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Bonding Essay

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Pure chemical substances are classified as ionic, metallic, covalent molecular and covalent network. In this essay I will describe the nature of each bonding present in these different types of substances and use this to explain the physical properties they exhibit and their structures.
Ionic compounds are compounds that are composed of positive and negative ions. An ionic compound is a chemical compound in which ions are held together in a lattice structure by ionic bonds. Usually, the positively charged portion consists of metal (cations) and the negatively charged portion is an (anion) or polyatomic ion. Ions in ionic compounds are held together by the electrostatic forces between oppositely charged bodies.

The positive and negative ions in these compounds are thought to be arranged in an orderly three-dimensional lattice. For example, the structure of sodium chloride is shown. In the lattice, each positive sodium ion is surrounded by six negative chloride ions and each negative chloride ion is surrounded by six positive sodium ions. The position of the ions is fixed and apart from vibration about these fixed positions no other movement of the ions occurs in the solid compound.

Each ion in an ionic solid is held in the crystal lattice by strong electrostatic attractions to the oppositely charged ions around it. These electrostatic forces between the positive and negative ions are called ionic bonds. Because ionic compounds have high melting points, in other words considerable energy is required to disrupt the attractive forces between the positive and negative ions; ionic bonding is regarded as an example of strong bonding. It’s also a bond in which on atom loses an electron to form a positive ion (cation) and the other atom gains an electron in order to have a full outer shell of electrons to make them stable. One atom pulls an electron from another atom.

For example in sodium chloride, sodium loses one electron to form a sodium ion, electron is not loss but transferred to chlorine to form chlorine ion. They are oppositely charged ions and will be electrostatically attracted to one another. This electrostatic attraction is an ionic bond. Sodium chloride (NaCl) then represents ionic compound because it is held intact by ionic bond.

Another example is magnesium oxide; magnesium transfers its covalent electrons to oxygen forming the magnesium ion and an oxide ion. Then, the resulting ion is electrostatically attracted to one another forming an ionic bond.

Calcium loses two electrons to form a calcium ion, one of these electrons is transferred to chlorine atom and another is transferred to another chlorine atom forming two chloride ions. The calcium ion is electrostatically attracted to the two chloride ions because the opposite charges must fully balance each other up in all three examples, the resulting ionic compound have an overall neutral charges because all charges are full balance out.

In summary an ionic bond is defined as the electrostatic attraction between oppositely charged ions, when an ionic compound is formed the charges must be balance so that the resulting ionic compound has an overall neutral charge.

As with metals there are variations in the properties of individual ionic compounds, but some generalisations can be made about the physical properties characteristic of ionic compounds. Ionic compounds have high melting and boiling points and are all solids at room temperature. They are made up of hard crystals and they are neither malleable nor ductile, but are brittle. In the solid state they are non-conductors of electricity but they are good conductors of electricity in the liquid state. The solutions formed from soluble ionic compounds are good conductors of electricity. Their solubilities in water vary from very soluble to insoluble. They are not soluble in non-polar solvents, such as oil. Ionic solids formed from group 1 and 2 metals are white or colourless while those formed from transition metals are usually coloured The strong, electrostatic attractions between the oppositely charged ions in the three-dimensional lattice result in ionic solids being hard and difficult to cut. The brittleness of ionic compound results from the orderly arrangement of ions in a crystal being disturbed after a layer of ions is forced to slide past another layer. Because of the displacement, ions of similar charge are forced closer to one another with an increase in repulsive forces and a decrease in attractive forces. As a result, the crystal will shatter. The forces of attraction between the oppositely charged ions in ionic compounds are so strong that large quantities of heat energy must be supplied to disrupt the crystal lattice and separate the ions. Consequently, ionic solids have high melting points. In the liquid state, the ions have sufficient energy to move around randomly, but they are still close together. As a result, the attractive forces between the positive and negative ions in the liquid state are still strong. These attractive forces will be broken when the ionic compound boils to form a gas. In other words, molten ionic compounds have high boiling points (even though many decompose before they reach their boiling temperatures). When a soluble ionic compound is added to water, the ions break away from the ionic lattice and mix with the water molecules.

If an insoluble ionic compound is added to water, the ions essentially remain bonded together in the ionic lattice. There is no simple explanation for why some ionic compounds are soluble and others are insoluble in water.
Ionic solids do not conduct electricity because, although they contain charged particles, these ions occupy fixed positions and are not free to move through the solid lattice. Molten ionic compounds do conduct electricity because in the molten state the positive and negative ions are no longer strongly bonded in fixed positions in a lattice but are able to move through the liquid. Although molten ionic compounds conduct an electric current, they do not conduct as well as metals. This suggests that electrons in metals are much more mobile than ions in the molten state.
Covalent bonding is a chemical bond that involves sharing a pair of electrons between atoms in a molecule. The two atoms have positive nuclei. They are both attracted to the same electron. In this way the two nuclei are held close to each other. Each atom will follow the octet rule – A full outer shell (of 8 electrons). In some cases this may involve sharing 4 electrons (a double bond) or 6 electrons (a triple bond).

For example in the hydrogen molecule, the shared pair of electrons would be in the region between the two nuclei most of the time, that is, there is a greater probability of finding the bonding electrons in this region. The density of the electron cloud is therefore greater between the two nuclei in the molecule. The two positively charged nuclei are held together by their mutual attraction to the two shared electrons, or in other words, they are held together by a covalent bond. Because the hydrogen molecule is stable, the electrostatic attractive forces between the two positive nuclei and the shared electrons must balance the repulsive forces between the nuclei and between the electrons. The attractive forces between the nuclei and shared electrons balance the repulsive forces between the nuclei and between the electrons in the stable hydrogen molecule. Covalent bond formation in a hydrogen molecule can also be illustrated using the Bohr-model representation of atoms and electron dot diagrams.

A covalent bond in which 2 electrons are shared is called a single covalent bond and it can be represented by a line drawn between the atoms. Hence a hydrogen molecule can be represented as H–H. Another example is when two chlorine atoms, with electron configurations of 2, 8, 7, combine to form a chlorine molecule, the chlorine atoms share a pair of electrons. In this way, each chlorine atom obtains a share in 8 valence electrons and acquires the electron configuration of the noble gas argon. As with ionic compounds, the octet rule is ‘obeyed’ in many covalent molecular substances.

The two chlorine atoms in the molecule are held together by the attractive forces between their nuclei and the pair of shared electrons. That is, there is a single covalent bond between the atoms in the Cl2 molecule. In the chlorine molecule, 2 electrons are shared, but the other valence electrons are not involved in the bonding. The electron pairs forming covalent bonds in molecules are called bonding electron pairs. The remaining valence electron pairs, if any, are called non-bonding electron pairs or lone pairs. The one bonding pair and six non-bonding pairs of valence electrons in a chlorine molecule are shown below.

Here are some examples of electron dot diagram:
Experimental evidence indicates that covalent molecular substances consist of neutral molecules. As defined earlier, a molecule is a group of two or more atoms held together by covalent bonds. For a covalent molecular substance in the solid state, the neutral molecules are organised in an orderly lattice structure. In the liquid and gas states, the molecules are arranged randomly. In the gas state, however, the molecules are much further apart from each other than in the liquid state. It made up of discrete molecules. The forces that hold the molecules together are comparatively weak and the strength of the covalent bond is very strong.

Substances made up of molecules have the following properties in common. Covalent molecular have low melting and boiling points and many are liquids or gases at room temperature. They are non-conductors of electricity in both the solid and liquid states and the aqueous solutions of some compounds are weak conductors of electricity however few compounds in aqueous solution are good conductors of electricity. Covalent molecular form solids that are generally quite soft are easily scratched and often have a waxy appearance. Many solids are malleable when bent and do not shatter when hit. And many are insoluble in water, but are soluble in non-polar solvents such as petrol and kerosene. Also many have an odour.
Generally, the attractive forces between the molecules of covalent molecular substances are weak, as indicated by the low melting points of these substances. For example, ice has the low melting point of 0°C because the attractive forces between the water molecules, the intermolecular forces, are relatively weak. However, a large amount of energy is required to break up water molecules into individual hydrogen and oxygen atoms. For the process H2O(g) → 2H(g) 1 O(g) to occur to a significant extent, a temperature of over 2000°C is required. This indicates that the covalent bonds between the hydrogen and oxygen atoms within water molecules must be strong. As with melting, when a covalent molecular substance is boiled, it is the weak intermolecular forces between the molecules that are disrupted, not the covalent bonds between the atoms in the molecules.
Covalent molecular solids such as wax, chocolate and butter are soft and can be cut easily. As with the low melting and boiling points, this property can also be explained in terms of the weak attractive forces between the molecules. Very little force is required to push the neutral molecules in the solid lattice past one another. Even with ice, which is possibly one of the harder molecular substances, it can be crushed fairly easily, compared with, for example, the ionic substance marble and the metal steel.
Many covalent molecular substances, particularly those made up of larger molecules such as petroleum jelly and some of the plastic materials, can be bent or pushed into another shape without the solid shattering. When a force is exerted on a molecular substance, the molecules can be pushed past one another easily because of the weak forces between them. But once the molecules reach their new positions, weak forces still exist between them and they remain held together in the new shape.
No charged particles exist. The molecules are neutral. As a result in an electrical current cannot be carried. Covalent molecular acids are the exception. They form ions when they are dissolved in water which is able to conduct electricity (like an ionic solution). For a substance to conduct electricity, mobile charged particles capable of conducting electricity must be present. These charged particles include electrons and positive and negative ions. In covalent molecular substances, there are no ions because the molecules are neutral. Also, in the molecules the valence electrons are localised between or around atoms, as either shared or unshared pairs, and so are not free to move throughout the solid lattice or the liquid. As a result, covalent molecular substances are non-conductors of electricity in the solid and liquid states. When dissolved in water some covalent molecular substances conduct electricity. These substances form or break up into ions (ionise) when added to water. The mobile ions in the aqueous solutions are then able to act as the charge carriers of the electric current.
The solubility properties of covalent molecular compounds vary considerably. Some are very soluble in water, others slightly soluble, but many are insoluble. For example, ethanol and sugar are soluble in water but butter, chocolate and olive oil is insoluble in water. Numerous covalent molecular substances are soluble in non-polar solvents such as oil and kerosene. For example, cooking oil and petrol are soluble in kerosene.

Metallic bonding refers to the interaction between the delocalised electrons and the metal nuclei. The physical properties of metals are the result of the delocalisation of the electrons involved in metallic bonding. Chemists have developed various models to explain the characteristic properties of metals. In one of these models, it is assumed that the outermost or valence electrons of metal atoms move about freely within a three-dimensional arrangement or lattice of positively charged metal ions. That is, the metal consists of a lattice of positive ions surrounded by a ‘sea’ of mobile electrons. The valence electrons are said to be delocalised as they are not associated with a particular metal ion but can move freely through the lattice of metal ions. The positively charged metal ions in the lattice are attracted to the negatively charged delocalised electrons and these electrostatic attractions hold the metallic lattice together. This type of bonding, that is, the electrostatic attractions between the delocalised electrons and the positive metal ions, is called metallic bonding.

Solid metals are good conductors of electricity in both the solid and liquid states. In the solid state, conductivity decreases with increasing temperature. Metals are good conductors of heat in both the solid and liquid states. As a consequence, the solid metals feel cold to touch. Metals are also lustrous (shiny) when freshly cut or cleaned. Metals are malleable (can be hammered or pressed out of shape without breaking) and ductile (able to be drawn into a wire). They are hard, tough and dense at room temperature. Silver in colour (except for copper and gold). Mostly high melting points (there is a great variety within this property though). Sonorous, make a pleasant ringing tone when tapped (if manufactured as a shape that is allowed to vibrate).
The ‘electron sea’ model of metallic bonding can be used to explain the common properties of metals. Metals are good conductors of electricity because of the mobility of the delocalised electrons within the lattice of positive ions. When a metal is used in an electrical circuit, electrons entering one end of the metal cause a similar number of electrons to be displaced from the other end, and the metal conduct. In the solid state, the positive ions do not act as charge carriers. They remain vibrating about fixed positions within the lattice, as the delocalised electron move. However, when a molten metal conducts electricity, both the delocalised electrons and the positive ions are able to move and act as charge carriers.

The delocalised electrons are also responsible for the rapid transmission of heat energy in metals. When one end of a piece of metal is heated, the kinetic energy of the positive ions and the delocalised electrons in the heated region increases, that is, their rate of movement increases. The heat energy is conducted along the piece of metal by the more energetic electrons and positive ions colliding with less energetic electrons and ions. However, because the electrons are delocalised, they are able to move quite freely through the lattice and so cause a transfer of the heat energy along the metal that is more rapid than for most other materials.
The delocalised electrons in metals do not ‘belong’ to any particular positive ion in the lattice, so metallic bonding is said to be non-directional. Therefore, if sufficient force is applied to the metal, one layer of positive ions can slide, or slip, over another without disrupting the metallic bonding. This means that, after the layer of positive ions has moved, there are still attractive forces between the delocalised electrons and the positive ions holding the lattice together in the newly deformed metal. As a result, metals can be hammered readily into sheets or drawn into wires without breaking.

The relatively high melting and boiling points of most metals are due to the strong electrostatic attraction between the positive metal ions and the delocalised electrons. The stronger the metallic bonds, the greater the amount of heat required to move the ions out of their fixed positions, that is, the higher the melting point. It takes a large amount of heat energy to disrupt these forces. Melting points tend to increase across a period. This is partly due to the number of valence electrons increasing – meaning there will be more delocalised electrons as you go across the table and therefore stronger bonding forces.

Most substances formed when non-metal atoms combine are covalent molecular substances. That is, substances in which the atoms are covalently bonded in small groups called molecules. However, a few non-metallic substances have quite different properties compared with those of covalent molecular substances. For example, both carbon and silicon, which are group 14 elements, form oxides with formulas of CO2 and SiO2 respectively, yet the physical properties of these two compounds, are somewhat different. An exception to the rule – where covalent substances usually form discrete molecules. In these few cases these substances form a lattice or network of covalent bonds. Examples include silicon dioxide, carbon (graphite, diamond and fullerenes).
Covalent network substances have a very high melting and boiling points. It is non-conductors of electricity in the solid and liquid states and it’s extremely hard and brittle. Covalent network is reasonably chemically inert and insoluble in water and most other solvents. Covalent network substances have very high melting points and are very hard because each atom is held in the rigid lattice by the strong covalent bonds. To break these bonds requires a very large amount of energy. However, if some of the covalent bonds are broken, the network lattice is placed under stress and rather than it changing shape or becoming distorted, the solid shatters. In a covalent network substance, the electrons are either inner-shell electrons or localised in the bonding and non-bonding pairs around each atom. As a result, these electrons cannot act as charge carriers in the conduction of electricity. There are also no ions present in the lattice. It therefore follows that pure substances, with atoms covalently bonded in a three-dimensional lattice, are not expected to be conductors of electricity.
As there are no charged particles free to move – covalent network substances cannot conduct electricity Graphite is the exception. Each carbon atom has 3 bonds to another carbon all in the same plane. The fourth exists in between two layers of this planar formation – holding the layers together. These electrons are free to move and can thus conduct electricity.
Graphite

Silicon Dioxide

DIAMOND
FULLENERES AND NANOTUBES

SILICON

SILICON CARBIDE

In conclusion, we talked about the different properties, structures and the bonding of metallic compound, ionic compound, covalent network and covalent molecular.

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...Jacinta Houng Comparing the Solubility of Chemicals in Water “Water is known as the “universal solvent” because so many different substances dissolve in it and we rely on this for many of our daily needs.” Introduction: Water is known as the ‘universal solvent’ as it is capable of dissolving a variety of different substances and dissolves more substances than any other liquid. However the ability to be soluble depends on a substances polarity and bonding. This then contributes to the various ways that different types of chemicals interact in water. Solubility is crucial to every living thing on earth as water can carry along valuable chemicals, minerals, and nutrients necessary for survival. In fact Water covers 70% of the Earths surface and composes 55-70% of the human body. Water is an excellent solvent due to its chemical composition and physical attributes. According to USGS (http://water.usgs.gov/edu/qa-solvent.html ) Water molecules have a polar arrangement of the oxygen and hydrogen atoms—one side (hydrogen) has a positive electrical charge and the other side (oxygen) has a negative charge. A polar bond is a covalent bond between two atoms where the electrons forming the bond are unequally distributed. This causes the molecule to have a slight electrical dipole moment where one end is slightly positive and the other is slightly negative. "Like dissolves like" is an expression used by chemists to help them remember how solvents work. The expression refers to...

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