Li, K, Ca, Na, Mg, Al, Zn, Cr, Fe, Pb, H 2 , Cu, Ag, Hg, Au
The further to the left the metal is in the series of standard electrode potentials, the stronger the reducing agent it is, the strongest reducing agent is metallic lithium, gold is the weakest, and, conversely, the gold (III) ion is the strongest oxidizing agent, lithium (I) is the weakest .
Each metal is able to restore from salts in solution those metals that are in a series of voltages after it, for example, iron can displace copper from solutions of its salts. However, it should be remembered that alkali and alkaline earth metals will interact directly with water.
Metals, standing in the series of voltages to the left of hydrogen, are able to displace it from solutions of dilute acids, while dissolving in them.
The reducing activity of a metal does not always correspond to its position in the periodic system, because when determining the place of a metal in a series, not only its ability to donate electrons is taken into account, but also the energy expended on the destruction of the metal crystal lattice, as well as the energy expended on the hydration of ions.
Interaction with simple substances
FROM oxygen most metals form oxides - amphoteric and basic:
4Li + O 2 \u003d 2Li 2 O,
4Al + 3O 2 \u003d 2Al 2 O 3.
Alkali metals, with the exception of lithium, form peroxides:
2Na + O 2 \u003d Na 2 O 2.
FROM halogens metals form salts of hydrohalic acids, for example,
Cu + Cl 2 \u003d CuCl 2.
FROM hydrogen the most active metals form ionic hydrides - salt-like substances in which hydrogen has an oxidation state of -1.
2Na + H 2 = 2NaH.
FROM gray metals form sulfides - salts of hydrosulfide acid:
FROM nitrogen some metals form nitrides, the reaction almost always proceeds when heated:
3Mg + N 2 \u003d Mg 3 N 2.
FROM carbon carbides are formed.
4Al + 3C \u003d Al 3 C 4.
FROM phosphorus - phosphides:
3Ca + 2P = Ca 3 P 2 .
Metals can interact with each other to form intermetallic compounds :
2Na + Sb = Na 2 Sb,
3Cu + Au = Cu 3 Au.
Metals can dissolve in each other at high temperature without interaction, forming alloys.
Alloys
Alloys are called systems consisting of two or more metals, as well as metals and non-metals that have characteristic properties inherent only in the metallic state.
The properties of alloys are very diverse and differ from the properties of their components, for example, in order to make gold harder and more suitable for making jewelry, silver is added to it, and an alloy containing 40% cadmium and 60% bismuth has a melting point of 144 °С, i.e. much lower than the melting point of its components (Cd 321 °С, Bi 271 °С).
The following types of alloys are possible:
Molten metals are mixed with each other in any ratio, dissolving in each other without limit, for example, Ag-Au, Ag-Cu, Cu-Ni and others. These alloys are homogeneous in composition, have high chemical resistance, conduct electricity;
The straightened metals are mixed with each other in any ratio, however, when cooled, they delaminate, and a mass is obtained, consisting of individual crystals of components, for example, Pb-Sn, Bi-Cd, Ag-Pb and others.
Objective: get acquainted by experience with the dependence of the redox properties of metals on their position in the electrochemical series of voltages.
Equipment and reagents: test tubes, test tube holders, spirit lamp, filter paper, pipettes, 2n. solutions HCl and H2SO4, concentrated H2SO4, diluted and concentrated HNO3, 0.5M solutions CuSO 4 , Pb(NO 3) 2 or Pb(CH 3 COO) 2; pieces of metal aluminum, zinc, iron, copper, tin, iron paper clips, distilled water.
Theoretical explanations
The chemical character of any metal is largely determined by how easily it oxidizes, i.e. how easily its atoms are able to pass into the state of positive ions.
Metals that exhibit an easy ability to oxidize are called base metals. Metals that oxidize with great difficulty are called noble metals.
Each metal is characterized by a certain value of the standard electrode potential. For standard capacity j0 of this metal electrode, the EMF of a galvanic cell is taken, composed of a standard hydrogen electrode located on the left, and a metal plate placed in a solution of a salt of this metal, and the activity (in dilute solutions, you can use the concentration) of metal cations in the solution should be equal to 1 mol/l; T=298 K; p=1 atm.(standard conditions). If the reaction conditions differ from the standard ones, the dependence of the electrode potentials on the concentrations (more precisely, activities) of metal ions in solution and temperature must be taken into account.
The dependence of electrode potentials on concentration is expressed by the Nernst equation, which, as applied to the system:
Me n + + n e -→Me
AT;
R is the gas constant, 
;
F- Faraday's constant ("96500 C/mol);
n-
a Me n + - mol/l.
Taking the value T=298TO, we get
mol/l.
j 0 , corresponding to the reduction half-reaction, a series of metal voltages is obtained (a series of standard electrode potentials). The standard electrode potential of hydrogen, taken as zero, is placed in the same row for the system in which the process takes place:
2H + + 2e - \u003d H 2
In this case, the standard electrode potentials of non-noble metals have a negative value, and noble ones - positive.
Electrochemical series of voltages of metals
Li; K; Ba; Sr; Ca; Na; Mg; Al; Mn; Zn; Cr; Fe; CD; Co; Ni; sn; Pb; ( h) ; Sb; Bi; Cu; Hg; Ag; Pd; Pt; Au
This series characterizes the redox ability of the "metal - metal ion" system in aqueous solutions under standard conditions. The further to the left in the series of stresses is the metal (the less it j0), the stronger the reducing agent it is, and the easier it is for metal atoms to give up electrons, turning into cations, but the cations of this metal are more difficult to attach electrons, turning into neutral atoms.
Redox reactions involving metals and their cations go in the direction in which a metal with a lower electrode potential is a reducing agent (i.e., is oxidized), and metal cations with a high electrode potential are oxidizers (i.e., are reduced). In this regard, the following regularities are typical for the electrochemical series of voltages of metals:
1. each metal displaces from the salt solution all other metals to the right of it in the electrochemical series of metal voltages.
2. all metals that are to the left of hydrogen in the electrochemical series of voltages displace hydrogen from dilute acids.
Experimental methodology
Experience 1: Interaction of metals with hydrochloric acid.
Pour 2-3 into four test tubes ml hydrochloric acid and place in them a piece of aluminum, zinc, iron and copper separately. Which of the given metals displaces hydrogen from the acid? Write reaction equations.
Experience 2: Interaction of metals with sulfuric acid.
Drop a piece of iron into a test tube and add 1 ml 2n. sulfuric acid. What is observed? Repeat the experiment with a piece of copper. Does the reaction take place?
Check the effect of concentrated sulfuric acid on iron and copper. Explain observations. Write all reaction equations.
Experience 3: Interaction of copper with nitric acid.
Put a piece of copper into two test tubes. Pour 2 into one of them ml dilute nitric acid, the second - concentrated. If necessary, warm the contents of the test tubes on an alcohol lamp. What gas is formed in the first test tube, and what in the second? Write the reaction equations.
Experience 4: Interaction of metals with salts.
Pour into test tube 2 – 3 ml copper (II) sulfate solution and lower a piece of iron wire. What's happening? Repeat the experiment, replacing the iron wire with a piece of zinc. Write reaction equations. Pour into test tube 2 ml a solution of acetate or lead (II) nitrate and lower a piece of zinc. What's happening? Write the reaction equation. Specify the oxidizing agent and reducing agent. Will the reaction proceed if zinc is replaced with copper? Give an explanation.
11.3 Required level of student preparation
1. Know the concept of standard electrode potential, have an idea about its measurement.
2. Be able to use the Nernst equation to determine the electrode potential under conditions other than standard.
3. Know what a series of metal stresses is, what it characterizes.
4. Be able to use a number of voltages of metals to determine the direction of redox reactions involving metals and their cations, as well as metals and acids.
Tasks for self-control
1. What is the mass of technical iron containing 18% impurities required to displace nickel sulfate from solution (II) 7.42 g nickel?
2. A copper plate with a mass of 28 g. at the end of the reaction, the plate was taken out, washed, dried and weighed. Its mass turned out 32.52 g. What mass of silver nitrate was in the solution?
3. Determine the value of the electrode potential of copper immersed in 0.0005 M copper nitrate solution (II).
4. Electrode potential of zinc immersed in 0.2 M solution ZnSO4, is equal to 0.8V. determine the apparent degree of dissociation ZnSO4 in a solution of the specified concentration.
5. Calculate the potential of the hydrogen electrode if the concentration of hydrogen ions in the solution (H+) is 3.8 10 -3 mol/l.
6. Calculate the potential of an iron electrode immersed in a solution containing 0.0699 g FeCI 2 in 0.5 l.
7. What is called the standard electrode potential of the metal? What equation expresses the dependence of electrode potentials on concentration?
Lab #12
Subject: Electroplating Cell
Objective: familiarization by experience with the principles of operation of a galvanic cell, mastering the calculation methodology EMF galvanic elements.
Equipment and reagents: copper and zinc plates attached to conductors, copper and zinc plates connected by conductors to copper plates, sandpaper, voltmeter, 3 chemical beakers 200-250 ml, measuring cylinder, tripod with a U - shaped tube fixed in it, salt bridge, 0.1 M solutions of copper sulfate, zinc sulfate, sodium sulfate, 0,1 % phenolphthalein solution in 50% ethyl alcohol.
Theoretical explanations
A galvanic cell is a chemical current source, that is, a device that generates electrical energy as a result of direct conversion chemical energy of the redox reaction.
Electric current (directed movement of charged particles) is transmitted through current conductors, which are divided into conductors of the first and second kind.
Conductors of the first kind conduct electric current with their electrons (electronic conductors). These include all metals and their alloys, graphite, coal, and some solid oxides. The electrical conductivity of these conductors is in the range from 10 2 to 10 6 Ohm -1 cm -1 (for example, coal - 200 Ohm -1 cm -1, silver 6 10 5 Ohm -1 cm -1).
Conductors of the second kind conduct electric current with their ions (ionic conductors). They are characterized by low electrical conductivity (for example, H 2 O - 4 10 -8 Ohm -1 cm -1).
When the conductors of the first and second kind are combined, an electrode is formed. This is most often a metal dipped in a solution of its own salt.
When a metal plate is immersed in water, the metal atoms in its surface layer are hydrated under the action of polar water molecules. As a result of hydration and thermal motion, their bond with the crystal lattice is weakened and a certain number of atoms pass in the form of hydrated ions into the liquid layer adjacent to the metal surface. The metal plate becomes negatively charged.
Me + m H 2 O \u003d Me n + n H 2 O + ne -
Where Me is a metal atom; Me n + n H 2 O is a hydrated metal ion; e-- electron, n is the charge of the metal ion.
The state of equilibrium depends on the activity of the metal and on the concentration of its ions in solution. In the case of active metals ( Zn, Fe, Cd, Ni), the interaction with polar water molecules ends with the detachment of positive metal ions from the surface and the transition of hydrated ions into solution (Fig. 1 a). This process is oxidative. As the concentration of cations near the surface increases, the rate of the reverse process, the reduction of metal ions, increases. Ultimately, the rates of both processes are equalized, an equilibrium is established, in which a double electric layer with a certain value of the metal potential appears at the solution-metal interface.
| + + + + |
| – – – – |
Zn 0 + mH 2 O → Zn 2+ mH 2 O+2e - + + – – Cu2+ nH 2 O + 2e - → Cu 0 + nH 2 O
+ + + – – –
Rice. 1. Scheme of the appearance of the electrode potential
When a metal is immersed not in water, but in a solution of a salt of this metal, the equilibrium shifts to the left, that is, in the direction of the transition of ions from the solution to the surface of the metal. In this case, a new equilibrium is established already at a different value of the potential of the metal.
For inactive metals, the equilibrium concentration of metal ions in pure water is very low. If such a metal is immersed in a solution of its salt, then the metal cations will be released from the solution at a higher rate than the rate of transition of ions from the metal to the solution. In this case, the metal surface will receive a positive charge, and the solution will receive a negative charge due to an excess of salt anions (Fig. 1. b).
Thus, when a metal is immersed in water or in a solution containing ions of this metal, a double electric layer is formed on the metal-solution interface, which has a certain potential difference. The electrode potential depends on the nature of the metal, the concentration of its ions in the solution, and the temperature.
The absolute value of the electrode potential j individual electrode cannot be determined experimentally. However, it is possible to measure the potential difference of two chemically different electrodes.
We agreed to take the potential of a standard hydrogen electrode equal to zero. The standard hydrogen electrode is a platinum plate coated with spongy platinum, immersed in an acid solution with a hydrogen ion activity of 1 mol/l. The electrode is washed with gaseous hydrogen at a pressure of 1 atm. and temperature 298 K. This establishes an equilibrium:
2 H + + 2 e \u003d H 2
For standard capacity j0 of this metal electrode is taken EMF a galvanic cell composed of a standard hydrogen electrode and a metal plate placed in a salt solution of this metal, and the activity (in dilute solutions, you can use the concentration) of metal cations in the solution should be equal to 1 mol/l; T=298 K; p=1 atm.(standard conditions). The value of the standard electrode potential is always referred to as the reduction half-reaction:
Me n + +n e - → Me
Arranging metals in ascending order of their standard electrode potentials j 0 , corresponding to the reduction half-reaction, a series of metal voltages is obtained (a series of standard electrode potentials). The standard electrode potential of the system, taken as zero, is placed in the same row:
H + + 2e - → H 2
The dependence of the electrode potential of the metal j on temperature and concentration (activity) is determined by the Nernst equation, which, as applied to the system:
Me n + + n e -→Me
It can be written in the following form:
where is the standard electrode potential, AT;
R is the gas constant, ;
F- Faraday's constant ("96500 C/mol);
n- the number of electrons involved in the process;
a Me n + - activity of metal ions in solution, mol/l.
Taking the value T=298TO, we get
moreover, the activity in dilute solutions can be replaced by the concentration of ions, expressed in terms of mol/l.
EMF any galvanic cell can be defined as the difference between the electrode potentials of the cathode and anode:
EMF = j cathode -j anode
The negative pole of the element is called the anode, the oxidation process takes place on it:
Me - ne - → Me n +
The positive pole is called the cathode, it is undergoing a recovery process:
Me + + ne - → Me
A galvanic cell can be written schematically, while following certain rules:
1. The electrode on the left should be written in the metal-ion sequence. The electrode on the right is written in the sequence ion - metal. (-) Zn/Zn 2+ //Cu 2+ /Cu (+)
2. The reaction occurring at the left electrode is recorded as an oxidative one, and the reaction at the right electrode as a reduction one.
3. If EMF element > 0, then the work of the galvanic cell will be spontaneous. If a EMF< 0, то самопроизвольно будет работать обратный гальванический элемент.
Experiment methodology
Experience 1: Compilation of copper-zinc cell
Get the necessary equipment and reagents from the laboratory assistant. In a chemical beaker 200 ml pour 100 ml 0.1 M copper sulfate solution (II) and lower into it a copper plate connected to a conductor. Pour the same volume into the second glass 0.1 M zinc sulfate solution and lower the zinc plate connected to the conductor into it. The plates must be pre-cleaned with sandpaper. Obtain a salt bridge from the laboratory assistant and connect two electrolytes with it. The salt bridge is a glass tube filled with gel (agar-agar), both ends of which are closed with a cotton swab. The bridge is kept in a saturated aqueous solution of sodium sulfate, as a result of which the gel swells and exhibits ionic conductivity.
With the help of a teacher, connect a voltmeter to the poles of the resulting galvanic cell and measure the voltage (if the measurement is carried out with a voltmeter with a small resistance, then the difference between the value EMF and stress is low). Using the Nernst equation, calculate the theoretical value EMF galvanic element. Voltage less EMF galvanic cell due to the polarization of the electrodes and ohmic losses.
Experience 2: Electrolysis of sodium sulfate solution
In the experiment, due to the electrical energy generated by the galvanic cell, it is proposed to carry out the electrolysis of sodium sulfate. To do this, pour a solution of sodium sulfate into the U - shaped tube and place copper plates in both knees of it, cleaned with sandpaper and connected to the copper and zinc electrodes of the galvanic cell, as shown in Fig. 2. Add 2-3 drops of phenolphthalein to each elbow of the U-tube. After some time, in the cathode space of the electrolyzer, the solution is colored pink due to the formation of alkali during the cathodic reduction of water. This indicates that the galvanic cell works as a current source.
Make up the equations of the processes occurring at the cathode and at the anode during the electrolysis of an aqueous solution of sodium sulfate.
(-) CATHODE ANODE (+)
salt bridge
→ Zn2+ Cu2+→
ZnSO4CuSO4
ANODE (-) CATHODE (+)
Zn - 2e - → Zn 2+ Cu 2+ + 2e - → Cu
oxidation reduction
12.3 Required level of student preparation
1. Know the concepts: conductors of the first and second kind, dielectrics, electrode, galvanic cell, anode and cathode of a galvanic cell, electrode potential, standard electrode potential. EMF galvanic element.
2. Have an idea about the causes of the occurrence of electrode potentials and methods for their measurement.
3. Have an idea about the principles of operation of a galvanic cell.
4. Be able to use the Nernst equation to calculate electrode potentials.
5. Be able to write circuits of galvanic cells, be able to calculate EMF galvanic elements.
Tasks for self-control
1. Describe conductors and dielectrics.
2. Why does an anode have a negative charge in a galvanic cell, and a positive charge in an electrolytic cell?
3. What is the difference and similarity between cathodes in an electrolyzer and a galvanic cell?
4. A magnesium plate was lowered into a solution of its salt. In this case, the electrode potential of magnesium turned out to be equal to -2.41V. Calculate the concentration of magnesium ions in mol/l. (4.17x10 -2).
5. At what concentration of ions Zn 2+ (mol/l) the potential of the zinc electrode will become 0.015 V smaller than its standard electrode? (0.3 mol/l)
6. Nickel and cobalt electrodes are lowered into solutions, respectively. Ni(NO 3) 2 and Co(NO 3) 2. In what ratio should the concentration of these metal ions be in order for the potentials of both electrodes to be the same? (C Ni 2+ :C Co 2+ = 1:0.117).
7. At what concentration of ions Cu2+ in mol/l the value of the potential of the copper electrode becomes equal to the standard potential of the hydrogen electrode? (1.89x 10 -6 mol/l).
8. Draw a diagram, write the electronic equations of electrode processes and calculate EMF a galvanic cell consisting of plates of cadmium and magnesium, lowered into solutions of their salts with a concentration = = 1.0 mol/l. Will the value change EMF if the concentration of each ion is reduced to 0.01 mol/l? (2.244 V).
Lab #13
Electrochemical activity series of metals (voltage range, a range of standard electrode potentials) - the sequence in which the metals are arranged in order of increasing their standard electrochemical potentials φ 0 corresponding to the metal cation reduction half-reaction Me n+ : Me n+ + nē → Me
A number of stresses characterize the comparative activity of metals in redox reactions in aqueous solutions.
Story
The sequence of the arrangement of metals in the order of change in their chemical activity in general terms was already known to alchemists. The processes of mutual displacement of metals from solutions and their surface precipitation (for example, the displacement of silver and copper from solutions of their salts by iron) were considered as a manifestation of the transmutation of elements.
Later alchemists came close to understanding the chemical side of the mutual precipitation of metals from their solutions. So, Angelus Sala in his work "Anatomy Vitrioli" (1613) came to the conclusion that the products of chemical reactions consist of the same "components" that were contained in the original substances. Subsequently, Robert Boyle proposed a hypothesis about the reasons why one metal displaces another from solution, based on corpuscular representations.
In the era of the formation of classical chemistry, the ability of elements to displace each other from compounds became an important aspect of understanding reactivity. J. Berzelius, on the basis of the electrochemical theory of affinity, built a classification of elements, dividing them into "metalloids" (now the term "non-metals" is used) and "metals" and putting hydrogen between them.
The sequence of metals according to their ability to displace each other, long known to chemists, was especially thoroughly and comprehensively studied and supplemented by N. N. Beketov in the 1860s and subsequent years. Already in 1859, he made a report in Paris on the topic "Research on the phenomena of the displacement of some elements by others." In this work, Beketov included whole line generalizations about the relationship between the mutual displacement of elements and their atomic weight, linking these processes with " the original chemical properties of the elements - what is called chemical affinity» . Beketov's discovery of the displacement of metals from solutions of their salts by hydrogen under pressure and the study of the reducing activity of aluminum, magnesium and zinc at high temperatures (metallothermy) allowed him to put forward a hypothesis about the relationship between the ability of some elements to displace others from compounds with their density: lighter simple substances are able to displace more heavy (therefore, this series is often also called Beketov displacement series, or simply Beketov series).
Without denying the significant merits of Beketov in the development of modern ideas about the number of activity of metals, it should be considered erroneous that prevails in the domestic popular and educational literature the idea of him as the sole creator of this series. Numerous experimental data obtained at the end of the 19th century disproved Beketov's hypothesis. Thus, William Odling described many cases of "activity reversal". For example, copper displaces tin from a concentrated acidified solution of SnCl 2 and lead from acid solution PbCl 2 ; it is also capable of dissolving in concentrated hydrochloric acid with the release of hydrogen. Copper, tin and lead are in the row to the right of cadmium, however, they can displace it from a boiling slightly acidified CdCl 2 solution.
The rapid development of theoretical and experimental physical chemistry pointed to another reason for the differences in the chemical activity of metals. With the development of modern concepts of electrochemistry (mainly in the works of Walter Nernst), it became clear that this sequence corresponds to a "series of voltages" - the arrangement of metals according to the value of standard electrode potentials. Thus, instead of a qualitative characteristic - the “tendency” of a metal and its ion to certain reactions - Nerst introduced an exact quantitative value characterizing the ability of each metal to pass into solution in the form of ions, and also to be reduced from ions to metal on the electrode, and the corresponding series was named a number of standard electrode potentials.
Theoretical basis
The values of electrochemical potentials are a function of many variables and therefore show a complex dependence on the position of metals in the periodic system. Thus, the oxidation potential of cations increases with an increase in the atomization energy of a metal, with an increase in the total ionization potential of its atoms, and with a decrease in the hydration energy of its cations.
In the most general form, it is clear that metals at the beginning of periods are characterized by low values of electrochemical potentials and occupy places on the left side of the voltage series. At the same time, the alternation of alkali and alkaline earth metals reflects the phenomenon of diagonal similarity. Metals located closer to the middle of the periods are characterized by large potential values and occupy places in the right half of the series. A consistent increase in the electrochemical potential (from -3.395 V for a pair of Eu 2+ /Eu [ ] to +1.691 V for the Au + /Au pair) reflects a decrease in the reducing activity of metals (the ability to donate electrons) and an increase in the oxidizing ability of their cations (the ability to attach electrons). Thus, the strongest reducing agent is europium metal, and the strongest oxidizing agent is gold cations Au+.
Hydrogen is traditionally included in the voltage series, since the practical measurement of the electrochemical potentials of metals is carried out using a standard hydrogen electrode.
Practical use of a range of voltages
A number of voltages are used in practice for a comparative [relative] assessment of the chemical activity of metals in reactions with aqueous solutions of salts and acids and for assessing cathodic and anodic processes during electrolysis:
- Metals to the left of hydrogen are stronger reducing agents than metals to the right: they displace the latter from salt solutions. For example, the interaction Zn + Cu 2+ → Zn 2+ + Cu is possible only in the forward direction.
- Metals in the row to the left of hydrogen displace hydrogen when interacting with aqueous solutions of non-oxidizing acids; the most active metals (up to and including aluminum) - and when interacting with water.
- Metals in the row to the right of hydrogen do not interact with aqueous solutions of non-oxidizing acids under normal conditions.
- During electrolysis, metals to the right of hydrogen are released at the cathode; the reduction of metals of moderate activity is accompanied by the release of hydrogen; the most active metals (up to aluminum) cannot be isolated from aqueous solutions of salts under normal conditions.
Table of electrochemical potentials of metals
| Metal | Cation | φ 0 , V | Reactivity | Electrolysis (at the cathode): |
|---|---|---|---|---|
| Li + | -3,0401 | reacts with water | hydrogen is released | |
| Cs + | -3,026 | |||
| Rb+ | -2,98 | |||
| K+ | -2,931 | |||
| F+ | -2,92 | |||
| Ra2+ | -2,912 | |||
| Ba 2+ | -2,905 | |||
| Sr2+ | -2,899 | |||
| Ca2+ | -2,868 | |||
| EU 2+ | -2,812 | |||
| Na+ | -2,71 | |||
| Sm 2+ | -2,68 | |||
| Md2+ | -2,40 | reacts with aqueous solutions of acids | ||
| La 3+ | -2,379 | |||
| Y 3+ | -2,372 | |||
| Mg2+ | -2,372 | |||
| Ce 3+ | -2,336 | |||
| Pr 3+ | -2,353 | |||
| Nd 3+ | -2,323 | |||
| Er 3+ | -2,331 | |||
| Ho 3+ | -2,33 | |||
| Tm3+ | -2,319 | |||
| Sm 3+ | -2,304 | |||
| Pm 3+ | -2,30 | |||
| Fm 2+ | -2,30 | |||
| Dy 3+ | -2,295 | |||
| Lu 3+ | -2,28 | |||
| Tb 3+ | -2,28 | |||
| Gd 3+ | -2,279 | |||
| Es 2+ | -2,23 | |||
| AC 3+ | -2,20 | |||
| Dy 2+ | -2,2 | |||
| Pm 2+ | -2,2 | |||
| cf2+ | -2,12 | |||
| Sc 3+ | -2,077 | |||
| Am 3+ | -2,048 | |||
| cm 3+ | -2,04 | |||
| Pu3+ | -2,031 | |||
| Er 2+ | -2,0 | |||
| Pr 2+ | -2,0 | |||
| EU 3+ | -1,991 | |||
| Lr 3+ | -1,96 | |||
| cf 3+ | -1,94 | |||
| Es 3+ | -1,91 | |||
| Th4+ | -1,899 | |||
| Fm 3+ | -1,89 | |||
| Np 3+ | -1,856 | |||
| Be 2+ | -1,847 | |||
| U 3+ | -1,798 | |||
| Al 3+ | -1,700 | |||
| Md 3+ | -1,65 | |||
| Ti 2+ | -1,63 | competing reactions: both hydrogen evolution and pure metal evolution | ||
| hf 4+ | -1,55 | |||
| Zr4+ | -1,53 | |||
| Pa 3+ | -1,34 | |||
| Ti 3+ | -1,208 | |||
| Yb 3+ | -1,205 | |||
| no 3+ | -1,20 | |||
| Ti 4+ | -1,19 | |||
| Mn2+ | -1,185 | |||
| V2+ | -1,175 | |||
| Nb 3+ | -1,1 | |||
| Nb 5+ | -0,96 | |||
| V 3+ | -0,87 | |||
| Cr2+ | -0,852 | |||
| Zn2+ | -0,763 | |||
| Cr3+ | -0,74 | |||
| Ga3+ | -0,560 | |||
Metals that react easily are called active metals. These include alkali, alkaline earth metals and aluminium.
Position in the periodic table
The metallic properties of the elements weaken from left to right in Mendeleev's periodic table. Therefore, elements of groups I and II are considered the most active.
Rice. 1. Active metals in the periodic table.
All metals are reducing agents and easily part with electrons at the external energy level. Active metals have only one or two valence electrons. In this case, the metallic properties are enhanced from top to bottom with an increase in the number of energy levels, because. the farther an electron is from the nucleus of an atom, the easier it is for it to separate.
Alkali metals are considered the most active:
- lithium;
- sodium;
- potassium;
- rubidium;
- cesium;
- francium.
The alkaline earth metals are:
- beryllium;
- magnesium;
- calcium;
- strontium;
- barium;
- radium.
You can find out the degree of activity of a metal by the electrochemical series of metal voltages. The more to the left of hydrogen an element is located, the more active it is. The metals to the right of hydrogen are inactive and can only interact with concentrated acids.
Rice. 2. Electrochemical series of voltages of metals.
The list of active metals in chemistry also includes aluminum, located in group III and to the left of hydrogen. However, aluminum is located on the border of active and medium active metals and does not react with certain substances under normal conditions.
Properties
Active metals are soft (can be cut with a knife), light, and have a low melting point.
The main chemical properties of metals are presented in the table.
|
Reaction |
The equation |
Exception |
|
Alkali metals ignite spontaneously in air, interacting with oxygen |
K + O 2 → KO 2 |
Lithium reacts with oxygen only at high temperatures. |
|
Alkaline earth metals and aluminum form oxide films in air, and spontaneously ignite when heated. |
2Ca + O 2 → 2CaO |
|
|
React with simple substances to form salts |
Ca + Br 2 → CaBr 2; |
Aluminum does not react with hydrogen |
|
React violently with water, forming alkalis and hydrogen |
|
The reaction with lithium proceeds slowly. Aluminum reacts with water only after the removal of the oxide film. |
|
React with acids to form salts |
Ca + 2HCl → CaCl 2 + H 2; 2K + 2HMnO 4 → 2KMnO 4 + H 2 |
|
|
React with salt solutions, first reacting with water and then with salt |
2Na + CuCl 2 + 2H 2 O: 2Na + 2H 2 O → 2NaOH + H 2; |
Active metals easily react, therefore, in nature they are found only in mixtures - minerals, rocks.
Rice. 3. Minerals and pure metals.
What have we learned?
Active metals include elements of groups I and II - alkali and alkaline earth metals, as well as aluminum. Their activity is due to the structure of the atom - a few electrons are easily separated from the external energy level. These are soft light metals that quickly react with simple and complex substances, forming oxides, hydroxides, salts. Aluminum is closer to hydrogen and requires additional conditions for its reaction with substances - high temperatures, the destruction of the oxide film.
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