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CHAPTER 3 CHEMICAL BONDS. The world around us is composed almost entirely of compounds and mixture of compounds. Most of the pure elements also contain many atoms bound together, for example, diamond is a native form of carbon, in which a large number of carbon atoms are bound.
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CHAPTER 3 CHEMICAL BONDS
The world around us is composed almost entirely of compounds and mixture of compounds. Most of the pure elements also contain many atoms bound together, for example, diamond is a native form of carbon, in which a large number of carbon atoms are bound. In compounds, atoms are held together by forces known as chemical bonds. Electrons play a key role in chemical bonding.
There are three ideal types of chemical bonds: - ionic bond (between metals and nonmetals); - covalent bond (between nonmetals); - metallic bond (between metallic atoms).
The ionic bond is a type of chemical bond based on the electrostatic attraction forces between ions havingoppositecharges. It can form between electropositive and electronegative elements, e.g. between metal and non-metal ions. The metal, with a few electrons on the last shell, donates one or more electrons to get a stable electron configuration and forms positively charged ions (cations). These electrons are accepted by the non-metal to form a negatively charged ion (anion) also with a stable electron configuration. The electrostatic attraction between the anions and cations causes them to come together and form a bond.
Example: the formation of ionic bond between Na and Cl. For the sodium atom the electron configuration is: 1s22s22p63s1 The first and second shells of electrons are full, but the third shell contains only one electron. When this atom reacts, it gains the configuration of the nearest rare gas in the periodic table: Ne 1s22s22p6 Na atom loses one electron from its outer shell: Na → Na+ + e-
The chlorine atom has the configuration 1s22s22p63s23p5 It gains one electron and realizes the stable electron configuration of Ar: 1s22s22p63s23p6 Cl + e- Cl- When sodium and chlorine react, the outer electron of the sodium atoms are transferred to the chlorine atoms to produce sodium ions Na+ and chlorine ions Cl- , which are held together by the electrostatic force of their opposite charges. NaCl is an ionic compound.
1s22s22p63s1 1s22s22p63s23p5 NaCl formation 1s22s22p6 1s22s22p63s23p6
NaCl formation may be illustrated showing the outer electrons only (Lewis symbol): In a similar way, a calcium atom may lose two electrons to two chlorine atoms forming a calcium ion Ca2+ and two chloride ions Cl-, that is calcium chloride CaCl2 :
In sodium chloride, the ionic bonds are not only between a pair of sodium ion Na+ an chlorine ion Cl-, but also between all the ions. These electrostatic interactions have as a result the formation of NaCl crystal. We write the formula of sodium chloride as NaCl, but this is the empirical formula. The sodium chloride crystal contains huge and equal numbers of Na+ and Cl- ions pocket together in a way that maximizes the electrostatic forces of the oppositely charged ions. Figures 3.2 and 3.3 show the crystal lattice of NaCl and LiBr.
Covalent bonds The covalent bond is a type of chemical bond formed by sharing pairs of electrons between atoms. When two electronegative atoms react together, ionic bonds are not formed because both atoms have a tendency to gain electrons. In such cases, an stable electronic configuration may be obtained only by sharing electrons. First, consider how chlorine atoms Cl react to form chlorine molecules Cl2 :
Each chlorine atom shares one of its electrons with the other atom. The electron is shared equally between both atoms, and each atom in the molecule has in its outer shell 8 electrons – a stable electronic configuration corresponding to that of Ar. In a similar way a molecule of carbon tetrachloride CCl4 is made up of carbon and four chloride atoms:
The carbon atom shares all its four electrons and the chlorine atoms share one electron each. The carbon atom forms 4 covalent bonds with 4 chlorine atoms. In this way, both the carbon and all four chlorine atoms attain a stable electronic structure. The sharing of a single pair of electrons results in a single covalent bond, often represented by a dash sign, so chlorine molecule may be written as follow: Cl — Cl carbon tetrachloride
For oxygen molecule O2, there are two pairs of electrons shared between the O atoms (double covalent bond): O ═ O In nitrogen molecule (N2) each nitrogen atom shares three electrons. The sharing of three pairs of electrons between two atoms leads to a triple covalent bond N ≡ N Coordinate bond A molecule of ammonia NH3 is made up of one nitrogen and three hydrogen atoms:
The nitrogen atom forms three bonds and the hydrogen atoms one bond each. In this case, one pair of electrons is not involved in bond formation and this is called a lone pair of electrons. It is possible to have a shared electron pair in which the pair of electrons comes just from one electron and not from both. Such bond is called coordinate covalent bond. Even though the ammonia molecule has a stable configuration, it can react with hydrogen H+ by donating the lone pair of electrons, forming the ammonium ion NH4+:
Partial ionic character of covalent bonds In the chlorine molecule Cl – Cl the pair of electrons of the covalent bond is shared equally between both chlorine atom. Because there is not a charge separation toward one of the Cl atoms, Cl2 molecule is nonpolar.
On the contrary, in HCl molecule, there is a shift of electrons toward the chlorine atom which is more electronegative than hydrogen one. Such molecules, in which a charge separation exists is called polar moleculeor dipole molecule The polar molecule of hydrochloric acid
The magnitude of the effect described above is denoted through the dipole moment μ. The dipole moment is the product of the magnitude of the charges (δ) and the distance separating them (d): μ = δ · d The symbol δ suggests small magnitude of charge, less than the charge of an electron ( 1.602 · 10-19 C ). The magnitude of 3.34 · 10-30 Cm means Debye ( D ): 1D = 3.34 · 10-30 C m
The hydrochloric acid molecule has a dipole moment μ=1.03 D and the distance between H and Cl atoms is 136 pm ( 136 · 10-12 m ). A charge δ will be: The charge δ is about 16% of the electron charge (1.602 · 10-19 C ). We can say therefore that the covalent H – Cl bond has about 16% ionic character.
Metallic bond Metals tend to have high melting and boiling points suggesting strong bonds between the atoms. Sodium has the electronic structure 1s22s22p63s1. When sodium atoms come together, the electron in the 3s atomic orbital of one sodium atom shares space with the corresponding electron of a neighbouring atom to form a molecular orbital in the same way that a covalent bond is formed. The difference, however, is that each sodium atom is touched by eight other sodium atoms, and the sharing occurs between the each atom and 3s orbitals of all the eight other atoms. And each oh these eight is in turn touched by eight sodium atoms and so on.
All of the 3s orbitals of all the atoms overlaps to give a vast number of molecular orbitals which extend over the whole piece of metal. There have to be huge numbers of molecular orbitals because any orbital can only hold two electrons. The electron can move freely within these molecular orbitals and so each electron becomes detached from its parent atom. The electrons are said to be delocalised. The metal is held together by the strong forces of attraction between the positive nuclei and the delocalised electron. This may be described as “ an array of positive ions in a sea of electrons “.
The “ free “ electrons of the metal are responsible for the characteristic metallic properties: ability to conduct electricity and heat, malleability (ability to be flattened into sheets), ductility (ability to be drawn into wires) and lustrous appearance. Intermolecular bonds Van der Waals forces Intermolecular forces are attractions between one molecule and neighboring molecules. All molecules are under the influence of intermolecular attractions, although in some cases those attractions are very weak. These intermolecular interactions are known as van der Waals forces. Even in a gas like hydrogen (H2), if you slow the molecules down by cooling the gas, the attractions are large enough for the molecules to stick together in order to form a liquid and then a solid.
In hydrogen’ s case the attractions are so weak that the molecules have to be cooled to 21 K (-252C) before the attractions, are enough to consider the hydrogen as a liquid. Helium’ s intermolecular attractions are even weaker – the molecules won’ t stick together to form a liquid until the temperature drops to 4 K ( -269 C). Attractions are electrical in nature. In a symmetrical molecule like hydrogen, however, these doesn’t seem to be any electrical distortion to produce positive or negative parts. But that’ s only true in average. In the next figure the symmetrical molecule of is represented.
The even shading shows that on average there is no electrical distraction. But the electrons are mobile and at any one instant they might find them selves towards one out if the molecule. This end of the molecule becomes slightly negative (charge -). The other end will be temporarily short of electrons and so becomes slightly positive (+ ) as we can see in the next figure. An instant later the electrons may well have moved up to the other end, reversing the polarity of the temporary dipole of molecule.
Temporary dipole of H2 This phenomena even happens in monoatomic molecules of rare gases, like helium, which consists of a simple atom. If both the helium electrons happen to be on one side of the atom at the same time, the nucleus is no longer properly covered by electrons for that instant.
The question is how temporary dipoles give intermolecular bonds? Imagine a molecule which has a temporary polarity being approached by one which happens to be non- polar just at that moment. Induced dipole
As the right hand molecule approaches, its electrons will tend to be attracted by the slightly positive end of the left hand one. This sets up on induced dipole in the approaching molecule, which is orientated in such a way that the + end of one is attached to the - end of the other. Dipole-dipole attraction
An instant later the electrons in the left hand molecule may well have up the other end. In doing so, they will repel the electrons in the right hand one. The polarity of both molecules reverses, but there is still attraction between - end and + end. As long as the molecules stay close to each other, the polarities will continue to fluctuate in synchronization so that the attraction is always maintained. This phenomena can occur over huge numbers of molecules. The following diagram shows how a whole lattice of molecules could be held together in a solid.
Molecular distribution in a solid The interactions between temporary dipoles and induced dipoles are known as van derWaals dispersion forces .
Now, let us consider a molecule like HCl. Such a molecule has a permanent dipole because chlorine is more electronegative than hydrogen. These permanent dipoles will cause the HCl molecules to attract each other rather than if they had to rely only on dispersion forces. It’s important to realize that all molecules experience dispersion forces. Dipole-dipole interactions are not an alternative to dispersion forces. They occur in addition to them. Molecules which have permanent dipoles will have boiling points higher than molecules which have only temporary fluctuating dipoles. Surprisingly, dipole-dipole attractions are fairly minor compared with dispersion forces, and their effect can be seen if we compare two molecules with the same number of electrons and the same size.
For example, the boiling points of ethane (CH3-CH3) and fluoromethane (CH3F) are: 184.5 K (-88.7C), respectively 194.7 K (-78.5C). The molecule of ethane is symmetric while that of fluoromethane has permanent dipole.
Hydrogen bond If we plot the boiling points the hydride of the elements of groups 15, 16 end 17 we find that the boiling point of the first elements in each group is abnormally high.
In the cases of NH3, H2O and HF there must be some additional intermolecular forces of attraction, requiring significantly more heat energy to break. These relatively powerful intermolecular forces are described as hydrogen bonds.
We can observe that in each of these molecules: • the hydrogen is attached directly to one of the most electronegative elements, causing the hydrogen to acquire a significant amount of positive charge; • each of the elements to which the hydrogen atom is attached is not only significant, but also has one “active“ lone pair of elements. Lone pairs of the 2nd level have the elements contained in a relatively small volume of space which therefore has a high density of negative charge. Lone pair at higher levels are more diffuse and not so attractive to positive particles.
The slightly + charge of hydrogen is strongly attracted to the lone pair end as a result a coordinate bond is formed. This is a hydrogen bond . Hydrogen bond is significantly stronger than a dipole- dipole interaction, but has about a tenth of the strength of an average covalent bond. In liquid water, hydrogen bonds are constantly broken and reformed. In solid water each water molecule can form hydrogen bond surrounding water molecules as we can see in the next figure.
This is why the boiling pint of water is higher than of ammonia or hydrogen fluoride. In the case of ammonia, the amount of hydrogen bonding is limited by the fact that each nitrogen atom has only one lone pair. As well, in hydrogen fluoride, the number of hydrogen atoms is not enough to form a three- dimensional structure.
CHAPTER 4 GAS LAWS
GAS LAWS In a gas the molecules are in a permanent and chaotic motion. Each particle travels in random directions at high speed until it reaches another one, when it is deflected, or until it collides with the wall of the vessel. This movement is called Brownian motion and the gas phase is a completely disordered state. The thermodynamic state of a gas is characterized by its pressure, its volume, and its temperature. The relationship between the pressure, volume, temperature and amount of gas are called gas laws.
Pressure is measured as force per unit area. The SI unit for pressure is Pa (Pascal). However, several other units are commonly used. The table below shows the conversion between these units:
Volume is related between all gases by Avogadro’s hypothesis, which states: Equal volumes of gases, at the same temperature and pressure contain equal numbers of molecules. From this, one can derive the molar volume of a gas, that is the volume occupied by one mole of gas under certain conditions. This values, at 1atm and 0°C is: VM = 22.41 L·mole-1 Temperature is a measure of how much energy the particles have in a gas.
1. Boyle’s law This law was discovered by Robert Boyle (1662) and describes the relationship between the gas pressure and volume. The volume occupied by a given amount of gas is inversely proportional to the pressure at constant temperature: where: p – is the pressure (Pa); V – is the volume (m3); k – is a constant.
Boyle’s law may be written as the relationship: where p1 and V1 are the pressure and the volume in another state, at the same temperature. If we represent this relationship we obtain a set of curves with a shape called equilateral hyperbola, corresponding to a particular temperature.
The explanation of Boyle’s law is based on the fact that the pressure exerted by a gas arises from the impact of its molecules to the walls of the vessel. If the volume is halved, the density of molecules is doubled. In a given interval of time twice as many molecules strike the walls and so, the pressure is doubled in accord with Boyle’s law. This law is universal in the sense that it applies to all gases without reference to their chemical composition.
