INTRODUCTION TO ALKALI METALS, HALOGENS AND INERT GASES ~ BY AFFAN BIN USMAN 8TH GREEN
The alkali metals consist of the chemical elements lithium, sodium, potassium, rubidium, caesium, and francium. Together with hydrogen they constitute group 1, which lies in the s-block of the periodic table. Alkali Metals, sorted by Atomic Number: Atomic Symbol Atomic Mass Name Number 3 Li 6.941 Lithium 11 Na 22.98977 Sodium 19 K 39.0983 Potassium Rb 85.4678 Rubidium 37 Cs 132.9054 Cesium Fr (223) Francium 55 87
The Group 1 elements have similar properties because of the electronic structure of their atoms - they all have one electron in their outer shell. The alkali metals consist of the chemical elements lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). All alkali metals have their outermost electron in an s-orbital.
Relationships between different Alkali metals: Alkali metals share common characteristics. They are all soft, silver metals. Due to their low ionization energy, these metals have low melting points and are highly reactive. The reactivity of this family increases as you move down the table. The Group 1 elements have similar properties because of the electronic structure of their atoms - they all have one electron in their outer shell. Reactivity:- Increases or decreases down the group 1? ~ Yes the reactivity of group 1 elements increases as you go down the group because: the atoms become larger. the outer electron becomes further from the nucleus. the force of attraction between the nucleus and the outer electron decreases.
Why do alkali metals get more reactive going down group 1? Alkali metals from lithium to potassium get more reactive because the force of attraction between the nucleus (core) and the outer electron gets weaker as you go down group 1 elements. The distance \"c\" is greater than \"a\" and the force of attraction between the nucleus and the outer shell (rings) diminishes with distance.
The more electron shells (rings) between the nucleus and outer electron also creates shielding and again this weakens the nuclear attraction. The outer electron is more easily transferred to say an oxygen atom, which needs electrons to complete its full outer shell.
REACTIONS OF THE GROUP 1 ELEMENTS WITH WATER Lithium Lithium's density is only about half that of water so it floats on the surface, gently fizzing and giving off hydrogen. It gradually reacts and disappears, forming a colourless solution of lithium hydroxide. The reaction generates heat too slowly and lithium's melting point is too high for it to melt Sodium Sodium also floats on the surface, but enough heat is given off to melt the sodium (sodium has a lower melting point than lithium and the reaction produces heat faster) and it melts almost at once to form a small silvery ball that dashes around the surface. A white trail of sodium hydroxide is seen in the water under the sodium, but this soon dissolves to give a colourless solution of sodium hydroxide. The sodium moves because it is pushed around by the hydrogen which is given off during the reaction. If the sodium becomes trapped on the side of the container, the hydrogen may catch fire to burn with an orange flame. The colour is due to contamination of the normally blue hydrogen flame with sodium compounds. Potassium Potassium behaves rather like sodium except that the reaction is faster and enough heat is given off to set light to the hydrogen. This time the normal hydrogen flame is contaminated by potassium compounds and so is coloured a faintly bluish pink.
Rubidium Rubidium is denser than water and so sinks. It reacts violently and immediately, with everything spitting out of the container again. Rubidium hydroxide solution and hydrogen are formed. Caesium Caesium explodes on contact with water, quite possibly shattering the container. Caesium hydroxide and hydrogen are formed Atomic Number, Mass, and Melting and Boiling Points of the Alkali Metals: -Lithium (Li) Melting Point: 180.5°C | Boiling Point: 1,347°C -Sodium (Na) Melting Point: 97.8°C | Boiling Point: 883° -Potassium (K) Melting Point: 63.28°C | Boiling Point: 759°C -Rubidium (Rb) Melting Point: 39.31°C | Boiling Point: 688°C -Cesium (Cs) Melting Point: 28.4°C | Boiling Point: 671°C -Francium (Fr) Melting Point: 27°C | Boiling Point: 677°C
HALOGENS (7) The halogens are a group in the periodic table consisting of five chemically related elements: fluorine, chlorine, bromine, iodine (I), and astatine. Key Points ○ Halogens are nonmetals in group 17 (or VII) of the periodic table. Down the group, atom size increases. As a diatomic molecule, fluorine has the weakest bond due to repulsion between electrons of the small atoms.
○ Due to increased strength of Van der Waals forces down the group, the boiling points of halogens increase. Therefore, the physical state of the elements down the group changes from gaseous fluorine to solid iodine. ○ Due to their high effective nuclear charge, halogens are highly electronegative. Therefore, they are highly reactive and can gain an electron through reaction with other elements. Halogens can be harmful or lethal to biological organisms in sufficient quantities.
Physical Properties Atoms get bigger down the group as additional electron shells are filled. When fluorine exists as a diatomic molecule, the F–F bond is unexpectedly weak. This is because fluorine atoms are the smallest of the halogens—the atoms are bonded close together, which leads to repulsion between free electrons in the two fluorine atoms. The boiling points of halogens increase down the group due to the increasing strength of Van der Waals forces as the size and relative atomic mass of the atoms increase. This change manifests itself in a change in the phase of the elements from gas (F2 , Cl2) to liquid (Br2) , to solid (I2). The halogens are the only periodic table group containing elements in all three familiar states of matter (solid, liquid, and gas) at standard temperature and pressure.
Chemical Properties Electronegativity is the ability of an atom to attract electrons or electron density towards itself within a covalent bond. Electronegativity depends upon the attraction between the nucleus and bonding electrons in the outer shell. This, in turn, depends on the balance between the number of protons in the nucleus, the distance between the nucleus and bonding electrons, and the shielding effect of inner electrons. In hydrogen halides (HX, where X is the halogen), the H-X bond gets longer as the halogen atoms get larger. This means the shared electrons are further from the halogen nucleus, which increases the shielding of inner electrons. This means electronegativity decreases down the group. Halogens are highly reactive, and they can be harmful or lethal to biological organisms in sufficient quantities. This reactivity is due to high electronegativity and high effective nuclear charge. Halogens can gain an electron by reacting with atoms of other elements.
The halogens have the following properties: ○ They are non-metals stable as diatomic molecules (this means at room temperature and pressure, they exist as molecules made of two atoms, e.g. Cl2). ○ They have a valence of 1 and form covalent bonds with nonmetals atoms, or ionic bonds with metal atoms. H alogen ions will usually have a single negative charge (X- ), where they are known as halides. ○ They are colored. ○ They have relatively l ow melting and boiling points compared to other non-metals (except the noble gases). ○ Halogens in elemental form are relatively toxic, reactive substances. ○ They do not conduct heat or electricity. ○ They are b rittle as solids. As we go down the group, the properties of the elements c hange in the following ways: ○ The melting and boiling point gets higher – starting as gases, bromine is a liquid while iodine is a solid. ○ The color of the halogens gets darker – fluorine is pale yellow, followed by green chlorine, brown/purple bromine and purple iodine. ○ Electronegativity decreases down the group. The smallest halogen, fluorine, is the most electronegative element in the periodic table. ○ The halogens get l ess reactive – fluorine, top of the group, is the most reactive element known. Iodine is the least reactive halogen (besides astatine which is often ignored because it is extremely rare). ○ As one of the more reactive groups of elements, there are a variety of reactions the halogens take part in: The halogens react well with group 1 and 2 metals because these have electron configurations that complement the halogens. The metals react by losing electrons; the halogens react by gaining them. These are vigorous, exothermic reactions.
INERT GASES (8) The noble gases, also known as the inert gases or rare gases, are located in Group VIII or International Union of Pure and Applied Chemistry (IUPAC) group 18 of t he periodic table. This is the column of elements along the far right side of the periodic table. This group is a subset of the nonmetals. Collectively, the elements are also called the helium group or the neon group. The noble gases are: 1. Helium (He) 2. Neon (Ne) 3. Argon (Ar) 4. Krypton (Kr) 5. Xenon (Xe) 6. Radon (Rn) 7. Oganesson (Og) With the exception of oganesson, all of these elements are gases at ordinary temperature and pressure. There haven't been enough atoms produced of oganesson to know its phase for certain, but most scientists predict it will be a liquid or solid. Both radon and oganesson consist only of radioactive isotopes. Noble Gas Properties The noble gases are relatively non-reactive. In fact, they are the least reactive elements on the periodic table. This is because they have a complete v alence shell. They have little tendency to gain or lose electrons. In 1898, Hugo Erdmann coined the phrase \"noble gas\" to reflect the low reactivity of these elements, in much the same way as the noble metals are less reactive than other metals. The noble gases have high ionization energies and negligible
electronegativities. The noble gases have low boiling points and are all gases at room temperature. The Atomic and Physical Properties ● Atomic mass, boiling point, and atomic radii INCREASE down a group in the periodic table. ● The first ionization energy DECREASES down a group in the periodic table. ● The noble gases have the largest ionization energies, reflecting their chemical inertness. ● Down Group 18, atomic radius and interatomic forces INCREASE resulting in an INCREASED melting point, boiling point, enthalpy of vaporization, and solubility. ● The INCREASE in density down the group is correlated with the INCREASE in atomic mass. ● Because the atoms INCREASE in atomic size down the group, the electron clouds of these non polar atoms become increasingly polarized, which leads to weak van Der Waals forces among the atoms. Thus, the formation of liquids and solids is more
easily attainable for these heavier elements because of their melting and boiling points. ● Because noble gases’ outer shells are full, they are extremely stable, tending not to form chemical bonds and having a small tendency to gain or lose electrons. ● Under standard conditions all members of the noble gas group behave similarly. ● All are monoAtomic gases under standard conditions. ● Noble gas atoms, like the atoms in other groups, INCREASE steadily in atomic radius from one period to the next due to the INCREASING number of electrons. ● The size of the atom is positively correlated to several properties of noble gases. The ionization potential DECREASES with an INCREASING radius, because the valence electrons in the larger noble gases are further away from the nucleus; they are therefore held less tightly by the atom. ● The attractive force INCREASES with the size of the atom as a result of an INCREASE in polarizability and thus a DECREASE in ionization potential. ● Overall, noble gases have weak interatomic forces, and therefore very low boiling and melting points compared with elements of other groups. Summary of Common Properties 1. Fairly nonreactive 2. Complete outer electron or valence shell (oxidation number = 0) 3. High ionization energies 4. Very low electronegativities 5. Low boiling points (all monatomic gases at room temperature) 6. No color, odor, or flavor under ordinary conditions (but may form colored liquids and solids) 7. Nonflammable
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