Electron Chemistry PDF

Electron Chemistry If a beam of white light passes through a prism, the light will disperse into its component colours, as indicated the by Figure 1. Figure 1 However, if the light is passed through a flask having hydrogen gas in it, an unexpected result occurs. Although light still disperses i...

Electron Chemistry If a beam of white light passes through a prism, the light will disperse into its component colours, as indicated the by Figure 1. Figure 1 However, if the light is passed through a flask having hydrogen gas in it, an unexpected result occurs. Although light still disperses into its composing colours, at certain wavelengths (meaning at certain energies) black lines appear, see Figure 2. Figure 2 What is the meaning of this? An atom has a nucleus surrounded with ‘zones’ where electrons can exist. These ‘zones’ are called energy levels or orbitals, as shown in Figure 3. Furthermore, this means that:   Each orbital has a fixed energy  The higher the energy level the further it is away from the nucleus  Electrons can only exist within each orbital (they cannot exist between orbitals).  Electrons need an exact amount of energy to go to a higher orbital and emit an exact amount of energy when going to lower orbitals (see Figure 4). This amount of energy is called a quantum of energy or photon. In other words, a quanta is the amount of energy needed to move from one energy level to another. Figure 3 Note that when presenting energy levels, we don’t need to draw the orbitals around the nucleus as ‘planets around the sun’. Instead we can use a schematic representation using straight lines. Referring to Figure 3, (right hand side) each straight line represents an energy level. n=1, represents the orbit which is nearest to the nucleus. This orbit has the lowest energy. Orbit 2 is further away from the nucleus and has a higher energy and so on for the other orbitals. Figure 4 This model, which was proposed by Bohr and thus it is called Bohr’s theory. The following rules may be used to find the electron configuration (the arrangement of electrons in an atom) of an element. 1. Electrons occupy the lowest possible energy levels 2. The first shell can hold up to 2 electrons 3. The second shell can hold up to 8 electrons 4. The third shell can also hold up to 8 electrons. Example:  Argon (atomic number has the electronic configuration : 2,8,8  Magnesium has the electronic configuration : 2,8,2 Note: The periodic table of elements and electron configurations. The Periodic table is a tabular arrangement of the chemical elements according to their atomic number. An example of the periodic table is given below. Note: it is very important to examine the Key of the periodic table in order to know what the numbers means. In this case, the number at the top is the atomic number while the number at the bottom is the Mass number Key Atomic Number Symbol Mass Number Figure 5 The periodic table can be subdivided into three main groups, the metals, the metalloids and the non-metals, as shown below. Figure 6 The vertical columns of the periodic table are called groups. All the elements in a group have the same number of electrons in their outer shells. These are called Valency electrons. The group number is the same number of outer shell electrons. For example:  All of the group 1 elements have one valence electron. This causes these elements to react in the same ways as the other members of the family. The elements in group 1 are all very reactive and form compounds in the same ratios with similar properties with other elements. Group 1 elements are also known as the alkali metals.  Group 2 elements are known as the alkaline earth metals. They have two electrons in the valence shell.   Group 7 elements are known as the halogens. They have seven electrons in the valence shell.  Group 8 elements are known as the Noble gases. They have eight electrons in the valence shell. The figure below shows the position of the above groups in the periodic table. Figure 7 The horizontal rows of elements in the periodic table are called periods. For example, the elements sodium (Na) and magnesium (Mg) are both in period 3. The elements astatine (At) and radon (Rn) are both in period 6. The period number gives the number of occupied shells of an element. For example, sodium and Magnesium have three occupied shells whilst astatine and radon have six occupied shells. Thus, if you can locate an element on the Periodic Table, you can use the element’s position to figure out the energy level of the element’s valence electrons. Bonding Most elements form compounds because they want a full outer shell and to achieve that they must react with other atoms. There are various types of bonds, a number of which are described below. Ionic bond The ionic bond involves the electrostatic attraction between ions of opposite charge. Ionic compounds normally form from the reaction between a metal and a non-metal. For example: Sodium (Group 1) has just one electron in its outer shell. It can obtain a full outer shell by losing this electron to anther atom. If sodium losses an electron, it will have 10 electrons but the number of protons will still be 11. Therefore, sodium will form an ion (sodium ion) having 1 positive charge On the other hand, chlorine (group 7) has seven electrons in its outer shell. It can obtain a full outer shell by gaining one electron from another atom. If chlorine gains an electron, it will have 18 electrons but the number of protons will still be 17. Therefore, chlorine will form an ion (chlorine ion) having 1 negative charge. Thus chlorine and sodium can react together. In this case, sodium will lose 1 electron whilst chlorine will gain 1 electron. This will result in the formation of ions (positive ion in the case of sodium and a negative ion in the case of the chlorine). Thus, the sodium ion and chlorine ions will attract each other, forming an ionic bond. We can show this through a dot-cross diagram Figure 8 When sodium reacts with chlorine, billions and billions of sodium and chlorine ions form and they attract each other. They cluster together so that they form a crystal structure (see Figure 9). The overall charge of the structure is 0 since 1 positive charge and 1 negative charge neutralize each other. Figure 9 Note:  Hydrogen and the metals form positive ions  Non-metals form negative ions, and their names end in ide  Group 4 and 5 do not usually form ions because they would have to lose or gain several electrons and that takes too much energy  Group 8 elements do not form ions; they already have full outer shells  Some of the transition metals form more than one ion. Covalent Bond A covalent bond involves the sharing of an electron pair between two atoms each of which contributes one electron to the bond. Non-metals undergo Covalent bonding since they need to gain electrons. For example, chlorine atoms have seven valance electrons. If there are two chlorine atoms, each chlorine atom needs 1 electron to have a full outer shell. Thus, the chlorine atoms bond together by sharing an electron pair, as shown below. Figure 10 Note that  When a pair of electrons is shared between two atoms, a single covalent bond (single bond) is formed.  When 2 pairs of electrons are shared between two atoms, a double covalent bond (double bond) is formed.  When 3 pairs of electrons are shared between two atoms, a triple covalent bond (triple bond) is formed. The molecules in a covalent compound are not flat because electrons repel each other and thus try to get as far apart from each other as possible. Two examples are shown in Figure 11 below. H N H H Figure 11 Dipole-Dipole attractions In a covalent bond, the electron pair found in that bond is not always equally shared between the two atoms of that bond. The electron pair tends to be shifted towards the most electronegative atom. Electronegativity can be defined as a measure of the tendency of an atom to attract a bonding pair of electrons. For instance, taking hydrogen Chloride (HCl) as an example, Chlorine is more electronegative electronegative than hydrogen and therefore, the electrons will be found closer to the chlorine atom Figure 12 This will cause the hydrogen (H) and chlorine (Cl) to have partial charges, a partial positive charge in the case of hydrogen and a partial negative charge in the case of chlorine. Therefore, these molecules will have an electrostatic attraction to each other as shown in Figure 11. Figure 13 A particular strong dipole-dipole attraction is called the hydrogen bond; this involves the attractive force between molecules containg hydrogen and a very elctro- negative element, such as oxygen (O). The small hydrogen atom in one molecule can approach very closely the electro-negative atom in a neighbouring molecule. Hydrogen bonding is responsible for the abnormally high boiling and melting temperatures in water. Hydrogen bonding is also responsible to high viscosity and high surface tension in water. Polar molecules tend to interact together but not with non-polar molecules. Since water is a polar molecule, and an important solvent, polar molecules which dissolve in water are caller hydrophilic, which literally means water loving. On the other hand, non polar molecules which do not dissolve in water are called hydrophobic or lipophilic. These molecules will usually dissolve in fats, oils and organic solvents. Van der Waal’s Forces. Electrons in an atom or molecule are in constant motion and at any moment, the centres of positive charge and negative charge are displaced. Due to this electron motion we have an oscillating electron dipole. Oscillating electric dipoles attract one another, with such forms of attraction being called Wan der Waals’s forces. The larger an atom or a molecule is the larger the number of electrons present and the large the magnitude of this oscillating electric dipole. This results in a stronger Van der Waals’ force. Figure 14 Metallic Bonding Metallic Bonding consists of an array of cations in a ‘sea of delocalised electrons’. This theory explains why metals are good thermal and electrical conductors. For example, when one side of a metal bar is heated, the kinetic energy of electrons in the vicinity increases. These electrons transfer the energy to neighbouring ones. Consequently, the kinetic energy of the electrons at the opposite end of the bar also increases meaning that the temperature in this part is raised. Figure 15