Notes to a video lecture on http://www.unizor.com
Isotopes
Isotopes are atoms that have the same number of protons (atomic number Z) but different number of neutrons (N).
Since the number of electrons and their distribution among shells and subshells are the same for different isotopes of atoms with the same atomic number, these isotopes have practically the same chemical properties.
At the same time, since the number of neutrons of different isotopes of atoms with the same atomic number is different, certain physical properties of these atoms might be quite different.
Identification of different isotopes involves specifying their atomic mass.
For example, carbon has 6 protons and 6 electrons, but the number of neutrons can be 6, 7 or 8 in its nucleus with atomic masses of these isotopes, correspondingly,
A=Z+N=6+6=12,
A=Z+N=6+7=13 and
A=Z+N=6+8=14.
These isotopes are named, according to their atomic mass, carbon-12, carbon-13 and carbon-14.
Another example: uranium-235 (symbol 92235U) with 92 protons, 92 electrons (atomic number is Z=92) and 143 neutrons (N=143), which has atomic mass A=235, and uranium-238 (symbol 92238U) with the same number of protons, electrons and the same atomic number Z=92, but with 146 neutrons (N=146) and atomic mass A=238.
In most cases one particular isotope of an atom is sufficiently stable and can be found in nature, while other isotopes of the same atom might be more or less radioactive, that is their atoms break with time into other (smaller) atoms - a process called radioactive decay.
The role of neutrons in a nucleus is to stabilize it. Since protons are positively charged and repel each other, neutrons serve as a buffer between them and, therefore, their number is usually equal or greater than the number of protons.
Obviously, so-called strong forces inside a nucleus that bind together all particles (protons to protons, protons to neutrons, neutrons to neutrons) are the main forces that hold a nucleus as a whole.
Interestingly, most nuclei with even atomic number Z and even number of neutrons N demonstrate more stability than those with either Z or N or both are odd.
Radiocarbon Dating
As an example of the usage of isotopes, let's describe the determination of age of certain samples found by archeologists using radiocarbon dating.
The most stable and abundant isotope of carbon is the one with 6 protons and 6 neutrons in a nucleus - carbon-12 612C.
This carbon isotope plays extremely important role in organisms, as carbon is one of the main building blocks of all living beings on Earth.
Firstly, let's examine where carbon-14 comes from.
High energy cosmic rays bombard the Earth and break atoms of the upper layer of atmosphere, releasing some elementary particles, electrons, protons and neutrons.
Next, some neutrons (01n) hit the atoms of nitrogen (714N) in the air causing production of carbon-14 (614C) and hydrogen (11H) in the following nuclear reaction:
714N + 01n → 614C + 11H
The reaction above describes the process of a neutron hitting an atom of nitrogen, kicking out a proton and an electron that connect to each other in the form of an atom of hydrogen, while replacing a missed proton with itself, thereby reducing the atomic number by one, while retaining the atomic mass, which transforms it into an isotope carbon-14.
Now carbon-14 is formed in the air in some small quantity and participates in the cycle of life. Any living organism absorbs it the same way it absorbs regular carbon-12 during its life time and, therefore, it's present in some quantity inside this organism in proportion to regular carbon-12 similar to its proportion that exists in nature.
As we know, carbon-14 is not a stable element, it decays with a half-life about 5730 years.
The process of decaying is rather complex, but can be described in an oversimplified form as follows.
A neutron in the nucleus of unstable carbon-14 transforms into a pair proton+electron. This does not change the electric neutrality of the atom, does not change the atomic mass, it remains 14, but atomic number increases by 1, thereby creating an atom of stable nitrogen.
This nuclear reaction can be described (in an oversimplified form) as
614C → 714N + e− + ?
where ? signifies additional participants in this transformation that we cannot discuss at this point because it requires knowledge of other elementary particles beyond the main ones (electron, proton and neutron).
While an organism lives, the relative amount of carbon-14 in it is maintained on the same level. Whatever is decayed is replenished from the outside world as part of the organism's existence.
As soon as life stops, absorption of carbon-14 stops and, whatever is present in the organism is no longer renewed, but decays according to the laws of half life.
If the amount of carbon-14 in the dead tree is 1/2 of whatever is normal, the tree died approximately 5730·1=5730 years ago.
If the amount of carbon-14 in the dead tree is 1/4 of whatever is normal, the tree died approximately 5730·2=11460 years ago.
If the amount of carbon-14 in the dead tree is 1/8 of whatever is normal, the tree died approximately 5730·3=17190 years ago.
etc.
In general, if the amount of carbon-14 in the dead tree is 1/N of whatever is normal, the tree died approximately 5730·log2N years ago.
Sunday, February 12, 2023
Saturday, February 11, 2023
Main particles: UNIZOR.COM - Physics4Teens - Atoms - Nucleus and Electrons
Notes to a video lecture on http://www.unizor.com
Main Particles
Three particles constitute the main building blocks of any atom:
electrons,
protons and
neutrons.
Let's examine their properties and functions in maintaining the integrity of an atom.
The basic atom model assumes that electrons are negatively charged particles flying on some orbits around an atom's nucleus that, in turn, contains certain number of positively charged protons and electrically neutral neutrons.
We discussed in details the distribution of electrons in shells (#1, #2, #3 etc.) and subshells (s, p, d etc.) around a nucleus.
These electrons (primarily, those on higher orbits) get involved in chemical reactions among atoms, facilitating creation of different molecules by either ionic or covalent bonding (see previous lectures on this topic).
The number of protons in a nucleus should be equal to the number of electrons circulating on orbits around this nucleus to maintain atom's neutrality.
This number is called an atomic number Z of an atom.
The electrostatic attraction between positive protons and negative electrons keeps the electrons on their stationary orbits.
But electrostatic forces repel from each other particles charged the same way (positive or negative). So, what keeps the protons in the nucleus and holds the atom's integrity?
Apparently, elementary particles inside a nucleus are held together by other forces, much stronger than electrostatic. They are called (not surprisingly) strong or nuclear forces.
These strong forces act only on a very small distance between particles inside a nucleus and have no noticeable influence on electrons around a nucleus.
Strong forces exist between protons, between protons and neutrons and between neutrons. Since neutrons are electrically neutral, greater number of them in the nucleus keeps the nucleus stronger, preventing electrostatic forces or external forces (like bombarding the nucleus with other particles) to break up an atom.
The total number of protons (Z) and neutrons (N) is called atomic mass of an atom (A=Z+N).
Nucleus of an atom takes a very small amount of space relatively to an atom's size.
For example, a nucleus of an atom of hydrogen, which consists of only one proton, has a diameter of the order of 10−15 meter, while a diameter of an atom of hydrogen is more than 100,000 greater (of the order of 10−10 meter).
Heavier elements that contain hundreds of particles (protons and neutrons, commonly called nucleons) in a nucleus have larger diameter of a nucleus, about 10 times larger than hydrogen, and their atoms are thousands times greater than their nuclei.
At the same time, the mass of an atom is concentrated in its nucleus. The nucleus' share in the mass of an atom is, approximately, 99.9%. The rest of mass is the mass of electrons.
These very approximate numbers do not take into consideration the effect of mass-energy relationship, according to Einstein Theory of Relativity formula E=m·c².
Since chemical properties of an element are defined by its electrons, the atomic number (that is, the number of electrons or equal to it the number of protons) defines chemical properties of an atom.
Different number of neutrons with the same number of protons in a nucleus of an element can occur, but these different compositions of a nucleus do not change the chemical properties of an element.
Atoms with the same number of protons but different number of neutrons are called isotopes.
We will address it in the next lecture.
Main Particles
Three particles constitute the main building blocks of any atom:
electrons,
protons and
neutrons.
Let's examine their properties and functions in maintaining the integrity of an atom.
The basic atom model assumes that electrons are negatively charged particles flying on some orbits around an atom's nucleus that, in turn, contains certain number of positively charged protons and electrically neutral neutrons.
We discussed in details the distribution of electrons in shells (#1, #2, #3 etc.) and subshells (s, p, d etc.) around a nucleus.
These electrons (primarily, those on higher orbits) get involved in chemical reactions among atoms, facilitating creation of different molecules by either ionic or covalent bonding (see previous lectures on this topic).
The number of protons in a nucleus should be equal to the number of electrons circulating on orbits around this nucleus to maintain atom's neutrality.
This number is called an atomic number Z of an atom.
The electrostatic attraction between positive protons and negative electrons keeps the electrons on their stationary orbits.
But electrostatic forces repel from each other particles charged the same way (positive or negative). So, what keeps the protons in the nucleus and holds the atom's integrity?
Apparently, elementary particles inside a nucleus are held together by other forces, much stronger than electrostatic. They are called (not surprisingly) strong or nuclear forces.
These strong forces act only on a very small distance between particles inside a nucleus and have no noticeable influence on electrons around a nucleus.
Strong forces exist between protons, between protons and neutrons and between neutrons. Since neutrons are electrically neutral, greater number of them in the nucleus keeps the nucleus stronger, preventing electrostatic forces or external forces (like bombarding the nucleus with other particles) to break up an atom.
The total number of protons (Z) and neutrons (N) is called atomic mass of an atom (A=Z+N).
Nucleus of an atom takes a very small amount of space relatively to an atom's size.
For example, a nucleus of an atom of hydrogen, which consists of only one proton, has a diameter of the order of 10−15 meter, while a diameter of an atom of hydrogen is more than 100,000 greater (of the order of 10−10 meter).
Heavier elements that contain hundreds of particles (protons and neutrons, commonly called nucleons) in a nucleus have larger diameter of a nucleus, about 10 times larger than hydrogen, and their atoms are thousands times greater than their nuclei.
At the same time, the mass of an atom is concentrated in its nucleus. The nucleus' share in the mass of an atom is, approximately, 99.9%. The rest of mass is the mass of electrons.
These very approximate numbers do not take into consideration the effect of mass-energy relationship, according to Einstein Theory of Relativity formula E=m·c².
Since chemical properties of an element are defined by its electrons, the atomic number (that is, the number of electrons or equal to it the number of protons) defines chemical properties of an atom.
Different number of neutrons with the same number of protons in a nucleus of an element can occur, but these different compositions of a nucleus do not change the chemical properties of an element.
Atoms with the same number of protons but different number of neutrons are called isotopes.
We will address it in the next lecture.
Thursday, February 9, 2023
Valence Electrons: UNIZOR.COM - Physics4Teens - Atoms - Interaction of A...
Notes to a video lecture on http://www.unizor.com
Valence Electrons
Previous lecture was about chemical reaction between elements - bonding between atoms that does not affect their nuclei.
We have seen that the most important role in these reactions is played by electrons in the outer most shell of an atom.
In particular, chemical reactions are expected between elements that have incomplete outer subshells, but complementary to each other to form complete subshells by either borrowing electrons from each other, forming electrovalent or ionic bonding between positively and negatively charged ions or sharing electrons, forming a covalent bonding between atoms.
Recall the chemical bonding between atoms of sodium and chlorine.
Sodium has atomic number 11, and its electrons are arranged as
Na: 1s2 2s2 2p6 3s1
Chlorine has atomic number 17, and its electrons are arranged as
Cl: 1s2 2s2 2p6 3s2 3p5
The mechanism of bonding can be represented in such a case as follows.
1. Sodium atom releases the only electron from the subshell 3s, becoming a positive ion
Na → Na+ + e−
2. Chlorine atom captures this electron in its subshell 3p, becoming negative ion
Cl + e− → Cl−
3. Now all subshells are filled to capacity and both atoms are electrically charged with opposite charges, which makes them stick to each other, forming a molecule of salt.
That was a typical example of electrovalent or ionic bonding.
Valence electrons are those exact electrons that are capable of participating in bonding with other electrons of other atoms.
In most cases they belong to the outer electronic shell of an atom.
Elements with electron structure ending in 1 electron in the outer subshell s, like sodium (Na), have that 1 electron as the only valence electron. The valency of such elements is 1. Atoms of these elements often bond by giving away this electron to complete other atom's outer subshell, like ionic bonding with chlorine (Cl) that needs exactly one electron to complete its outer p subshell. The valency of the counterpart atom in such a case is also 1.
The structure of a molecule that results in bonding of these two atoms can be symbolically expressed as
Na−Cl
where a single bar between symbols of elements corresponds to the valency of each.
Considering this is an ionic bonding, we might as well use more precise symbolics
Na+−Cl−
Elements with electron structure ending in 2 electrons in the outer subshell s, like barium (Ba), have these 2 electrons as the only valence electrons. Atoms of these elements often bond by giving away both of these electrons to complete other atom's outer subshell, like ionic bonding with two atoms of chlorine (Cl), each of which needs one electron to complete its outer p subshell, forming a molecule BaCl2.
The structure of a molecule that results in bonding of these three atoms can be symbolically expressed as
Cl−Ba−Cl
where bars correspond to the valency of each atom, 2 for barium and 1 for chlorine.
This process can be described as follows.
1. Barium atom (atomic number 56) frees two electrons from the outer subshell 6s, becoming a positive ion
Ba → Ba2+ + 2e−
2. Two chlorine atoms capture one electron each in their outer subshell 3p, becoming negative ions
2Cl + 2e− → 2Cl−
3. Now all subshells are filled to capacity and three atoms are electrically charged. One atom of barium is positive, two atoms of chlorine are negative, which makes these two atoms of chorine stick to an atom of barium, forming a molecule of barium chloride BaCl2.
So, more precise representation of this molecule would be
Cl−−Ba+−Cl−
Let's consider a more complicated case of forming aluminum oxide - a molecule combining atoms of aluminum (Al) and oxygen (O).
The configuration of 13 electrons in the atom of aluminum is:
Al: 1s2 2s2 2p6 3s2 3p1
The configuration of 8 electrons in the atom of oxygen is:
O: 1s2 2s2 2p4
As we see, aluminum has one valence electron in subshell 3p. If it loses this electron, its 3rd subshell would still have two electrons in subshell 3s. While technically complete, it's still exposed to be captured by some other atom.
So, an atom of aluminum can lose either (only under special conditions) 1 electron, the 3p1, or (under normal conditions) 3 electrons, the 3p1 and 3s2.
The valency of aluminum is, therefore, either 1 (only under special conditions) or, typically, 3.
A possible molecular bonding between 2 atoms of aluminum and 1 atom of oxygen Al2O that can be established in special laboratory conditions would look like this:
Al−O−Al
This is a very unstable molecule because the 3rd shell of the atom of aluminum with only one remaining subshell 3s and two electrons in it still remains exposed.
Much more stable molecule is formed, when 2 atoms of aluminum contribute 6 electrons (3 electrons from each atom, emptying subshells 3s and 3p) and 3 atoms of oxygen capture them (2 electrons are captured by each oxygen atom to fill its 2p subshell). The formula of this aluminum oxide molecule is Al2O3 with the structure that can be pictured as
O=Al−O−Al=O
where bars correspond to the valency of each atom, 3 for aluminum and 2 for oxygen.
From the position of an exchange of electrons this can be represented as follows.
1. Two aluminum atoms free 3 electrons each from their outer shell #3 (1 electron from subshell 3p and 2 electrons from subshell 3s), becoming positive ions
2Al → 2Al3+ + 6e−
2. Three oxygen atoms capture 2 electrons each in their outer subshell 2p, becoming negative ions
3O + 6e− → 3O2−
3. Now all subshells are filled to capacity and all 5 atoms are electrically charged. Two atoms of aluminum are positive, three atoms of oxygen are negative, which makes these three atoms of oxygen stick to two atoms of aluminum, forming a molecule of aluminum oxide Al2O3.
The atoms in this molecule are bonded ionically, so we can represent the molecule's structure as
O2−=Al3+−O2−−Al3+=O2−
As you saw from the previous example, in some cases the valency of an atom can be different under different circumstances. While for a bonding between aluminum and oxygen only one stable configuration is observed, in other cases more than one configuration occurs under normal conditions.
The common property of the examples above is that we analyzed the molecules comprised of a metal and non-metal.
Metals have their outer electrons relatively easier departing from their nuclei, that's why in all cases above the bonding between a metal and non-metal was ionic, when metals gave their electrons to non-metals.
As a final example, consider the formation of a molecule of carbon monoxide CO.
The electron configuration of carbon (atomic number 6) is
C: 1s2 2s2 2p2
The configuration of 8 electrons in the atom of oxygen, as we already discussed, is:
O: 1s2 2s2 2p4
None of the elements in a molecule of carbon monoxide is a metal, so their electrons are not capable, ready and willing to depart from the their nuclei, as in metals. As a result, the electrons are not transferred from atom to atom forming ions, but are shared between the atoms to form a covalent bonding.
Two electrons from 2p subshell of carbon and four electrons from 2p subshell of oxygen, totaling 6 electrons (exactly the maximum number for 2p subshell), become a common property of both atoms, completing their corresponding outer subshells 2p.
The act of sharing electrons between atoms of carbon and oxygen is the glue holding the molecule together in covalent bonding.
Valence Electrons
Previous lecture was about chemical reaction between elements - bonding between atoms that does not affect their nuclei.
We have seen that the most important role in these reactions is played by electrons in the outer most shell of an atom.
In particular, chemical reactions are expected between elements that have incomplete outer subshells, but complementary to each other to form complete subshells by either borrowing electrons from each other, forming electrovalent or ionic bonding between positively and negatively charged ions or sharing electrons, forming a covalent bonding between atoms.
Recall the chemical bonding between atoms of sodium and chlorine.
Sodium has atomic number 11, and its electrons are arranged as
Na: 1s2 2s2 2p6 3s1
Chlorine has atomic number 17, and its electrons are arranged as
Cl: 1s2 2s2 2p6 3s2 3p5
The mechanism of bonding can be represented in such a case as follows.
1. Sodium atom releases the only electron from the subshell 3s, becoming a positive ion
Na → Na+ + e−
2. Chlorine atom captures this electron in its subshell 3p, becoming negative ion
Cl + e− → Cl−
3. Now all subshells are filled to capacity and both atoms are electrically charged with opposite charges, which makes them stick to each other, forming a molecule of salt.
That was a typical example of electrovalent or ionic bonding.
Valence electrons are those exact electrons that are capable of participating in bonding with other electrons of other atoms.
In most cases they belong to the outer electronic shell of an atom.
Elements with electron structure ending in 1 electron in the outer subshell s, like sodium (Na), have that 1 electron as the only valence electron. The valency of such elements is 1. Atoms of these elements often bond by giving away this electron to complete other atom's outer subshell, like ionic bonding with chlorine (Cl) that needs exactly one electron to complete its outer p subshell. The valency of the counterpart atom in such a case is also 1.
The structure of a molecule that results in bonding of these two atoms can be symbolically expressed as
Na−Cl
where a single bar between symbols of elements corresponds to the valency of each.
Considering this is an ionic bonding, we might as well use more precise symbolics
Na+−Cl−
Elements with electron structure ending in 2 electrons in the outer subshell s, like barium (Ba), have these 2 electrons as the only valence electrons. Atoms of these elements often bond by giving away both of these electrons to complete other atom's outer subshell, like ionic bonding with two atoms of chlorine (Cl), each of which needs one electron to complete its outer p subshell, forming a molecule BaCl2.
The structure of a molecule that results in bonding of these three atoms can be symbolically expressed as
Cl−Ba−Cl
where bars correspond to the valency of each atom, 2 for barium and 1 for chlorine.
This process can be described as follows.
1. Barium atom (atomic number 56) frees two electrons from the outer subshell 6s, becoming a positive ion
Ba → Ba2+ + 2e−
2. Two chlorine atoms capture one electron each in their outer subshell 3p, becoming negative ions
2Cl + 2e− → 2Cl−
3. Now all subshells are filled to capacity and three atoms are electrically charged. One atom of barium is positive, two atoms of chlorine are negative, which makes these two atoms of chorine stick to an atom of barium, forming a molecule of barium chloride BaCl2.
So, more precise representation of this molecule would be
Cl−−Ba+−Cl−
Let's consider a more complicated case of forming aluminum oxide - a molecule combining atoms of aluminum (Al) and oxygen (O).
The configuration of 13 electrons in the atom of aluminum is:
Al: 1s2 2s2 2p6 3s2 3p1
The configuration of 8 electrons in the atom of oxygen is:
O: 1s2 2s2 2p4
As we see, aluminum has one valence electron in subshell 3p. If it loses this electron, its 3rd subshell would still have two electrons in subshell 3s. While technically complete, it's still exposed to be captured by some other atom.
So, an atom of aluminum can lose either (only under special conditions) 1 electron, the 3p1, or (under normal conditions) 3 electrons, the 3p1 and 3s2.
The valency of aluminum is, therefore, either 1 (only under special conditions) or, typically, 3.
A possible molecular bonding between 2 atoms of aluminum and 1 atom of oxygen Al2O that can be established in special laboratory conditions would look like this:
Al−O−Al
This is a very unstable molecule because the 3rd shell of the atom of aluminum with only one remaining subshell 3s and two electrons in it still remains exposed.
Much more stable molecule is formed, when 2 atoms of aluminum contribute 6 electrons (3 electrons from each atom, emptying subshells 3s and 3p) and 3 atoms of oxygen capture them (2 electrons are captured by each oxygen atom to fill its 2p subshell). The formula of this aluminum oxide molecule is Al2O3 with the structure that can be pictured as
O=Al−O−Al=O
where bars correspond to the valency of each atom, 3 for aluminum and 2 for oxygen.
From the position of an exchange of electrons this can be represented as follows.
1. Two aluminum atoms free 3 electrons each from their outer shell #3 (1 electron from subshell 3p and 2 electrons from subshell 3s), becoming positive ions
2Al → 2Al3+ + 6e−
2. Three oxygen atoms capture 2 electrons each in their outer subshell 2p, becoming negative ions
3O + 6e− → 3O2−
3. Now all subshells are filled to capacity and all 5 atoms are electrically charged. Two atoms of aluminum are positive, three atoms of oxygen are negative, which makes these three atoms of oxygen stick to two atoms of aluminum, forming a molecule of aluminum oxide Al2O3.
The atoms in this molecule are bonded ionically, so we can represent the molecule's structure as
O2−=Al3+−O2−−Al3+=O2−
As you saw from the previous example, in some cases the valency of an atom can be different under different circumstances. While for a bonding between aluminum and oxygen only one stable configuration is observed, in other cases more than one configuration occurs under normal conditions.
The common property of the examples above is that we analyzed the molecules comprised of a metal and non-metal.
Metals have their outer electrons relatively easier departing from their nuclei, that's why in all cases above the bonding between a metal and non-metal was ionic, when metals gave their electrons to non-metals.
As a final example, consider the formation of a molecule of carbon monoxide CO.
The electron configuration of carbon (atomic number 6) is
C: 1s2 2s2 2p2
The configuration of 8 electrons in the atom of oxygen, as we already discussed, is:
O: 1s2 2s2 2p4
None of the elements in a molecule of carbon monoxide is a metal, so their electrons are not capable, ready and willing to depart from the their nuclei, as in metals. As a result, the electrons are not transferred from atom to atom forming ions, but are shared between the atoms to form a covalent bonding.
Two electrons from 2p subshell of carbon and four electrons from 2p subshell of oxygen, totaling 6 electrons (exactly the maximum number for 2p subshell), become a common property of both atoms, completing their corresponding outer subshells 2p.
The act of sharing electrons between atoms of carbon and oxygen is the glue holding the molecule together in covalent bonding.
Thursday, February 2, 2023
Chemical Bonding: UNIZOR.COM - Physics4Teens - Atoms - Interaction of Atoms
Notes to a video lecture on http://www.unizor.com
Chemical Bonding
From the atomic viewpoint, classical chemical bonding between atoms are interactions that do not change the composition of these atoms' nuclei.
Only electrons that surround nuclei interact among themselves, breaking some old connections and establishing new ones.
Chemical bonding can occur between atoms of simple elements, like a single atom of sodium Na and a single atom of chlorine Cl forming their compositions, a molecule of sodium chloride NaCl, which is a molecule of regular salt.
Scientists thought about why some elements bond with each other, if mixed in certain proportion and at some external conditions, while some others are inert, that is refuse to combine with other elements. So called noble gases, like helium, neon, argon or krypton are inert.
Let's examine the electron configuration of above mentioned noble gases.
Helium with atomic number 2 has the following electron configuration
He: 1s2
Neon with atomic number 10 has the following electron configuration
Ne: 1s2 2s2 2p6 =
= [He] + 2s2 2p6
Argon with atomic number 18 has the following electron configuration
Ar: 1s2 2s2 2p6 3s2 3p6 =
= [Ne] + 3s2 3p6
Krypton with atomic number 36 has the following electron configuration
Kr: 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 =
= [Ar] + 3d10 4s2 4p6
As you see, all subshells of noble gases are filled to capacity:
max(s) = 2
max(p) = 6
max(d) = 10
and this is very important for stability of atoms.
Those elements that do not have this characteristic of its subshells will be easier to get into chemical bonding that facilitates sharing electrons to complete top most subshells of all involved components.
Let's examine the chemical process between elements using the above example of producing sodium chloride from sodium and chlorine
Na + Cl → NaCl
First of all, let's analyze the electronic structure of both substances.
Sodium has atomic number 11, and its electrons are arranged as
Na: 1s2 2s2 2p6 3s1
Chlorine has atomic number 17, and its electrons are arranged as
Cl: 1s2 2s2 2p6 3s2 3p5
As you see, the first subshell s of the shell #1 (that is, subshell 1s), which can hold up to 2 electrons, is completely filled for both elements, and there can be no more subshells within this shell #1.
The first subshell s of the shell #2 (that is, subshell 2s) is also completely filled for both elements, but shell #2 can have two subshells, s and p, and subshell p has capacity for 6 electrons. As we see, both 2p subshells are filled to capacity in both elements.
The first subshell s of the shell #3 (that is, subshell 3s) has only 1 electron in an atom of sodium.
The atom of chlorine, on the other hand, has this subshell filled to capacity of 2 electrons, but the next subshell 3p has only 5 electrons out of maximum 6.
Now the sharing of electrons comes to play.
The top subshell of sodium 3s contains 1 electron out of maximum 2. If it loses this electron, all its subshells will be complete, but losing electron means becoming positively charged.
The top subshell of chlorine 3p contains 5 electrons out of maximum 6. If it captures one electron, all its subshells will be complete, but capturing an electron means becoming negatively charged.
The mechanism of bonding can be represented in such a case as follows.
1. Sodium atom frees the only electron from the subshell 3s, becoming a positive ion
Na → Na+ + e−
2. Chlorine atom captures this electron in its subshell 3p, becoming negative ion
Cl + e− → Cl−
3. Now all subshells are filled to capacity and both atoms are electrically charged with opposite charges, which makes them stick to each other, forming a molecule of salt.
Let's make a few conclusive statements about this process of forming molecules.
I. Atoms like their subshells to be completely filled up to a maximum of 4m−2, where m is the subshell number (subshell #1 is labeled s, subshell #2 is labeled p etc.)
II. Some atoms are ready to accept electrons to fill up their top subshell from other atoms that are willing to give up their electrons.
III. Some other atoms are ready to give up electrons from their incomplete top subshell, leaving with themselves only complete subshells, if there is a recipient atom of these extra electrons that needs them to complete its top subshell.
IV. The necessity for an atom to have its shells filled up is very strong. Atoms with incomplete top subshells are looking for partners to get into chemical combination. If there is a fit between two atoms to transfer electrons from an incomplete top subshell of one of them to incomplete top subshell of another that completes both, electrons are transferred.
V. After the transfer of electrons that completes both atoms, they become oppositely charged ions that are attracted to each other, which establishes a stable combination.
VI. The complex combinations of more than 2 atoms can also be created based on the same principle. The mechanism is a little more complex, but follows the same idea to complete top subshells of all participants.
Chemical Bonding
From the atomic viewpoint, classical chemical bonding between atoms are interactions that do not change the composition of these atoms' nuclei.
Only electrons that surround nuclei interact among themselves, breaking some old connections and establishing new ones.
Chemical bonding can occur between atoms of simple elements, like a single atom of sodium Na and a single atom of chlorine Cl forming their compositions, a molecule of sodium chloride NaCl, which is a molecule of regular salt.
Scientists thought about why some elements bond with each other, if mixed in certain proportion and at some external conditions, while some others are inert, that is refuse to combine with other elements. So called noble gases, like helium, neon, argon or krypton are inert.
Let's examine the electron configuration of above mentioned noble gases.
Helium with atomic number 2 has the following electron configuration
He: 1s2
Neon with atomic number 10 has the following electron configuration
Ne: 1s2 2s2 2p6 =
= [He] + 2s2 2p6
Argon with atomic number 18 has the following electron configuration
Ar: 1s2 2s2 2p6 3s2 3p6 =
= [Ne] + 3s2 3p6
Krypton with atomic number 36 has the following electron configuration
Kr: 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 =
= [Ar] + 3d10 4s2 4p6
As you see, all subshells of noble gases are filled to capacity:
max(s) = 2
max(p) = 6
max(d) = 10
and this is very important for stability of atoms.
Those elements that do not have this characteristic of its subshells will be easier to get into chemical bonding that facilitates sharing electrons to complete top most subshells of all involved components.
Let's examine the chemical process between elements using the above example of producing sodium chloride from sodium and chlorine
Na + Cl → NaCl
First of all, let's analyze the electronic structure of both substances.
Sodium has atomic number 11, and its electrons are arranged as
Na: 1s2 2s2 2p6 3s1
Chlorine has atomic number 17, and its electrons are arranged as
Cl: 1s2 2s2 2p6 3s2 3p5
As you see, the first subshell s of the shell #1 (that is, subshell 1s), which can hold up to 2 electrons, is completely filled for both elements, and there can be no more subshells within this shell #1.
The first subshell s of the shell #2 (that is, subshell 2s) is also completely filled for both elements, but shell #2 can have two subshells, s and p, and subshell p has capacity for 6 electrons. As we see, both 2p subshells are filled to capacity in both elements.
The first subshell s of the shell #3 (that is, subshell 3s) has only 1 electron in an atom of sodium.
The atom of chlorine, on the other hand, has this subshell filled to capacity of 2 electrons, but the next subshell 3p has only 5 electrons out of maximum 6.
Now the sharing of electrons comes to play.
The top subshell of sodium 3s contains 1 electron out of maximum 2. If it loses this electron, all its subshells will be complete, but losing electron means becoming positively charged.
The top subshell of chlorine 3p contains 5 electrons out of maximum 6. If it captures one electron, all its subshells will be complete, but capturing an electron means becoming negatively charged.
The mechanism of bonding can be represented in such a case as follows.
1. Sodium atom frees the only electron from the subshell 3s, becoming a positive ion
Na → Na+ + e−
2. Chlorine atom captures this electron in its subshell 3p, becoming negative ion
Cl + e− → Cl−
3. Now all subshells are filled to capacity and both atoms are electrically charged with opposite charges, which makes them stick to each other, forming a molecule of salt.
Let's make a few conclusive statements about this process of forming molecules.
I. Atoms like their subshells to be completely filled up to a maximum of 4m−2, where m is the subshell number (subshell #1 is labeled s, subshell #2 is labeled p etc.)
II. Some atoms are ready to accept electrons to fill up their top subshell from other atoms that are willing to give up their electrons.
III. Some other atoms are ready to give up electrons from their incomplete top subshell, leaving with themselves only complete subshells, if there is a recipient atom of these extra electrons that needs them to complete its top subshell.
IV. The necessity for an atom to have its shells filled up is very strong. Atoms with incomplete top subshells are looking for partners to get into chemical combination. If there is a fit between two atoms to transfer electrons from an incomplete top subshell of one of them to incomplete top subshell of another that completes both, electrons are transferred.
V. After the transfer of electrons that completes both atoms, they become oppositely charged ions that are attracted to each other, which establishes a stable combination.
VI. The complex combinations of more than 2 atoms can also be created based on the same principle. The mechanism is a little more complex, but follows the same idea to complete top subshells of all participants.
Monday, January 30, 2023
Energy Levels: UNIZOR.COM - Physics4Teens - Atoms - Electronic Structure...
Notes to a video lecture on http://www.unizor.com
Energy Levels
In the previous lecture we have addressed how electrons are positioned around nucleus.
We mentioned that, according to Bohr's model, they are only located on stationary orbits within shells and subshells.
We had a nice formula for the maximum number of electrons that can be located on each subshell, depending on its order number, traditionally designated by a letter (subshell #1 is designated letter s, the next #2 is p, then d, f, g etc.) For subshell #m the maximum number of electrons on it is 4m−2.
The number of subshells within shell #N is N.
Consequently, the maximum number of electrons held in each shell is 2·N².
This mathematically perfect picture gets more complicated, when we consider the energy level associated with each subshell.
The main principle of distribution of energy among subshells, that is, the greater radius of a subshell - the greater energy electrons within this subshell possess, is absolutely true.
The problem is, the further we go from the nucleus - the closer to each other are the shells and, subshells of greater radius of a shell #N might overlap with subshells of smaller radius of shell #N+1.
Let's consider a few elements in a sequence of their atomic numbers, that is the number of protons in their nucleus or electrons within shells on different levels.
Hydrogen has only one proton in the nucleus and one electron within the first subshell (designated by a letter s) of the only shell #1.
So, the electron structure of the atom of hydrogen, where the number specifies the shell number and a letter specifies the subshell, where the superscript at the subshell letter indicates the number of electrons within the same subshell, is:
1s1
Helium has two protons in the nucleus and two electrons within the first subshell (designated by a letter s) of the only shell #1.
The subshell s (the first one) can hold maximum of 2 electrons, so everything is fine.
The electron structure of the atom of helium is:
1s2
Lithium has three protons in the nucleus and three electrons.
The first (and only) subshell s of the shell #1 can hold only 2 electrons. Therefore, a new shell #2 should exist and one electron should be in its first subshell s.
The subshell s of the second shell can hold maximum of 2 electrons, so everything is fine.
The electron structure of the atom of lithium is:
1s2 2s1
Let's skip a few elements and consider carbon with atomic number 6.
Its 6 electrons should fill the shell #1 with its only subshell s holding 2 electrons, then the shell #2 with its subshell s holding another 2 electrons, then, considering shell #2 has two subshells, s and p, the second subshell p will hold the remaining 2 electrons.
The subshell p, as the second subshell, can hold maximum of 6 electrons, so we still have room.
The electronic structure of an atom of carbon is, therefore,
1s2 2s2 2p2
Skip a few more elements to silicon with atomic number 14.
We already know that the first shell can hold 2 electrons in one and only subshell s.
The second shell can hold 2 electrons in subshell s and 6 electrons in subshell p.
That totals 10 electrons. To accommodate 4 other electrons the third shell should be used.
Its subshell s will hold 2 electron and subshell p will hold remaining 2 ones.
The electronic structure of silicon is
1s2 2s2 2p6 3s2 3p2
So far, increasing in atomic number is synchronous with sequential filling of shells and subshells within shells.
But let's consider element argon. Its atomic number is 18 and it completely fills one subshell s of the first shell, both subshells of the second shell and two (out of three) subshells of the third shell.
Its electronic structure is
1s2 2s2 2p6 3s2 3p6
The next element is potassium with atomic number 19.
It would be reasonable to expect that, since the second shell's last subshell p is already filled as well as two first subshells of shell #3, and the third shell allows three subshells, the extra electron #19 should be in the third subshell d (that is the third subshell's letter) of the shell #3.
However, the experiments showed that the last electron #19 goes to the fourth shell's first subshell s instead. Why?
Here we face the fact that shells overlap, and higher energy level of the third subshell d of shell #3 exceeds the lower energy level of the first subshell s of shell #4.
That's why the location of electron #19 of potassium is not 3d1 but 4s1.
That is the reason why the electronic structure of potassium is
1s2 2s2 2p6 3s2 3p6 4s1
The order of filling the electronic subshells, as the number of electrons in the atom grows, under normal conditions corresponds to an order of increasing their energy levels. This is known as the Aufbau principle or Madelung or Klechkovsky rule.
According to this principle, and taking into consideration that shells get closer and closer to each other as we move away from a nucleus, while the number of subshells is increasing linearly with the shell number, the order of filling the electronic structure of an atom is not the numerical order of shell/subshell numbers, but as follows: 1s→2s→2p→3s→3p→4s→
→3d→4p→5s→4d→5p→6s→
→4f→5d→6p→7s→5f→
→6d→7p→...
In the lecture "Orbiting Electron" of the chapter "Building Blocks of Matter" of this course we have derived a formula for a total energy of an orbiting electron:
E = −k·e²/(2·r)
where
k is Coulomb constant
e is the electric charge of an electron
r is the radius of an orbit.
From the above formula we see that absolute value of the energy is decreasing with increasing of a radius of an orbit, but, since it's negative, the energy itself is increasing.
As a radius of an orbit of an electron is increasing, its (negative) energy is increasing, getting asymptotically closer to zero. At the same time the number of subshells of each shell is increasing with the shell number.
That explains the overlapping energy characteristic of the shells.
Energy Levels
In the previous lecture we have addressed how electrons are positioned around nucleus.
We mentioned that, according to Bohr's model, they are only located on stationary orbits within shells and subshells.
We had a nice formula for the maximum number of electrons that can be located on each subshell, depending on its order number, traditionally designated by a letter (subshell #1 is designated letter s, the next #2 is p, then d, f, g etc.) For subshell #m the maximum number of electrons on it is 4m−2.
The number of subshells within shell #N is N.
Consequently, the maximum number of electrons held in each shell is 2·N².
This mathematically perfect picture gets more complicated, when we consider the energy level associated with each subshell.
The main principle of distribution of energy among subshells, that is, the greater radius of a subshell - the greater energy electrons within this subshell possess, is absolutely true.
The problem is, the further we go from the nucleus - the closer to each other are the shells and, subshells of greater radius of a shell #N might overlap with subshells of smaller radius of shell #N+1.
Let's consider a few elements in a sequence of their atomic numbers, that is the number of protons in their nucleus or electrons within shells on different levels.
Hydrogen has only one proton in the nucleus and one electron within the first subshell (designated by a letter s) of the only shell #1.
So, the electron structure of the atom of hydrogen, where the number specifies the shell number and a letter specifies the subshell, where the superscript at the subshell letter indicates the number of electrons within the same subshell, is:
1s1
Helium has two protons in the nucleus and two electrons within the first subshell (designated by a letter s) of the only shell #1.
The subshell s (the first one) can hold maximum of 2 electrons, so everything is fine.
The electron structure of the atom of helium is:
1s2
Lithium has three protons in the nucleus and three electrons.
The first (and only) subshell s of the shell #1 can hold only 2 electrons. Therefore, a new shell #2 should exist and one electron should be in its first subshell s.
The subshell s of the second shell can hold maximum of 2 electrons, so everything is fine.
The electron structure of the atom of lithium is:
1s2 2s1
Let's skip a few elements and consider carbon with atomic number 6.
Its 6 electrons should fill the shell #1 with its only subshell s holding 2 electrons, then the shell #2 with its subshell s holding another 2 electrons, then, considering shell #2 has two subshells, s and p, the second subshell p will hold the remaining 2 electrons.
The subshell p, as the second subshell, can hold maximum of 6 electrons, so we still have room.
The electronic structure of an atom of carbon is, therefore,
1s2 2s2 2p2
Skip a few more elements to silicon with atomic number 14.
We already know that the first shell can hold 2 electrons in one and only subshell s.
The second shell can hold 2 electrons in subshell s and 6 electrons in subshell p.
That totals 10 electrons. To accommodate 4 other electrons the third shell should be used.
Its subshell s will hold 2 electron and subshell p will hold remaining 2 ones.
The electronic structure of silicon is
1s2 2s2 2p6 3s2 3p2
So far, increasing in atomic number is synchronous with sequential filling of shells and subshells within shells.
But let's consider element argon. Its atomic number is 18 and it completely fills one subshell s of the first shell, both subshells of the second shell and two (out of three) subshells of the third shell.
Its electronic structure is
1s2 2s2 2p6 3s2 3p6
The next element is potassium with atomic number 19.
It would be reasonable to expect that, since the second shell's last subshell p is already filled as well as two first subshells of shell #3, and the third shell allows three subshells, the extra electron #19 should be in the third subshell d (that is the third subshell's letter) of the shell #3.
However, the experiments showed that the last electron #19 goes to the fourth shell's first subshell s instead. Why?
Here we face the fact that shells overlap, and higher energy level of the third subshell d of shell #3 exceeds the lower energy level of the first subshell s of shell #4.
That's why the location of electron #19 of potassium is not 3d1 but 4s1.
That is the reason why the electronic structure of potassium is
1s2 2s2 2p6 3s2 3p6 4s1
The order of filling the electronic subshells, as the number of electrons in the atom grows, under normal conditions corresponds to an order of increasing their energy levels. This is known as the Aufbau principle or Madelung or Klechkovsky rule.
According to this principle, and taking into consideration that shells get closer and closer to each other as we move away from a nucleus, while the number of subshells is increasing linearly with the shell number, the order of filling the electronic structure of an atom is not the numerical order of shell/subshell numbers, but as follows: 1s→2s→2p→3s→3p→4s→
→3d→4p→5s→4d→5p→6s→
→4f→5d→6p→7s→5f→
→6d→7p→...
In the lecture "Orbiting Electron" of the chapter "Building Blocks of Matter" of this course we have derived a formula for a total energy of an orbiting electron:
E = −k·e²/(2·r)
where
k is Coulomb constant
e is the electric charge of an electron
r is the radius of an orbit.
From the above formula we see that absolute value of the energy is decreasing with increasing of a radius of an orbit, but, since it's negative, the energy itself is increasing.
As a radius of an orbit of an electron is increasing, its (negative) energy is increasing, getting asymptotically closer to zero. At the same time the number of subshells of each shell is increasing with the shell number.
That explains the overlapping energy characteristic of the shells.
Thursday, January 26, 2023
Electrons and Shells: UNIZOR.COM - Physics4Teens - Atoms - Electronic St...
Notes to a video lecture on http://www.unizor.com
Electrons and Shells
First of all, there are some very complicated theories of the electronic structure of an atom. They are experimentally confirmed and, therefore, are considered as true representation of how electrons are arranged around a nucleus.
These theories are above the level of this course, but certain facts based on them we will mention without any proof.
Recall the Bohr's model of an atom and electrons that are supposed to be on stationary orbits (or, rather, within stationary shells), where they do not emit any energy.
Shells of a larger radius can hold more electrons, and electrons in those higher orbit shells have higher level of energy.
Shells are numbered in order of increasing radius as shell #1, shell #2, ...shell #N.
Any shell can have one or more subshells. The number of subshells in each shell corresponds to this shell's number, that is:
shell #1 has 1 subshell,
shell #2 has 2 subshell,
shell #3 has 3 subshell,
shell #4 has 4 subshell,
shell #5 has 5 subshell,
shell #6 has 6 subshell,
shell #7 has 7 subshell,
etc.
For historical reasons subshells within each shell are not enumerated, but rather assigned a letter. The first 4 subshells are called s, p, d and f, then the letter is assigned in alphabetical order.
So, the first few subshell names are:
s for subshell #1,
p for subshell #2,
d for subshell #3,
f for subshell #4,
g for subshell #5,
h for subshell #6,
i for subshell #7,
etc.
To bring a bit of math into this structure, let's use letter N for a shell's number and letter m for a subshell number within a shell.
Using these symbols, we can state the following:
shell #N has N subshells:
#1, #2,...,#N.
Each subshells has its own capacity to hold electrons. There are theoretical reasons for this based on Quantum Theory. We just state the result of this theory confirmed by experiments:
subshell #m has 4·m−2 electrons.
So,
subshell #1 (s) has 2 electrons,
subshell #2 (p) has 6 electrons,
subshell #3 (d) has 10 electrons,
subshell #4 (f) has 14 electrons,
subshell #5 (g) has 18 electrons,
subshell #6 (h) has 22 electrons,
subshell #7 (i) has 26 electrons,
etc.
Now let's calculate the maximum number of electrons in each shell.
Shell #1 has 1 subshell #1(s) and, therefore, can hold no more than 2 electrons.
Shell #2 has 2 subshells #1(s), #2(p) and, therefore, can hold no more than 2+6=8 electrons.
Shell #3 has 3 subshells #1(s), #2(p), #3(d) and, therefore, can hold no more than 2+6+10=18 electrons.
These calculations can be generalized in a formula for the maximum number of electrons in shell #N:
Σm∈[1,N](4m−2) = 2·N²
This formula can be easily proven by induction.
Indeed, it's correct for N=1 because
4·1−2 = 2 = 2·1²
Assuming the formula is correct for some number N, let's check it for N+1.
Σm∈[1,N+1](4m−2) =
=Σm∈[1,N](4m−2)+4(N+1)−2 =
= 2·N² + 4·(N+1)−2 =
= 2·N² + 4·N + 2 =
= 2·(N² + 2·N + 1) =
= 2·(N+1)²
which is the same formula, but for N+1.
Electrons and Shells
First of all, there are some very complicated theories of the electronic structure of an atom. They are experimentally confirmed and, therefore, are considered as true representation of how electrons are arranged around a nucleus.
These theories are above the level of this course, but certain facts based on them we will mention without any proof.
Recall the Bohr's model of an atom and electrons that are supposed to be on stationary orbits (or, rather, within stationary shells), where they do not emit any energy.
Shells of a larger radius can hold more electrons, and electrons in those higher orbit shells have higher level of energy.
Shells are numbered in order of increasing radius as shell #1, shell #2, ...shell #N.
Any shell can have one or more subshells. The number of subshells in each shell corresponds to this shell's number, that is:
shell #1 has 1 subshell,
shell #2 has 2 subshell,
shell #3 has 3 subshell,
shell #4 has 4 subshell,
shell #5 has 5 subshell,
shell #6 has 6 subshell,
shell #7 has 7 subshell,
etc.
For historical reasons subshells within each shell are not enumerated, but rather assigned a letter. The first 4 subshells are called s, p, d and f, then the letter is assigned in alphabetical order.
So, the first few subshell names are:
s for subshell #1,
p for subshell #2,
d for subshell #3,
f for subshell #4,
g for subshell #5,
h for subshell #6,
i for subshell #7,
etc.
To bring a bit of math into this structure, let's use letter N for a shell's number and letter m for a subshell number within a shell.
Using these symbols, we can state the following:
shell #N has N subshells:
#1, #2,...,#N.
Each subshells has its own capacity to hold electrons. There are theoretical reasons for this based on Quantum Theory. We just state the result of this theory confirmed by experiments:
subshell #m has 4·m−2 electrons.
So,
subshell #1 (s) has 2 electrons,
subshell #2 (p) has 6 electrons,
subshell #3 (d) has 10 electrons,
subshell #4 (f) has 14 electrons,
subshell #5 (g) has 18 electrons,
subshell #6 (h) has 22 electrons,
subshell #7 (i) has 26 electrons,
etc.
Now let's calculate the maximum number of electrons in each shell.
Shell #1 has 1 subshell #1(s) and, therefore, can hold no more than 2 electrons.
Shell #2 has 2 subshells #1(s), #2(p) and, therefore, can hold no more than 2+6=8 electrons.
Shell #3 has 3 subshells #1(s), #2(p), #3(d) and, therefore, can hold no more than 2+6+10=18 electrons.
These calculations can be generalized in a formula for the maximum number of electrons in shell #N:
Σm∈[1,N](4m−2) = 2·N²
This formula can be easily proven by induction.
Indeed, it's correct for N=1 because
4·1−2 = 2 = 2·1²
Assuming the formula is correct for some number N, let's check it for N+1.
Σm∈[1,N+1](4m−2) =
=Σm∈[1,N](4m−2)+4(N+1)−2 =
= 2·N² + 4·(N+1)−2 =
= 2·N² + 4·N + 2 =
= 2·(N² + 2·N + 1) =
= 2·(N+1)²
which is the same formula, but for N+1.
Monday, January 23, 2023
Nucleus of Atom: UNIZOR.COM - Physics4Teens - Atoms - Building Blocks of...
Notes to a video lecture on http://www.unizor.com
Nucleus of Atom
In 1897 J.J.Thomson, experimenting with cathode ray tubes, came to a conclusion that atoms contain tiny negatively charged particles.
He had demonstrated that cathode rays consist of negatively charged particles - electrons.
He then suggested a plum pudding model of an atom with these tiny negatively charged particles embedded into positively charged "soup".
In 1911 Rutherford experimented with a beam of positively charged particles (called alpha particles) directed toward a thin foil. He found that some particles go through a foil, while some are reflected back.
As a result, he came up with a planetary model of an atom with positively charged nucleus and negatively charged electrons rotating around a nucleus on a substantial (relatively to a size of a nucleus) distance, so an atom is substantially empty.
In 1917 Rutherford proved experimentally that nucleus of hydrogen atom is present in atoms of all substances he dealt with.
Later on the nucleus of hydrogen atom was called proton - another building block of an atom together with electron.
A few years later Rutherford suggested that another type of particle - an electrically neutral combination of tightly coupled together proton and electron, called by him neutron - must be present inside an atom's nucleus. While this hypothesis was not confirmed experimentally, the term "neutron" remained and used later on.
In 1932 James Chadwick discovered electrically neutral particles emitted from beryllium bombarded by alpha particles. These electrically neutral particles, in turn, were used to bombard paraffin wax and liberated hydrogen nuclei - protons.
That was a proof of existence of electrically neutral neutrons.
Right after that, in 1932, Dmitry Ivanenko and Werner Heisenberg proposed a proton-neutron structure of an atom's nucleus. This model of a nucleus together with electrons outside of a nucleus moving within stationary shells, each with a specific energy level, remains as the main atom's model.
The number of protons in a nucleus of any element under normal conditions is equal to the number of electrons around a nucleus to maintain electric neutrality of an atom. The properties of any element very much depend on this number and, actually, characterize the element's properties. This number is called an atomic number of an element.
The atomic number of a hydrogen is 1, its nucleus contains 1 proton and there is 1 electron outside a nucleus.
The atomic number of a helium is 2, its nucleus contains 2 proton and there are 2 electrons outside a nucleus.
The atomic number of a carbon is 6, its nucleus contains 6 proton and there are 6 electrons outside a nucleus.
The atomic number of a gold is 79, its nucleus contains 79 proton and there are 79 electrons outside a nucleus.
The atomic number of a uranium is 92, its nucleus contains 92 proton and there are 92 electrons outside a nucleus.
The number of neutrons inside a nucleus is also very important, but can vary for the same element.
Thus, a nucleus of an atom of hydrogen, besides one proton, can have no neutrons, one neutron and two neutrons. Their atomic numbers are the same. To differentiate them, another characteristic is used - a sum of the number of protons and the number of neutrons called atomic mass.
Different kinds of the same element with different numbers of neutrons (that is, different atomic mass) are called isotopes. So, there are three isotopes of hydrogen - with atomic masses of 1 (1 proton, no neutrons), 2 (1 proton, 1 neutron) and 3 (1 proton, 2 neutrons).
Isotopes of the same element have close but not identical properties. To fully identify an element, including its isotope, the following notation is used.
The element is identified by its abbreviated Latin name.
For example,
He for helium,
Au for gold ("aurum" in Latin),
Fe for iron ("ferrum" in Latin) etc.
To fully identify an element, in-front of this abbreviated name two indices are used: top for atomic mass (sum of the numbers protons and neutrons) and bottom for atomic number (the number protons).
Examples:
24He for helium
(2 protons and 2 neutrons),
92238U for uranium
(92 protons and 146 neutrons),
79197Au for gold
(79 protons and 118 neutrons),
2656Fe for iron
(26 protons and 30 neutrons).
Positively charged nucleus with certain number of protons keeps the same number of negatively charged electrons inside an atom because opposite charges attract.
Why a nucleus is held together, when similarly charged protons are lumped together and repel each other?
Apparently, there other attractive forces acting specifically on very small distances between particles inside a nucleus, which are stronger than repelling electrical forces. These forces are called nuclear or strong forces.
They are short range forces of attraction between any particles inside a nucleus, protons and neutrons (collectively called nucleons), and they are millions of times stronger than electric forces. They keep nucleons together inside a nucleus.
Nucleus of Atom
In 1897 J.J.Thomson, experimenting with cathode ray tubes, came to a conclusion that atoms contain tiny negatively charged particles.
He had demonstrated that cathode rays consist of negatively charged particles - electrons.
He then suggested a plum pudding model of an atom with these tiny negatively charged particles embedded into positively charged "soup".
In 1911 Rutherford experimented with a beam of positively charged particles (called alpha particles) directed toward a thin foil. He found that some particles go through a foil, while some are reflected back.
As a result, he came up with a planetary model of an atom with positively charged nucleus and negatively charged electrons rotating around a nucleus on a substantial (relatively to a size of a nucleus) distance, so an atom is substantially empty.
In 1917 Rutherford proved experimentally that nucleus of hydrogen atom is present in atoms of all substances he dealt with.
Later on the nucleus of hydrogen atom was called proton - another building block of an atom together with electron.
A few years later Rutherford suggested that another type of particle - an electrically neutral combination of tightly coupled together proton and electron, called by him neutron - must be present inside an atom's nucleus. While this hypothesis was not confirmed experimentally, the term "neutron" remained and used later on.
In 1932 James Chadwick discovered electrically neutral particles emitted from beryllium bombarded by alpha particles. These electrically neutral particles, in turn, were used to bombard paraffin wax and liberated hydrogen nuclei - protons.
That was a proof of existence of electrically neutral neutrons.
Right after that, in 1932, Dmitry Ivanenko and Werner Heisenberg proposed a proton-neutron structure of an atom's nucleus. This model of a nucleus together with electrons outside of a nucleus moving within stationary shells, each with a specific energy level, remains as the main atom's model.
The number of protons in a nucleus of any element under normal conditions is equal to the number of electrons around a nucleus to maintain electric neutrality of an atom. The properties of any element very much depend on this number and, actually, characterize the element's properties. This number is called an atomic number of an element.
The atomic number of a hydrogen is 1, its nucleus contains 1 proton and there is 1 electron outside a nucleus.
The atomic number of a helium is 2, its nucleus contains 2 proton and there are 2 electrons outside a nucleus.
The atomic number of a carbon is 6, its nucleus contains 6 proton and there are 6 electrons outside a nucleus.
The atomic number of a gold is 79, its nucleus contains 79 proton and there are 79 electrons outside a nucleus.
The atomic number of a uranium is 92, its nucleus contains 92 proton and there are 92 electrons outside a nucleus.
The number of neutrons inside a nucleus is also very important, but can vary for the same element.
Thus, a nucleus of an atom of hydrogen, besides one proton, can have no neutrons, one neutron and two neutrons. Their atomic numbers are the same. To differentiate them, another characteristic is used - a sum of the number of protons and the number of neutrons called atomic mass.
Different kinds of the same element with different numbers of neutrons (that is, different atomic mass) are called isotopes. So, there are three isotopes of hydrogen - with atomic masses of 1 (1 proton, no neutrons), 2 (1 proton, 1 neutron) and 3 (1 proton, 2 neutrons).
Isotopes of the same element have close but not identical properties. To fully identify an element, including its isotope, the following notation is used.
The element is identified by its abbreviated Latin name.
For example,
He for helium,
Au for gold ("aurum" in Latin),
Fe for iron ("ferrum" in Latin) etc.
To fully identify an element, in-front of this abbreviated name two indices are used: top for atomic mass (sum of the numbers protons and neutrons) and bottom for atomic number (the number protons).
Examples:
24He for helium
(2 protons and 2 neutrons),
92238U for uranium
(92 protons and 146 neutrons),
79197Au for gold
(79 protons and 118 neutrons),
2656Fe for iron
(26 protons and 30 neutrons).
Positively charged nucleus with certain number of protons keeps the same number of negatively charged electrons inside an atom because opposite charges attract.
Why a nucleus is held together, when similarly charged protons are lumped together and repel each other?
Apparently, there other attractive forces acting specifically on very small distances between particles inside a nucleus, which are stronger than repelling electrical forces. These forces are called nuclear or strong forces.
They are short range forces of attraction between any particles inside a nucleus, protons and neutrons (collectively called nucleons), and they are millions of times stronger than electric forces. They keep nucleons together inside a nucleus.
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