How Many Bonds Can Chlorine Form – Electric forces are responsible for the chemical bonding of atoms, ions, and ionic groups that form solid crystals. The physical and chemical properties of minerals are largely determined by the types and strength of these binding forces; hardness, cleavage, fusibility, electrical and thermal conductivity, as well as the coefficient of thermal expansion are examples of such properties. In general, the hardness and melting point of the crystal increase in proportion to the bond strength, while the coefficient of thermal expansion decreases. The extremely strong forces that bind the carbon atoms of diamonds, for example, are responsible for the pronounced hardness. Periclase (MgO) and halite (NaCl) have similar structures; however, periclase has a melting point of 2,800 °C (5,072 °F), while halite melts at 801 °C (1,474 °F). This discrepancy reflects the difference in the bond strength of the two minerals: because the periclase atoms are held together by a stronger electrical force, more heat is required to separate them.
Electrical forces, called chemical bonds, can be divided into five types: ionic, covalent, metallic, van der Waals, and hydrogen bonds. Classification in this way is largely optional; in fact, the chemical bonds in a given mineral may have the characteristics of more than one type of bond. For example, the forces binding silicon and oxygen atoms in quartz exhibit characteristics of both ionic and covalent bonds almost equally. As mentioned above, the electrical interaction between the atoms of a crystal determines its physical and chemical properties. Thus, the classification of minerals according to their electrical capacity will provide a grouping of species with similar properties. This fact justifies the classification by type of bond.
How Many Bonds Can Chlorine Form
Atoms tend to gain or lose electrons to make their outer orbitals stable; this is usually achieved by filling these orbitals with the maximum allowable number of valence electrons. For example, sodium metal has one valence electron in its outermost orbital; it ionizes by easily losing this electron and exists as the Na cation
Why The Formation Of Ionic Compounds Is Exothermic
Ions together there is an attraction between two opposite charges. This bonding mechanism is called ionic or electrovalent (
Crystals associated with an ionic bond usually exhibit moderate hardness and specific gravity, fairly high melting points, and poor thermal and electrical conductivity. The ion’s electrostatic charge is evenly distributed over its surface, so the cation tends to be surrounded by the maximum number of anions that can be located around it. Because ionic bonding is nondirectional, crystals bonded in this way usually exhibit high symmetry.
In the discussion of ionic bonding, it was noted that chlorine readily gains an electron to achieve a stable electronic configuration. The incomplete outer orbit puts the chlorine atom in a very reactive state, so it tries to bond with almost every atom in its vicinity. Since its nearest neighbor is usually another chlorine atom, they can bond together by sharing a pair of electrons. Due to this extremely strong bond, each chlorine atom enters a stable state.
An electronic or covalent bond is the strongest of all types of chemical bonds. Minerals bound in this way show general insolubility, high stability and high melting point. Crystals of covalently bonded minerals tend to show lower symmetry than their ionic counterparts because the covalent bond is highly directional, localized near the shared electrons.
Ch104: Chapter 3
Molecules formed by binding two adjacent chlorine atoms are stable and do not combine with other molecules. However, atoms of some elements have more than one electron in their outer orbit, so they can bond with several neighboring atoms to form groups, which in turn can combine into larger combinations. Carbon in the polymorphic form of diamond is a good example of this type of covalent bond. A carbon atom has four valence electrons, so each atom bonds with four others in a stable tetrahedral configuration. Connecting each carbon atom in this way creates a continuous network. The rigid diamond structure results from strong localization of binding energy near shared electrons; this makes diamond the hardest of all natural substances. Diamond does not conduct electricity, because all the valence electrons of the atoms that make up its composition form bonds between themselves and therefore are not mobile.
Bonding in metals differs from bonding in their salts, as evidenced by significant differences between the properties of the two groups. Unlike salts, metals exhibit high plasticity, viscosity, plasticity and conductivity. Many are characterized by lower hardness and higher melting and boiling points than, for example, covalently bonded materials. All of these properties result from the metal’s bonding mechanism, which can be thought of as a collection of positively charged ions immersed in a cloud of valence electrons. The attraction between cations and electrons holds the crystal together. Electrons are not bound to any particular cation and are therefore free to move through the structure. In fact, the radiant energy of light in the metals sodium, cesium, rubidium, and potassium can cause complete removal of electrons from their surfaces. (This result is known as the photoelectric effect.) The mobility of electrons is responsible for the ability of metals to conduct heat and electricity. Native metals are the only minerals that exhibit pure metallic bonding.
Neutral molecules can be held together by a weak electrical force known as van der Waals bonding. This is the result of deforming the molecule so that a small positive charge develops at one end and a corresponding negative charge at the other end. A similar effect is caused in neighboring molecules, and this dipole effect spreads throughout the structure. Then an attraction occurs between the oppositely charged ends of the dipoles. Van der Waals bonding is common in gases, organic liquids, and solids, but rare in minerals. Its presence in the mineral defines a weak region with good cleavage and low hardness. In graphite, the carbon atoms lie in covalently bonded sheets with van der Waals forces acting between the layers.
In addition to the four basic types of bonds described above, there is an interaction called a hydrogen bond. This occurs when a hydrogen atom bonded to an electronegative atom such as oxygen, fluorine, or nitrogen is also attracted to the negative end of a neighboring molecule. A strong dipole-dipole interaction occurs, forming a bond between the two molecules. Hydrogen bonds are common in hydroxides and many layered silicates, such as micas and clay minerals.
Question Video: Naming The Product Of The Addition Reaction Of Ethene And Chlorine Gas
The physical properties of minerals are a direct result of the structural and chemical properties of minerals. Some properties can be determined by inspection of a sample taken from the hand or by relatively simple tests on such a sample. Others, such as those determined by optical and X-ray diffraction methods, require specialized and often complex equipment and may involve lengthy sample preparation. The following discussion focuses on those properties that are easiest to assess with simple tests. The Lewis structure of CCL4 is a diagram depicting the electronic configuration of covalently bonded compounds. Lewis structures are designed to visualize the atomic structure and distribution of electrons in a particular chemical compound.
Carbon tetrachloride (CCl4) is a covalently bonded compound consisting of a central carbon surrounded by 4 chlorine atoms in a tetrahedral structure. The Lewis diagram of carbon tetrachloride looks like this:
A normal carbon atom has 4 unshared electrons in its outer shell. Chlorine has 7 electrons, so 1 electron is missing to completely fill the outer shell. Thus, each of the 4 outer electrons of the carbon atom will be shared with one chlorine atom, giving the individual carbon atoms and 4 chlorine atoms a full outer shell of electrons. in the resulting compound, each element has reached a stable electronic configuration, having 8 electrons in the outer shell.
Lewis structures were first presented by the American chemist H.N. Lewis in 1916. Since then, they have become ubiquitous in high school and college chemistry courses as an easy way to understand chemical bonding.
A) Polyvinyl Chloride Structure And Possible Defects, Where The R…
Lewis structures are designed to represent the atomic and electronic structure of a chemical compound. Each element of a compound is represented in the Lewis structure by its chemical symbol, so H for hydrogen, C for carbon, O for oxygen, and so on. The electron shell configuration of an element is represented by a pattern of dots around the chemical symbol. Shared electron pairs are shown as a single line connecting two bonded elements. Lone pairs (electrons not involved in a chemical bond) appear as a pair of loose dots next to the chemical symbol.
—the number of electrons in the outer shell. For example, oxygen has a valence of 6 because it has 6 electrons in its outer shell. Most elements will try to fill their outer shell completely and will bond with other elements until their valence number is 8, which corresponds to a full outer shell of 8 electrons. Tendency to
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