The main type of bonding between two metals is so-called metal bond. The metal atoms give off all their valence electrons and thus reach the noble gas configuration.

The metal atoms become positively charged cations upon release of the electrons. Between these positively charged cations, the released electrons form the so-called electron gas, since the electrons can move freely in the atomic structure as in a gas so to speak. The cohesion of the atoms is due to the electrostatic attraction between the positively charged cations and the negatively charged electron gas. As the name implies, this type of binding has a special significance, especially in the case of metals.

The free mobility of the electrons in the electron gas is ultimately the cause of the generally good electrical and thermal conductivity of metals (the exception to this property is the group of so-called metalloids). The mutual repulsive forces of the metal cations and the simultaneous attracting force of the electron gas lead to a regular lattice structure.

In contrast to the lattice structure of an ionic bond, which consist of anions or cations, the atomic structure of the metal bond is completely identical. When single atoms or entire atomic series are displaced, there are basically no changes in the atomic structure in a metal.

In contrast to this, opposed charged ions encounter one another when shifting the ionic lattice. The repulsive forces between the identical ions finally “shatter” the material. This is the reason why ceramics are much more brittle due to their ionic bonding and can not be deformed like metals.

The ionic bond is the predominant type of bonding between a metal and a nonmetal. The metal atoms involved in the binding release their valence electrons, which are taken up by the nonmetal atoms. In both cases, the noble gas configuration for the respective atoms is achieved.

The metal atom becomes a positively charged ion (cation) after the release of the electrons. The non-metal atom becomes a negatively charged ion (anion) after the electrons are taken up. The cohesion between the metal and non-metal atoms is due to the electrostatic attraction of the resulting ions. The ionic bond has special significance for ceramics.

In an ionic bond, the metal atoms release their valence electrons, which are taken up the nonmetal atoms in order to reach the noble gas configuration for each atom!

Such crystalline solid compounds of anions and cations are often referred to as salts. A typical example of an ionic compound is therefore common salt, also known as table salt (NaCl). In this compound, the sodium atoms (Na) as alkali metals give off their only valence electrons. The metal atoms thus lose their third M shell, so that on the underlying L shell, the noble gas configuration with eight outer electrons is obtained. The emitted electrons of the sodium atoms are taken up by the nonmetallic chlorine atoms. The chlorine atoms with their seven outer electrons thus now bind eight outer electrons around each other, thus achieving the noble gas configuration. As a result, a crystal structure (ionic lattice structure) ist obtained by the attractive forces between the positive sodium atoms and the negative chlorine atoms.

Salts are ionic compounds consisting of anions and cations!

Basically, the tendency of an atom to bind electrons to itself is particularly great when only a few external electrons are missing to achieve the noble gas configuration. This applies in particular to the elements of the group of halogens with seven valence electrons (e.g. chlorine). Conversely, the tendency to take up electrons is low for those atoms which have only a small number of valence electrons. For these atoms it is usually energetically cheaper to donate the few electrons instead of take up so many. This is especially true for the group of alkali metals whose atoms have only one external electron each (e.g. Na).

Such a more or less pronounced tendency of atoms to attract additional electrons in the binding case is referred to electronegativity. For the above reasons, the electronegativity increases from left to right within a period of the periodic table. This can also be explained by the fact that the number of protons and thus the positive charge of the nucleus increases with increasing atomic number. This increases the atom’s ability to bind electrons as well. The values ​​for the electronegativity of the chemical elements are shown in the figure below (the darker the red the higher the electronegativity).

On the other hand the electronegativity usually decreases from top to bottom within a group. The reason for this is the greater distance of the outermost shell from the atomic nucleus, since with each period a new shell is added. Due to the greater distance, the attractive force between the nucleus and the valence electrons is reduced. Accordingly, the ability of the atoms to attract more electrons decreases. Note that the noble gases can not be assigned electronegativity because they do not bind or tend to donate electrons. The artificially generated elements can as well not be assigned electronegativity for practical reasons, since they are not stable and thus cannot be examined.

The tendency of atoms to attract electrons in the binding case is referred to as electronegativity. In the periodc table the electronegativity usually increases from left to right but decreases from top to bottom!

Whether an ionic bond wil be occur or not depends on the property of the atoms involved to release or take up electrons. The non-metal atoms should have the highest possible tendency to take up electrons (high electronegativity), while the metal atoms should tend to electron donation (low electronegativity). For this reason, it can be concluded of their ionic character from the difference of the electronegativity values of the two elements.

If the difference in the electronegativities of two chemical elements is greater than 1.8, then mainly an ionic bond will be present. On the other hand, if the difference is less than 1.8, a covalent bond is predominantly to be expected. In principle, however, there is always a certain amount of covalent behavior in ionic bonding. For example, sodium chloride has an ionic binding character of about 75%. The remaining 25% is accounted for a covalent part.

If the difference in the electronegativities of two chemical elements is greater than 1.8, then mainly an ionic bond will be present!

The covalent bond mainly occurs between two non-metals. The atoms involved in the bond, share together (“co”) valence electrons (“valent”) to reach the noble gas configuration. For that reason, this type of binding is also called covalent bond. Somewhat imprecisely this is often referred to as molecular bond. In addition to gases, this type of bonding has a special significance for plastics and ceramics.

In a covalent bond, the involved atoms use common valence electrons to reach the noble gas configuration!

With the understanding of covalent bonding, the reason why hydrogen does not occur in nature as a single H atom but always as an H₂ molecule (elementary hydrogen) is finally revealed. Two hydrogen atoms can share their outer electrons. This achieves the noble gas configuration of an helium atom with its two valence electrons.

For the same reason, chlorine particles always occur at a microscopic scale not as a single atom but as a Cl₂ molecule (elementary chlorine). In addition to these examples, the figure above shows the covalent bonding of four hydrogen atoms (4 H) and one carbon atom (C) to one methane molecule (CH₄). Also shown is the covalent bond of a water molecule (H₂O) consisting of one oxygen atom (O) and two hydrogen atoms (2 H). Note that in the figure, only the valence electrons of the atoms are shown, as only these are responsible for the chemical bond.

The representation of covalent bonds shown in the figure above is indeed very clear, but in many cases too costly or especially for double bonds not even possible. For this reason, one uses the so-called structural formula. The valence electrons of the individual atoms are represented by a dot next to the element symbol. For a hydrogen atom with only one valence electron, there is thus a point to the right of the H symbol. In the case of four external electrons, such as the carbon atom, on the other hand, there is one point to the right, left, above and below the element symbol. For atoms with more than four valence electrons, one electron is added to each side. Such a notation is also referred to as Lewis dot structure.

The Lewis dot structure is a structural formula which discribes the covalent bond. Binding electrons are represented by a dash between the atoms.

If atoms now enter into covalent bonds, the electrons involved in the binding are combined with a dash. These electrons are referred to as bonding electrons (bonding pairs or shared pairs). In the same way, the remaining pairs of electrons of an atom are also connected by a dash. However, these are no longer bonding electrons, since they have nothing to do with the bond itself. With the help of the Lewis dot structure one obtains a very clear information about the binding of the molecule.

The figure above shows very clear the noble gas configuration, which is formed by the surrounding electron pairs around an element. Thus, in all cases shown, either the noble gas configuration of helium with two electrons (or an electron pair) or the noble gas configuration with eight valence electrons (or four electron pairs). With the Lewis dot structure, double bonds can also be illustrated very easily, as is the case with carbon dioxide (CO₂) or ethylene (C₂H₄).

The notation describing only the type and number of atoms occurring in a molecule (for example H₂, CO₂, C₂H₄), etc.) is called the molecular formula!

In nature, substances rarely appear as pure elements. Much more, different elements bond for energetic reasons with each other and form chemical compounds. A typical example of this is water (H₂O). In this case, two hydrogen atoms (2H) combine with one oxygen atom (O) to form a stable water molecule:

\begin{align}
\label{wasser}
& 2 H ~+~ O ~\rightarrow ~ H_2O \\[5px]
\end{align}

On the other hand, bringing the two hydrogen atoms into contact with an argon atom (Ar) will not result in a stable bonding between these elements. Rather, the argon atom remains for itself and the two hydrogen atoms combine to form elementary hydrogen (H₂):

\begin{align}
\label{argon}
& 2 H ~+~ Ar ~\rightarrow ~ H_2 ~+~ Ar \\[5px]
\end{align}

If one considers the chemical elements with regard to their bonding behavior, it is noticeable that the elements of the 8th main group in the periodic table, are particularly stable. They take virtually no chemical reactions with other atoms and therefore do not form molecules. In nature they occur only monatomic (that means as single atoms).

For this reason, the argon atom does not bind chemically with the two hydrogen atoms. The elements of the 8th main group are all gaseous at room temperature, which gives this group the name noble gases. Based on their “sluggish” chemical behavior, these gases are also referred to as inert gases. Since these noble gases practically do not react with other substances, some of them are used as shielding gases against unwanted oxidation during welding.

Since the number of outermost electrons decisively influence the chemical behavior of an atom, the number of eight so-called valence electrons (or two in case of helium) means a particularly stable electron occupation. This electron configuration is obviously very favorable in terms of energy. The experimental investigations of the chemical bonding behavior of different atoms also confirm this assumption. This shows that atoms always try to form chemical bonds in such a way that eight or two valence electrons form around the atoms involved.

This state of an atom within a chemical bond with eight or two external electrons is also referred to as an noble gas configuration. Thus, an important rule for the chemical bonding behavior can be derived:

Each atom strives to reach the closest noble gas configuration in the periodic table (octet rule).

Based on the noble gases with their eight valence electrons (exception: helium), the effort to achieve the noble gas configuration is also called octet rule. The noble gas state is achieved by the fact that the atoms form chemical bonds and thereby

• absorb or release electrons (ionic bond, metal bond), or
• use it together with other atoms (covalent bond).

The respective chapters will briefly explain these most important types of bonding. It should always be noted that in reality, bonds can not be sharply limited to a certain type of bond. Rather, chemical compounds have features of different types of bonds.

Important note

in the context of Bohr’s atomic model, it is often incorrectly claimed that the noble gas configuration means a fully occupied outermost shell. This statement is wrong! For in the Bohr model, the maximum number of electrons on the \(n\)^th shell results from the following equation:

\begin{equation}
N_{max} = 2 \cdot n^2
\end{equation}

With \(n\) = 3 argon offers space on its outermost third shell for a maximum of \(N\) = 18 electrons. However, argon only has 8 valence electrons on this shell. The outermost shell is therefore far from being fully occupied! At this point it is argued with the wrong atomic model. Rather, the statement of the fully occupied (sub)shell is to be seen in connection with the orbitals that Sommerfeld introduced as an enhancement of the Bohr model (Bohr-Sommerfeld model):

The noble gas configuration means a fully occupied orbital (subshell) in the Bohr-Sommerfeld model!

More specifically, the noble gas configuration means a fully occupied p orbital. The exception is helium with a fully occupied s orbital.

The weaknesses of the Bohr model could be partially eliminated by the physicist Arnold Sommerfeld. In addition to the already introduced shells by Bohr, Sommerfeld further introduced subshells (also referred to as orbitals). With the introduction of these subshells, it was finally possible to explain the distribution of the electrons within the shells. The distribution of electrons in an atom is referred to as electron configuration. It will be explained in more detail below.

One imagines the main shells introduced by Bohr subdivided into subshells. The number of subshells depends on the main shell. The number of the main shell indicates the number of subshells. In the figure above, not all subshells of the higher main shells are shown, as these usually have no relevance. The lower shells are not labeled with numbers but with lowercase letters (s, p, d and f). A g-subshell exists only for theoretical elements with atomic numbers greater than 121 (called superactinoids), which is why this orbital has only theoretical meaning.

1^th main shell (K): 1 lower shell (s), identical to the main shell
2^nd main shell (L): 2 lower shells (s, p)
3^rd main shell (M): 3 lower shells (s, p, d)
4^th main shell (N): 4 lower shells (s, p, d, f)

For example, the shell designation 3p means the subshell p (“2nd subshell”) of the third main shell and the designation 4s the subshell s (“1st subshell”) of the fourth main shell. The shell designation 2d, however, does not exist because the second main shell has only one s and one p subshell! Subshells can only be occupied by a certain number of electrons:

s-subshell: 2 electrons
p-subshell: 6 electrons
d-subshell: 10 electrons
f-subshell: 14 electrons

Here, a subshell of a lower main shell number may well have a higher energy level than the subshell of a higher shell number (Sommerfeld explained this with elliptical orbits of electrons instead of circular orbits after Bohr)! For example, the subshell 3d has a higher energy state than the subshell 4s! The graphical representation by shells with their subdivisions in subshells is therefore no longer possible. Instead atomic orbitals are used, which will not be discussed further here.

The energetic distribution of the shells is shown in the figure above. The subshells are divided into white-framed blocks, each providing space for a total of two electrons.

In order to better remember the energetic order of the orbitals, you can first create a table. Therein, the line numbering corresponds to the main shell number and the column numbering corresponds to the subshell. Thus, the subshell with associated main shell is clearly defined for each field. The energetic order of the orbitals can now be obtained by going through the table diagonally line by line from top right to bottom left. This principle is also knows as the aufbau principle (“Aufbau” is a german word which means “configuration”).

The animation below shows the electron configuration with increasing atomic number of the atoms. Note that the number of electrons increases to the same extent as the number of protons and thus increases by one from element to element. The subshells belonging to a main shell are all marked in a uniform color. Also shown are the outer electrons (valence electrons) on the outermost main shell (valence shell), since these are decisive for the chemical behavior of an element. The animation also explains the order of the chemical elements in the periodic table.

The occupation of the shells with electrons always starts from the lowest energy state, only then are higher energy levels occupied. Each block of a subshell is initially filled with only one electron. This is symbolized by an ascending arrow. Only if all blocks of a subshell are solitary occupied by an electron, then the each block will be filled with one more electron. These second electrons are represented by a descending arrow. This symbolic distinction is due to the so-called Pauli exclusion principle of quantum mechanics. According to this principle no two identical electron states can exist. The different arrow directions take this principle into account (to be more precise: each arrow represents one of two spin quantum numbers).

The Pauli exclusion principle forbids that two electrons share the same state!

In the subshell 3p, it is striking that after it has been completely filled with electrons, it is energetically more favorable to start the fourth main shell and fill it with electrons (4s shell) instead of the subshell 3d! Note, that some subshells of a smaller main shell number are obviously of a higher energy level than subshells of a higher main shell number. Thus, the occupation of electrons of the main shells, which in the Bohr model at first appears a little bit strange, can now be explained.

Only after the 4s orbital is fully occupied, the third main shell is filled up with the energetically higher 3d orbital. This jump is also noticeable in the chemical behavior. It marks the transition from the so-called main group elements to the  transition group (transition metals). The transition metals are characterized by the fact that each of them have an incompletely occupied d-orbital (called d-block)!

Another jump where even one main shell is being skiped can be seen in the transition from the element barium (Ba) to cerium (Ce). After the 6s orbital of barium has been completely filled, two main shells are “jumped back” and the 4f orbital is filled up (the lanthanum in between of those elements is an exception to the Aufbau principle). This jump introduces a subgroup of transition metals, called lanthanides or actinides (f-block). The lanthanides or actinides are characterized by the fact that the f-orbital is gradually filled with electrons! Strictly speaking, the elements lanthanum and actinium do not belong to the group of lanthanides or actinides, although they are very often counted for practical reasons (after all, the suffix “ide” means similar to). Therefore, these elements also fall into the f-block.

Depending on which orbital an electron is added to, one can divide the periodic table into blocks, which correspond to s-, p-, d- or f-block.

Note, that there are other exceptions to the aufbau principle, for example, for the metals copper and chromium. There, an electron changes from the 4s subshell to the 3d subshell and thus remains occupied by only one electron. Such exceptions to the regular principle can be found especially at higher atomic numbers, as the electrons influence each other more and more. In addition, relativistic effects come into play, which are not taken into account by this model.

The Rutherford model in many cases provides a very good explanation of physical processes in matter. However, some phenomena can not be explained with this atomic model. For example, some atoms can only be excited to glow when bombarded with particles of specific energy. If the energy is only slightly lower, suddenly there is no illumination (for example Franck-Hertz experiment).

The physicist Niels Bohr suspected that this behavior must have something to do with the electron shell. Therefore, he expanded the atom model of Rutherford especially with regard to the atomic shell. He postulated that the electrons can only move in certain orbits around the atomic nucleus, comparable to the planetary motion around the sun. He called these discrete orbits shells. For this reason, the Bohr model is also referred to as shell model.

Each shell corresponds to a specific energy value of the electron (also called energy state or energy level). An electron can not assume an energy state that lies between two shells, since there can not be any electron there. The further away the shell is from the atomic nucleus, the more energetic is the state of an electron located there (higher energy level). This is the first innovation in the atomic concept that Bohr formulated as a postulate:

Electrons move only on discrete shells around the nucleus, each representing a certain energy level!

A postulate is a principle on which a theory is based!

With Bohr’s shell model can now finally be understood why atoms absorb only certain amounts of energy. Such an energy intake is also referred to as absorption. The absorption of energy can only happen if the energy supply is at least as large as an electron can be “lifted” from its current shell to the next higher one. Since there are no energy states between two shells, at lower energy levels, no electron can be brought to a next higher shell. The amount of energy supplied is not absorbed by the atom or by the electrons. The atom then remains in its lowest-energy state, which is also called the ground state. The state of an atom after one or more electrons have been brought to a higher energy level is called an excited state.

Conversely, when an electron “falls” to a lower energy level (towards an inner shell), only discrete energy packets can be released. The process of releasing energy is referred to as emission. The emitted energy corresponds exactly to the difference in the energy level of the two shells which are invovled in this process. This energy is emitted in terms of radiation (photons).

In this way, it can be explained why mercury, for example, emits a specific energy spectrum to which specific wavelengths (colors) in the light spectrum belong. The figure below shows the emitted spectrum of a mercury-vapor lamp. It can be seen that only certain wavelengths are emitted. So only discrete energy leaps occur. This corresponds to the leaps of the electrons from an energetically higher to an energetically lower shell. Because of the sharply defined lines in the spectrum, it is also called a line spectrum.

This provides another important insight into Bohr’s new nuclear concept:

When an electron “falls” from an outer shell to an inner shell, a photon is emitted. The energy of the photon corresponds to the energy difference of the two shells.

Note: The discretely portioned energy in the transition of an electron between two shells is referred to as quantum and the process of releasing a quant is called quantum leap. The Bohr atom model thus already contains basic features of quantum physics.

Although the Bohr model of the atom is a further development of Rutherford’s atomic model, it also contains some weak points. Thus, the imaginary circular motion of an electron around the nucleus is an accelerated motion. However, such an accelerated movement of charged particles would have to lead to an energy dissipation. Thus, after a short time the electrons should have no more energy to stably orbit around the nucleus. The consequence would be that the electrons fall into the nucleus and the atom decays. Since this is obviously not the case, Bohr had to postulate another postulate, which, however, contradicts everyday knowledge:

The electrons orbit the nucleus without emitting radiation!

Based on the different energy states of the shells Bohr also made a statement about the distribution of the electrons on the respective shells. Thus, on the innermost shell, which he called K-shell, there can be a maximum of two electrons. On the following shell, the L-shell, a maximum of 8 electrons can be found. The following M-shell contains a maximum of 18 electrons and the N-shell a maximum of 32 electrons, etc .:

• 1st shell (K-shell): 2 electrons
• 2nd shell (L-shell): 8 electrons
• 3rd shell (M-shell): 18 electrons
• 4th shell (N-shell): 32 electrons
• 5th shell (O -shell) Shell): 50 electrons
• 6th shell (P shell): 72 electrons
• 7th shell (Q shell): 98 electrons

The maximum number \( N_{max} \) of electrons on a particular shell can be determined by the following equation, where \( n \) is the shell number:
\begin{equation}
\boxed{N_{max} = 2 \cdot n^2 } \\[5px]
\end{equation}

The occupation of the shells with electrons always takes place from the lowest-energy state or the lowest-energy shell. The magnesium atom with its total of 12 electrons thus occupies 2 electrons on the K shell and 8 electrons on the L shell. These shells are now fully occupied so that the last two electrons can fit on the M shell. The electrons on an unfilled shell (in this case: the two electrons on the M shell) are also called valence electrons. The shell itself is called valence shell. The valence electrons on the outermost shell decisively determine the chemical properties of the atom and are also responsible for the position of the element in the periodic table.

The outermost shell is called valence shell and the electrons located in there are referred to as valence electrons. Chemical properties as mostly influenced by the number of valence electrons!

Note that the maximum number of electrons on a shell does not mean that an atom can have as many valence electrons! Because not always such a simple occupation rule shows up as in the case of the magnesium atom. In some cases a new shell is being filled (which then forms the outer electrons) although the underlying one is not yet fully occupied. This is evident, for example, in the case of the calcium atom. While the valence shell contains two electrons, the underlying M shell is filled only with 8 electrons and not with the maximum number of 18.

The order of occupation of the shells with electrons must therefore also be based on further influences that can not yet be explained by the Bohr model. In addition, experimental findings show that the classification of the electron orbits into the shells was too simple. For in some experiments, one also found energetic radiation transitions, which were also discrete, but located between the energy levels of the above-mentioned shells. So there had to be a finer division of the shells. The question of how and why elements form chemical compounds can not be explained by Bohr model as well. The physicist Sommerfeld provided an important development of Bohr’s atomic model (Bohr-Sommerfeld model).

In 1910, the physicist Ernest Rutherford found that when a thin gold foil was bombarded with α-particles (twice positively charged helium nuclei with two neutrons \( ^4_2\text{He}^{2+}  \)), only very few of these particles collided with the atomic nuclei of the gold atoms. Almost all α-particles traveled on a straight trajectory through the foil, while only a few were deflected.

Obviously, very few α-particles were close enough to the positive nucleus of the gold atoms that they could be deflected to a significant degree by the repulsive forces. In most cases, the α-particles traversed the gold foil at quite a distance from the respective atomic nuclei and were scarcely affected in their trajectory. This experiment concluded that the nucleus would have to be much smaller compared to the rest of the atom or rather to its atomic shell.

Today we know that the atomic nucleus has a diameter which is 10,000 to 100,000 times smaller than the atomic shell! If the atomic nucleus had the size of a dollar coin, the diameter of the atomic shell would amount to about 2 km!

The gold foil experiment further showed that some α-particles were reflected back to the gold foil with almost no energy loss. They obviously had to hit something very massive and heavy (analogous to a tennis ball that hits a massive concrete wall and flies back at almost the same speed). From this, Rutherford concluded that nearly the entire mass of an atom must be concentrated in the nucleus to produce such a strong reflection effect. And indeed, nearly 99.9% of the total mass of an atom is contained in its nucleus. Only 0.1% of the mass is therefore attributable to the atomic shell. Today we know that a proton (as well as a neutron) has a mass about 1800 times as large as an electron.

These findings formed the basis for Rutherford’s atomic model (Rutherford model), whose quintessences are summarized below:

• an atom consists of an atomic nucleus and an atomic shell,
• the nucleus is positively charged and the atomic shell carries a negative charge,
• in the nucleus are positively charged protons (and neutrons),
• in the atomic shell are the negatively charged electrons,
• the nucleus is much smaller than the atomic shell and
• almost the entire mass of an atom is concentrated in its nucleus.

With the Rutherford model, the results of scattering experiments (such as those of the gold foil experiment) could be correctly explained. The basic mass and size ratios as well as the corresponding division into atomic nucleus and electron shell also reflect this atomic model.

For example, the question of why atoms can only be excited with certain energies can not be answered by this model. Or why atoms emit characteristic line spectra. Likewise, Rutherford’s atomic model gives no explanation why an atom is stable, because the circular motion of the electrons around the nucleus would actually lead to an energy dissipation. Accordingly, the electrons ought to fall into the nucleus after only a short time and no atom should therefore be stable!

Some of the weaknesses of Rutherford’s atomic model could be corrected by the physicist Niels Bohr in his model (Bohr model).

Note

In principle, models (such as the atomic models or the particle model) never claim to give a complete explanation of reality. Models are always attempts to depict reality within certain limits and make it explainable.

The Rutherford model is not fundamentally “wrong” but has only limits of validity. Therefore, the Rutherford is not obsolete but it depends on the phenomena to be described and explained. To explain, for example, the gold foil experiment, the Rutherford model is completely sufficient; this does not require an unnecessarily complex quantum mechanical model.

Models are attempts to describe observable phenomena within certain validity limits.

In the periodic table all chemical elements are classified according to their atomic number and their chemical properties in main group elements (columns IA to VIII A) and transition group elements (IB to VIII B). The number of protons increases continuously from left to right. In addition to this horizontal classification, the periodic table is divided vertically into periods. These periods are not chosen randomly but correspond in the shell model to the electron shell introduced by Bohr (K, L, M, N, O, P and Q).

Elements in a certain group are all showing similar chemical behavior due to their identical amount of electrons in their outermost shell (this accounts only for elements in the main group).

The periodic table arranges elements into groups with similar chemical properties and periods with identical number of shells!

As the number of period increases a new electron shell is added. Therefore  the atoms within a certain group grow larger from top to bottom of the periodic table. On the other hand the atomic radius decreases from left to right oft the periodic table. This is due to the increasing number of protons that comes along with the atomic number. The number of electrons in the shell will increase as well . The more protons an atomic nulceus contains the higher its charge and the higher the charge of the electron shell. However, a higher charge results in an stronger force of attraction between the nucleus and the shell. Since the number of shells will not increase within a period the stronger force of attraction will bind the shell much stronger to the nucleus.

The size of an atom will increase within a group from top to bottom but will decrease within a period from left to right!

Within each period the element on the rightmost side of the periodic table will have the highest force of attraction between its nucleus and its shell. This configuration makes the element extremely stable. Since the elements on the rightmoste side are gaseous they are referred to as noble gases (or inert gases).

The classification of the periodic table into main groups and transition groups is due to their different distribution of electrons in atomic orbitals (electron configuration). For the same reason a further division can be made into lanthanides and actinides (the term actionide derives from the fact that all these elements are radioactive).

In the main group the s and p orbitals of the respective atoms are occupied by electrons (“s-block” or “p-block”), while in the transition group an electron is added in the d-orbital of the respective atom (“d-block”). In case of lanthanides and actinides the occupation of the f-orbital (“f-block”) takes place.

The main group elements can be further subdivided according to their physical and chemical behavior. This is usually done as follows:

• non-metals
• alkali metals
• alkaline earth metals
• metals
• metalloids (sometimes misleading called semimetals)
• halogens and
• noble gases.

Note that alkali metals and alkaline earth metals are “metals” in the true sense. Between the group of the alkaline earth metals and the metals is the transition group which is not shown in this figure. The reason for the word “transition” now becomes clear and since all the elements within the transition group are metals they are also referred to as transition metals.  Thus, about 80 % of the existing elements are metals!

A few elements have proporties of metals as well as of nonmetals. Those are referred to as metalloids. However, there is no cear definition of a metalloid! Usually the metalloids include:

• boron (B)
• silicon (Si)
• germanium (Ge)
• arsenic (As)
• antimony (Sb)
• bismuth (Bi)
• selenium (Se)
• tellurium (Te)
• polonium (Po)

The number of outer electrons of an atom (also referred to as valence electrons) significantly determines the chemical properties of the respective element. For main group elements, the number of valence electrons corresponds directly to the main group number. For example, potassium (Ka) belongs to the main group number 1 and therefore has one electron in its outer shell; so does  sodium (Na) and cesium (Cs). Accordingly to the fifth main group, elements like nitrogen (N), phosphorus (P) and arsenic (As) do have five valence electrons.

The number of the main group corresponds to the number of valence electrons of the elements assigned therein! The chemical behavier is mostly influenced by the number of valence electrons!

The chemical similarity due to the common number of valence electrons is particularly evident in the case of the alkali metals (1st main group, with the exception of hydrogen), the alkaline earth metals (2nd main group), the halogens (7th main group) and the noble gases (8th main group) ,

However, this relatively simple determination of the valence electrons based on the group number only works for the main group elements. In the transition group, however, this principle fails. Thus, the entire transition metals have only one or two outer electrons. Consequently, all transition elements have similar chemical properties.

First, you can distinguish between pure substances and mixtures. Pure substances contain only one type of particle. In the simplest case, particle means a chemical element. For example hydrogen (H), pure iron (Fe) or graphite (carbon C). Not only single atoms but also whole molecules can form pure substances. These substances are characterized by a certain atomic ratio and are referred to as chemical compounds. Pure compounds for example are water (H₂O), carbon dioxide (CO₂), acetone (C₃H₆O) or cementite (Fe₃C).

In contrast to mixtures, pure substances have fixed atomic ratios (only in the specific case of elements does a pure substance consist of one type of atom)!

If matter consist of several types of particles (atoms or molecules) with no specific atomic ratio due to chemical bonds, this will be referred to as mixtures. Such mixtures can be further classified into heterogeneous mixtures and homogeneous mixtures.

Homogeneous mixtures do have a uniform distribution of the different particles. Homogeneous mixtures for example are gasmixtures like air as well as solutions. But in contrast to a gas mixture, the state of matter of a solution is liquid. Therefore homogeneous mixtures can further by classified into gasmixtures and solutions. The group of solutions includes, for example, sugar water, salt water or carbonated soda. In addition to gases and liquids, solids can also form homogeneous mixtures. This is the case with some alloys such as copper-nickel alloys. So this could be a third classification of homogeneous mixtures.

Mixtures of substances with an diverse distribution of the containing particles are referred to as heterogeneous mixtures. A mixture of a solid and a liquid is called a suspension, for example, iron sludge, quicksand or liquid concrete. Heterogeneous mixtures of different liquids, which can not be mixed homogeneously, will be referred to as emulsions (e.g. oil-water-mixture, milk, mayonnaise, etc.). For heterogeneous mixtures of two or more solids, such as iron ore, granite or marble, one speaks of a solid sol. The last group of heterogeneous mixtures are the aerosols. These are mixtures of solids or liquids in gases. Examples of aerosols are cigarette smoke, water fog or car exhaust.

A suspension is a mixture of solid particles dissolved in a liquid. A mixture of two different liquids, which can not be dissolved homogeneously, is called emulsion.

Atomic structure

Matter is made up of microscopic units called atoms. Chemical elements are composed of atoms of a certain type. The classification of the elements takes place according to the periodic table. If several atoms (chemical elements) react with each other and form a stable unit by chemical bonds, these are called molecules.

Molecules are stable compounds of atoms by using chemical bonds.

For example, water consists of the elements hydrogen and oxygen. In each case, two hydrogen atoms (H) and one oxygen atom (O) join together in order to form a stable H₂O molecule. Atomic units such as molecules, atoms, protons, neutrons, electrons, etc. are simply referred to as particles. In this connection one often speaks of the so-called particle model of matter.

The particle model of matter means that matter is made up of particles without making a difference between atom, molecules, etc.

According to Rutherford’s atomic model, an atom consist of an positively charged atomic nucleus (lat. nucleus = “core”) and an negatively charged electron shell. The nucleus contains positively charged protons. These protons are the reason for the positive charge of the atomic nucleus.

The repulsive force between the protons due to their identical charges is compensated by the strong attraction of the neutrons,  who are also present in the atomic nucleus. The neutrons themselves are electrically neutral, but they still exert a strong attraction to the protons. In this way the protons are held together stably in the nucleus.

The particles in the nucleus (neutrons and protons) are also referred to as nucleons.

The attraction between the protons an the neutrons can not be caused by a electrostatic field because neutrons do not carry electric charges and theirfore can not be affected by such a field. It is rather another type of force. This force is called strong nuclear force or strong interaction. In addition to the electromagnetic force, the gravitational force (gravity) and the weak interaction (weak nuclear force), the strong interaction is one of the four fundamental forces of physics.

The interaction of the strong nuclear force is very limited in range, but at small distances as in atomic nuclei , that force can be extremely strong. The strong interaction between the protons and neutrons is the reason why this nuclear force outweighs the repulsive electrostatic forces of the protons and thus holds the nucleus together. The strong interaction is the “clue” for the nucleus so to speak.

The strong nuclear force (strong interaction) between the nucleons holds the atomic nucleus together.

The electron shell is located around the positively charged nucleus of an atom. It is formed by the negatively charged electrons. In a simplistic notion, the electrons in this imaginary shell orbit the positive nucleus. The electrostatic forces of attraction between the positive nucleus and the negative electrons ensure that the orbiting electrons are held stably on their path around the atomic nucleus, so that the atom does not fall apart.

The electron shell is an imaginary shell where the electrons orbit the nucleus.

Atomic number

Characteristic for a particular type of atom or for a chemical element is the number of protons in the nucleus! The number of protons essentially determines the chemical behavior of the element and is responsible for the order in the periodic table. Therefore, the number of protons is often referred to as atomic number. For example, a hydrogen atom always has one proton in its nucleus. If it contained two or three protons in the nucleus, it would no longer be a hydrogen atom but a helium atom (2 protons) or a lithium atom (3 protons).

The number of protons in an atom’s nucleus (atomic number) determines the element.

Isotopes

In contrast to the number of protons, the number of neutrons is not characteristic for a chemical element. For example, a lithium atom usually has four neutrons in its nucleus. However, this only applies to 92.5% of all lithium atoms. The remaining 7.5% contain only three neutrons in the nucleus. Such modifications of atoms with different numbers of neutrons but of course still having the same number of protons (otherwise it would be another element) are referred to as isotopes. The lithium atom thus has two (stable) isotopes.

Isotopes do have the same number of protons but different number of neutrons.

The hydrogen atom even has three isotopes. The most abundant hydrogen isotope (99.98%) has only one proton in its nucleus and no neutron. Becaus of only having one proton that hydrogen isotope is also called protium (symbol: H).

If the hydrogen atom has a neutron in addition to the proton, this isotope is called deuterium (symbol: D). Deuterium is represented by only 0.015% of all naturally occurring hydrogen atoms.

Another hydrogen isotope even has two neutrons in the nucleus and is called tritium (symbol: T). This isotope accounts for only a tiny fraction of total hydrogen in nature. However, unlike protium and deuterium, tritium is not stable and decomposes with a half-life of approximately 12 years. Due to the decomposition tritium is radioactive.

Ions

In the electrically neutral state, there are just as many positively charged protons in the nucleus as there are electrons in shell. The electric charge of an electron and a proton is identical in magnitude, but with the opposite sign. In the macroscopic point of view, the electrostatic effects cancel each other out. In this state, the particle is electrically neutral. However, if this neutral state is disturbed by absorbing or removing electrons, the atom is called an ion. The process of absorbing or losing electron is called ionization.

With excess of electrons, the atom is negatively charged. A negatively charged ion is referred to as an anion. An electrically positively charged ion is  called a cation. Since the number of protons determines the element, an ion can only be obtained by donation of electrons, but not through the release of a proton (changing the number of protons would create a completely different element).

An ion is an electrically charged atom (or a group of atoms). A negatively charged atom is called anion and a positively charged atom is called cation.

Note, that anions are larger in size than their respective atom due to the excess of electron and cations are smaller in size because of the loss of electrons.