home Success stories Chemical properties of metals, corrosion of metals. Metal corrosion

Chemical properties of metals, corrosion of metals. Metal corrosion

Elements with metal properties are located in IA–VIA groups of the Periodic Table (Table 7).

Metals are also all elements located in IB – VIIIB‑groups ( transition metals).

There are currently 92 metals in the Periodic Table.

Typical metals are s‑elements (elements of group IA from Li to Fr, elements of group IIA from Mg to Ra). The general electronic formula of their atoms is ns 1–2. They are characterized by oxidation states +I and +II, respectively.

The small number of electrons (1–2) at the outer energy level of typical metal atoms suggests an easy loss of these electrons and the manifestation of strong reducing properties, which reflect low values electronegativity. This implies the limited chemical properties and methods of obtaining typical metals.

Characteristic feature typical metals is the tendency of their atoms to form cations and ionic chemical bonds with non-metal atoms. Compounds of typical metals with nonmetals are ionic crystals of “metal cation and nonmetal anion,” for example K + Br ‑, Ca 2+ O 2‑. Cations of typical metals are also included in compounds with complex anions - hydroxides and salts, for example Mg 2+ (OH ‑) 2, (Li +) 2 CO 3 2‑.

The A-group metals that form the amphoteric diagonal in the Periodic Table Be-Al-Ge-Sb-Po, as well as the metals adjacent to them (Ga, In, Tl, Sn, Pb, Bi) do not exhibit typical metallic properties. General electronic formula of their atoms ns 2 np 0–4 involves a greater variety of oxidation states, a greater ability to retain their own electrons, a gradual decrease in their reducing ability and the appearance of oxidizing ability, especially in high degrees oxidation (typical examples are compounds Tl III, Pb IV, Bi v). Similar chemical behavior is characteristic of most d-elements, i.e. elements of the B-groups of the Periodic Table ( typical examples– amphoteric elements Cr and Zn).

This manifestation of duality (amphoteric) properties, both metallic (basic) and non-metallic, is due to the nature of the chemical bond. In the solid state, compounds of atypical metals with nonmetals contain predominantly covalent bonds (but less strong than bonds between nonmetals). In solution, these bonds are easily broken, and the compounds dissociate into ions (in whole or in part). For example, the metal gallium consists of Ga 2 molecules; in the solid state, the chlorides of aluminum and mercury(II) AlCl 3 and HgCl 2 contain strongly covalent bonds, but in solution AlCl 3 dissociates almost completely, and HgCl 2 - to a very small extent (and even then into HgCl + and Cl ‑ ions).

In their free form, all metals are solids, except for one - mercury Hg, which under normal conditions is liquid. In metal crystals it predominates special kind communications ( metal connection); valence electrons are weakly bound to a particular atom in the lattice, and inside the metal there is a so-called electronic gas. All metals have high electrical conductivity (the highest is Ag, Cu, Au, Al, Mg) and thermal conductivity. There are low-melting metals (cesium Cs with a melting point of 28.7 °C melts from the heat of the hand) and, on the contrary, very refractory (tungsten W melts only at 3387 °C). A distinctive property of metals is their plasticity (malleability), as a result of which they can be rolled into thin sheets - foil (Sn, Al, Au) or drawn into wire (Cu, Al, Fe), however, very brittle metals (Zn, Sb) are also found , Bi).

In industry, they often use not pure metals, but mixtures of them - alloys, in which the beneficial properties of one metal are complemented by the beneficial properties of another. Thus, copper has low hardness and is unsuitable for the manufacture of machine parts, while alloys of copper and zinc ( brass) are already quite hard and are widely used in mechanical engineering. Aluminum has high ductility and sufficient lightness (low density), but is too soft. Based on it, an alloy with magnesium, copper and manganese is prepared - duralumin (duralumin), which, without losing useful properties aluminum, acquires high hardness and becomes suitable for aircraft construction. Alloys of iron with carbon (and additives of other metals) are widely known cast iron And steel.

Free metals are restorers. However, the reactivity of some metals is low due to the fact that they are coated surface oxide film, to varying degrees, resistant to the action of chemical reagents such as water, solutions of acids and alkalis.

For example, lead is always covered with an oxide film; its transition into solution requires not only exposure to a reagent (for example, dilute nitric acid), but also heating. The oxide film on aluminum prevents its reaction with water, but is destroyed by acids and alkalis. Loose oxide film (rust), formed on the surface of iron in moist air, does not interfere with further oxidation of iron.

Under the influence concentrated acids form on metals sustainable oxide film. This phenomenon is called passivation. So, in concentrated sulfuric acid metals such as Be, Bi, Co, Fe, Mg and Nb are passivated (and then do not react with acid), and in concentrated nitric acid– metals Al, Be, Bi, Co, Cr, Fe, Nb, Ni, Pb, Th and U.

When interacting with oxidizing agents in acidic solutions, most metals transform into cations, the charge of which is determined by the stable oxidation state of a given element in compounds (Na +, Ca 2+, Al 3+, Fe 2+ and Fe 3+).

The reducing activity of metals in an acidic solution is transmitted by a series of stresses. Most metals are transferred into solution with hydrochloric and dilute sulfuric acids, but Cu, Ag and Hg - only with sulfuric (concentrated) and nitric acids, and Pt and Au - with “regia vodka”.

An undesirable chemical property of metals is their corrosion, i.e. active destruction (oxidation) upon contact with water and under the influence of oxygen dissolved in it (oxygen corrosion). For example, the corrosion of iron products in water is widely known, as a result of which rust forms and the products crumble into powder.

Corrosion of metals also occurs in water due to the presence of dissolved gases CO 2 and SO 2; an acidic environment is created, and H + cations are displaced by active metals in the form of hydrogen H 2 (hydrogen corrosion).

The area of contact between two dissimilar metals can be especially corrosive. (contact corrosion). A galvanic couple is created between one metal, such as Fe, and another metal, such as Sn or Cu, placed in water. The flow of electrons goes from the more active metal, which is to the left in the voltage series (Fe), to the less active metal (Sn, Cu), and the more active metal is destroyed (corroded).

This is what causes the tinned surface to rust. tin cans(iron coated with tin) when stored in a humid atmosphere and handled carelessly (iron quickly deteriorates after even a small scratch appears, allowing the iron to come into contact with moisture). On the contrary, the galvanized surface of an iron bucket does not rust for a long time, since even if there are scratches, it is not the iron that corrodes, but the zinc (a more active metal than iron).

The corrosion resistance of a given metal increases when it is coated with a more active metal or when they are fused; Thus, coating iron with chromium or making an alloy of iron and chromium eliminates corrosion of iron. Chromed iron and steel containing chromium (stainless steel), have high corrosion resistance.

Are common ways to get metals in industry:

electrometallurgy, i.e., obtaining metals by electrolysis of melts (for the most active metals) or salt solutions;

pyrometallurgy, i.e., the recovery of metals from ores at high temperatures (for example, the production of iron in the blast furnace process);

hydrometallurgy, i.e., the separation of metals from solutions of their salts by more active metals (for example, the production of copper from a solution of CuSO 4 by the action of zinc, iron or aluminum).

Sometimes found in nature native metals(typical examples are Ag, Au, Pt, Hg), but more often metals are in the form of compounds ( metal ores). Metals vary in abundance in the earth’s crust: from the most common – Al, Na, Ca, Fe, Mg, K, Ti to the rarest – Bi, In, Ag, Au, Pt, Re.

DEFINITION

When in contact with the environment, many metals, as well as metal-based alloys, can be subject to destruction due to chemical interaction (ORR with substances in environment). This process is called corrosion.

There are corrosion in gases (gas corrosion), which occurs when high temperatures in the absence of exposure to moisture on metal surfaces, and electrochemical corrosion (corrosion in electrolyte solutions, as well as corrosion in a humid atmosphere). As a result of gas corrosion, oxide, sulfide, etc. are formed on the surface of metals. films. Furnace fittings, parts of internal combustion engines, etc. are subject to this type of corrosion.

As a result of electrochemical corrosion, metal oxidation can lead to both the formation of insoluble products and the transition of the metal into solution in the form of ions. This type of corrosion affects pipelines located in the ground, underwater parts of ships, etc.

Any electrolyte solution is an aqueous solution, and water contains oxygen and hydrogen that are capable of reduction:

O 2 + 4H + +4e = 2H 2 O (1)

2H + +2e=H 2 (2)

These elements are oxidizing agents that cause electrochemical corrosion.

When writing about the processes occurring during electrochemical corrosion, it is important to take into account standard electrode potentials (EP). Thus, in a neutral environment, the EC of process 1 is equal to 0.8B, therefore, metals whose EC is less than 0.8B (metals located in the activity series from its beginning to silver) are subject to oxidation by oxygen.

The EP of process 2 is -0.41V, which means that only those metals whose potential is lower than -0.41V (metals located in the activity series from its beginning to cadmium) are subject to oxidation with hydrogen.

The rate of corrosion is greatly influenced by impurities that a particular metal may contain. Thus, if a metal contains non-metallic impurities, and their EC is higher than the EC of the metal, then the corrosion rate increases significantly.

Types of corrosion

There are several types of corrosion: atmospheric (corrosion in humid air at zero altitude), corrosion in the soil, corrosion with uneven aeration (oxygen access to different parts of a metal product in solution is not the same), contact corrosion (contact of 2 metals with different EP in an environment where moisture is present).

During corrosion, electrochemical reactions occur on the electrodes (anode and cathode), which can be written by the corresponding equations. Thus, in an acidic environment, electrochemical corrosion occurs with hydrogen depolarization, i.e. Hydrogen is released at the cathode (1). In a neutral environment, electrochemical corrosion occurs with oxygen depolarization—water is reduced at the cathode (2).

K (cathode) (+): 2H + +2e=H 2 - reduction (1)

A (anode) (-): Me – ne →Me n + – oxidation

K (cathode) (+): O 2 + 2H 2 O + 4e → 4OH - - reduction (2)

In the case of atmospheric corrosion, the following electrochemical reactions occur on the electrodes (and at the cathode, depending on the environment, various processes can occur):

A (anode) (-): Me→Me n + +ne

K (cathode) (+): O 2 + 2H 2 O + 4e → 4OH - (in alkaline and neutral environments)

K (cathode) (+): O 2 + 4H + + 4e → 2H 2 O (in acidic medium)

Corrosion protection

Used for corrosion protection following methods: use of chemically resistant alloys; protection of the surface of metals with coatings, which most often use metals that are coated in air with oxide films that are resistant to the effects of the external environment; treatment of corrosive environments; electrochemical methods (cathodic protection, protector method).

Examples of problem solving

EXAMPLE 1

EXAMPLE 2

Exercise

The part consists of an alloy of iron and nickel. Which metal will corrode faster? Write down the equations of the anodic and cathodic processes during atmospheric corrosion. The values of standard electrode potentials are E(Fe 2+ /Fe) = - 0.444V, E(Ni 2+ /Ni) = -0.250V.

Solution

First of all, active metals (those with the most negative values of standard electrode potentials) are subject to corrosion; in this case, it is iron.

Corrosion is usually called the spontaneous destruction of metals as a result of their chemical and electrochemical interaction with external environment and converting them into stable compounds (oxides, hydroxides, salts).

Strictly speaking, corrosion is a set of redox processes that occur when metals come into contact with an aggressive environment, which leads to destruction metal products. By aggressive environment we mean an oxide atmosphere (the presence of oxygen in the Earth’s atmosphere makes it oxide), especially in the presence of water or electrolyte solutions.

Based on the mechanism of the process, a distinction is made between chemical and electrochemical corrosion of metals. Chemical corrosion is a common chemical reaction between metal atoms and various oxidizing agents. Examples of chemical corrosion are high-temperature oxidation of metals with oxygen, oxidation of the surface of aluminum in air, interaction of metals with chlorine, sulfur, hydrogen sulfide H 2 S, etc.

Electrochemical corrosion occurs in solutions, that is, mainly when metals come into contact with electrolyte solutions, especially in cases where the metals are in contact with less active metals. The rate of corrosion depends significantly on the activity of the metals, as well as on the concentration and nature of impurities in the water. In pure water, metals almost do not corrode, and in contact with more active metals, even in electrolyte solutions, they do not corrode.

Cause of metal corrosion

Many metals, including Iron, are found in the earth's crust in the form of oxides. The transition from metal to oxide is an energetically favorable process, in other words, oxides are more stable compounds than metals. In order to reverse the process and extract the metal from the ore, it is necessary to expend a lot of energy, so iron tends to turn back into oxide - as they say, iron rusts. Rusting is the term for corrosion, which is the process of oxidation of metals under the influence of the environment.

The cycle of metals in nature can be depicted using the following diagram:

Metal products rust because the steel from which they are made reacts with oxygen and water contained in the atmosphere. During corrosion of iron or steel, hydrated forms of ferum(III) oxide of various compositions are formed(Fe 2 O 3 ∙ xH 2 O). The oxide is permeable to air and water and does not form a protective layer on the metal surface. Therefore, metal corrosion continues under the layer of rust that has formed.

When metals come into contact with moist air, they are always subject to corrosion, but many factors influence the rate of rusting. Among them are the following: the presence of impurities in the metal; the presence of acids or other electrolytes in solutions that come into contact with the surface of the iron; oxygen contained in these solutions.

The mechanism of electrochemical corrosion of a metal surface

In most cases, corrosion is an electrochemical process. Electrochemical cells are formed on the surface of the metal, in which different areas act as areas of oxidation and areas of reduction.

Below are two reactions of the redox rusting process:

The overall equation for the iron corrosion reaction can be written as follows:

Schematically, the processes that occur on the surface of iron or steel upon contact with water can be represented as follows:

The concentration of oxygen dissolved in a drop of water determines which areas on the metal surface are the site of reduction and which are the site of oxidation.

At the edges of the drop, where the concentration of dissolved oxygen is higher, oxygen is reduced to hydroxide ions.

The electrons needed to reduce oxygen move from the center of the drop, where they are released during the oxidation of Iron and where the concentration of dissolved oxygen is low. Iron ions go into solution. The released electrons move along the metal surface to the edges of the drop.

The above explains why corrosion is most severe in the center of a drop of water or under a layer of paint: these are areas where oxygen supply is limited. Here, so-called “shells” are formed, in which Iron goes into solution.

Rust as such arises as a result of a sequence of secondary processes in a solution into which Iron ions and hydroxide ions diffuse from the metal surface. A protective layer is not formed on the surface.

The activity of the Oxygen reduction reaction depends on the acidity of the environment, so corrosion accelerates in an acidic environment. Any salt impurities, such as sodium chloride in seawater spray, contribute to the formation of rust because they increase the electrical conductivity of the water.

The corrosion problem may never be completely solved, and the best that can be hoped for is to slow it down, not stop it.

Corrosion protection methods

Today there are several ways to prevent corrosion.

Separation of metal from an aggressive environment - painting, lubrication with oils, coating with inactive metals or enamel (I), Bringing the surface of metals into contact with more active metals (II). Use of corrosion inhibitors (corrosion inhibitors) and corrosion resistant alloys (III).

I. The simplest way to protect steel from corrosion is to isolate the metal from atmospheric air. This can be done using oil, grease or applying a protective layer of paint.

Protective coatings made from organic polymers are now widely used. The coating can be made in different colors, and this is enough flexible solution corrosion problems. Even a quick glance at the things that surround us in everyday life provides a lot of examples of such a solution: a refrigerator, a dish dryer, a tray, a bicycle, etc.

II. Sometimes the iron is coated with a thin layer of another metal. Some manufacturers make car bodies from galvanized zinc coated steel. With this treatment, a layer of zinc oxide firmly adhered to the base is formed, and if the galvanic coating is not damaged, it protects well from serviceberry.

Even if such a coating has flaws, the steel body of the machine is still protected from rapid destruction because in this system zinc preferentially corrodes rather than iron, since zinc is a more reactive metal than iron. In this case, zinc is sacrificed. One of the earliest proposals for the use of sacrificial ("sacrificial") metals was made in 1824 to protect the metal plating of the hulls of sea boats from corrosion.

Today zinc blocks are used for corrosion protection oil platforms in the seas: corrosion from expensive, complex steel structures is transferred to pieces of metal that are easy to replace. What is the principle of such protection? Let's illustrate it using a diagram.

At certain intervals along the entire support that is in the sea, zinc blocks are attached. Since zinc is more active than iron (located to the left in the electrochemical voltage series), zinc is predominantly oxidized, and the iron surface predominantly remains untouched. In principle, any metal located to the left of iron in the electrochemical voltage series can be used to protect steel products.

A similar principle is used to protect reinforced concrete structures of residential buildings, in which all the iron rods are connected to each other and connected to a piece of magnesium buried in the ground.

III. A very common solution to the problem of corrosion protection is the use of corrosion-resistant alloys. Many steel products used in everyday life, especially those that are in constant contact with water: cookware, spoons, forks, knives, washing machine tank, etc. - made of stainless steel, which does not require additional protection.

Hard steel was invented in 1913 by Sheffield chemist Harry Brearley. He investigated the rapid wear of the rifling of gun barrels and decided to try steel with a high chromium content to see if it was possible to prolong the life of the weapon in this way.

Typically, when analyzing steel, the sample was dissolved in acid. Brearley, conducting such an analysis, encountered unexpected difficulties. His steel, with its high chromium content, did not dissolve. He also noticed that the samples left in the laboratory retained their original shine. Brearley immediately realized that he had invented a steel that was resistant to corrosion.

The invention of Harry Briarley encountered some prejudices. One of Sheffield's leading metal utensil makers thought Briarley's idea was "against nature" and another said that "resistance to corrosion is not a great advantage in knives which require cleaning after each use." Today we take it for granted that cookware retains its shine and is not affected by the acids contained in food.

Stainless steel does not corrode because a film of chromium(III) oxide forms on its surface. Unlike rust, this oxide is not affected by water, and it adheres tightly to the metal surface. With a thickness of only a few nanometers, the oxide film is invisible to the naked eye and does not hide the natural shine of the metal. At the same time, it is impenetrable to air and water and protects the metal. Moreover, if you scrape off the surface film, it will quickly recover.

Unfortunately, stainless steel is expensive, and we have to take this into account when choosing which steel to use. In modern technology, highly resistant steel is most often used with the following composition: 74% iron, 18% chromium, 8% nickel.

Since the use of stainless steel is not always economically justified, as is the use of protective layers of lubricants and paints, today quite often they use a thin layer of zinc (galvanized iron) or tin (tinned iron) to coat iron products. The latter is very often used in the manufacture of canned food.

The method of protecting canned food by coating the inner metal surface with tin was proposed by the Englishman Peter Durand. With such protection, canned food remains edible for a long time. Unfortunately, the canned food and beverage industry is not without its challenges. Different products create different environments inside the can, which have different effects on the metal and can cause corrosion.

At the beginning of the 20th century, canned beer began to be produced. However New Product was not an immediate success, and the reason for this was that the banks were being destroyed from the inside. The thin layer of tin that was used to cover the jars very rarely came out solid. Most often it had minor flaws. In an aqueous solution, iron oxidizes faster than tin (due to its higher activity). Iron ions Fe 2+ dissolved in beer (which is generally a good remedy for anemia) and gave the drink a metallic taste, and in addition, reduced its transparency. This reduced the popularity of canned beer. However, manufacturers managed to overcome this problem after they began to coat the inside of the cans with a special inert organic varnish.

Canned fruit contains organic acids, such as citric acid. In solution, these acids promote the binding of tin ions Sn 2+ and thereby increase the rate of dissolution of the tin coating, so in canned fruits (peaches, etc.) the tin mainly corrodes. Tin ions that enter food in this way are non-toxic. They do not significantly change the taste of canned fruits, except that they provide them with an islandy aftertaste. However, if such a jar is stored for too long, problems may arise. The thin layer of tin, which is oxidized, will eventually collapse under the influence of organic acids and begin to corrode the iron layer quite quickly.

Omsk State Technical University
Department of Chemistry
Novgorodtseva L.V.
GENERAL CHEMISTRY
Chemical properties metals
Corrosion of metals.
Lecture
Multimedia slide lecture
©OmSTU, 2014

Distribution of metals in nature

METALS IN NATURE. CLARK.

Claark number (or clarke of elements, even more often they say
simply the clarke of the element) - numbers expressing the average
content chemical elements in the earthly
crust, hydrosphere, Earth, space
bodies, geochemical or cosmochemical systems, etc.,
in relation to the total mass of this system.
Expressed in % or g/kg.
Most
distributed from
metals in the earth
bark
Aluminum
Al –
8.45% (wt.)
Iron
Fe-
4.4% (wt.)
Calcium
Ca-
3.3% (wt.)
Sodium
Na-
2.6% (wt.)
Magnesium
Mg-
2.1% (wt.)
Titanium
Ti-
0.61% (wt.)

THE MOST COMMON METAL COMPOUNDS IN NATURE

Of the natural metal compounds, the most
Oxides are common.
Fe2O3 - hematite; Fe3O4 – magnetic iron ore, magnetite;
Cu2O - cuprite; Al2O3 - corundum; TiO2 – rutile, anatase, brookite;
MnO2 - pyrolusite; SnO2 – cassiterite, etc.
Low-active sulfides are widespread
metals: NiS; CuS; ZnS; PbS; FeS2.
In the form of halides: - fluorides, chlorides - alkaline and
alkaline earth metals.
In the form of carbonates – light metals - Mg, Ca (CaCO3).
In the form of sulfates – active metals Na, Ca, Ba, Mg (Na2SO4).
Soluble metal salts are found in the water of oceans, seas, and lakes.

Obtaining metals

Extraction of metals from ores

Most metals occur in nature in the form
compounds with other elements, mainly in the form of ores.
In a free state (nuggets)
gold and platinum are found, and
silver and copper - partly; Sometimes
native mercury comes across
some other metals.
Au and Pt are mined through
mechanical separation from the rock,
in which they are contained (for example
by washing), or by removing
them from the rock with various reagents
followed by separation from
solution

PYROMETALLURGY

Obtaining metals from their ores
recovery at high
temperatures
Restorers
Carbon (coke)
PbO + C = Pb + CO
Carbon monoxide
(II)
Fe2O3 + 3CO = 2Fe + 3CO2
Hydrogen
MnO2 + 2H2 = Mn + 2H2O
hydrothermy
More active
metal
(metallothermy)
Fe2O3 + 2Al = 2Fe + Al2O3
aluminothermy
TiCl4 + 2Mg = Ti + 2MgCl2
magnetothermy
carbothermy

HYDROMETALLURGY

Dissolution of natural
compounds in the form of aqueous
solutions using various
reagents followed by
release of metal from
solution. The process goes on
normal temperatures.
Reducing agents - active
metals or electrons
Gold is extracted from ores using potassium cyanide, and
then reduced with powdered zinc
2K + Zn→ K2 + 2Au
The metal is obtained in a finely crushed state

ELECTROMETALLURGY

Electrometallurgy - obtaining metals from water
solutions or melts using electric current
(by electrolysis)
Electrolysis aqueous solutions: for the production of low-active metals
CuSO4 + H2O→ Cu0 + H2SO4 + O2
Cathode (-): Cu2+ + 2e- → Cu0
Anode (+): 2H2O - 4e- → O2 + 4H+
Electrolysis of melts: for
obtaining active metals
2NaClmelt→ 2Na0 + Cl20
Cathode (-): Na+ + 1e- → Na0
Anode (+): 2Cl- - 2e- → Cl20

FLOTATION METHOD

Flotation is a method based on
different surface wettability
mineral water.
Example: ore consisting of sulfur metal and empty
rocks, crushed, filled with water, adding low-polarity
organic matter (for foam formation) and small
amount of "collector" reactant that is adsorbed
surface of the mineral. A stream is passed through the mixture from below
air. The result is a mineral particle with a layer of molecules
"collectors" stick to air bubbles, and the particles are empty
rocks moistened with water sink to the bottom. Then foam
collected, pressed and obtained ore with high content
metal

MAGNETIC METHOD

Magnetic separation is used to enrich ores containing
minerals with relatively high magnetic susceptibility. To them
include magnetite, franklinite, ilmenite and pyrrhotite, as well as
some other iron minerals, the surfaces of which may be
imparted the desired properties by low-temperature firing.
Separation is carried out as in water,
as well as in dry environments. Dry separation
more suitable for large grains,
wet - for fine-grained sands and
sludge. Conventional magnetic separator
is a device in which
layer of ore several grains thick
moves continuously in magnetic
field. Magnetic particles are pulled out
from the flow of grains by tape and collected for
further processing; non-magnetic
the particles remain in the flow.

The nature of chemical bonds in metals

GENERAL PHYSICAL PROPERTIES

High electrical conductivity, high thermal conductivity,
plasticity, i.e. ability to undergo deformation
normal and elevated temperatures without collapsing.
Due to this property, metals
amenable to forging, rolling,
drawing into a wire (drawing),
stamping.
Metals are inherent
also metal
shine due to
their ability
reflects light well.

METALS. SEMICONDUCTORS. DIELECTRICS. ZONE THEORY.

METAL LINK

The ability of electrons to move freely throughout a crystal
and serves to transfer energy from one part of it to another
the reason for the high thermal and electrical conductivity of metals
Valence electrons that carry out chemical bonds
do not belong to two or more specific atoms,
and the entire metal crystal. In this case, the valence electrons
are able to move freely in the volume of the crystal.
Educated in this way
chemical bond is called
metal connection.
The totality of "free"
electrons in metal electron gas

Chemical properties of metals

Standard hydrogen electrode

To construct a numerical scale of electrode
potentials you need the potential of some electrode
process to accept equal to zero. As a reference for
To create such a scale, an electrode process is adopted:
2Н+ +2е- = Н2
Hydrogen electrode
Platinum plate,
electrolytically coated
spongy platinum and immersed
in 1M sulfuric acid solution,
through which it bubbles
hydrogen gas by
pressure 1 atmosphere.
On the contact surface
platinum with acid solution
equilibrium is established
process:
2Н+ +2е- ⇆ Н2

ELECTROCHEMICAL VOLTAGE SERIES OF METALS

The potential of the hydrogen electrode is reproduced with a very high
accuracy. Therefore, the hydrogen electrode is adopted as
standard when creating a scale of electrode potentials.
To determine the potential
or other electrode process
need to create a galvanic
element
from
subject
And
standard
hydrogen
electrode and measure its EMF.
Since the potential of the standard
hydrogen electrode is zero,
That
measurement
EMF
will
introduce
by yourself
potential
electrode process.
In this way, the electrochemical voltage series is obtained
metals Because measurements are carried out relative to hydrogen
electrode, this series is called the hydrogen scale.

RESTORING PROPERTIES. IONIZATION ENERGY

Weakening of restorative properties and activity
This series is called the electrochemical voltage series
metals Ionization energy is determined by the position
metal in the periodic table. In electrochemical
in the voltage series, the metal located to the left can displace
from solutions or molten salts, the metal to the right.

Using this series, you can predict how the metal will be
behave in pairs with another.
The electrochemical voltage series also includes
hydrogen. This allows us to draw a conclusion about what
metals can displace hydrogen from acid solutions.
For example, iron displaces hydrogen from solutions
acids, since it is to the left of it;
copper does not displace hydrogen, since it is located to the right
it in the series of metal stresses.

ACTIVITY OF METALS IN ACCORDANCE WITH METALS VOLTAGE RANGE

Li, K, Ba, Na, La, Mg, Lu, Be, Sc, Ti, Hf, Al, Zr, V, Mn, Cr, Zn, Fe, Cd,
Co, Mo, Sn, W, Pb, H
Ge, Sb, Bi, Cu, Re, Ag, Pd, Hg, Pt, Au.
All metals can be divided into groups:
active metals are in the activity series before Cd;
medium activity - are in the series from Cd to H;
low-active metals appear in the activity series after N.

GENERAL CHEMICAL PROPERTIES OF METALS

The chemical properties of metals are determined by:
the structure of their atoms,
type of crystal lattice.
The main and most general property of metals is good
reducing agents, i.e. easily give up electrons:
Ме0 - ne-→ Мen+
Based on a range of standard electrode potentials
it is possible to draw a conclusion about the chemical activity of metals
With salts, the more active metal (standing to the left in the row
voltage of metals) displaces the less active one from it
salts: Zn + CuCl2 → ZnCl2 + Cu

GENERAL PROPERTIES OF METALS. INTERACTION WITH SIMPLE SUBSTANCES

METAL
oxygen
oxides, peroxides,
superoxides
halogens
fluorides, chlorides,
bromides, iodides
sulfur
sulfides
nitrogen
nitrides
phosphorus
phosphides
hydrogen
hydrides
carbon
carbides
silicon
silicides

Interaction of metals with water

INTERACTION OF METALS WITH WATER FROM THE THERMODYNAMIC POINT OF VIEW

The interaction of metals with water proceeds according to the reaction:
Me0 + H2O = MeOH + 1/2 H2

Ϥ0Red
Oxidizer:
2H+ + 2e- → H2
Ϥ0Oh
E= Ϥ0Ох - Ϥ0Red > 0
Ϥ0Ох > Ϥ0Red
Standard electrode potential of hydrogen ions in water
(pH = 7):
Ϥ0Ох = -0.59.рН = -0.41 V
Therefore, the condition for the reduction of metal with water can be
write it in the form:
Ϥ0Red< -0,41 В
Those. All interact with water, displacing hydrogen from it.
metals up to Cd, standard electrode
whose potential is below -0.41 V.

INTERACTION OF METALS WITH WATER

Active metals (metals from the beginning of the activity series to Mg) with
Hydroxides and hydrogen are produced with water:
2Na + 2H2O = 2NaOH + H2
Metals of medium activity (from Mg to H2) give oxides and
hydrogen (when heated):
WITH hot water metals in the series from Mg to Cd react:
Mg + 2H2O = Mg(OH)2 + H2
3Fe + 4H2O = Fe3O4 + 4H2
Reaction temperature t = 100 °C
Reaction temperature t = 700 °C
Some of the metals located
between Mg and Cd, for example, Zn, Al are covered with protective oxide
films (ZnO, Al2O3) and do not dissolve in water, i.e. metal is not
active (passive). The phenomenon is called metal passivation.
Inactive metals do not react with water.

Interaction of metals with acids

INTERACTION WITH ACIDS

Dilute acids are oxidizing agents due to
hydrogen
Reducing agent: Ме0 - ne-→ Мen+
Ϥ0Red
Oxidizer:
2H+ + 2e- → H2
Ϥ0Ох = Ϥ0 2H+/H2 = 0
E= Ϥ0Ох - Ϥ0Red > 0
Ϥ0Ох > Ϥ0Red
Ϥ0Red< 0 В
Mg0 + 2HCl → MgCl2 + H2
Metals in the metal voltage series up to
hydrogen displaces it from acids (exceptions:
concentrated sulfuric acid, any nitric acid
concentration).

PASSIVATION OF METALS WITH ACIDS

Sometimes insoluble or slightly soluble
foods that inhibit the reaction.
For example, lead Pb does not dissolve in dilute sulfuric acid.
acid and hydrochloric acid, because PbSO4 and PbCl2 are formed, which are not
dissolve in water and inhibit oxidation.
Pb + 2HCl = PbCl2 + H2
Passivation effect due to the formation of a protective film on
surfaces, leading to a slower reaction,
observed in some other metals.
Most often, products are formed when interacting with
the following acids: H3PO4, H2SO3, H2CO3, HCN, HF.

SOLUBILITY TABLE

INTERACTION OF METALS WITH CONCENTRATED SULFURIC ACID

In concentrated sulfuric acid as an oxidizing agent
sulfur appears in the oxidation state +6, which is included in
composition of the sulfate ion SO42-.
Concentrated sulfuric acid oxidizes everything
metals whose standard electrode potential
less than 0.36 V, the maximum value of the electrode
potential in electrode processes involving SO42- sulfation.
Concentrated sulfuric acid is reduced to
following products
H2S+6O4 (k) → S+4O2 → S0 → H2S2-

INFLUENCE OF ACTIVITY OF METALS WHEN INTERACTING WITH CONCENTRATED SULFURIC ACID

Active metals react with acid, reducing it
to hydrogen sulfide
5H2S6+O4(k) + 4Zn = 4ZnSO4 + H2S +4H2O
Low-active metals react with acid, reducing
it to SO2
2H2S6+O4(k) + Cu0 = CuSO4 + SO2 + 2H2O
Concentrated sulfuric acid passivates metals
medium activity: Fe, Be, Cr, Co, Al. On a surface
metal, dense oxide films are formed:
3H2SO4(k) +2Fe = Fe2O3 + 3H2O +3SO2
Metals Re, Mo, Tc, Ti, V interact in accordance with
equation
2V +5H2SO4(k) =2HVO3+5SO2 + 4H2O

OXIDATING POWER OF NITRIC ACID

Acid residue of nitric acid (any concentration)
has high oxidizing ability.
In nitric acid, nitrogen acts as an oxidizing agent.
oxidation state +5.
The acid is reduced to the following products:
HN5+O3 → N4+O2 → N2+O → N+2O → N20 → N3-H3
Recovery rate is growing
The more concentrated the acid, the deeper
she is recovering.
The nature of the reaction products depends on both the concentration
acid and metal activity

HEAVY METALS

Known
near
magpie
various
definitions
term heavy metals, and it is impossible to indicate one of
them as the most accepted.
The criterion used may be atomic weight above
50,
Heavy metals include more than 40 metals
periodic system D.I. Mendeleev: V, Cr, Mn, Fe, Co,
Ni, Cu, Zn, Mo, Cd, Sn, Hg, Pb, Bi, etc.
Another commonly used criterion is density,
approximately equal to or greater than the density of iron (8 g/cm3),
N. Reimers classification: Pb, Cu, Zn, Ni, Cd, Co, Sb, Sn,
Bi, Hg.
There are classifications based on other
values of threshold density or atomic weight. Some
classifications make exceptions for noble and
rare metals, not classifying them as heavy, some
exclude non-ferrous metals (iron, manganese).

INTERACTION OF HEAVY METALS WITH NITRIC ACID

In case of interaction heavy metals With
concentrated nitric acid most often
nitrogen oxide (IV) NO2 is released, with dilute - oxide
nitrogen (II) NO.
HNO3(dil) + Cu0 → Cu(NO3)2 + NO + H2O
HNO3(conc) + Cu0 → Cu(NO3)2 + NO2 + H2O
In the case of concentrated nitric acid, most often
nitrogen oxide (IV) NO2 is released, in the case of dilute -
nitric oxide (II) NO.

INTERACTION OF ALKALI AND ALKALINE EARTH METALS WITH NITRIC ACID

Concentrated nitric acid when reacting with
alkaline (elements of group 1 of the main subgroup: Li, Na,
K, Rb, Cs, Fr) and alkaline earth metals (elements 2
groups of the main subgroup (except Be, Mg): Ca, Sr, Ba, Ra)
is reduced to nitric oxide (I) N2O
HNO3(conc) + Ca0 → Ca+2(NO3)2 + N+2O + H2O
HNO3(dil) + Ca0 → N-3H4NO3 + Ca(NO3)2+ H2O
Dilute nitric acid reacts with
alkali, alkaline earth metals, Zn, Fe
is reduced to ammonium nitrate NH4NO3.

INFLUENCE OF METAL OXIDATION STATE. PASSIVATION

Depending on the chemical nature of the metal, the following are noted:
patterns:
metals that have a stable low oxidation state,
form the corresponding ions:
Mg0 + HNO3 (dil) → Mg(NO3)2 + NO + H2O
metals (W, Ti, V, Re, Tc), for which the most characteristic
high degree
acids:
oxidation
are formed
oxygen-containing
W0 + 2HNO3 (dil) → H2WO4 + NO
W0 + 6HNO3 (conc) → H2WO4 + 6NO2 + 2H2O
3Tc + 7HNO3 (dil) → 3HTcO4 +7NO + 2H2O
Nitric acid of any concentration passivates
metals: Fe, Cr, Al, Be, Bi, Ni in the cold.

The oxidizing ability of nitric acid is enhanced
adding hydrogen fluoride to it
or hydrochloric acids.
These mixtures dissolve the most inactive metals.
21HF + 5HN+5O3 + Ta → 3H2-2 + 5NO + 10H2O

"Royal vodka" - mixture
concentrated acids HNO3
and HCl in a ratio of 1:3.
Is a liquid
yellow in color with the smell of chlorine and
nitrogen oxides.

INTERACTION WITH ACID MIXTURES

"Royal vodka" dissolves gold and platinum. Its action
due to the fact that nitric acid oxidizes hydrochloric acid
With
release of free chlorine and formation of chloride
nitrosyl N+3OCl:
HN+5O3 + 3HCl = Cl2 + N+3OCl + 2H2O
Nitrosyl chloride is a reaction intermediate and
decomposes:
2N+3OCl = 2NO + Cl2
Chlorine at the moment of release consists of atoms, which determines

High oxidizing ability of aqua regia.
Au+ HN+5O3 + 3HCl →AuCl3 + NO + 2H2O
3Pt+ 4HN+5O3 + 12HCl →3PtCl3 + 4NO + 8H2O
With excess HCl, gold(III) chloride and platinum(IV) chloride
form complex compounds H and H2
Au+ HN+5O3 + 4HCl → H + NO + 2H2O

Interaction of metals with aqueous solutions of alkalis

INTERACTION OF METALS WITH AQUEOUS SOLUTIONS OF ALKALI

Metals interact with alkali solutions
prone to the formation of anionic complexes, i.e. those
metals whose oxides and hydroxides have
amphoteric character:
These are amphoteric metals - Zn, Al, Be, Ga, Sn, Pb.
The mechanism of the reaction (oxidation occurs
due to water molecules):
Zn +2 H2O = Zn(OH)2↓ + H2
Zn(OH)2↓ +2 NaOH = Na2-2.

STABILITY OF ANIONIC COMPLEXES

The more stable complex anions like [E(OH)n]x-,
the easier the reaction is. It is noted that such anions
most
stable
at
such
metals
How
zinc,
aluminum, beryllium, so they dissolve easily
in aqueous solutions of alkalis. For iron, cobalt,
titanium,
manganese
complexes
slowly.
Not
And
row
stable
others
And
metals
interaction
such
coming

INTERACTION OF METALS WITH HIGH OXIDATION DEGREES WITH AQUEOUS ALKALI SOLUTIONS

Some d-elements also react with alkalis,
which in the presence of oxidizing agents form compounds with
high oxidation states. Vanadium, tungsten, chromium,
For example,
V
melted
alkalis
oxidize
oxygen into vanadates – Me3VO4, into tungstates – Me2WO4
and into chromates – Me2СrO4, respectively:
2W + 4NaOH + 3O2 = 2Na2WO4 + 2 H2O

DISSOLUTION OF ALUMINUM IN AN AQUEOUS ALKALI SOLUTION

Aluminum does not interact with water, although it is active
metal. The reason for the inertness of aluminum is the formation
on its surface under the influence of atmospheric oxygen in ordinary
conditions oxide film Al2O3, which has a very strong
protective effect. The added alkali dissolves the oxide
film with the formation of hydroxoaluminate and creates the possibility
direct interaction of aluminum with water.
The reaction proceeds according to the following scheme:
1.Al2O3 + 2NaOH + 3H2O → 2Na
2. 2Al + 6H2O → 2Al(OH)3 + 3H2
3. Al(OH)3 + NaOH → Na

Metal corrosion

CORROSION OF METALS.

Corrosion is
spontaneously
leaky
destruction process
metal in
result
interaction with
environment.

MATERIAL LOSSES.

Material losses when
corrosion:
Destruction of pipelines,
metal machine parts,
ship hulls, marine
structures (more than 10%
annual metal smelting
lost due to corrosion).
Cost of lost
product via
corroded system
pipes
Downtime of enterprises
replacement period
metal structures
corroded.

MECHANISM OF THE CORROSION PROCESS.

The process involves the release of energy and dispersion of matter
(the entropy of the system increases ∆S > 0).
The production of metals in their pure form is always accompanied by
energy expenditure.
This energy accumulates in them as free energy
Gibbs, and makes them chemically active substances.
Metallurgical process:
_
Corrosion process:
_
Мen+ + ne → Me0
Me0 - ne → Мen+
∆G0х.р. >0
∆G0х.р.< 0
(comes with energy cost)
(spontaneous process)

CLASSIFICATION OF CORROSION PROCESSES.

CORROSION OF METALS

Chemical corrosion of metals

ESSENCE OF CHEMICAL CORROSION.

Chemical corrosion is typical for non-conductive environments
electricity.
The essence of the chemical corrosion process comes down to
redox
reactions,
and
direct transition of metal electrons is observed
to the oxidizer.
It represents spontaneous destruction
metals in an oxidizing gas environment (O2, SO2, H2S,
halogens) or in liquid non-electrolytes (organic
liquids – sour oil).

CHEMICAL CORROSION IN GAS ENVIRONMENT

IN general view for gas corrosion:
1.2Me0 (t) + O2 (g) ⇄ 2 Me+2O (t)
Oxide formation on the metal surface in
as a result of interaction with atmospheric oxygen.
2. MeO (t) → [MeO] (p)
Dissolution of the oxide film in the metal itself,
the balance will be shifted to the right, because majority oxides
metals are capable of dissolving in the metal and leaving
equilibrium systems.
The mechanism of such corrosion is reduced to ion diffusion
metal through a film of corrosion products on one side,
and on the other side of the counter diffusion of oxygen atoms
deep into the film.

EXAMPLE OF CHEMICAL CORROSION IN A LIQUID ENVIRONMENT.

OXIDATION RATE

Determined by the properties of the oxide film on
metal surface:
film continuity;
film diffusion ability;
structure of the oxide film.
Film continuity () is estimated by the volume ratio
of the formed oxide to the volume of metal,
spent on the formation of this oxide (factor
Pilling-Badwords)
Values for
metals are given in reference books.

FILM CONTINUITY

If
< 1, то образующаяся пленка не сплошная. Толщина
oxide film grows in proportion to the oxidation time.
With increasing temperature, the corrosion process sharply
accelerates due to poor heat dissipation, the metal is heating up
and the reaction rate increases.
If
= 1.2 – 1.6, then the resulting oxide film
solid. Such a film inhibits the diffusion of the oxidizing agent. And by
As the film thickens, its further growth will continue
slow down. Continuous films form on the surface
metals such as Co, Ni, Mn, Ti.
At
> 1.6 oxide films
also not solid and easy
are separated
from
surfaces
metal (iron scale)

BLUEING

Voronea
steel (oxidation, blackening,
bluing) -
process
receiving
on
surfaces
carbon
or
low alloy steel or cast iron layer
oxides
iron with a thickness of 1-10 microns. The thickness of this layer determines its
color - the so-called tarnish colors, replacing each other
friend as the film grows (yellow, brown, cherry,
purple, blue, gray).
The structure of the coating is fine-crystalline,
microporous. To add shine and
improving the protective properties of the oxide film
it is also impregnated with oil (mineral
or vegetable).
Nowadays bluing is used
mainly as decorative
finishing, and before - mainly - for
reducing metal corrosion.

INFLUENCE OF THE STRUCTURE OF OXIDE FILMS ON THE RATE OF CORROSION

For metals with variable oxidation states, the structure
film thickness will be different, for example, when
When steel blued, the following oxide layers are formed:
Fe|diffuse layer|FeO|diffuse layer|Fe3O4
This structure of the oxide film
provides a strong connection
oxide layer with surface
metal
It has been experimentally proven that
oxide films with structure
spinels RO.R2O3 (FeO.Cr2O3 or
NiO.Cr2O3) serve as reliable
protection against corrosion.

QUANTITATING CORROSION RATE

Quantitatively, the rate of any type of corrosion is measured in
units of mass of lost metal (∆m) per unit
area (S) per unit time (t):
The corrosion rate can also be measured by layer thickness
lost metal per unit of time.
To determine the corrosion rate, weights are used,
volumetric and physical methods.

AGGRESSIVE ENVIRONMENTS AND STABILITY OF VARIOUS METALS

In addition to oxygen, they have a strong aggressive property
and other gases. The most active are fluorine (F),
sulfur dioxide (SO2), chlorine (Cl2), hydrogen sulfide (H2S). Their
aggressiveness towards metals, and therefore
the corrosion rate is not the same.
For example, aluminum and its alloys,
chromium and high content steels
chromium, unstable in the atmosphere,
containing
chlorine,
Although
By
attitude
To
oxygen
They
stable.
Nickel is not stable in the atmosphere
sulfur dioxide (SO2), and copper is quite
stable.

Electrochemical corrosion of metals

ELECTROCHEMICAL CORROSION

Electrochemical corrosion is typical for environments with
ionic conductivity, i.e. for electrolytes
In this case, the reaction between the metal and
oxidizing agent occurs in several stages:
1. Anodic
oxidation
metal
The metal in the form of ions goes into
solution,
A
equivalent
the number of electrons remains in
_
metal: Me0 - ne → Мen+
2. Cathode process - assimilation
(holding)
redundant
electrons in metal.
3. Movement of ions in solution.

CONDITIONS CONTINUING ELECTROCHEMICAL CORROSION

The position of the metal in the activity series of metals: than
The farther apart they are, the faster
corrosion occurs.
Metal purity: impurities accelerate corrosion.
Irregularities in the metal surface, cracks.
Groundwater, seawater, electrolyte solution
(ions: H+, Cl-, Br-, I-, for amphoteric metals OH-).
Temperature increase.
Action
microorganisms
(mushrooms,
bacteria, bacteria
lichens): affect metals with high
corrosion resistance.

MECHANISM OF ELECTROCHEMICAL CORROSION. ANODIC PROCESS.

Mechanism
electrochemical
corrosion
determined
the potential difference between the cathode and anode sections and
comes down to the operation of a gas galvanic cell.
The main difference between electrochemical corrosion processes
from galvanic cell processes is the absence
external circuit. In this case, electrons do not leave
corroding metal, but move inside the metal.
Since any metal always contains impurities of other
metals, then on its surface in an electrolyte environment
many short-circuited microgalvanic
elements.
The anode in them will be the base metal, which is oxidized by
_
reactions: Me0 - ne → Мen+

CATHOdic DEPOLARIZATION PROCESS.

The cathodic process most often occurs with oxygen or
hydrogen depolarization.
A depolarizer is a substance that holds electrons.
Oxygen depolarization occurs with the participation
dissolved oxygen, which is
depolarizer:
at
: O2 + 2H2O + 4e- → 4OHat
: O2 + 2H+ + 4e- → 2H2O
Hydrogen depolarization occurs with the participation of cations
hydrogen medium (depolarizer - hydrogen):
at
: 2Н+ + 2е- → Н20

ELECTROCHEMICAL CORROSION

Anode (-) Fe / O2, H2O, NaCl / Sn (+) Cathode
A: Fe
0
_
- 2e → Fe
2+
_

DIAGRAM OF CORROSIVE MICROGALVANO COUPLE AT Zn - Cu CONTACT

Anode (-) Zn / medium / Cu (+) Cathode
Anode (-) Zn / O2, H2O, NaCl / Cu (+) Cathode
Zn
0
_
- 2e → Zn
2+
_
O2 + 2H2O + 4e → 4OH oxygen depolarization
Anode (-) Zn / H2SO4 / Cu (+) Cathode
_
2Н+ + 2е → Н20
hydrogen depolarization

DIAGRAM OF MICROGALVANO COUPLE OF ELECTROCHEMICAL CORROSION IN CONTACT OF ZINC IRON

Anode (-) Zn / O2, H2O, NaCl / Fe (+) Cathode

AERATION CORROSION

The processes of oxidation and reduction occur at different
areas of the metal surface and are accompanied by the appearance
electric current.
With unequal access of oxygen
To
surfaces
metal
on
her
arise
galvanic
pair
special
kind:
plot
more
adsorbing oxygen is
cathode, and less adsorbing
anode.
Due to
spherical
flattened water drops circular
the zone under its edges will be the cathode,
and under the central part - the anode.
Anode (-) Fe (center) / O2, H2O, NaCl / Fe (edge) (+) Cathode
A: Fe2+ + K3 = 3K+ + Fe2+
Potassium hexacyanoferrate(III)
"Turnbull blue"

PASSIVATION OF METALS DURING CORROSION

Sometimes the corrosion rate can be limited by the anodic
process. This is typical for metals that can
passivate (Cr, Al, Ti, Ni, Zr, Ta, etc.)
The passivity of a metal is called
his state of heightened
corrosion resistance,
caused by inhibition
anodic process.
Passivation is associated with the formation on the surface
metal of adsorbed or phase layers (sometimes those and
others), which inhibit the process of metal dissolution.
Strong oxidizing agents usually promote passivation
metal

∆G = -nFE< 0
E= Ϥ0Ох - Ϥ0Red > 0
Ϥ0Ох > Ϥ0Red
Ϥ0Ох > Ϥ0Me+n/Me0

CONDITIONS FOR CORROSION WITH OXYGEN AND HYDROGEN DEPOLARIZATION

If E0Me+n/Me0< E0Н+/Н2 меньше потенциала водородного
electrode (area 1), corrosion is possible and with absorption
oxygen, and with the release of hydrogen (alkali and alkaline earth metals, zinc, aluminum)
If E0Me+n/Me0 is less than the potential of the oxygen electrode,
but greater than the potential of the hydrogen electrode (region 2),
then corrosion is possible only with the participation of oxygen.
_
E0H+/H2< E0Me+n/Me0 < E0О2/ОН-
A (-): Me0 - ne → Me+n
_
K (+): O2 + 2H2O + 4e → 4OH-

CONDITIONS FOR CORROSION WITH OXYGEN AND HYDROGEN DEPOLARIZATION

If E0Me+n/Me0 > E0O2/OH- potential of the oxygen electrode
(area 3), then metal corrosion is impossible.
Example: gold – in the absence of a complexing agent it will not
corrodes with oxygen absorption or release
hydrogen.
The potentials of many metals lie in the second region.

Methods for protecting metals from corrosion

METHODS FOR PROTECTING METALS FROM CORROSION

All protection methods are conditionally divided into the following groups:
1. Alloying of metals;
2. Protective coatings (metallic and non-metallic);
3.Electrochemical protection;
4.Change in the properties of the corrosive environment.
Choosing one or another method of corrosion protection
defined:
on the one hand, its effectiveness,
on the other hand, its economic feasibility.

ALLOYING OF METALS

This is a security method associated with
changing the properties of the corrosive
metal Effective though usually
expensive method of protection. At
alloying into the alloy composition is usually
introduce components that cause
metal passivation (introduction
chromium, nickel, tungsten, etc.)
For heat-resistant alloys with alloying
additives include chromium, aluminum,
nickel, silicon - they improve properties
protective films formed when
oxidation of metals.

PROTECTIVE COATINGS. METAL COATINGS

Protective coatings are layers artificially created
on the metal surface to protect against corrosion.
Metal coatings
Coating materials can be
be like pure metals (Zn, Cd,
Al, Ni, Cu, Cr, Ag), and their alloys
(bronze, brass).
By the nature of their behavior
coatings are divided
to cathode and anode

CATHOdic METAL COATINGS

Cathodic coatings include coatings, electrode
whose potentials in a given environment have more
positive value than the potential of the base metal.
For steel (Fe), the cathode coating will be copper,
nickel, silver.
Anode (-) Fe / O2, H2O, NaCl / Cu (+) Cathode
_
A: Fe 0 - 2e → Fe 2+
_
K: O2 + 2H2O + 4e → 4OH-
If the cathode is damaged
coverage occurs
oxygen depolarization
galvanic cell in
_
+
0
2Н + 2е → Н2
which oxidation occurs
hydrogen depolarization
main material.
Therefore, cathodic coating can protect products
only in the absence of pores and cracks, i.e. when not violated
integrity of the coating.

ANODIC METAL COATINGS

Anodic coatings have a more negative potential,
than the potential of the base metal.
For example: coating
steel (Fe) with zinc – the base metal in this case will be
cathode and will not corrode.
Anode (-) Zn / O2, H2O, NaCl / Fe (+) Cathode
_
A: Zn 0 - 2e → Zn 2+
_
K: O2 + 2H2O + 4e → 4OH oxygen depolarization
_
2Н+ + 2е → Н20
hydrogen depolarization
At
damage
anode
coatings
arises
galvanic cell in which oxidation occurs
coating, and the base material remains unchanged until
complete dissolution of the coating.

METHODS FOR PRODUCING METAL PROTECTIVE COATINGS

Electrochemical (electroplating).
Metallization (Immersion in molten metal).
Thermal diffusion
way
(For
receiving
heat-resistant coatings: Al – aluminizing, Si –
siliconization, Cr – chrome plating, Ti – titanation).
At elevated temperature the product is immersed in
powder
metal,
which
is
coating.
Diffusion of the applied metal occurs into the base
metal.
Chemical.
Product
place
V
solution,
containing coating metal ions and a reducing agent. IN
result
redox
reactions
metal ions are being reduced to free
metal Thus, it is applied to metals and non-metals
coatings of silver, copper, nickel and palladium.

NON-METALLIC COATINGS

The protective properties of non-metallic coatings boil down to
isolation of metal from the environment.
Such coatings can be
be:
inorganic enamels,
paint coatings,
resin coatings,
plastics,
polymer films,
rubber.

ELECTROCHEMICAL PROTECTION. PROTECTORS

The method is based on the inhibition of anodic or cathodic
corrosion processes.
Cathodic protection – the product is connected to (-) external
current source, it becomes the cathode, and the anode usually
An auxiliary electrode (usually steel) serves as an auxiliary electrode.
If the auxiliary anode is made of metal,
having
more
negative
potential,
how
the metal being protected, then the current is not connected. IN
the resulting galvanic cell dissolves the anode, and
the product is not subject to corrosion.
Such electrodes are called protectors (magnesium and its
alloys, zinc, aluminum).
Anodic protection - consists of creating anodic polarization
due to externally applied current (stainless steel protection
steel in sulfuric acid).

CHANGES IN THE PROPERTIES OF A CORROSIVE ENVIRONMENT

WITH
purpose
reduction
corrosive
activity
environment
carry out its processing.
For example:
deletion
oxygen
(boiling
solution;
bubbling with inert gas; restoring it using
corresponding reducing agents - sulfites, hydrazine);
decrease in the concentration of H+ ions
- alkalization
solution
IN last years for corrosion protection widely
inhibitors are used.

INHIBITORS

Inhibitor – a substance that reduces speed is called
corrosion.
Inhibitors are used in systems that work with
constant or slightly changing volume.
The inhibitory effect is most clearly expressed in
the following types of compounds: amines, nitrogen-containing
heterocyclic compounds, sulfides, aldehydes,
mercaptans.
According to the conditions of use, inhibitors are divided into:
inhibitors for aqueous solutions (acid, alkaline
and for neutral environments);
“volatile inhibitors” - for protection from atmospheric
corrosion (compounds of amines with nitrous, carbon or
chromic acids.)

MECHANISM OF ACTION OF INHIBITORS

The mechanism of action of inhibitors is adsorption
them on the corroding surface and subsequent
inhibition of cathodic and anodic processes.
Anodic inhibitors – oxidizers (NO2-, NO3-, CrO42-, PO43-).
At the same time, the metal becomes stable passive
state.
Cathode inhibitors - reduce the speed of cathode
process or reduce the area of the cathode sections.
Organic substances containing sulfur, nitrogen and
oxygen (diethylamine, methenamine, hydrazine).

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