Bonding Lecture 2 PDF

**Which out of NH~3~ and NF~3~ has higher dipole moment and why?** **The dipole moment of NH~3~ is higher in comparison to NF~3~. REASON The F is more electronegative than N and pulls the electrons. As the number NF bonds is 3 the 3 F atoms pull the electrons. The resulting electronegative forces d...

**Which out of NH~3~ and NF~3~ has higher dipole moment and why?** **The dipole moment of NH~3~ is higher in comparison to NF~3~. REASON The F is more electronegative than N and pulls the electrons. As the number NF bonds is 3 the 3 F atoms pull the electrons. The resulting electronegative forces decrease the dipole moment. NF3 DIPOLE MOMENT 0.24DThe NH3 molecule having 3 Hydrogen molecules that form bond with N do not have electronegativity to over come the pull of N molecule unlike the NF3 which results in higher dipole moment. NH3 DIPOLE MOMENT 1.24D.** **Coordinate Covalent Bond** Co-ordinate bond is a type of alternate covalent bond that is formed by sharing of electron pair from a single atom. Both shared electrons are donated by the same atom. It is also called dative bond or dipolar bond. Co-ordinate [[covalent bonds]](https://byjus.com/jee/covalent-bond/) are usually formed in reactions that involve two non-metals such as a hydrogen atom or during bond formation between metals ions and ligands. Characteristics Of Coordinate Covalent Bond 1. In this type of bonding, the atom that shares an electron pair from itself is termed as the donor. 2. The other atom which accepts these shared pair of electrons is known as a receptor or acceptor. 3. The bond is represented with an arrow →, pointing towards acceptor from the donor atom. 4. After sharing of electron pain each atom gets stability. 5. This type of bonding is central to the Lewis theory. 6. Getting a good understanding of co-ordinate covalent bonds can help in properly designing complex organic molecules. Coordinate Bond Diagram ----------------------- Below we have given a simple diagram of a co-ordinate bond. The bond is shown by an arrow which points in the direction where an atom is donating the lone pair to the atom that is receiving it. ###  ### ### ### ### ### Formation Of Ammonium Ion The nitrogen atom in Ammonia donates its electron pair to the empty orbital of H^+ ^ion thus nitrogen is donor, H^+^ is acceptor and a co-ordinate bond is formed Coordinate Covalent Bond Definition , Examples ,Formation and Properties \| Covalent bonding, Coordinates, Bond EXAMPLES: --------- Carbon monoxide is linear in structure As we know that one of the bonds in CO is a coordination bond, which is formed by the donation of lone pairs of electrons from Oxygen to Carbon atoms. ------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------------- 2). ------------------------ 3) -- 4) ----------------------- What is Hydrogen Bonding? ------------------------- Hydrogen bonding refers to the formation of Hydrogen bonds, which are a special class of attractive intermolecular forces that arise due to the dipole-dipole interaction between a hydrogen atom that is bonded to a highly electronegative atom and another highly electronegative atom which lies in the vicinity of the hydrogen atom. For example, in water molecules (H~2~O), hydrogen is covalently bonded to the more electronegative oxygen atom. Therefore, hydrogen bonding arises in water molecules due to the dipole-dipole interactions between the hydrogen atom of one water molecule and the oxygen atom of another H~2~O molecule. Here, the location of the bond pair of electrons in the O-H bond is very close to the oxygen nucleus (due to the large difference in the electronegativities of oxygen and hydrogen). Therefore, the oxygen atom develops a partial negative charge (-δ) and the hydrogen atom develops a partial positive charge (+δ). Now, hydrogen bonding can occur due to the electrostatic attraction between the hydrogen atom of one water molecule (with +δ charge) and the oxygen atom of another water molecule (with -δ charge). Thus, hydrogen bonds are a very special class of intermolecular attractive forces that arise only in compounds featuring hydrogen atoms bonded to a highly electronegative atom. Hydrogen bonds are mostly strong in comparison to normal dipole-dipole and dispersion forces. However, they are weak compared to true covalent or ionic bonds. What are the Conditions for Hydrogen Bonding? In a molecule, when a hydrogen atom is linked to a highly electronegative atom, it attracts the shared pair of electrons more and so this end of the molecules becomes slightly negative while the other end becomes slightly positive.  The negative end of one molecule attracts the positive end of the other and as a result, a weak bond is formed between them. This bond is called the **hydrogen bond.** As a result of hydrogen bonding, a hydrogen atom links the two [electronegative atoms](https://byjus.com/chemistry/electronegativity/) simultaneously, one by a covalent bond and the other by a hydrogen bond. The conditions for hydrogen bonding are: 1. The molecule must contain a highly electronegative atom linked to the hydrogen atom. The higher the electronegativity more is the polarization of the molecule. 2. The size of the electronegative atom should be small. The smaller the size, the greater is the electrostatic attraction. Examples of Hydrogen Bonding ---------------------------- ### Hydrogen Bonding in Water A water molecule contains a highly electronegative oxygen atom linked to the [hydrogen atom](https://byjus.com/chemistry/hydrogen/). Oxygen atom attracts the shared pair of electrons more and this end of the molecule becomes negative whereas the hydrogen atoms become positive. ### ### ### ### ### ### Hydrogen Bonding in Ammonia It contains highly electronegative atom nitrogen linked to hydrogen atoms.  ### ### Strength of the Hydrogen bond ----------------------------- The hydrogen bond is a weak bond. The strength of hydrogen bond is in-between the **weak [van der Waals forces](https://byjus.com/chemistry/van-der-waals-forces/) **and the strong covalent bonds. The dissociation energy of the hydrogen bond depends upon the attraction of the shared pair of electrons and hence on the electronegativity of the atom.\ Properties of Hydrogen Bonding ------------------------------ - **Solubility: **Lower alcohols are soluble in water because of the hydrogen bonding which can take place between water and alcohol molecule. - **Volatility: **As the compounds involving hydrogen bonding between different molecules have a higher boiling point, so they are less volatile. - **Viscosity and surface tension: **The substances which contain hydrogen bonding exists as an associated molecule. So their flow becomes comparatively difficult. They have higher viscosity and high surface tension. - **The lower density of ice than water: **In the case of solid ice, the hydrogen bonding gives rise to a cage-like structure of water molecules. As a matter of fact, each water molecule is linked tetrahedral to four water molecules. The molecules are not as [closely packed](https://byjus.com/chemistry/close-packing-three-dimensions/) as they are in a liquid state. When ice melts, this case like structure collapses and the molecules come closer to each other. Thus for the same mass of water, the volume decreases and density increases. Therefore, ice has a lower density than water at 273 K. That is why ice floats**.** Types of Hydrogen Bonding ------------------------- There are two types of H bonds, and it is classified as the following: - Intermolecular Hydrogen Bonding - Intramolecular Hydrogen Bonding ### Intermolecular Hydrogen Bonding When hydrogen bonding takes place between different molecules of the same or different compounds, it is called** intermolecular hydrogen bonding.** For example -- hydrogen bonding in water, alcohol, [ammonia](https://byjus.com/chemistry/ammonia/) etc. ### Intramolecular Hydrogen Bonding The hydrogen bonding which takes place within a molecule itself is called **intramolecular hydrogen bonding.** It takes place in compounds containing two groups such that one group contains hydrogen atom linked to an electronegative atom and the other group contains a highly electronegative atom linked to a lesser electronegative atom of the other group. The bond is formed between the hydrogen atoms of one group with the more [electronegative atom](https://byjus.com/chemistry/electronegativity/) of the other group n water, each hydrogen nucleus is covalently bound to the central oxygen atom by a pair of electrons that are shared between them. In H~2~O, only two of the six outer-shell electrons of oxygen are used for this purpose, leaving four electrons which are organized into two non-bonding pairs. The four electron pairs surrounding the oxygen tend to arrange themselves as far from each other as possible in order to minimize repulsions between these clouds of negative charge. This would ordinarly result in a tetrahedral geometry in which the angle between electron pairs (and therefore the H-O-H *bond angle*) is 109.5°. However, because the two non-bonding pairs remain closer to the oxygen atom, these exert a stronger repulsion against the two covalent bonding pairs, effectively pushing the two hydrogen atoms closer together. The result is a distorted tetrahedral arrangement in which the H---O---H angle is 104.5°. ### Water\'s large dipole moment leads to hydrogen bonding The H~2~O molecule is electrically neutral, but the positive and negative charges are not distributed uniformly. This is illustrated by the gradation in color in the schematic diagram here. The electronic (negative) charge is concentrated at the oxygen end of the molecule, owing partly to the nonbonding electrons (solid blue circles), and to oxygen\'s high nuclear charge which exerts stronger attractions on the electrons. This charge displacement constitutes an *electric dipole*, represented by the arrow at the bottom; you can think of this dipole as the electrical \"image\" of a water molecule. Opposite charges attract, so it is not surprising that the negative end of one water molecule will tend to orient itself so as to be close to the positive end of another molecule that happens to be nearby. The strength of this ***dipole-dipole attraction*** (described in more detail [here](http://www.chem1.com/acad/webtext/states/interact.html#3A)) is less than that of a normal chemical bond, and so it is completely overwhelmed by ordinary thermal motions in the gas phase. #### Hydrogen bonding in water But when the H~2~O molecules are crowded together in the liquid, these attractive forces exert a very noticeable effect, which we call (somewhat misleadingly) *hydrogen bonding*. And at temperatures low enough to turn off the disruptive effects of thermal motions, water freezes into ice in which the hydrogen bonds form a rigid and stable network. Notice that the hydrogen bond (shown by the dashed green line) is somewhat longer than the covalent O---H bond. It is also *much weaker*, about 23 kJ mol^--1^ compared to the O--H covalent bond strength of 492 kJ mol^--1^. One factor that makes water unique even among other hydrogen-bonded liquids is its very small mass in relation to the large number of hydrogen bonds it can form. Owing to disruptions of these weak attractions by thermal motions, the lifetime of any single hydrogen bond is very short --- on the order of a picosecond. At any instant, the average H~2~O molecule is bound to somewhat fewer than four neighbors --- estimates vary from 2.4 to 3.6. #### Why ice floats on water The most energetically favorable configuration of H~2~O molecules is one in which each molecule is hydrogen-bonded to four neighboring molecules. Owing to the thermal motions described above, this ideal is never achieved in the liquid, but when water freezes to ice, the molecules settle into exactly this kind of an arrangement in the ice crystal. This arrangement requires that the molecules be somewhat farther apart then would otherwise be the case; as a consequence, ice, in which hydrogen bonding is at its maximum, has a more open structure, and thus a lower density than water. Here are three-dimensional views of a typical local structure of ice **Ice**, like all solids, has a well-defined structure; each water molecule is surrounded by four neighboring H~2~Os. two of these are hydrogen-bonded to the oxygen atom on the central H~2~O molecule, and each of the two hydrogen atoms is similarly bonded to another neighboring H~2~O.  When ice melts, the more vigorous thermal motion disrupts much of the hydrogen-bonded structure, allowing the molecules to pack more closely. Water is thus one of the very few substances whose solid form has a lower density than the liquid at the freezing point. Localized clusters of hydrogen bonds still remain, however; these are continually breaking and reforming as the thermal motions jiggle and shove the individual molecules. As the temperature of the water is raised above freezing, the extent and lifetimes ofthese clusters diminish, so the density of the water increases. At higher temperatures, another effect, common to all substances, begins to dominate: as the temperature increases, so does the amplitude of thermal motions. This more vigorous jostling causes the average distance between the molecules to increase, reducing the density of the liquid; this is ordinary thermal expansion. Because the two competing effects (hydrogen bonding at low temperatures and thermal expansion at higher temperatures) both lead to a decrease in density, it follows that there must be some temperature at which the density of water passes through a maximum. This temperature is 4° C; this is the temperature of the water you will find at the bottom of an ice-covered lake in which this most dense of all water has displaced the colder water and pushed it nearer to the surface..