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Encyclopedia about Chemistry


Enthalpy and reaction heat

Every chemical reaction results in a change in energy. This change in energy leads to a temperature change of the participating reactants and the surrounding area. Reaction energy is released or taken in, but in any case transferred in the form of heat energy.

Just like in physics, the law of conservation of energy is a cornerstone of chemistry. This law states that in a closed system, the amount of energy before and after a reaction remains unchanged. The existing energy can, however, be transferred from one reactant to another, or to a product of the reaction. The surrounding area, be it a beaker or other vessel, can also be affected, either heating up or cooling down. In this case, this surrounding area is considered a part of the system.


Enthalpy is the amount of energy contained in a certain material, the inner heat a material contains (from Greek, thalpos = heat, en = inner).

If energy is released to the surroundings in a reaction, it is called an exothermic reaction. If, however, energy is required for the reaction to occur, it can be transferred from the surroundings to the reaction itself (for example, if a reaction seems to produce "cold"), an endothermic reaction has taken place. Exothermic reactions take place with no outside help, while endothermic reactions need a little help from outside the system, often with the help of heat.

If reactants are in the same state of matter, we say a reaction is homogeneous. If they are in different states of matter, a reaction is said to be heterogeneous.

Bond energy and reaction energy

Bonds between atoms are also a function of the amount of inner bond energy they contain. The energy of the participating atoms in this bond are assigned a bond energy. To break a bond, it is necessary to add a sufficient amount of energy. Then, when new bonds are formed, a new bond energy is the result. The difference between the energy expended to break the bonds, and the energy of the bonds formed is called the reaction energy of a reaction.

Heat energy can be released through burning, or combustion, reactions (oxidation), in neutralisation reactions between acids and bases (though in the reaction of a weak acid with a weak base, only a small amount of energy will be released), in the dissolving of salts, during melting, and in the creation of new bonds, or the breaking of old ones.

The interaction of energy with its surroundings is a natural phenomenon for us, and often we do not even realise that a chemical reaction has taken place. We should be clear, however, that any type of burning reaction must, by definition, be an exothermic reaction.

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Galvanisation and Galvanic cells

In electrolysis, ions in an electrical field are attracted to an oppositely charged electrode. These ions then donate excess electrons to the electrode, or accept them. In this way, new neutral atoms or, as the case may be, molecules are formed, these lodging onto the electrodes in their solid state. This effect of electrolysis is used in a number of ways - for example in metal coating, or electric plating, which has as its goal the protection of metals from corrosion. It can also improve a metal’s outer appearance.

The process whereby one body is covered by a metal is called galvanisation. The body which we want to galvanise is dissolved in an electrolyte - a solution which has the ability to conduct an electrical current, thanks to the presence of free ions in that solution.

Two bodies are dissolved in solution. These, thanks to their ionic nature, allow the solution to form an electrical circuit. We want to electroplate one of the bodies, so we choose the second body to be one which is made up of the metal that we want to coat onto the first body. Often an ore is used, its ions dissociating into solution.

The material that we want to electroplate, or galvanise, is added in as the cathode. In its charged state, it has an excess of electrons, thereby an overall negative charge. If a current is allowed to run through solution and the bodies included in it, a layer of ore begins to form on the surface of the second body. The oppositely charged electrode (anode) accepts electrons from the ore dissolved in solution. Positively charged ions of the ore form, these moving to the cathode. There, the negative charge which was present is neutralised by the alloyed body, bonding with the ore ions. In this way, the surface of the coated body is covered with a perfect covering of ore.

The Galvanic Cell

In the course of a reaction, energy is often released in the form of an electrical voltage differential. If two electrodes of differing metals are placed in an acidic solution, so that they do not touch each other,

a voltage difference results between the two of them. A current of electrons begins to traverse towards the metal (a conductor of the first degree) and solution, with its free ions (conductors of the second degree). If both electrodes are joined by a line wire, the solution conducts an electrical current, and a closed circuit is the result. This is a galvanised cell. This change in chemical energy was first discovered by the Italian scientist Alessandro Volta (1745 - 1827).

The principle of the galvanised cell is one which results in a method to produce a source of electrical energy which can produce electrical current over the long term. Volta used his findings to invent the very first battery. All of today’s batteries are based on the same principle of the galvanised cell. Volta’s original cell was composed of two pieces of metal. These were submerged in an electrolytic solution, the result of a salt or an acid being dissolved in water (to a solution which contained free ions). Such batteries have a short lifetime, because polarisation forces appear, and the tension between poles gradually decreases.

Polarisation forces are caused by polarisation. Polarisation is the formation of bubbles of hydrogen on the copper coated anode. These bubbles isolate the surface of the anode, preventing any further movement of ions to the anode.

The chemist Leclanche was able to stop these polarisation effects by using a manganese oxide compound, which attracts electrons strongly. The reaction is still, however, very slow.

Dry cells are very similar to Leclanche’s first batteries, but they are portable, because the original ammonium chloride that was dissolved in solution is replaced by another electrolyte - an ammonium chloride paste. Similar cells yield electrical energy at a maximum voltage of 1.5 V. More energy can be gained by connecting multiple batteries.

Secondary batteries, or storage batteries, are also used. These can be recharged. The best example is the common car battery. In this case, lead discs are placed next to each other, these immersed in sulphuric acid. The discs have a large surface area which helps them to hold a high voltage - two volts on each disc. If we want to increase the voltage, more such discs must be connected in series. In the production of an electrical current, these lead discs are slowly transformed into lead sulphate, and the concentraion of sulphuric acid falls. When a battery is charged, the chemical process which results in the battery releasing electrical voltage can be reversed, and a storage battery can be recharged.

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Metals and their compounds

Chemical elements can be divided into metals and non-metals. But in reality, there are metals, metalloids and non-metals. Metals have a similar crystal structure to non-metals, but the two groups differ in their characteristic properties. Metals are good conductors. They are malleable and elastic. They have high boiling points and melting points. and, metals show a certain lustre, or shine. They are almost always found in the solid state of matter.

In the periodic table of the elements, metals are found in the various periods on the left side of the table. From left to right, metallic character decreases. On the right side of the table are the non-metals. Between these two groups, the metals and non-metals, are the metalloids.

Metalloids are similar to metals, but they do have a number of properties which set them apart from the metals. Some of the most important elements are metalloids. Iron, zinc and copper are three of the most important metalloids.

Metals and metalloids are used for a number of purposes in industry, surgery and other technology related fields. Naming all of the uses of these important elements Is beyond the scope of this document, but suffice it to say that life as we know it would not be possible without the processes that these elements participate in. Importantly, where metalloids are not able to be employed, we can usually use metals. Their expandability and high melting point are very important.

A big problem with metals, however, is corrosion, in layman’s terms, rusting. But this can be reduced by a variety of techniques and chemical tricks, or by using an alloy.

Metals have few electrons in their outermost shells, so they must combine with other elements which have more electrons in their outer shells to attain a noble gas configuration. Most of the metals are able to combine with oxygen. The combination of iron with oxygen is common rust.

Corrosion occurs when reactive metals come into contact with less stable compounds containing oxygen. Most metals react with oxygen. When this occurs, electrons are transferred from the metal to an oxygen-containing compound or to oxygen itself (reducing agent, oxidising agent). The metal oxidises. The production of corrosion causes metals to gradually degenerate, with their conductivity dropping. Alloys, however, can slow the process of rusting. There are several metals which corrode on their surfaces. This layer of corrosion is so air-tight that even oxygen cannot penetrate to lower levels of the metal which have not been corroded (passivation of metals). If these metals are treated with iron to produce a steel alloy, these protective properties are transferred to the steel (stainless steel).

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Metallic bonding

Atoms of metals have very few electrons in their outer electron shells (1-4 electrons), so their electrons cannot form electron pairs. There is therefore no possibility to form covalent bonds in order to gain a stable electron configuration. The atoms of metals are arranged in crystal lattices. The outer electrons are so loosely held on to that they can easily jump between atoms of a metal. In essence, they can move freely, and are only associated with an atom of a metal. This mass of electrons which moves back and forth between atoms can be considered a swarm of electrons. The movement of these electrons, in such a free way, gives the torso of the atom, the nucleus of a metallic atom, a positive charge.

Metallic bonds are based on the fact that this electron swarm, negatively charged, is attracted to the positively charged torso of other nuclei of other atoms. Individual atoms in metals are arranged in a so-called metallic lattice. This can also be called a metallic crystal, because it looks somewhat similar to the ionic lattice of a crystal.

Characteristics of metals

Metals occur in cubic lattice structures, face-centred (copper), hexagonally (magnesium) or cubic face-centred (alkaline metals). The good electrical conductivity of metals is a result of the movement of a metal’s electrons – the electron swarm. The presence of this swarm also gives metals their characteristic lustre, or shine. At high pressure, metals can be shaped, thanks to the individual torsos of metallic atoms being able to push toward each other, if the force is great enough. The atoms are pushed into the electron swarm, so the electrostatic attractions which characterise metallic bonds

remains unchanged. When heated, oscillations in the atomic nuclei of metals increase. For that reason, metals are good conductors of heat. When heated, particle movement can become so frenetic that the attractive forces between atoms can be broken, and the nuclei of metallic atoms can move towards and away from each other. This is the point when a metal starts to melt.

In the first group of the periodic table are the alkaline metals. They have one electron in their outer shells. With water, they react to create basic solutions (alkaline bases, alkaline hydroxides). Their chemical properties are identical, while their physical properties differ according to some basic rules. From top to bottom of the periodic table, they react with increasing intensity.

The alkaline earth metals (second main group, IIA) also react with water, but not as fiercely. This group has two electrons in its outer shell. They are often used in fireworks because they can colour flames quite efficiently.

Half of the metals belong to the groups which are called transition metals. They are in the eight groups in the centre of the periodic table, because their characteristics are not so similar to those of other metals. They are found in nature in a number of oxidation states. Iron, copper and zinc are some of the most important of the transition metals. Transition metals have numerous uses, and are often found in nature in states other than the metallic state.

All metals can be differentiated according to their reactivity. Metals are in general quite reactive due to the ease with which they can give up their electrons. When combined, a more reactive metal can replace a less reactive one. This characteristic is used to protect against corrosion. We simply take the metals we want to protect from corrosion and bring them into contact with a very reactive metal. The corrosion will be attacked by the more reactive metal, and through this contact, electrons will be transferred to the more pure metal. This excess of electrons will protect the metal from corrosion. The more reactive metal is, in essence, sacrificed to save the second metal, the less reactive one.

The relative reactivities of individual metals can be determined by allowing them to react with weak acids. Magnesium is very reactive, so that it reacts violently with acids. Metals with little or no reactivity are called pure metals. Mercury, silver, gold and platinum are considered to be pure metals. These find uses in areas where it is necessary for a material to last a long time, like in electronics or in space flight.

Pure and impure metals

Generally, metals can be divided into pure and impure. Purity and impurity was in the past mostly concerned with regards to a metal’s image and uses, which determined the price of that metal. Chemically, these terms are more concerned with how metals behave. If we dissolve a metal salt into a solution with electrolytes, electrons are released into solution. These electrons remain associated with the metal, while at the same time, metal ions are released. The negative charge of the metal prevents further dissociation of metal ions into solution and leads to the attraction of the dissolved metal ioins, which again accept electrons. A certain equilibrium is reached. At the interface of phases between the metal and electrolyte, a layer is formed which is a dividing line between voltages. This voltage is called the potential of the electrode or its redox potential.

If we compare metals, we see that less pure metals donate their electrons easily, while more pure metals hold onto their electrons more strongly. According to their electrochemical characteristics, we can rate metals according to their electrochemical potential. Another possibility is to arrange them according to electrode potential. To do this, a galvanic cell can be produced from the interaction of two metals. The metal which donates electrons into solution of electrolyte is less pure. The difference between metals is quantified with the help of the voltage difference produced between metals during the experiment. Although the differences between metals can be tested only relatively, one in comparison with another, a list of most pure to least pure can be created from individual metals. The redox potential of metals can be measured in comparison with a standard hydrogen electrode whose potential is set at zero.

The Nernst Equation

The electrode potential of a metal is dependent on the concentration of a solution of electrolyte. This relation was discovered by W.H. Nernst (1864-1941). In a diluted solution of electrolyte, evidently, more metal ions are released than in a concentrated solution. If we combine these half cells with differing concentration of electrolyte solution, we observe an electrical current between cells from the cell with lower concentration of electrolyte to the cell with higher concentration of electrolyte. These types of galvanised cell are called concentration cells. The Nernst equation gives a quantitative relationship between the concentration of a solution of ions (an electrolyte solution) and electrode potential.

Redox reactions dependent on pH

Electrode potential of a hydrogen half cell is dependent on the pH of a solution of electrolytes. At a known electrode potential we can calculate the pH of a solution using the Nernst equation. This implies the possibility of calculating pH using the electrode potential. Today’s chemists use various glass electrodes and measuring devices which make speedy and accurate determination of pH possible.

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To be able to understand the oxidation of metals more easily, it is necessary to understand the meaning of the word oxidation in a clear and straightforward manner.

In past times, the word "oxidation" was used for processes in which some atom or atoms were combined with oxygen. Today, we use the word for processes in which the oxidation number of a substance changes. During oxidation, the oxidation number of an atom in a compound we are interested in rises.

Today we can describe oxidation with the help of the transfer of electrons. In an oxidation reaction, electrons are released by either an element as it stands alone or one in a compound. This loss of electrons is the oxidation of a material.

When elements are found in their natural state, they have the same amount of protons as electrons, that is, they have an oxidation number of 0. If an atom donates its electrons to a chemical bond in order to form a compound, this results in its oxidation. Its oxidation state rises in relation to the number of electrons it donates (+1, +2, +3,…). If an atom in a reaction with other atoms gains electrons (either partially or otherwise), its oxidation state has been reduced, and a reduction has taken place. Again, depending on the number of electrons received, the atom which receives the electrons can show a negative oxidation state (-1, -2, -3,…). There are some elements which can only donate electrons and cannot receive any at all. Their oxidation states can only be positive, never negative. They cannot hold electrons in their outer shell; it is energetically more advantageous for these elements to give up their electrons, thereby reaching the previous energy level’s noble gas configuration, a filled electron shell. Among the elements which do this are hydrogen (H) and all of the metals.

On the other hand, there are naturally existing metals which can only accept electrons and cannot donate any at all. In addition, oxygen (O) and flourine (F) are such elements, which can only accept electrons, never donating. Then there are some elements which can go both ways – either accepting or donating. If we want to determine whether an element in a compound accepts or donates electrons, it is necessary to begin with a comparison of elements whose behaviour is known (for example oxygen, hydrogen and some metals).

The oxidation number is a measure of how many electrons an element can donate or receive. In the equation of a reaction, the number of positive and negative charges (oxidation numbers of the individual elements) has to be the same on both sides of the arrow. The oxidation number is written following the chemical symbol of an element using Roman numerals, like this: Fe III Cl3. It is also necessary to know that one element can be assigned more than one oxidation number. That is, one element can lose, or gain, either one, two or more electrons.

An oxidation can only take place if a corresponding reduction takes place at the same time. Because both processes have to occur concurrently, we call these types of reactions redox reactions. A material which oxidises spontaneously functions as a reducing agent. At the same time, a material which reduces spontaneously is an oxidising agent.

Generally it can be said that a strong reducing agent has a high negative redox potential, and a strong oxidising agent always has a positive redox potential. Individual elements can be rated according to their redox potential, similar to the rating of the reactivity of metals. All materials have the ability to oxidise elements which are in front of them in this list, as well as being able to reduce the ones which come behind. The further apart two elements are from each other in redox potential, the more easily and violently they react with each other.

The burning of materials is nothing more than a violent oxidation process. The reaction takes place so quickly that the energy released can actually be seen, in the form of light and heat. All of the modern firefighting agents are based on one simple principle – denying oxygen to the fire, or taking it away. If there is not enough oxygen, an oxidation reaction cannot continue, and the flames of a fire die out.

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