Can you convert one metal to another



The chemistry is the doctrine of the structure, behavior and transformation of chemical elements and their compounds as well as the laws that apply to them.

In chemical reactions, bonds between atoms are separated and newly formed, so there is a change in substance. Since the properties of atoms that are relevant to chemistry are almost exclusively based on their electronic structure (electron shell), fundamental areas of chemistry can also be used as "Physics of the outer electron shell" to be viewed as.

All processes in which at least one new substance is created therefore belong to the field of chemistry. All interventions that leave the composition of the substance (substance = substance) unchanged with regard to the elements it consists of (e.g. melting, solidification), on the other hand, belong to physics. Nuclear physics includes changes to the atomic nucleus.

See on this topic also that in Wikipedia.

Atoms

An atom consists of the nucleus with protons and neutrons and the electron shell. In simplified terms, one can assume that the electrons circle around the nucleus on certain orbits (orbitals). Since an atom is uncharged, the number of protons is equal to the number of electrons. Most atoms only appear in compounds. The only exception are the noble gases, of which only a few noble gas compounds are known.

Valence electrons

Valence electrons are those electrons that are involved in a chemical bond. They are located on the outermost shell, but with the D block (transition metals) the electrons from the second outermost shell also have an effect, with lanthanides and actinides even those from the third outermost shell. The outer shell is in the ideal state when it is fully occupied, so the octet rule applies to all elements of the main groups (except hydrogen and helium, since the innermost shell is already fully occupied with 2 electrons), because the outer shell requires a total of 8 electrons . From the third period onwards, the main group elements can "expand their octets" in connections, such as B. the sulfur in sulfuric acid proves. The noble gases, which they already have what they call the “noble gas state” with a full outer shell, rarely react with other atoms.

The number of valence electrons for main group elements can be determined by looking at the periodic table, the elements of the 1st main group (alkali metals) have one outer electron, the elements of the 2nd main group (alkaline earth metals) two, etc.

With the metals in the D-block and the higher periods, however, half-filled shells can also occur, since this state is also very stable. The exact number of valence electrons can be seen in the D block from the electron configuration.

elements

Elements consist of atoms with the same atomic number (also proton number), but not necessarily the same number of neutrons (elements with the same atomic number, but different neutron and mass numbers are called isotopes). Each element reacts in a certain way, but here, too, periodicities can be recognized, according to which the elements can be sorted. In this periodic table, elements that are standing among each other behave very similarly and are grouped together. All elements of a group have the same number of outer electrons. Horizontal lines in the periodic table show the elements with the same number of shells, with the ordinal number increasing from left to right.

Each element in the periodic table is assigned a specific abbreviation based on the Latin form of the name. This short notation is particularly important for empirical formulas and reaction equations, as it represents an enormous simplification.

Examples: iron, lat .:ferrum, has the abbreviation Fe; Hydrogen, lat .:hydrogenium, has the abbreviation H

Metals

Metals have only a few and / or far from the core and therefore weakly bound valence electrons, which they can easily give off (low ionization energy, low electronegativity). As a result, they often form positive ions by releasing their valence electrons and thus dissolving the outermost shell. The new, lower-lying outer shell is always fully occupied and the resulting cation thus fulfills the octet rule.

Examples of cations with a noble gas state: Na+, Mg2+, Al3+

The metals of the D block as well as those of the higher periods can form several different ions (e.g .: Fe2+ and Fe3+).

In their elementary state, metals are metallic, shiny solids that conduct electricity. They are also good heat conductors and malleable (ductile).

Non-metals

Non-metals have many valence electrons that are tightly bound to the nucleus (high ionization energy, high electronegativity). They reach the noble gas state by accepting electrons and thereby form negatively charged anions.

Examples of anions with a noble gas state: N3−, O2−, F

Another way for non-metals to meet the octet rule is covalent bonding.

In the elementary state, non-metals do not conduct electricity (exception: graphite, a modification of carbon).

Semi-metals

The semi-metals cannot really be classified here, because they combine typical properties of both classes (for example, many of them have a metallic sheen, but are poor conductors). Their electronegativity and ionization energy lie between the corresponding values ​​for metals and non-metals. Many elements of this class of substances exist in several modifications.

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Types of binding

The atoms that make up all the materials around us can combine to form chemical compounds. In general, a distinction is made between the following 3 basic types of ties, although there are also mixed forms:

Metal bond

If the sum and the difference of the electronegativities of the binding partners are small, metal binding occurs. The atoms give off their outer electrons and form a metal lattice in which the free outer electrons can move freely. These free external electrons are the reason for the good thermal and electrical conductivity of metals and metal alloys. Because there are no clearly delineated parts, metals and their alloys usually have high melting and boiling points. Examples of metal bonding: iron (Fe), bronze (Cu + Sn), brass (Cu + Zn)

Ionic bond

If there is a large difference in electronegativities (non-metal + metal), ionic bonding (= salt) occurs. The element with the lower electronegativity releases its outer electrons, which are absorbed by the binding partner. The resulting charges create strong electrostatic forces that hold cations and anions together. Salts are electrically neutral, so there is a certain ratio of cation number to anion number for each salt. A crystal lattice forms, which leads to poor thermal conductivity and a high melting and boiling point and is also responsible for the brittleness of salts, since charges of the same name are next to each other and repel if the structure is deformed. As solids, salts are insulators, but as a melt they conduct electricity because there are freely moving charge carriers (ions). Examples of ionic bonding: table salt (Na+Cl), Magnesium oxide (Mg2+O2−)

In the case of transition metals, their oxidation number must be stated in the name: Iron (II) chloride (Fe2+Cl2), Iron (III) chloride (Fe3+Cl3)

Polyatomic Ions

An ion can also contain several atoms. In general, a distinction is mainly made here between negatively charged acid residues and positively charged ions, which are differentiated by the addition of H+ to form bases. The charge can be recognized by the number of H.+-Ions.

Examples:

Salts can also contain polyatomic ions: ammonium chloride (NH4+Cl), Sodium hydrogen sulfate (Na+HSO4)

Covalent bond (atomic bond)

If the sum of the electronegativities is large but the difference is small, the covalent bond is formed (non-metal + non-metal). The binding partners share one or more electron pairs, which then belong to both. This is how they reach the noble gas state. There are single and multiple bonds according to the number of bonding electron pairs, which is reflected in the Lewis formula. There are also fixed numerical ratios for the atomic bond, since each atom involved is given an octet.

For example, one atom each of chlorine and hydrogen combine to form hydrogen chloride.

Because the strong binding forces only act between two atoms, clearly demarcated structures develop, these are called molecules. Because they only interact weakly with one another, the melting and boiling points of substances with an atomic bond are usually relatively low. Example: carbon dioxide (CO2) Since there are also slight differences in electronegativity in the covalent bond, polarization of the molecule can occur. The more electronegative binding partner attracts the binding electrons, which leads to partial charges in the molecule. Substances with a covalent bond consist of molecules, are non-conductors (insulators) and - depending on polarity and molecular size - relatively volatile, unless they are plastic or diamond-like giant molecules. Hydrogen sulfide and ammonia are relatively volatile (weakly polar), hydrogen gas is extremely volatile (non-polar) and water and hydrogen chloride (HCl) are less volatile (strongly polarized bond) - non-volatile substances with atomic bonds, on the other hand, consist of giant molecules (examples: diamond, polyethylene, starch) .

Structural chemistry

For the chemical properties of a compound, however, it is not only decisive which atoms it contains, but also how they are connected to one another (see chemical bond). In the case of certain chemical compounds, especially proteins and other organic compounds, not only the bonds between the atoms are decisive for the chemical properties, but also their spatial orientation (see isomerism).

The challenge in chemical synthesis is therefore usually to selectively break and / or create bonds between individual atoms of the reactant molecules in order to produce a desired substance (reaction product).

Basic substances

Basic substances always consist of only one type of compound. They have a fixed melting and boiling point and certain physical and chemical properties. They can be described with the aid of sum formulas that indicate the numerical ratios. All elements can also be assigned to this category at the same time.

Examples of basic substances (with empirical formula): Water (H2O), table salt (NaCl), oxygen (O2)

Substances can only be broken down further using chemical agents.

mixture

Mixtures (mixtures of substances) contain different types of pure substances, they have a melting and boiling range. A distinction is made here between homogeneous and heterogeneous mixtures.

Homogeneous mixture

In the case of a homogeneous mixture, the components are completely mixed at the molecular level, so that they are no longer e.g. E.g. they can be separated from one another by simple filtration. From the outside, a homogeneous mixture appears like a whole.

Examples of homogeneous mixtures: salt solutions, air, alloys

Homogeneous mixtures can, however, due to the different boiling and melting points of the components, by physical processes such as E.g. separate distillation.

Heterogeneous mix

In the case of a heterogeneous mixture, the mixing is not complete. There are always clearly separated phases that can be separated from one another relatively easily.

Examples of heterogeneous mixtures: blown dust, tobacco smoke, bath foam, mud, earth, suspensions, emulsions

Chemical reaction

Main article: Chemical reactions

Chemistry knows a multitude of different types of chemical reactions. However, they all have in common that at least one substance is converted (substance conversion = chemical reaction). A distinction is made between the substances involved Educts (Starting materials) and Products (Final substances). Generally speaking, a response is made using a Reaction equation written (also as Reaction scheme designated):

In most reactions, however, this does not only take place in the direction of the products, there is also a reverse reaction. To emphasize this, one writes:

All others can be derived from this simplified basic type of reaction scheme. Numbers in front of the reactants (coefficients) also indicate the molar ratio, whereby “1” is not written. In exceptional cases, fractional numbers are also possible as coefficients.

Example: Formation of water from the elements:

Example of a reversible (reversible) reaction: Boudouard equilibrium:

Chemical equilibrium

Main article: Chemical equilibrium

Chemical equilibrium is a state of dynamic equilibrium in a system, which means that no reaction can be observed from outside, although it is still taking place. The reason that the external state of the system remains unchanged is that the back and forth reactions proceed at the same pace. This dynamic equilibrium can be proven by marking some molecules. I would like to explain this using a specific example:

We take the following equilibrium reaction:

The percentages in equilibrium can be calculated as described in the paragraph "Constants and formulas" and are filled into a container, but instead of the normal hydrogen isotope protium with a proton in the nucleus, the isotope deuterium with one proton and one neutron is used. Outwardly no reaction can be seen, since the system is in equilibrium, but if you examine the contents of the vessel again after a few days, you will find that the deuterium atoms can be found in both ammonia and hydrogen .

Chemical equilibrium calculations

For any equilibrium one can define the so-called thermodynamic equilibrium constant K, which describes the position of the equilibrium. The following applies:

with activities a and stoichiometric coefficients ν. This equilibrium constant is related to the free enthalpy of reaction

coupled.

For a simple equilibrium reaction

then applies, for example

In many cases the activity can approximately be replaced by the concentration c of the substance. The equilibrium constant is then given a K for better identificationc designated. It is also possible to set up equilibrium constants with the molar fraction x or the partial pressure p. However, these then have a different numerical value and can therefore not be used directly in the equations of thermodynamics. So they have to be converted into "K's" beforehand. For example, the constant K appliesX in our example

.

Similarly, for Kp:

Using the previously established reaction equation for the formation of ammonia, one can now calculate z. B. Kx to demonstrate:

Influencing the balance

According to the principle of LeChatelier (principle of the smallest compulsion), a system tries to evade external constraints. If a reduction in volume occurs during a reaction, an increase in pressure will accelerate the forward reaction. Conversely, a reduction in pressure accelerates the reverse reaction. If the enthalpy increases with an endothermic reaction, a rise in temperature can accelerate it. This principle is often used by industry to increase the yield.

Another special feature of the chemical equilibrium is that it is always re-established when part of the reaction mixture is removed, which sometimes increases the yield enormously:

An esterification with a low equilibrium constant is given:

If dehydrating reagents are now added, the reaction will proceed completely to the right, despite the low equilibrium constant, since the reverse reaction is no longer possible without water.

Catalysts

Catalysts lower the activation energy for a reaction by directing the reaction path through intermediate stages, whereby the energy expenditure for this path is far lower than for the direct variant. Catalysts are not converted, i.e. H. the amount of substance of catalyst before the reaction is equal to the amount of substance after the reaction. Catalysts influence the reaction rate and thus also the time until equilibrium is established, but not the position of the equilibrium.

Important inorganic reactions

Education from the elements

Examples:

Formation of methane:

Formation of hydrogen sulfide:

Redox reactions

Oxidation and reduction are summarized under the term “redox reactions”. These completely opposite reactions are electron release in the case of oxidation (outdated: oxygen uptake / hydrogen release) and in reduction electron absorption (outdated: oxygen release / hydrogen uptake). Rapid oxidation is also known as "combustion". If a substance accepts electrons, another substance must have given them up beforehand.Therefore, reduction and oxidation always occur together. This must be taken into account when creating a reaction scheme for redox reactions. Substances that can easily oxidize other substances are called oxidizing agents, substances that can easily reduce other substances are called reducing agents. Reducing agents give off electrons, oxidizing agents take them up. (Acid-base reactions also fall into this category in principle, whereby the acid is oxidized and the base is reduced, although many see this as a separate reaction mechanism, since protons instead of electrons are exchanged here).

Examples:

Burning sulfur:

Reaction of fluorine with sodium:

Acid-base reactions

According to Brønsted's acid-base theory, an acid is a proton donor and a base is a proton acceptor. In acid-base reactions, the acid gives off a hydrogen ion, which is taken up by a base. When an acid reacts with a base, this process is called neutralization, as the resulting salt has a less acidic / basic pH than the starting substances. If the base is a metal hydroxide, the by-product is water. They differ from the redox reactions only in the sign of the transferred charge, which in this case is an H.+-Ion is.

Examples:

Formation of table salt:

Formation of potassium bromide:

Formation of ammonium chloride:

The Lewis acid-base theory corresponds to oxidation, with a Lewis acid being an electron acceptor and a Lewis base being an electron donor. The only difference is that with Lewis acids / bases always electron pairs instead of single electrons are released / accepted.

Reaction of oxides with water

Non-metal oxides react with water to form an acid, metal oxides form a metal hydroxide (base) with water.

Examples:

Formation of sulfuric acid:

Formation of silica:

Formation of aluminum hydroxide:

Formation of potassium hydroxide:

Reaction of acids with metals

Acids form salts with metals, the acid replacing the acidic H atoms with a metal. This reaction produces hydrogen as a by-product. Example of a monoprotonic acid and an alkaline earth metal (2nd main group).

Me: metal; R: acid residue; H: hydrogen atom

This reaction is all the more lively, the stronger and the more concentrated the acid and the less noble the metal. If the strength of the acid is too weak and / or the metal is too noble, there will be no or at best a weak reaction. The formation of a passivation layer can also prevent this reaction.

Examples:

Reaction of sulfuric acid with sodium:

Reaction of nitric acid with potassium:

But: Reaction of carbonic acid with copper: (Acid too weak and metal too noble)

Complex formation reaction

In inorganic chemistry, colored complexes are very often formed. Many such reactions are good detection reactions.

The general principle of construction of complexes is very simple:

In a complex formation reaction according to a certain scheme, ligands attach to a central atom (usually a metal ion) and thus ensure a massive improvement in solubility, for example. These ligands enter into a coordinative bond with the central atom. Depending on the coordination number (number of ligands on a central atom), size, polarity and charge, different geometric bodies are formed.

Detection reactions

Detection means are added to identify certain substances in unknown material samples. These react with the desired substance in a known way (acid-base, redox, complex formation and precipitation reactions - see under detection reactions). Detection reactions are characterized by good visibility, so that z. B. characteristically colored solutions and precipitates as well as specific smells occur. Any substances that react similarly must be separated or masked beforehand (see cation separation process).

Important organic reactions

See name reactions

See also

Category: Chemistry