Alkali metals are a class of chemical elements found in Periodic Table Group 1. Lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr) are all members of this group. Alkali metals are extremely reactive and have unique features that distinguish them from other elements.
One distinguishing feature of alkali metals is their tendency to generate alkaline solutions when they react with water. This is why they are referred to as “alkali” metals. When alkali metals come into contact with water, they undergo a strong reaction that results in the formation of alkaline solutions by releasing hydrogen gas (H+) and producing hydroxide ions (OH-).
Chemistry of Alkali Metals
Alkali metals have a wide range of uses due to their strong reactivity and distinctive characteristics. They are used to make batteries, such as lithium-ion batteries, which power numerous modern gadgets. Alkali metals are also important in organic chemistry as powerful reducing agents and catalysts. They are also used in the production of pyrotechnics, glass, soaps, and fertilizers.
Here, let us delve more into the occurrence and extraction of alkali metals, as well as their characteristics, reactivity, and uses. We may obtain insights into the behavior of elements in the main group and grasp their relevance in numerous scientific and industrial sectors by knowing the chemistry of alkali metals.
Alkali Metal Occurrence
Alkali metals are rather plentiful in nature, albeit they are not found in their pure form. They are found in a variety of minerals and salts, typically as ionic complexes.
The most frequent alkali metal-containing minerals are:
Lithium: Spodumene (LiAlSi2O6) and lepidolite (KLi2Al(Al,Si)3O10(F,OH)2)
Potassium minerals include sylvite (KCl), carnallite (KMgCl36H2O), and langbeinite (K2Mg2(SO4)3).
Rubidium and cesium: These elements are found in trace amounts in minerals such as lepidolite and pollucite.
Alkali Metal Extraction
Alkali metals are generally extracted from their ores or salts using a variety of extraction processes. The extraction method used is determined on the kind of alkali metal and its source.
Extraction of Lithium
Lithium is usually recovered from lithium-bearing minerals such as spodumene and lepidolite. Several stages are involved in the extraction process:
1. Crushing and grinding of the ore: The ore is crushed and ground into fine particles.
2. Roasting: High temperatures in the presence of oxygen are used to transform lithium minerals into lithium oxide (Li2O).
3. Acid Treatment: To produce lithium sulfate (Li2SO4), the roasted ore is treated with sulfuric acid (H2SO4).
4. Lithium Carbonate Conversion: Lithium sulfate is then converted to lithium carbonate (Li2CO3) by precipitation with sodium carbonate (Na2CO3) or by the addition of soda ash.
5. Purification: To achieve pure lithium carbonate, the lithium carbonate is further refined using methods such as filtering and crystallization.
Extraction of Sodium and Potassium
Sodium and potassium are extremely reactive metals that are often produced through electrolysis of their salts.
The Downs’ Process, which utilizes electrolysis of molten sodium chloride (NaCl) or potassium chloride (KCl) using a Downs’ cell, is the most prevalent technique.
Sodium or potassium ions flow towards the cathode in the Downs’ cell, where reduction occurs and molten metal is created.
During electrolysis, the following reactions occur at the electrodes:
Cathode: 2Na+ + 2e– → 2Na (sodium metal)
Anode: 2Cl– → Cl2 + 2e– (for sodium extraction)
Extraction of Rubidium and Cesium
Rubidium and cesium are very uncommon alkali metals, and their extraction is normally done from ores containing these elements.
– One approach includes extracting rubidium and cesium from pollucite, a mineral with high quantities of both elements. The extraction procedure consists of the following steps:
1. Pollucite ore is crushed and processed into a fine powder.
2. Digestion: Sulfuric acid (H2SO4) is applied to the powdered ore to generate a soluble sulfate complex.
3. Filtration: To eliminate any solid contaminants, the resultant solution is filtered.
4. Ion Exchange: After the filtering stage, the solution is exposed to ion exchange. The solution is passed through a column filled with a resin that preferentially binds and separates the rubidium or cesium ions from other contaminants during ion exchange.
Typically, the resin utilized is a zeolite substance with a strong affinity for rubidium or cesium ions.
The rubidium or cesium ions are exchanged with other ions present on the resin as the solution goes through the column, thereby isolating them.
5. Elution: The rubidium or cesium ions are securely bonded to the resin after the ion exchange procedure.
An eluent solution is run through the column to release and recover the required alkali metal.
Because the eluent solution is more receptive to rubidium or cesium ions than the resin, the alkali metal ions are displaced from the resin and dissolved in the eluent. The required alkali metal ions are now present in the eluent solution.
6. Precipitation and Purification: The alkali metal ions in the eluent solution are then precipitated to produce the solid metal salts.
By adding hydrochloric acid (HCl) to the eluent solution, for example, rubidium chloride (RbCl) or cesium chloride (CsCl) can be created, causing the metal ions to react and form insoluble metal chloride salts.
The precipitate that forms is recovered by filtration and refined further by further washing and drying procedures.
7. Final Refining: Depending on the required purity level, the resulting alkali metal salts may be refined further.
To eliminate any residual impurities and create very pure alkali metals, techniques such as recrystallization, distillation, or fractional crystallization can be used.
Properties of Alkali Metals
General Characteristics of Alkali Metals
Alkali metals have some general properties that set them apart from other elements.
High reactivity, low density, suppleness, and low melting and boiling temperatures are among these properties.
Alkali metals are good heat and electrical conductors. They have a silvery look when freshly cut, but tarnish fast when exposed to air due to their reactivity.
Atomic Number and Symbol
The atomic number and symbol of alkali metals are used to identify them.
Alkali metals have atomic numbers ranging from 3 (lithium) to 87 (francium).
Alkali metal symbols are derived from their English or Latin names, for example, Li for lithium, Na for sodium, K for potassium, Rb for rubidium, Cs for cesium, and Fr for francium.
Alkali metals are classified as Group 1 on the Periodic Table, often known as the alkali metal group or Group IA.
They are found on the periodic table’s far left side.
On the periodic table, alkali metals are found in Period 2 (lithium), Period 3 (sodium), Period 4 (potassium), Period 5 (rubidium), Period 6 (cesium), and Period 7 (francium).
Electronic Configuration IV:
In their outermost energy level (s-orbital), alkali metals have a distinctive electronic arrangement.
They all have one valence electron, which is held loosely due to its location in the lowest energy level.
The electronic configurations of alkali metals follow the pattern [n]s1, where n indicates the energy level.
For example, lithium has an electronic configuration of 1s22s1, sodium has 1s22s22p63s1, and so on.
Oxidation State
Alkali metals rapidly lose their valence electron, resulting in the production of a unipositive ion (cation).
The loss of a valence electron results in the formation of a very stable noble gas electron configuration.
As a result, alkali metals frequently have an oxidation state of +1.
The production of alkali metal ions may be described by the following equation:
M → M+ + e–
where M represents an alkali metal.
Reactivity with Water
Alkali metals are extremely reactive with water, creating alkaline solutions and emitting hydrogen gas. The reaction between alkali metals and water may be expressed by the following equation (using sodium as an example):
2Na + 2H2O → 2NaOH + H2
The reaction results in the formation of sodium hydroxide (NaOH) and hydrogen gas (H2).
Flame Test
When burned in a flame, alkali metals emit various flame colors. The flame test may be used to detect and differentiate various alkali metals.
Lithium, for example, generates a crimson-red flame, sodium a brilliant yellow flame, potassium a violet flame, and so on.
| Element | Atomic Number | Symbol | Electronic Configuration | Oxidation State |
|---|---|---|---|---|
| Lithium | 3 | Li | 1s2 2s1 | +1 |
| Sodium | 11 | Na | 1s2 2s2 2p6 3s1 | +1 |
| Potassium | 19 | K | 1s2 2s2 2p6 3s2 3p6 4s1 | +1 |
| Rubidium | 37 | Rb | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 | +1 |
| Cesium | 55 | Cs | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s1 | +1 |
| Francium | 87 | Fr | 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s1 | +1 |
Physical Properties of Alkali Metals
Alkali metals have numerous common features that set them apart from other elements, including high reactivity, low density, softness, and low melting and boiling temperatures.
Alkali metals are good heat and electrical conductors.
Appearance and State at Room Temperature
When newly cut, alkali metals have a silvery look. However, due to their reactivity, they quickly tarnish when exposed to air, creating oxides or hydroxides on the surface.
The tarnished coating might make things look drab.
The Chemistry of Storing Alkali Metals.
Several storage strategies are used to keep alkali metals from tarnishing or interacting with oxygen, moisture, or other substances. The goal of these strategies is to keep the alkali metals isolated from their surroundings while maintaining a regulated environment. The following are some popular ways for storing alkali metals:
Inert Environment: Alkali metals are frequently kept in containers containing an inert gas, such as argon or nitrogen. These gases keep alkali metals from reacting with oxygen or moisture in the air, which might cause oxidation or corrosion.
Oil Immersion: Alkali metals can also be immersed in a non-reactive oil, such as mineral oil or kerosene. The oil acts as a barrier between the alkali metal and the surrounding environment, preventing direct contact.
Sealed Containers: To prevent exposure to air and moisture, alkali metals are kept in carefully sealed containers. These containers are often built of corrosion-resistant materials, such as glass or certain types of polymers.
Desiccants: To absorb any moisture that may react with the alkali metals, desiccants such as silica gel or molecular sieves may be added to the storage containers. The danger of oxidation and subsequent deterioration of alkali metals is reduced by limiting moisture content.
Low Temperatures: Keeping alkali metals at low temperatures helps to decrease their reactivity. Lithium, for example, is frequently held at temperatures below its melting point to keep it solid and limit its reactivity.
Density, Melting Point, and Boiling Point
In comparison to other metals, alkali metals have low densities.
From lithium to cesium, the densities usually decrease.
Alkali metals have low melting and boiling points that drop as one moves down the group.
For example, lithium has a melting point of 180.5°C and a boiling point of 671°C.
Reactivity with Water and Air
Alkali metals are extremely reactive, especially when exposed to water and air. Alkali metals react with water to form alkaline solutions and release hydrogen gas. The group’s reactivity with water rises, with the reactions becoming more violent.
2Na + 2H2O → 2NaOH + H2
The reaction results in the formation of sodium hydroxide (NaOH) and hydrogen gas (H2).
Another example, Potassium with Water. Potassium (K) is highly reactive and vigorously reacts with water, producing potassium hydroxide (KOH) and hydrogen gas (H2).
2K + 2H2O → 2KOH + H2
Alkali metals react with air as well, generating oxides or hydroxides on their surfaces. The interaction with air causes alkali metals to tarnish.
Chemical Properties of Alkali Metals
Alkali metals are extremely reactive elements found in Periodic Table Group 1. They have unique features such as low ionization energy, a high atomic radius, and a single valence electron in their outermost shell. Alkali metals rapidly lose their valence electron to create +1-charged cations.
Because of their low ionization energy and susceptibility to shed their outermost electron, alkali metals have a high chemical reactivity. They have strong reactions with a variety of elements and compounds, including oxygen, halogens, and water.
Reaction with Oxygen:
Alkali metals have a great attraction for oxygen and rapidly react with it to generate metal oxides. The interaction of alkali metals with oxygen is strongly exothermic, releasing a considerable quantity of heat.
The general equation for the reaction of an alkali metal (M) with oxygen (O2) is as follows:
4M + O2 → 2M2O
The reaction yields alkali metal oxide (M2O).
For example: 4K + O2 → 2K2O
Reaction with Halogens
Alkali metals react violently with halogens (elements from Group 17) to generate alkali metal halides. The reaction between alkali metals with halogens is extremely exothermic.
The following is the general equation for the reaction of an alkali metal (M) with a halogen (X2):
2M + X2 → 2MX
The reaction produces alkali metal halide (MX).
For example: 2Na + Cl2→ 2NaCl
Formation of Alkali Metal Salts
Alkali metals readily produce salts when they react with non-metallic elements or compounds. They contribute their valence electron to the non-metal, leading in the production of a cation and an anion, which combine to create an ionic compound.
When sodium interacts with chlorine, sodium chloride (NaCl) is formed:
Na + Cl2 → 2NaCl
The reaction leads to the formation of sodium chloride, a common salt.
Electronic Structure and Reactivity
Electronic Structure of Alkali Metals
Alkali metals have a unique electronic structure that contributes to their reactivity.
They all have one valence electron in their outermost shell (ns1), where n represents the principal energy level.
The valence electron of alkali metals is located in the outermost s orbital which is used for bonding.
Because of the existence of a single valence electron, alkali metals have a high chemical reactivity. Their reactivity is mostly due to how easily they may shed this valence electron in order to create a stable electron configuration.
Relationship between Electronic Structure and Reactivity
The reactivity of alkali metals is directly related to their electronic structure.
– Alkali metals have a strong tendency to lose their single valence electron to achieve a stable electron configuration.
– By losing this valence electron, alkali metals form cations with a +1 charge, as they have one fewer electron than protons in their nucleus.
– The formation of a cation allows alkali metals to achieve a stable noble gas electron configuration, similar to the nearest noble gas element in the periodic table.
The reaction of alkali metals with water provides an impressive example of their reactivity.
When alkali metals combine with water, an exothermic reaction occurs, releasing hydrogen gas and creating metal hydroxides.
The reaction of alkali metals with water demonstrates the trend in reactivity.
Alkali metals react more violently with water as you progress below Group 1, resulting in a more exothermic and explosive reaction.
For example, lithium interacts slowly with water, but sodium responds quickly and potassium reacts even faster.
The general equation for the reaction of an alkali metal (M) with water (H2O) when it results in the formation of metal hydroxide (MOH) and hydrogen gas (H2):
2M + 2H2O → 2MOH + H2
Observations during the Reaction
Several visible effects occur when alkali metals combine with water:
a) Effervescence: the generation of gas bubbles, specifically hydrogen gas.
b) Heat Production: The reaction is extremely exothermic, generating a large quantity of heat.
c) Alkaline Solution: An alkaline solution is generated, resulting in a basic pH solution.
Reaction Mechanism of Metal with Water
Ionization of Metal Atoms and Formation of Metal Ions
When an alkali metal comes into contact with water, the metal atoms undergo ionization.
The metal atoms lose their outermost valence electron, forming metal ions with a positive charge.
This ionization process is facilitated by the presence of water molecules, which act as a medium for ion transfer.
The ionization reaction can be represented using sodium (Na) as an example:
Na → Na+ + e–
Hydrogen Gas Evolution and Heat Release
As metal ions form, the liberated valence electron combines with water molecules, resulting in the production of hydrogen gas. The apparent effervescence is caused by hydrogen gas bubbles escaping from the solution. Furthermore, the interaction between metal ions and water is extremely exothermic, resulting in large heat generation.
The heat release contributes to the reaction’s total energy shift, making it extremely exothermic. The exothermic character of the reaction can be ascribed to the metal ions formed’s high reactivity and instability.
Reaction of Sodium with Water
– When sodium (Na) reacts with water (H2O), a vigorous reaction occurs.
– Sodium floats on the surface of the water and moves around rapidly due to the evolution of hydrogen gas.
– The reaction produces sodium hydroxide (NaOH), which dissolves in water, resulting in the formation of an alkaline solution.
– The equation for the reaction of sodium with water is:
2Na + 2H2O → 2NaOH + H2
Formation of Metal Hydroxides and Hydrogen Gas
– The reaction between alkali metals and water can be represented by the following general equation (taking sodium as an example):
2M + 2H2O → 2MOH + H2
The reaction yields metal hydroxide (MOH) and hydrogen gas (H2).
Electronic Structure and Reactivity of Group 1 Elements
Alkali metals become increasingly reactive from the top to the bottom of the group. As you move down Group 1, alkali metals exhibit a gradual increase in reactivity. The reactivity trend is primarily influenced by changes in atomic size and ionization energy.
Factors Influencing Reactivity
Atomic Size
– Alkali metals exhibit a gradual increase in atomic size from the top to the bottom of Group 1.
– The larger atomic size results in a greater distance between the nucleus and the valence electron.
– As a result, the valence electron is less strongly attracted to the nucleus, making it easier to remove and increasing reactivity.
Ionization Energy
– Ionization energy refers to the energy required to remove an electron from an atom or ion in the gaseous state.
– Alkali metals have low ionization energies due to the weak attraction between the valence electron and the nucleus.
– The low ionization energy allows alkali metals to readily lose their valence electron and form cations, enhancing reactivity.
Alkali Metals as Reducing Agents
– The reactivity of alkali metals with water makes them powerful reducing agents.
– They readily donate electrons, reducing other substances in chemical reactions.
– For example, alkali metals can be used as reducing agents in organic synthesis or metallurgical processes.
Hydrogen Gas Production
– The reaction between alkali metals and water is a significant source of hydrogen gas.
– Hydrogen gas is widely used in various industries, including fuel cells, hydrogenation processes, and as a clean energy source.
– The reaction with water provides a practical and efficient method for producing hydrogen gas.
Alkali Metals in Alkali Solutions
– The reaction of alkali metals with water produces alkaline solutions.
– These solutions, such as sodium hydroxide (NaOH) and potassium hydroxide (KOH), have numerous applications.
– Alkali solutions are used in industries such as manufacturing soap, detergents, and various chemical processes.
Reaction with Oxygen and Air
Alkali Metals and Oxygen
– Alkali metals readily react with oxygen present in the air, resulting in the formation of metal oxides.
– The reaction between alkali metals and oxygen is highly exothermic.
– The reaction with oxygen is influenced by the reactivity trend observed within Group 1 of the periodic table.
– The general equation for the reaction of an alkali metal (M) with oxygen (O2) is:
4M + O2 → 2M2O
– The reactivity trend is also observed in the reaction of alkali metals with oxygen.
– As you move down Group 1, alkali metals show an increasing tendency to react with oxygen and form metal oxides.
– For instance, lithium reacts slowly with oxygen, forming lithium oxide (Li2O), while potassium reacts vigorously, forming potassium oxide (K2O).
Formation of Metal Oxides
– When alkali metals react with oxygen, they form metal oxides, which are compounds consisting of a metal cation bonded to oxygen.
– The metal oxide formation is an oxidation reaction, where the alkali metal atoms lose electrons and are oxidized, while oxygen gains electrons and is reduced.
– The metal oxides produced have distinct properties and find various applications in different industries.
– Alkali metals readily react with the oxygen present in the air, even at room temperature.
– This reactivity leads to the formation of a thin layer of metal oxide on the surface of the alkali metal, which acts as a protective layer against further reaction with oxygen.
– Sodium (Na) reacts with oxygen and moisture in the air to form sodium oxide (Na2O) and sodium hydroxide (NaOH).
– The balanced chemical equation for the reaction is:
4Na + O2 + 2H2O → 2Na2O + 2NaOH
Gradual Increase in Reactivity with Oxygen Down the Group
– The reactivity of alkali metals with oxygen increases as you move down Group 1 of the periodic table.
– This trend is attributed to the decreasing ionization energy and increasing atomic size down the group.
Reaction of Potassium with Oxygen
– Potassium (K), being the most reactive alkali metal, reacts vigorously with oxygen in the air.
– The reaction results in the formation of potassium oxide (K2O), a white solid compound.
4K + O2 → 2K2O
Reactivity Trend
– Lithium (Li) reacts slowly with oxygen, forming lithium oxide (Li2O).
– Sodium (Na) reacts more vigorously, producing sodium oxide (Na2O).
– Potassium (K) reacts even more vigorously, forming potassium oxide (K2O).
Importance of Metal Oxides
– Potassium oxide (K2O) is used as a drying agent and in the production of other chemicals.
– Metal oxides are also utilized in the manufacturing of ceramics, glass, and catalysts.
Group 1 Elements Reaction with Halogens
Alkali Metals and Halogens
– Alkali metals readily react with halogens, which are elements from Group 17 of the periodic table.
– The reaction between alkali metals and halogens leads to the formation of alkali metal halides.
Formation of Alkali Metal Halides
– When alkali metals react with halogens, they transfer electrons to the halogen atoms, resulting in the formation of ionic compounds known as alkali metal halides.
– The alkali metal donates one electron to the halogen, forming a positively charged alkali metal ion (cation), while the halogen accepts one electron, forming a negatively charged halide ion (anion).
Gradual Decrease in Reactivity from Top to Bottom
– The reactivity of alkali metals with halogens decreases as you move down Group 1 of the periodic table.
– This trend is attributed to the increasing atomic size and decreasing ionization energy down the group.
Reaction of Sodium with Chlorine
– Sodium (Na) reacts with chlorine (Cl2) to form sodium chloride (NaCl), which is a white crystalline compound commonly known as table salt.
– The balanced chemical equation for the reaction is:
2Na + Cl2 → 2NaCl
Reactivity Trend
– Lithium (Li) reacts vigorously with halogens, such as chlorine (Cl2), bromine (Br2), and iodine (I2), to form lithium halides (LiX).
– Sodium (Na) also reacts readily with halogens, but its reactivity decreases compared to lithium.
– Potassium (K) reacts less vigorously with halogens compared to lithium and sodium.
Explanation of the Reaction Mechanism
– The reaction between alkali metals and halogens involves the transfer of electrons, leading to the formation of alkali metal halides.
– This reaction follows an ionic mechanism, where electrons are transferred from the alkali metal atoms to the halogen atoms.
Ionization of Metal Atoms and Formation of Metal Ions
– The alkali metal atoms, such as sodium (Na), have a single valence electron in their outermost energy level.
– During the reaction with halogens, the alkali metal atoms lose their valence electron, becoming positively charged ions.
– The ionization of metal atoms can be represented by the following general equation:
M → M+ + e-, where M represents the alkali metal atom.
Formation of Ionic Compounds through the Transfer of Electrons
– The halogen atoms, such as chlorine (Cl), have high electronegativity, meaning they have a strong tendency to attract electrons.
– When alkali metals react with halogens, the halogen atoms accept the electrons donated by the alkali metals, forming halide ions.
– The formation of alkali metal halides can be represented by the following general equation:
M++ X– → MX, where M+ represents the alkali metal cation and X– represents the halide anion.
Ionic Bonding in Alkali Metal Halides
– Alkali metal halides exhibit ionic bonding, where the positively charged alkali metal ions and negatively charged halide ions are held together by electrostatic forces.
– The transfer of electrons and the resulting formation of ionic compounds contribute to the stability of alkali metal halides.
Reaction of Potassium with Bromine
– Potassium (K) reacts with bromine (Br2) to form potassium bromide (KBr), which is an ionic compound.
2K + Br2 → 2KBr
| Alkali Metal Compound | Applications |
|---|---|
| Sodium Chloride (NaCl) |
|
| Potassium Chloride (KCl) |
|
| Lithium Fluoride (LiF) |
|
| Cesium Iodide (CsI) |
|
| Rubidium Chloride (RbCl) |
|
Importance of Alkali Metal Halides
Sodium chloride (NaCl) is used in food seasoning, water treatment, and chemical production.
– Alkali metal halides are also utilized in the synthesis of other chemicals, as catalysts, and in the production of specialty materials.
Compounds of Alkali Metals
– Alkali metals, located in Group 1 of the periodic table, have a strong tendency to form compounds due to their highly reactive nature.
Ionic Nature of Alkali Metal Compounds
– Alkali metals readily lose their outermost valence electron to form positively charged ions, known as cations.
– The compounds of alkali metals generally have a +1 charge due to the loss of one electron.
– These cations combine with negatively charged ions, known as anions, to form ionic compounds.
| Common Alkali Metal Compounds | |
|---|---|
| A. Alkali Metal Oxides |
|
| B. Alkali Metal Hydroxides |
|
| C. Alkali Metal Carbonates |
|
| D. Alkali Metal Chlorides |
|
| E. Alkali Metal Nitrates |
|
Properties and Applications of Alkali Metal Compounds
– Alkali metal oxides are commonly used as desiccants and in the production of glass and ceramics.
– Alkali metal hydroxides have applications in industries such as chemical manufacturing, soap production, and water treatment.
– Alkali metal carbonates are used in the production of glass, detergents, and as pH regulators.
– Alkali metal chlorides have applications in the production of salts, as electrolytes in batteries, and in the food industry.
– Alkali metal nitrates find uses in fertilizers, pyrotechnics, and the production of nitric acid.
Corrosion of Alkali Metals
– Corrosion refers to the degradation of materials due to chemical reactions with their surroundings.
– In the case of alkali metals, the presence of oxygen and moisture in the air leads to the corrosion of the metal surfaces.
– The formation of surface oxides and hydroxides can affect the appearance, integrity, and functionality of alkali metals.
Alkali Metals in Batteries
– Alkali metals, such as lithium and sodium, are widely used in batteries.
– In rechargeable lithium-ion batteries, the reaction between lithium and oxygen plays a crucial role during the discharge process.
– The chemical equation for the reaction in the positive electrode of a lithium-ion battery is:
2Li + O2 → Li2O
Alkali Metals in Air Purification Systems
– The reactivity of alkali metals with oxygen in the air is utilized in air purification systems.
– Alkali metals, like potassium and cesium, are incorporated into materials used for capturing and removing harmful gases and pollutants from the air.
– These materials react with oxygen, converting the pollutants into less harmful compounds.
– The chemical equations for the reactions depend on the specific pollutants being targeted.
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