isotopes of carbon are:

Carbon-12: 6 protons + 6 neutrons = mass number 12

Carbon-13: 6 protons + 7 neutrons = mass number 13

Carbon-14: 6 protons + 8 neutrons = mass number 14

Although they have different mass numbers, they are all carbon atoms because they contain 6 protons. The chemical properties remain unchanged between isotopes. But isotopes have slightly different physical properties like atomic mass and radioactivity stability.

Isotopic Notation

Isotopes are identified by their atomic mass number and represented in isotopic notation. This notation puts the mass number after the element’s name or symbol.

For example:

– Carbon-14 = 14C
– Uranium-235 = 235U
– Hydrogen-2 = 2H

The atomic mass number is superscripted before the element’s symbol. This notation specifies the isotope’s mass number (protons + neutrons) uniquely from just writing the element. It also distinguishes between an element’s different isotopes.

Isotopic notation allows us to represent the specific isotope rather than just the element generally. This is extremely useful across chemistry, physics, geology and a number of other scientific areas that utilize isotope analysis.

Key Applications of Isotopes

– Radiometric dating – Radioactive isotopes like carbon-14 can be used to accurately date ancient fossil, rocks, and artifacts. As radioactive isotopes decay at known rates, measuring isotope concentrations reveals time passed.

– Medical tracers – Safe radioactive isotopes like technetium-99 can track internal organ function and abnormal tissue in PET scans and similar diagnostic tests.

– Stable isotope analysis – Variations in stable isotope rations of carbon, nitrogen, oxygen and more can detect adulterated food products, study ancient diets, and trace pollutants.

– Food and health safety – Regulatory testing looks for radioactive isotopes from contaminants and toxins to ensure food and products are safe for human consumption and use.

Atomic structure defines the basic components and interactions that build up all the physical world around us. Understanding principles of atomic particles, elements, isotopes and radioactive decay is fundamental to the wider field chemistry and its many applications across science and industry. A solid grasp of these atomic properties and processes forms the foundation for delving deeper into the exciting world of chemistry.

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