Chemical Bonding and Intermolecular Forces
Under conditions prevailing on earth, most atoms do not exist as individual atoms but they combine with each other to form molecules or unite with other atoms to form compounds. Let us take oxygen as an example. An oxygen atom can combine with another oxygen atom to form diatomic molecule O2 and it can also combine with hydrogen atoms to form water, H2O. If we are to understand the chemistry of everyday matter, we need to understand the nature of chemical bonds that hold atoms together. In this unit, you will learn about the forces that hold atoms together to form molecules or compounds.You will also learn about forces that keep molecules together.
To understand organic chemistry, it is necessary to have some understanding of the chemical bond (the forces that hold atoms together within molecules). Review of Chemical Bonding Bonding is the joining of two atoms in a stable arrangement. Bonding may occur between atoms of the same or different elements.Bonding is a favorable process because it always leads to lowered energy and increased stability.
What is a chemical bond?
Only the noble gases exist as single atoms in nature while most of the other elements are linked together by an attractive force to form molecules or compounds.For example one hydrogen atom can link with another H atom to form H-H (H2) molecule. H can also link with a different type of atom such as carbon (C) to form a compound like CH4. The attractive force that hold atoms together in a molecule or compound is referred to as chemical bond.
One general rule governs the bonding process: Through bonding, atoms attain a complete outer shell of valence electrons. Alternatively, because the noble gases in column 8A of the periodic table are especially stable as atoms having a filled shell of valence electrons, the general rule can be restated as: Through bonding, atoms attain a stable noble gas configuration of electrons.
Noble gases have outer shells filled with electrons and this configuration makes the noble gases stable. Bonding occurs in an attempt to acquire the inert gas configuration so that the atoms are stable. Bonding, therefore, involves shells of atoms that are not completely filled - the outermost shells. The outermost shells are called valence shells and electrons in these shells are called valence electrons. The electrons involved in bonding are called bonding electrons while the valence electrons not involved in bonding are called non-bonding electrons. During bonding, atoms can acquire inert gas configuration by either losing, gaining or sharing of electrons.
Electronegativity and bonding
Bonding can involve atoms of the same type as in hydrogen (H-H) or it can involve two different atoms like hydrogen fluoride (H-F). Molecules formed from bonding of the same atoms are called homonuclear diatomic molecules while those formed from different atoms are called heteronucleardiatomic molecules. The nature of chemical bond is determined by the power with which each participating atom will attract electrons to itself in a molecule. We called this attracting power as electronegativity.
If two atoms have the same electronegativity values it is possible for the two atoms to share pairs of electrons equally in the region of space between their nuclei. On the other hand, if the two uniting atoms have different electronegativity values the two atoms will still share the electrons but unequally or there could be complete transfer of electron(s) from one atom to another. Different types of bonds are possible depending on the electronegativity of the uniting atoms.
Elements that behave in this manner are said to follow the octet rule. There are two different kinds of bonding:
- Ionic bonding and
- Covalent bonding.
Ionic bonds result from the transfer of electrons from one element to another. Elements that behave in this manner are said to follow the octet rule. Covalent bonds result from the sharing of electrons between two nuclei.
A chemical bond is the net result of attractive and repulsive electrostatic forces. When bringing together two atoms that are initially very far apart. Three types of interaction occur, one attractive and two repulsive force. The nucleus-electron attractions are greater than the nucleus-nucleus and electron-electron repulsions, resulting in a net attractive force that holds the atoms together in a molecule. The type of bonding is determined by the location of an element in the periodic table. An ionic bond generally occurs when elements on the far left side of the periodic table combine with elements on the far right side, ignoring the noble gases, which form bonds only rarely. The resulting ions are held together by extremely strong electrostatic interactions. A positively charged cation formed from the element on the left side attracts a negatively charged anion formed from the element on the right side.
Types of Chemical Bonds
There are two major classes of chemical bonds:
- Ionic bond
- Covalent bond.
The Two Extremes
IONIC BOND results from the transfer of electrons from a metal to a nonmetal.
COVALENT BOND results from the sharing of electrons between the atoms. Usually found between nonmetals.
1. Ionic bond
The ionic bond involves atoms that have different electronegativities. The bond forms between two atoms when one or more electrons are transferred from the valence shell of one atom (usually a metal) to the valence shell of the other (usually a non-metal). The two atoms become charged and are called ions. The atom that loses electrons becomes positively charged and is called a cation, and the atom that gains electrons becomes negatively charged and is called an anion.Since the two ions are oppositely charged they attract each other. The electrostatic attraction which hold two ions together is called the ionic bond.
2. Covalent bonding
For elements that have the same or near the same electronegativity values they can bond by sharing the electrons. A covalent bond results from sharing one or more electron pairs between two atoms. In sharing of the electrons, an orbital of one atomoverlaps with the orbital of the other from a different atom. Each orbital contains one electron and so, each atom contributes one electron to a bond. Covalent bonds are formed between non-metal atoms.
Polar covalent bond
So far we have looked at two extremes of bonding where electron completely transfers from one atom to another (ionic bonding) and where the sharing is equal due to same or nearly same electronegativity values of the uniting atoms. Between these extremes are intermediate cases in which the atoms are not so different that electrons are completely transferred but are different enough that unequal sharing results, forming what is called a polar covalent bond. An example of this type of bond occurs in the hydrogen fluoride (HF) molecule. Recall that F is more ele ctronegative than H and so the shared electrons spend more time closer to F than to H. This makes F to acquire a partial negative charge (δ-)and the H to acquire a partial positive charge (δ+). The representation of HF bonding is shown below:
H + F→ Hδ+Fδ-
Such a molecule with partial charges is said to be polarised. In a polarised molecule such as HF, the electrons are still shared as in H2, but unequally.
Coordinate or dative bond
Coordinate or dative bond is similar to covalent bond. However, in a dative bond, both shared electrons come from same atom termed the donor. The other atom with which they share the electrons is termed acceptor. This type of bond is sometimes called donor-acceptor bond. Usually the donor atom is a very electronegative atom such as oxygen and nitrogen.
Intermolecular forces (in order of decreasing strength) are: ion-ion, metallic, dipole- dipole and London dispersion (or induced dipole) forces. (Strictly speaking, covalent bonding, present in covalent network solids, is not an inter-molecular force since the solid in this case is a single giant molecule). ‘Hydrogen bonding’ is a special case of a dipole-dipole force, where an extra large dipole exists between the hydrogen covalently bonded to a small electronegative atom, such as N, O or F. Hydrogen bonding is an inter-molecular force between the hydrogen of one molecule and the lone pair of electrons on the nitrogen, oxygen or fluorine of a neighboring molecule. For molecules with a net dipole moment (or large individual bond dipole), the dominant interaction will be dipole-dipole interactions (such liquids are said to be polar). If the molecules have only weak dipoles (e.g., C-H bonds) then London dispersion (induced dipole) forces become important. If the molecules have no dipole moment, (e.g., H2, noble gases etc.) then the only interaction between them will be the weak London dispersion (induced dipole) force. Large atoms (or non-polar molecules) have larger London dispersion forces as there larger electron clouds are farther away from the nuclei and are therefore more polarizable. For liquids, stronger intermolecular forces result in higher viscosity, surface tension, boiling point and melting point and lower vapour pressure.
- The attractions that molecules have for each other are based on their charge distributions, which can be determined from their structures.
- Some intermolecular forces are stronger and some are weaker.
- A compound is water soluble if its IMFs allow it to effectively be surrounded by water molecules. “Like dissolves like”
- Enzymes recognize molecules by their unique shape and IMF profile
The two types of bonds discussed above involved forces between atoms – interatomic. Now we turn our focus on forces that keep molecules together -Intermolecular forces. Intermolecular forces are forces of attraction that hold molecules together. Note that heteronuclear molecules are either gases, liquids or solids at room temperature but almost all homonuclearmolecules are gases under the same conditions. The nature of intermolecular forces is therefore different in different molecules. One reason for the different sizes of intermolecular forces is the electronegativity of the atoms joined together.
Types of Intermolecular Forces
These forces occur between 2 polar covalent molecules. Shape determines polarity. These forces are weaker than H-bonding but stronger than London Dispersion Forces. So they tend to be very weakly held together solids, and liquids. Hence they tend to evaporate more readily with lower boiling/melting pts. we gave an example of HF molecule as a polar molecule. Polar molecules are called dipoles. The dipole moment in a molecule like HF is along the direction of the most electronegative atom, fluorine. Since the molecules are charged, they attract each other. The partially positively charged H in one HF will be attracted to the partially negatively charged F of the other HF molecule. These forces of attraction are called dipole-dipole attractions since they take place between two dipoles. Dipoles can be permanent or temporal. Polar molecules have permanent dipoles and always exhibit dipole-dipole interactions.
This type of IF only occurs if H is bonded to N,O, F (highly electronegative elements). When bonded the shared electrons are pulled so close to the N, O, or F atoms that it exposes Hydrogen's proton creating a very strong force between that proton and any non-bonded electron pairs occurring on another molecule. These forces are not as strong as ionic bonds and more of these tend to be liquids at room temperature, like H2O, HF, NH3 (ammonia), C2H5OH (ethanol)
When a hydrogen atom is attached to a very electronegative atom such as oxygen, nitrogen or fluorine, a highly polarised molecule is formed. Examples of such molecules or compounds are ammonia (NH3), water (H2O) or HF. The partially positively charged hydrogen in one water molecule is attracted to a partially negatively charged oxygen atom in another water molecule. This electrostatic attraction is the hydrogen bond, a type of dipole-dipole attraction.
Effects of hydrogen bonding on boiling points and melting points of compounds
When a substance melts or boils, intermolecular forces are broken and a high boiling or melting point indicates strong attractive forces. Hydrogen bonding always increases boiling points and melting points of compounds. For example, ethanol (CH3CH2OH) which exhibit hydrogen bonding has a boiling point of 78 °C while ethane (CH3CH3) which has same number of carbon atoms as ethanol has a much lower boiling point of -88.6 °C
Dispersion (London) forces
This type of IF occurs in all cpds but are the primary IF in non polar covalent cpds such as the diatomic elements (O2 , N2 , Cl2 ) as well as any non polar cpds like CH4 , or BCl3 .
These extremely weak forces are created two ways: by mass and by temporary induced dipoles. The more massive these element/cpds, the more London Dispersion Forces will be present.
These are the weakest of the intermolecular forces and therefore they have very low melting/boiling pts, low surface tension for liquids, high viscosity, high vapor pressure, high rates of evaporation, and volatility. Many of these because of their low forces are gases at room temperature, but some will be liquids and solids.
Polar molecules attract each other to form intermolecular forces. Non polar molecules do not have dipole and you may think that they have no attraction to each other. They do attract each other. We know this because all gases such as N2, O2 and even the noble gases can be liquefied or solidified under certain conditions. Liquefaction is brought about by some attractive forces between molecules. To understand the origin of these forces let us imagine two non polar N2, molecules next to each other As the electrons move about in one molecule, a momentary non-symmetrical electron distribution can develop that produces a temporary dipole arrangement of charge. The formation of this temporary dipole can, in turn, affect the electron distribution of a neighbouring molecule. That is, this instantaneous dipole that occurs accidentally in a given atom can then induce a similar dipole in a neighbouring atom.This kind of interaction produces dispersion forces, attractive forces that arise as a result of temporary dipoles induced in atoms or molecules. At very low temperatures, dispersion forces are strong enough to hold the N2 molecules together, causing the gas to condense. N2 gas liquefies at -200 °C
Dipole-dipole forces and dispersion forces are commonly referred to as van der Waals forces, named after a Danish physicist Johannes van der Waals.
Effect of dispersion forces on melting point and boiling point of molecules and compounds
Dispersion forces also called London forces are very weak forces and all molecules have them. These dispersion forces usually increase with molar mass because molecules with larger molar mass tend to have more electrons,and dispersion forces increase in strength with the number of electrons. Increase in dispersion forces increases boiling points of molecules.
These are the forces of attraction between ionic cpds and a polar covalent cpds (like water).These forces are strong enough to break apart the ions and suspend them in the polar solvent (usually water). (KCl (aq), CaCO3 (aq) , RbF (aq) HC2H3O2 (aq) )
The electrostatic forces between ions are greatest type of intermolecular forces. This only occurs between metal and nonmetallic atoms. Strong forces mean higher melting &boiling pts, higher viscosity and surface tension(if liquids), but very low rates of evaporation, vapor pressure, and volatility. Because of the high forces of attraction, the majority of these cpds will be solids at room temperature (KCl (s), CaCO3 (s) , RbF (s) HC2H3O2 (l)).