Structure of atom

Structure of Atom

The existence of different kinds of matter is due to different atoms constituting them. Two questions arise:

(i) What makes the atom of one element different from the atom of another element? and

(ii) Are atoms really indivisible, or are there smaller constituents inside the atom?

One of the first indications that atoms are not indivisible, comes from studying static electricity and the condition under which electricity is conducted by different substances.

Charged Particles in Matter

Many scientists contributed in revealing the presence of charged particles in an atom. It was known by 1900 that the atom was not a simple, indivisible particle but contained at least one sub-atomic particle – the electron identified by J.J. Thomson.

E. Goldstein in 1886 discovered the presence of new radiations in a gas discharge and called them canal rays. These rays were positively charged radiations which ultimately led to the discovery of another sub-atomic particle. This sub-atomic particle had a charge, equal in magnitude but opposite in sign to that of the electron. Its mass was approximately 2000 times as that of the electron. It was given the name of proton. In general, an electron is represented as ‘e–’ and a proton as ‘p+’. The mass of a proton is taken as one unit and its charge as plus one. The mass of an electron is considered to be negligible and its charge is minus one.

The Structure of an Atom

J.J. Thomson was the first one to propose a model for the structure of an atom.

THOMSON’S MODEL OF AN ATOM

Thomson proposed that:

(i) An atom consists of a positively charged sphere and the electrons are embedded in it.

(ii) The negative and positive charges are equal in magnitude. So, the atom as a whole is electrically neutral. Although Thomson’s model explained that atoms are electrically neutral, the results of experiments carried out by other scientists could not be explained by this model.

RUTHERFORD’S MODEL OF AN ATOM

Ernest Rutherford was interested in knowing how the electrons are arranged within an atom. Rutherford designed an experiment for this. In this experiment, fast moving alpha (α)-particles were made to fall on a thin gold foil.

• He selected a gold foil because he wanted as thin a layer as possible. This gold foil was about 1000 atoms thick.

• α-particles are doubly-charged helium ions. Since they have a mass of 4 u, the fast-moving α-particles have a considerable amount of energy.

• It was expected that α-particles would be deflected by the sub-atomic particles in the gold atoms. Since the α-particles were much heavier than the protons, he did not expect to see large deflections.

The following observations were made:

(i) Most of the fast moving α-particles passed straight through the gold foil.

(ii) Some of the α-particles were deflected by the foil by small angles.

(iii) Surprisingly one out of every 12000 particles appeared to rebound.

Rutherford concluded from the α-particle scattering experiment that–

(i) Most of the space inside the atom is empty because most of the α-particles passed through the gold foil without getting deflected.

(ii) Very few particles were deflected from their path, indicating that the positive charge of the atom occupies very little space.

(iii) A very small fraction of α-particles were deflected by 1800, indicating that all the positive charge and mass of the gold atom were concentrated in a very small volume within the atom.

On the basis of his experiment, Rutherford put forward the nuclear model of an atom, which had the following features:

(i) There is a positively charged centre in an atom called the nucleus. Nearly all the mass of an atom resides in the nucleus.

(ii) The electrons revolve around the nucleus in well-defined orbits.

(iii) The size of the nucleus is very small as compared to the size of the atom.

Drawbacks of Rutherford’s model of the atom

The orbital revolution of the electron is not expected to be stable. Any particle in a circular orbit would undergo acceleration. During acceleration, charged particles would radiate energy. Thus, the revolving electron would lose energy and finally fall into the nucleus. If this were so, the atom should be highly unstable and hence matter would not exist in the form that we know. We know that atoms are quite stable.

BOHR’S MODEL OF ATOM

In order to overcome the objections raised against Rutherford’s model of the atom, Neils Bohr put forward the following postulates about the model of an atom:

(i) Only certain special orbits known as discrete orbits of electrons, are allowed inside the atom.

(ii) While revolving in discrete orbits the electrons do not radiate energy. These orbits or shells are called energy levels.

These orbits or shells are represented by the letters K,L,M,N,… or the numbers, n=1,2,3,4,….

NEUTRONS

In 1932, J. Chadwick discovered another subatomic particle which had no charge and a mass nearly equal to that of a proton. It was eventually named as neutron. Neutrons are present in the nucleus of all atoms, except hydrogen. In general, a neutron is represented as ‘n’. The mass of an atom is therefore given by the sum of the masses of protons and neutrons present in the nucleus.

How are Electrons Distributed in Different Orbits (Shells)?

The distribution of electrons into different orbits of an atom was suggested by Bohr and Bury.

The following rules are followed for writing the number of electrons in different energy levels or shells:

(i) The maximum number of electrons present in a shell is given by the formula 2n2, where ‘n’ is the orbit number or energy level index, 1,2,3,…. Hence the maximum number of electrons in different shells are as follows:

first orbit or K-shell will be = 2 × 1² = 2,

second orbit or L-shell will be = 2 × 2² = 8,

third orbit or M-shell will be = 2 ×3² = 18,

fourth orbit or N-shell will be = 2 × 4²= 32, and so on.

(ii) The maximum number of electrons that can be accommodated in the outermost orbit is 8.

(iii) Electrons are not accommodated in a given shell, unless the inner shells are filled. That is, the shells are filled in a step-wise manner.

Valency

The electrons present in the outermost shell of an atom are known as the valence electrons. Valency or valency number, is a measure of the number of chemical bonds formed by the atoms of a given element.

According to Bohr-Bury, outermost shell of an atom can have two electrons in its outermost shell and all other elements have atoms with eight electrons in the outermost shell. The combining capacity of the atoms of other elements, that is, their tendency to react and form molecules with atoms of the same or different elements was thus explained as an attempt to attain a fully-filled outermost shell. An outermost-shell, which had eight electrons was said to possess an octet. Atoms would thus react, so as to achieve an octet in the outermost shell. This was done by sharing, gaining or losing electrons. The number of electrons gained, lost or shared so as to make the octet of electrons in the outermost shell, gives us directly the combining capacity of the element.

For example, hydrogen/lithium/sodium atoms contain one electron each in their outermost shell, therefore each one of them can lose one electron. So, they are said to have valency of one. If the number of electrons in the outermost shell of an atom is close to its full capacity, then valency is determined in a different way. For example, the fluorine atom has 7 electrons in the outermost shell, and its valency could be 7. But it is easier for fluorine to gain one electron instead of losing seven electrons. Hence, its valency is determined by subtracting seven electrons from the octet and this gives a valency of one for fluorine. Valency can be calculated in a similar manner for oxygen. Therefore, an atom of each element has a definite combining capacity, called its valency.

Atomic Number and Mass Number

ATOMIC NUMBER- The number of protons in the nucleus of an atom determines an element's atomic number. Each element has a unique number that identifies how many protons are in one atom of that element. For example, all hydrogen atoms, and only hydrogen atoms, contain one proton and have an atomic number of 1. All carbon atoms, and only carbon atoms, contain six protons and have an atomic number of 6. Oxygen atoms contain 8 protons and have an atomic number of 8. The atomic number of an element never changes, meaning that the number of protons in the nucleus of every atom in an element is always the same.

MASS NUMBER- mass of an atom is practically due to protons and neutrons alone. These are present in the nucleus of an atom. Hence protons and neutrons are also called nucleons. Therefore, the mass of an atom resides in its nucleus. For example, mass of carbon is 12 u because it has 6 protons and 6 neutrons, 6 u + 6 u = 12 u. Similarly, the mass of aluminum is 27 u (13 protons+14 neutrons). The mass number is defined as the sum of the total number of protons and neutrons present in the nucleus of an atom.

All atoms have a mass number which is derived as follows.

Number of Neutrons + Number of Protons = Mass Number

Atom

 

The atom is a basic unit of matter that consists of a dense central nucleus surrounded by a cloud of negatively charged electrons. The atomic nucleus contains a mix of positively charged protons and electrically neutral neutrons (except in the case of hydrogen-1, which is the only stable nuclide with no neutrons). The electrons of an atom are bound to the nucleus by the electromagnetic force. Likewise, a group of atoms can remain bound to each other, forming a molecule. An atom containing an equal number of protons and electrons is electrically neutral, otherwise it has a positive charge if there are fewer electrons (electron deficiency) or negative charge if there are more electrons (electron excess). A positively or negatively charged atom is known as an ion. An atom is classified according to the number of protons and neutrons in its nucleus: the number of protons determines the chemical element, and the number of neutrons determines the isotope of the element.^[1]

The name atom comes from the Greek ἄτομος (atomos, "indivisible") from ἀ- (a-, "not") and τέμνω (temnō, "I cut"),^[2] which means uncuttable, or indivisible, something that cannot be divided further.^[3] The concept of an atom as an indivisible component of matter was first proposed by early Indian and Greek philosophers. In the 17th and 18th centuries, chemists provided a physical basis for this idea by showing that certain substances could not be further broken down by chemical methods. During the late 19th and early 20th centuries, physicists discovered subatomic components and structure inside the atom, thereby demonstrating that the 'atom' was divisible. The principles of quantum mechanics were used to successfully model the atom.^[4][5]

Atoms are minuscule objects with proportionately tiny masses. Atoms can only be observed individually using special instruments such as the scanning tunneling microscope. Over 99.94% of an atom's mass is concentrated in the nucleus,^{[note 1]} with protons and neutrons having roughly equal mass. Each element has at least one isotope with an unstable nucleus that can undergo radioactive decay. This can result in a transmutation that changes the number of protons or neutrons in a nucleus.^[6] Electrons that are bound to atoms possess a set of stable energy levels, or orbitals, and can undergo transitions between them by absorbing or emitting photons that match the energy differences between the levels. The electrons determine the chemical properties of an element, and strongly influence an atom's magnetic properties.

Components

Subatomic particles

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 proton and the neutron. However, the hydrogen-1 atom has no neutrons and a positive hydrogen ion has no electrons.

The electron is by far the least massive of these particles at 9.11×10⁻³¹ kg, with a negative electrical charge and a size that is too small to be measured using available techniques.^[46] Protons have a positive charge and a mass 1,836 times that of the electron, at 1.6726×10⁻²⁷ kg, although this can be reduced by changes to the energy binding the proton into an atom. Neutrons have no electrical charge and have a free mass of 1,839 times the mass of electrons,^[47] or 1.6929×10⁻²⁷ kg. Neutrons and protons have comparable dimensions—on the order of 2.5×10⁻¹⁵ m—although the 'surface' of these particles is not sharply defined.^[48]

In the Standard Model of physics, electrons are truly elementary particles with no internal structure. However, both protons and neutrons are composite particles composed of elementary particles called quarks. There are two types of quarks in atoms, each having a fractional electric charge. Protons are composed of two up quarks (each with charge +²⁄₃) and one down quark (with a charge of −¹⁄₃). Neutrons consist of one up quark and two down quarks. This distinction accounts for the difference in mass and charge between the two particles.^[49][50]

The quarks 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 has somewhat different range-properties (see the article on the nuclear force for more). The gluon is a member of the family of gauge bosons, which are elementary particles that mediate physical forces.

Nucleus

The binding energy needed for a nucleon to escape the nucleus, for various isotopes

All the bound protons and neutrons in an atom make up a tiny atomic nucleus, and are collectively called nucleons. The radius of a nucleus is approximately equal to , where A is the total number of nucleons. This is much smaller than the radius of the atom, which is on the order of 10⁵ fm. The nucleons are bound together by a short-ranged attractive potential called the residual strong force. At distances smaller than 2.5 fm this force is much more powerful than the electrostatic force that causes positively charged protons to repel each other.

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 nuclide. The number of neutrons relative to the protons determines the stability of the nucleus, with certain isotopes undergoing radioactive decay.

The neutron and the proton are different types of fermions. The Pauli exclusion principle is a quantum mechanical effect that prohibits identical fermions, such as multiple protons, from occupying the same quantum physical state at the same time. Thus every proton in the nucleus must occupy a different state, with its own energy level, and the same rule applies to all of the neutrons. This prohibition does not apply to a proton and neutron occupying the same quantum state.

For atoms with low atomic numbers, a nucleus that has a different number of protons than 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 roughly matching numbers of protons and neutrons are more stable against decay. However, with increasing atomic number, the mutual repulsion of the protons requires an increasing proportion of neutrons to maintain the stability of the nucleus, which modifies this trend. Thus, there are no stable nuclei with equal proton and neutron numbers above atomic number Z = 20 (calcium); and as Z increases toward the heaviest nuclei, the ratio of neutrons per proton required for stability increases to about 1.5.

Illustration of a nuclear fusion process that forms a deuterium nucleus, consisting of a proton and a neutron, from two protons. A positron (e⁺)—an antimatter electron—is emitted along with an electron neutrino.

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 3–10 keV to overcome their mutual repulsion—the coulomb barrier—and fuse 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 the 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 = mc², 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.^[58]

The fusion of two nuclei that create larger nuclei with lower atomic numbers than iron and nickel—a total nucleon number of about 60—is usually an exothermic process that releases more energy than is required to bring them together.^[59] It is this energy-releasing process that makes nuclear fusion in stars a self-sustaining reaction. For heavier nuclei, the binding energy per nucleon in the nucleus begins to decrease. That means fusion processes producing nuclei that have atomic numbers higher than about 26, and atomic masses higher than about 60, is an endothermic process. These more massive nuclei can not undergo an energy-producing fusion reaction that can sustain the hydrostatic equilibrium of a star.^[54]

Properties

Nuclear properties

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 (hydrogen-1, by far the most common form, also called protium), one neutron (deuterium), two neutrons (tritium) 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 ununoctium. All known isotopes of elements with atomic numbers greater than 82 are radioactive.

About 339 nuclides occur naturally on Earth, of which 255 (about 75%) have not been observed to decay, and are referred to as "stable isotopes". However, only 90 of these nuclides are stable to all decay, even in theory. Another 165 (bringing the total to 255) have not been observed to decay, even though in theory it is energetically possible. These are also formally classified as "stable". An additional 33 radioactive nuclides have half-lives longer than 80 million years, and are long-lived enough to be present from the birth of the solar system. This collection of 288 nuclides are known as primordial nuclides. Finally, an additional 51 short-lived nuclides are known to occur naturally, as daughter products of primordial nuclide decay (such as radium from uranium), or else as products of natural 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 3.2 stable isotopes per element. Twenty-six elements have only a single stable isotope, while the largest number of stable isotopes observed for any element is ten, for the element tin. Elements 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 shell of 50 protons for tin, confers unusual stability on the nuclide. Of the 255 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. Also, only four naturally occurring, radioactive odd-odd nuclides have a half-life over a billion years: potassium-40, vanadium-50, lanthanum-138 and tantalum-180m. 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.

Mass

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. The mass number is a simple whole number, and has units of "nucleons." 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 using the unified atomic mass unit (u), which is also called a dalton (Da). This unit is defined as a twelfth of the mass of a free neutral atom of carbon-12, which is approximately 1.66×10⁻²⁷ kg.^[73] Hydrogen-1, the lightest isotope of hydrogen and the atom with the lowest mass, has an atomic weight of 1.007825 u.^[74] 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 mass of the atomic mass unit. However, this number will not be an exact whole number except in the case of carbon-12 (see below)^[75] The heaviest stable atom is lead-208,^[68] with a mass of 207.9766521 u.

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 6.022×10²³). This number was chosen so that if an element has an atomic mass of 1 u, 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 u, and so a mole of carbon-12 atoms weighs exactly 0.012 kg.^{[73][page needed]}

Shape and size

Atoms lack a well-defined outer boundary, so their dimensions are usually described in terms of an atomic radius. This is a measure of the distance out to which the electron cloud extends from the nucleus. However, this assumes the atom to exhibit a spherical shape, which is only obeyed for atoms in vacuum or free space. Atomic radii may be derived from the distances between two nuclei when the two atoms are joined in a chemical bond. The radius varies with the location of an atom on the atomic chart, the type of chemical bond, the number of neighboring atoms (coordination number) and a quantum mechanical property known as spin. On the periodic table of the elements, atom size tends to increase when moving down columns, but decrease when moving across rows (left to right). Consequently, the smallest atom is helium with a radius of 32 pm, while one of the largest is caesium at 225 pm.

When subjected to external fields, like an electrical field, the shape of an atom may deviate from that of a sphere. The deformation depends on the field magnitude and the orbital type of outer shell electrons, as shown by group-theoretical considerations. Aspherical deviations might be elicited for instance in crystals, where large crystal-electrical fields may occur at low-symmetry lattice sites. Significant ellipsoidal deformations have recently been shown to occur for sulfur ions in pyrite-type compounds.

Atomic dimensions are thousands of times smaller than the wavelengths of light (400–700 nm) so they can not be viewed using an optical microscope. However, individual atoms can be observed using a scanning tunneling microscope. To visualize the minuteness of the atom, consider that a typical human hair is about 1 million carbon atoms in width. A single drop of water contains about 2 sextillion (2×10²¹) atoms of oxygen, and twice the number of hydrogen atoms. A single carat diamond with a mass of 2×10⁻⁴ kg contains about 10 sextillion (10²²) atoms of carbon. If an apple were magnified to the size of the Earth, then the atoms in the apple would be approximately the size of the original apple.

Valence and bonding behavior

The outermost electron shell of an atom in its uncombined state is known as the valence shell, and the electrons in that shell are called valence electrons. The number of valence electrons determines the bonding behavior with other atoms. Atoms tend to chemically react with each other in a manner that fills (or empties) their outer valence shells. For example, a transfer of a single electron between atoms is a useful approximation for bonds that form between atoms with one-electron more than a filled shell, and others that are one-electron short of a full shell, such as occurs in the compound sodium chloride and other chemical ionic salts. However, many elements display multiple valences, or tendencies to share differing numbers of electrons in different compounds. Thus, chemical bonding between these elements takes many forms of electron-sharing that are more than simple electron transfers. Examples include the element carbon and the organic compounds.

The chemical elements are often displayed in a periodic table that is laid out to display recurring chemical properties, and elements with the same number of valence electrons form a group that is aligned in the same column of the table. (The horizontal rows correspond to the filling of a quantum shell of electrons.) The elements at the far right of the table have their outer shell completely filled with electrons, which results in chemically inert elements known as the noble gases.

States

Quantities of atoms are found in different states of matter that depend on the physical conditions, such as temperature and pressure. By varying the conditions, materials can transition between solids, liquids, gases and plasmas. ^[105] Within a state, a material can also exist in different phases. An example of this is solid carbon, which can exist as graphite or diamond.

At temperatures close to absolute zero, atoms can form a Bose–Einstein condensate, at which point quantum mechanical effects, which are normally only observed at the atomic scale, become apparent on a macroscopic scale.^[107][108] This super-cooled collection of atoms then behaves as a single super atom, which may allow fundamental checks of quantum mechanical behavior.^[109]

Origin and current state

Atoms form about 4% of the total energy density of the observable universe, with an average density of about 0.25 atoms/m³.^[116] Within a galaxy such as the Milky Way, atoms have a much higher concentration, with the density of matter in the interstellar medium (ISM) ranging from 10⁵ to 10⁹ atoms/m³.^[117] The Sun is believed to be inside the Local Bubble, a region of highly ionized gas, so the density in the solar neighborhood is only about 10³ atoms/m³. Stars form from dense clouds in the ISM, and the evolutionary processes of stars result in the steady enrichment of the ISM with elements more massive than hydrogen and helium. Up to 95% of the Milky Way's atoms are concentrated inside stars and the total mass of atoms forms about 10% of the mass of the galaxy.^[119] (The remainder of the mass is an unknown dark matter.)

Earth

Most of the atoms that make up the Earth and its inhabitants were present in their current form in the nebula that collapsed out of a molecular cloud to form the Solar System. The rest are the result of radioactive decay, and their relative proportion can be used to determine the age of the Earth through radiometric dating. Most of the helium in the crust of the Earth (about 99% of the helium from gas wells, as shown by its lower abundance of helium-3) is a product of alpha decay.

There are a few trace atoms on Earth that were not present at the beginning (i.e., not "primordial"), nor are results of radioactive decay. Carbon-14 is continuously generated by cosmic rays in the atmosphere.^[132] Some atoms on Earth have been artificially generated either deliberately or as by-products of nuclear reactors or explosions.^[133][134] Of the transuranic elements—those with atomic numbers greater than 92—only plutonium and neptunium occur naturally on Earth.^[135][136] Transuranic elements have radioactive lifetimes shorter than the current age of the Earth^[137] and thus identifiable quantities of these elements have long since decayed, with the exception of traces of plutonium-244 possibly deposited by cosmic dust.^[129] Natural deposits of plutonium and neptunium are produced by neutron capture in uranium ore.

The Earth contains approximately 1.33×10⁵⁰ atoms.^[139] In the planet's atmosphere, small numbers of independent atoms of noble gases exist, such as argon and neon. The remaining 99% of the atmosphere is bound in the form of molecules, including carbon dioxide and diatomic oxygen and nitrogen. At the surface of the Earth, atoms combine to form various compounds, including water, salt, silicates and oxides. Atoms can also combine to create materials that do not consist of discrete molecules, including crystals and liquid or solid metals.^[140][141] This atomic matter forms networked arrangements that lack the particular type of small-scale interrupted order associated with molecular matter.^[142]

Atomic Structure

        In the last lesson we learned that atoms were particles of elements, substances that could not be broken down further.  In examining atomic structure though, we have to clarify this statement.  An atom cannot be broken down further without changing the chemical nature of the substance.  For example, if you have 1 ton, 1 gram or 1 atom of oxygen, all of these units have the same properties.  We can break down the atom of oxygen into smaller particles, however, when we do the atom looses its chemical properties.  For example, if you have 100 watches, or one watch, they all behave like watches and tell time.  You can dismantle one of the watches: take the back off, take the batteries out, peer inside and pull things out.  However, now the watch no longer behaves like a watch.  So what does an atom look like inside?
        Atoms are made up of 3 types of particles electrons  , protons  and neutrons .  These particles have different properties.  Electrons are tiny, very light particles that have a negative electrical charge (-). Protons are much larger and heavier than electrons and have the opposite charge, protons have a positive charge.  Neutrons are large and heavy like protons, however neutrons have no electrical charge.  Each atom is made up of a combination of these particles.  Let's look at one type of atom:

The atom above, made up of one proton and one electron, is called hydrogen (the abbreviation for hydrogen is H).  The proton and electron stay together because just like two magnets, the opposite electrical charges attract each other.  What keeps the two from crashing into each other?  The particles in an atom are not still.  The electron is constantly spinning around the center of the atom (called the nucleus).  The centrigugal force of the spinning electron keeps the two particles from coming into contact with each other much as the earth's rotation keeps it from plunging into the sun. 

            Keep in mind that atoms are extremely small.  One hydrogen atom, for example, is approximately 5 x 10^-8 mm in diameter.  To put that in perspective, this dash - is approximately 1 mm in length, therefore it would take almost 20 million hydrogen atoms to make a line as long as the dash.  In the sub-atomic world, things often behave a bit strangely.  First of all, the electron actually spins very far from the nucleus.  If we were to draw the hydrogen atom above to scale, so that the proton were the size depicted above, the electron would actually be spinning approximately 0.5 km (or about a quarter of a mile) away from the nucleus.  In other words, if the proton was the size depicted above, the whole atom would be about the size of Giants Stadium.  Another peculiarity of this tiny world is the particles themselves.  Protons and neutrons behave like small particles, sort of like tiny billiard balls.  The electron however, has some of the properties of a wave.  In other words, the electron is more similar to a beam of light than it is to a billiard ball.  Thus to represent it as a small particle spinning around a nucleus is slightly misleading.  In actuality, the electron is a wave that surrounds the nucleus of an atom like a cloud. 

While you should keep in mind that electrons actually form clouds around their nucleii, we will continue to represent the electron as a spinning particle to keep things simple.
        In an electrically neutral atom, the positively charged protons are always balanced by an equal number of negatively charged electrons.  As we have seen, hydrogen is the simplest atom with only one proton and one electron.  Helium is the 2nd simplest atom.  It has two protons in its nucleus and two electrons spinning around the nucleus.  With helium though, we have to introduce another particle.  Because the 2 protons in the nucleus have the same charge on them, they would tend to repel each other, and the nucleus would fall apart.  To keep the nucleus from pushing apart, helium has two neutrons in its nucleus.  Neutrons have no electrical charge on them and act as a sort of nuclear glue, holding the protons, and thus the nucleus, together.

        As you can see, helium is larger than hydrogen.  As you add electrons, protons and neutrons, the size of the atom increases.  We can measure an atom's size in two ways: using the atomic number (Z) or using the atomic mass (A, also known as the mass number).  The atomic number describes the number of protons in an atom.  For hydrogen the atomic number, Z, is equal to 1.  For helium Z = 2.  Since the number of protons equals the number of electrons in the neutral atom, Z also tells you the number of electrons in the atom.  The atomic mass tells you the number of protons plus neutrons in an atom.  Therefore, the atomic mass, A, of hydrogen is 1.  For helium A = 4.

Atomic Structure

---

An atom is the smallest building block of matter. Atoms are made of neutrons, protons and electrons. The nucleus of an atom is extremely small in comparison to the atom. If an atom was the size of the Houston Astrodome, then its nucleus would be the size of a pea.

Introduction to the Periodic Table

Scientists use the Periodic Table in order to find out important information about various elements. Created by Dmitri Mendeleev (1834-1907), the periodic table orders all known elements in accordance to their similarities. When Mendeleev began grouping elements, he noticed the Law of Chemical Periodicity. This law states, "the properties of the elements are periodic functions of atomic number." The periodic table is a chart that categorizes elements by "groups" and "periods." All elements are ordered by their atomic number. The atomic number is the number of protons per atom. In a neutral atom, the number of electrons equals the number of protons. The periodic table represents neutral atoms. The atomic number is typically located above the element symbol. Beneath the element symbol is the atomic mass. Atomic mass is measured in Atomic Mass Units where 1 amu = (1/12) mass of carbon measured in grams. The atomic mass number is equal to the number of protons plus neutrons, which provides the average weight of all isotopes of any given element. This number is typically found beneath the element symbol. Atoms with the same atomic number, but different mass numbers are called isotopes. Below is a diagram of a typical cells on the periodic table.

There are two main classifications in the periodic table, "groups" and "periods." Groups are the vertical columns that include elements with similar chemical and physical properties. Periods are the horizontal rows. Going from left to right on the periodic table, you will find metals, then metalloids, and finally nonmetals. The 4th, 5th, and 6th periods are called the transition metals. These elements are all metals and can be found pure in nature. They are known for their beauty and durability. The transition metals include two periods known as the lanthanides and the actinides, which are located at the very bottom of the periodic table. The chart below gives a brief description of each group in the periodic table.

Group 1A	Known as Alkali Metals Very reactive Never found free in nature React readily with water
Group 2A	Known as Alkaline earth elements All are metals Occur only in compounds React with oxygen in the general formula EO (where O is oxygen and E is Group 2A element)
Group 3A	Metalloids Includes Aluminum (the most abundant metal in the earth) Forms oxygen compounds with a X2O3 formula
Group 4A	Includes metals and nonmetals Go from nonmetals at the top of the column to metals at the bottom All oxygen form compounds with a XO2 formula
Group 5A	All elements form an oxygen or sulfur compound with E2O3 or E2S3 formulas
Group 6A	Includes oxygen, one of the most abundant elements. Generally, oxygen compound formulas within this group are EO2 and EO3
Group 7A	Elements combine violently with alkali metals to form salts Called halogens, which mean "salt forming" Are all highly reactive
Group 8A	Least reactive group All elements are gases Not very abundant on earth Given the name noble gas because they are not very reactive

Charges in the Atom

The charges in the atom are crucial in understanding how the atom works. An electron has a negative charge, a proton has a positive charge and a neutron has no charge. Electrons and protons have the same magnitude of charge. Like charges repel, so protons repel one another as do electrons. Opposite charges attract which causes the electrons to be attracted to the protons. As the electrons and protons grow farther apart, the forces they exert on each other decrease.

 

Atomic Models and the Quantum Numbers

There are different models of the structure of the atom. One of the first models was created by Niels Bohr, a Danish physicist. He proposed a model in which electrons circle the nucleus in "orbits" around the nucleus, much in the same way as planets orbit the sun. Each orbit represents an energy level which can be determined using equations generated by Planck and others discussed in more detail below. The Bohr model was later proven to be incorrect, but provides a useful model for building an explanation.

The "accepted" model is the quantum model. In the quantum model, we state that the electron cannot be found precisely, but we can predict the probability, or likelihood, of an electron being at some location in the atom. You should be familiar with quantum numbers, a series of three numbers used to describe the location of some object (like an electron) in three-dimensional space:

1. n: the principal quantum number, an integer value (1, 2, 3...) that is used to describe the quantum level, or shell, in which an electron resides. The principal quantum number is the primary number used to determine the amount of energy in an atom. Using one of the first important equations in atomic structure (developed by Niels Bohr), we can calculate the amount of energy in an atom with an electron at some value of n:

En = -

Rhc n2

2. where:
   R = Rydberg constant, a value of 1.097 X 10⁷ m^-1
   c = speed of light, 3.00 X 10⁸ m/s
   h = Planck's constant, 6.63 X 10 ^-34 J-s
   n = principal quantum number, no unit

3.      For example, how much energy does one electron with a principal quantum number of n= 2 have?

En = - Rhc n2

or

En = - (1.097x107 m-1 ∗ (6.63x10-34 J•s)∗(3.0x108 m•s-1) 22

= 5.5x10-19 J

4.      You might ask, well, who cares? In addition to the importance of knowing how much energy is in an atom (a very important characteristic!), we can also derive, or calculate, other information from this energy value. For example, can we see this energy? The table below suggests that we can. For example, suppose that an electron starts at the n=3 level (we'll call this the excited state) and it falls down to n=1 (the ground state). We can calculate the change in energy using the equation:

ΔE = hv = RH

1 ni2

-

1 nf2

5.      Where:
ΔE = change in energy (Joules)
h = Planck's constant with a value of 6.63 x 10^-34 (J-s)
ν is frequency (s^-1)
R_H is the Rydberg constant with a value of 2.18 x 10^-18J.
n_i is the initial quantum number
n_f is the final quantum number

6.      Using the equation below, we can calculate the wavelength and the frequency of the energy. The wavelength and the frequency give us information about how we might "see" the energy:

vλ = c

7. Where:
   ν = the frequency of radiation (s^-1)
   λ = the wavelength (m)
   c = the speed of light with a value of 3.00 x 10⁸ m/s in a vacuum

Speed of light =	3.00E+08
Rydberg constant =	2.18E-18
Planck's constant =	6.63E-34
Excited state, n =	3	4	5
Ground state, n =	2	2	2
Excited state energy (J)	2.42222E-19	1.363E-19	8.72E-20
Ground state energy (J)	5.45E-19	5.45E-19	5.45E-19
ΔE =	-3.02778E-19	-4.09E-19	-4.58E-19
ν =	4.56678E+14	6.165E+14	6.905E+14
λ(nm) =	656.92	486.61	434.47

8. l ("el", not the number 1): the azimuthal quantum number, a number that specifies a sublevel, or subshell, in an orbital. The value of the azimuthal quantum number is always one less than the principal quantum number n. For example, if n=1, then "el"=0. If n=3, then l can have three values: 0,1, and 2. The values of l are typically not identified as "0, 1, 2, and 3" but are more commonly called by their historic names, "s, p, d, and f", respectively. Since the quantum numbers were discovered through the study of light and lines on an electromagnetic spectra, chemists identified the lines by their quality: sharp, principal, diffuse and fundamental. The table below shows the relationship:

Value of l	Subshell designation
0	s
1	p
2	d
3	f

9. m: the magnetic quantum number. Each subshell is composed of one or more orbitals. In the study of light, it was discovered that additional lines appeared in the spectra produced when light was emitted in a magnetic field. The magnetic quantum number has values between -l and +l. When l =1, for example, m can have three values: -1, 0, and +1. Because you know from the chart above that the subshell designation for l =1 is "p", you now know that the p orbital has three components. In your study of chemistry, you will be presented with p_x, p_y, and p_z. Notice how the subscripts are related to a three-dimensional coordinate system, x, y, and z. The chart below shows a summary of the quantum numbers:

Principal Quantum Number (n)	Azimuthal Quantum Number (l)	Subshell Designation	Magnetic Quantum Number (m)	Number of orbitals in subshell
1	0	1s	0	1
2	0 1	2s 2p	0 -1 0 +1	1 3
3	0 1 2	3s 3p 3d	0 -1 0 +1 -2 -1 0 +1 +2	1 3 5
4	0 1 2 3	4s 4p 4d 4f	0 -1 0 +1 -2 -1 0 +1 +2 -3 -2 -1 0 +1 +2 +3	1 3 5 7

---

Chemists care about where electrons are in an atom or a molecule. In the early models, we believed that electrons move like billiard balls, and followed the rules of classical physics. The graphic below attempts to show that earlier models thought that we could identify the exact path, position, velocity, etc. of an electron or electrons in an atom:

A more accurate picture is that the electron(s) reside in a "cloud" that surrounds the nucleus of the atom. This concept is shown in the graphic below:

Chemists are interested in predicting the probability that the electron will be at some particular part of this cloud. The cloud is better known as an orbital, and comes in several different types, or shapes. Atomic orbitals are known as s, p, d, and f orbitals. Each type of atomic orbital has certain characteristics, such as shape. For example, as the graphic below shows, an s orbital is spherical in shape:

On this graph, the horizontal (x) axis represents the distance from the nucleus in units of a₀, or atomic units. The value of a₀ is 0.0529 nanometers (nm). The vertical (y) axis represents the probability density. What you should notice is that as the electron moves farther away from the nucleus, the probability of its being found at that distance decreases. In other words, the electron prefers to hang around close to the nucleus.

The three graphics below show some other orbitals. The first graph (top left) is of a "2s" orbital. Each "s" orbital can hold two electrons in its cloud. Notice how there is a relatively high probability of an electron being near the nucleus, then some space where the probability is close to zero, then the probability increases substantially at some distance from the nucleus. The graphic at the top right shows a "2p" atomic orbital. Orbitals that are "p" orbitals can hold up to six (6) electrons in their cloud. Notice its "dumbbell" or "figure of eight" shape. At the bottom left is a "3s" orbital. Again, notice its spherical shape. Finally, at the bottom right, is a "3p" orbital.

 

Determining Electron Configuration

One of the skills you will need to learn to succeed in freshman chemistry is being able to determine the electron configuration of an atom. An electron configuration is basically an account of how many electrons there are, and in what orbitals they reside under "normal" conditions. For example, the element hydrogen (H) has one electron. We know this because its atomic number is one (1), and the atomic number tells you the number of electrons. Where does this electron go? The one electron of hydrogen goes into the lowest energy state it possibly can, which means it will start at "level" one and goes into "s" orbitals first. We say that hydrogen has a "[1s¹]" electron configuration. Looking at the next element on the Periodic Table --helium, or He -- we see it has an atomic number of two, so two electrons. Since " s" orbitals can hold up to two electrons, helium has an electron configuration of "[1s²]".

What about larger atoms? Let's look at carbon, with an atomic number of 6. Where do its 6 electrons go?

• First two: 1s²
• Next two: 2s²
• Last two: 2p²

We can therefore say that carbon has the electron configuration of "[1s²2s²2p²]".

The table below shows the subshells, the number of orbitals, and the maximum number of electrons allowed:

Subshell	Number of Orbitals	Maximum Number of Electrons
s	1	2
p	3	6
d	5	10
f	7	14

The Abridged (shortened) Periodic Table below shows the electron configurations of the elements. Notice for space reasons we sometimes leave off a portion of the electron configuration. For example, look at argon (Ar), element 18. The table below shows its electron configuration as "[3s²3p⁶]" (remembering that "p" orbitals can hold up to six (6) electrons). Its actual electron configuration is:

Ar = [1s²2s²2p⁶3s²3p⁶]

Sometimes you will see the notation: "[Ne]3s23p6", which means to include everything that is configuration of magnesium could be written [Ne]3s2, rather than writing out 1s22s22p63s2 Electronic Configuration of the Elements

Hydrogen through Krypton

Related Resources • Periodic Table of the Elements • Valences of Elements • Periodic Properties of Elements • Chemistry of Element Groups

Here's a useful table for your chemistry homework or general use! This is a compilation of the electron configurations of the elements up through number 104, broken into three pages (the table was too large for anything less). To arrive at the electron configurations of atoms, you must know the order in which the different sublevels are filled. Electrons enter available sublevels in order of their increasing energy. A sublevel is filled or half-filled before the next sublevel is entered. For example, the s sublevel can only hold two electrons, so the 1s is filled at helium (1s²). The p sublevel can hold six electrons, the d sublevel can hold 10 electrons, and the f sublevel can hold 14 electrons. Common shorthand notation is to refer to the noble gas core, rather than write out the entire configuration.

All matter consists of particles called atoms. Here are some useful facts about atoms:

• Atoms cannot be divided using chemicals. They do consist of parts, which include protons, neutrons, and electrons, but an atom is a basic chemical building block of matter.
• Each electron has a negative electrical charge.
• Each proton has a positive electrical charge. The charge of a proton and an electron are equal in magnitude, yet opposite in sign. Electrons and protons are electrically attracted to each other.
• Each neutron is electrically neutral. In other words, neutrons do not have a charge and are not electrically attracted to either electrons or protons.
• Protons and neutrons are about the same size as each other and are much larger than electrons.
• The mass of a proton is essentially the same as that of a neutron. The mass of a proton is 1840 times greater than the mass of an electron.
• The nucleus of an atom contains protons and neutrons. The nucleus carries a positive electrical charge.
• Electrons move around outside the nucleus.
• Almost all of the mass of an atom is in its nucleus; almost all of the volume of an atom is occupied by electrons.
• The number of protons (also known as its atomic number) determines the element. Varying the number of neutrons results in isotopes. Varying the number of electrons results in ions. Isotopes and ions of an atom with a constant number of protons are all variations of a single element.
• The particles within an atom are bound together by powerful forces. In general, electrons are easier to add or remove from an atom than a proton or neutron. Chemical reactions largely involve atoms or groups of atoms and the interactions between their electrons.

·         Elements are made up of atoms.

·         Each atom has a nucleus situated at the center. It contains positively charged particles called protons, and neutral particles called neutrons.

·         Electrons are negatively charged particles which move around the nucleus in definite circular paths called orbits, shells or energy levels.

·         The mass number of an element is equal t the sum of the number of protons and number of neutrons in its nucleus.

·         Number of protons equals the number of electrons in an atom; therefore, an atom is electrically neutral.

·         Atomic number is the number of protons of an atom.

·         Isotopes are atoms of the same element having different mass numbers.

·         The distribution of electrons in various shells or energy levels in an atom is called the electronic configuration of that atom.

·         According to Bohr and Bury, the maximum number of electrons that can be accommodated in any energy level of an atom is given by the formula 2n², where ‘n’ represents the number of the energy level.

·         In order to exist independently by itself an atom must have eight electrons in its outermost shell two electrons if there is only one shell. This is the octet rule.

·         Atoms try to attain stable configuration (completing their outermost shell) either by losing, gaining or sharing electrons.

·         The force of attraction that holds atoms together in a molecule is known as a chemical bond.

·         A bond between an anion and cation is called an ionic bond. Cations give electrons to the anions.

·         A covalent bond id a bond in which both the reacting atoms are short of electrons. Thus, they attain stable electronic configuration by sharing electrons.

·         Coordinate bond is a covalent bond in which the shared pair of electrons is contributed by only one of the two atoms.