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21 -- Part 1: 1 Periodic Trends and Charge Density
The group 1 atom in a given period is the largest that hydrogen is often placed and is the most easily ionized.
The group 1 elements have a low densi table, but some of them are not ties and are an alkali metal.
This metal is highly reactive.
The reactivity of the alkali metals can be seen in their reactions with water.
The group 2 has been measured.
The metals of group 2 are not as dense as a typical metal, but they are still highly reactive and less dense than the alkali metals.
The group 2 metals have densities that are greater than that of water, and they only react slowly with water.
Group 13 and group 14 will be discussed in this chapter.
Both metals and nonmetals are encountered in these groups.
Boron has interesting chemistry because it tends to form molecule with incomplete octets around the central boron atoms.
One of the most widely used metals is aluminum.
It is possible to get aluminum metal from its compounds.
Because aluminum production requires a lot of electricity, aluminum-production plants are located close to a lot of hydroelectricity.
Group 13--Gallium, indium, and thallium--are all metals.
The chemistry of group 13 is dominated by aluminum and boron, and we will only mention the heavier elements in this chapter.
Group 14 has a nonmetal, two metalloids, and two metals.
The chemistry of carbon is the most important in the group since it occurs in all living systems.
Silicon is found in many minerals.
Tin and lead can be obtained using methods that have been used for thousands of years.
There are many opportunities to relate new information to principles presented earlier in the text.
The ideas of atomic structure, periodic trends in atomic and ionic radii, chemical bonding, and thermodynamics will help us understand the chemical behavior of the elements.
The periodic trends that we have covered in this text can be rationalized in the chemistry of the elements.
The elements in a given group have similar electronic configurations, but not the same chemical properties, because the atoms of each group have similar electronic configurations.
The lightest member of a group has features that are different from the rest of the group.
Trends in atomic properties will be reviewed in this section.
We will begin to understand the trends in the chemistry of the elements with these ideas.
The atomic properties of an element are responsible for its chemistry.
The electronegativity of an atom is important.
The ground state electronic configuration of He is 1s2.
A summary of trends in atomic radius, first ionization energy, electron affinity, electronegativity, and atomic polarizability can be found here.
The shaded elements are the focus of the chapter.
Chapter 9 discussed atomic radii, ionising energies, and electron affinities.
Chapter 10 and Chapter 12 discussed electronegativities and polarizabilities.
We have discussed polarization of the anion trends before in this text.
A high charge density distorts the electron from top to bottom in a group.
The anion has first ion energies, electron affinities, and cloud around it.
The electronegativities show that the quantities increase across the electron cloud and decrease down a group.
cations are smaller than the parent atoms when shown with a dashed line and radii, and anions are larger, so it is important to remember.
Explanations for the trends have been shown.
Chapters 9 and 10 have been given elsewhere in the text.
The bond between the anion electron cloud and the internuclear region is distorted when a cation interacts with an anion.
As a result of the distortion, the bond between the cation and the anion has a variety of different character.
Charge density is deterred by the polarizability of the anion and the cation.
The electron cloud of the anion is distorted for me.
The larger and more polarizable anions derived from atoms that are lower down in a group are defined by some authors as charge density.
The I- ion is more polarizable than the F-.
The charge density of the Main- Group Elements I: Groups 1, 2, 13, and 14 is between 1 and 1000 Cmml -3.
The charge density increases as the charge on the cation increases.
The higher the charge density, the greater the ability of a cation to distort the electron cloud of an anion toward itself.
The charge density concept will be used to rationalize certain observations.
It will be used to help us understand why there are dramatic differences in the properties of elements in the same group.
It is not possible to use a single quantity of Group 1 Elements as a substitute for careful consideration of all contributing factors.
3 and AlI3 are used to measure the amount of solid crust.
Since prehistoric times, some of their compounds have been used.
The elements were isolated in pure form about 200 years ago.
The elements of the alkali metals are difficult to break down by ordinary chemical means, so discovery had to wait for new scientific developments.
The two substances were discovered through electrolysis.
It was discovered in 1817.
Cesium and rubidium were identified as new elements.
Francium was isolated from actinium.
Natural brines can be used to obtain a number of Li, Na, and K compounds.
Na Albert Russ/Shutterstock 2CO3 can be mined as a solid deposit.
The salt is obtained from the water.
LiAl1SiO322 is the name for Rubidium and cesium.
The group 1 elements are the most active metals.
Several of their properties are listed in Table 21.2, and a few of them are discussed next.
ion pairs are converted to gaseous atoms when NaCl is vaporized.
As excited atoms 1Na*2 return to their groundstate electron configurations, light with a wavelength of 589 nm emits as the excited atoms Na(g) are excited to higher energies.
Yellow2 Alkali metal compounds are used in fireworks.
The atomic radii of the group 1 elements increase from the top to the bottom, as was described in Chapter 9.
These large atoms make Chip Clark/Fundamental Photographs for a relatively low mass per unit volume.
The lighters of the alkali metals will float.
The property leads to soft met The sodium, an active metal with low melting points.
A bar of sodium is covered with a thick oxide coating.
The values given here assume a coordination number of 4 for Li+ and 6 for the others.
Ten minerals are ranked on the Mohs scale, ranging from that of talc to diamond.
Only substances with lower values can be scratched.
A good indicator of the extreme metallic character of the group 1 elements is their standard reduction potentials, which are large, negative quantities.
The metal M(s) is very easy to oxidize to M+1aq2 because it is difficult to reduce the ion M+.
The alkali metals can reduce water to H21g2.
We can use the values given in Table 21.2 to calculate them.
The Edegcell values show that the strongest reducing agent in the solution is lithium.
The strong reducing agents in the alkali metals are in the solution.
The equilibrium position for reac is always very far to the right, and so the reaction is controlled by completion regardless of which alkali metal is involved.
The time it takes for a reaction to happen and the rate of it.
It is necessary to consider what happens to the energy that is released by the controlled factors to explain this observation.
The energy released by the reaction is used to heat the system.
The energy released by the reaction is enough to melt the unreacted metal.
The reaction of lithium and water is not as vigorous as it is for the other alkali metals because it does not melt as the reaction proceeds.
The melting point of NaCl is too high a temperature to carry this economically.
The Downs cell is used for the reduction of molten KCl.
The Downs cell has NaCl1l2 and K1g2 in it.
At low temperatures, most of the KCl(l) remains is molten NaCl(l) to which CaCl2 has been added unreacted.
The equilibrium is displaced far to the right as to lower the melting point of K(g) from the molten mixture.
The K(g) is free of any Na(g) present when the liquid metals are fractionally distilled.
Rb and Cs can be produced with Ca metal as the reducing agent.
The most important use of the metal is because it is so easy to oxidize and because it can be kept apart by reducing agent.
Titanium metal can be obtained from the reduction of TiCl4 by Na.
In a nuclear reactor, sodium is used as a heat-transfer medium.
It has better thermal con elevated temperature.
It is easy to pump because of its low density and low viscosity.
sodium is used in outdoor lighting.
The total quantity consumed in this application is small because each lamp uses a small amount of Na.
It is possible to make high-strength, low density alloys with aluminum and magnesium using lithium metal.
These are used in the aircraft industry.
The ease of oxidation and the large number of electrons produced by a small mass of lithium make it an anode material in batteries.
It takes 6.94 g Li to produce one mole of electrons.
In cardiac pacemakers, the installed battery must have high reliability and a long lifetime to be useful.
There is an X-ray photograph of a pacemaker.
Li is the smallest of the alkali metal atoms.
It is a different matter when it comes to oxidation.
The result of a hypothetical three-step process is what we can think of.
We must compare tendencies in each of the three steps if we want to form M+1aq2 by oxidation of the metals.
The secondary hydration sphere is formed by the listing of electrode potentials holding other molecule, but more weakly.
The easiest substance to oxidize is Li.
A common method of summing up important reactions is from a compound of central importance.
This section describes some of the reactions.
Alternative methods may be used to prepare a number of these compounds.
The conversion of Na2CO3 to NaOH is no longer important.
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