CSIR-NET LIFE SCIENCE NOTE FOR UNIT-1 - C.STABILIZING INTERACTION AND WATER

UNIT-1 - C.STABILIZING INTERACTION AND WATER

                  A.STABILIZING INTERACTIONS-

1.CO-VALENT BONDS -

Atoms can become more stable is by sharing electrons (rather than fully gaining or losing them), thus forming covalent bonds. Covalent bonds are more common than ionic bonds in the molecules of living organisms.

For instance, covalent bonds are key to the structure of carbon-based organic molecules like our DNA and proteins. Covalent bonds are also found in smaller inorganic molecules, such as $\text H_2\text O$H, start subscript, 2, end subscript, O, $\text {CO}_2$C, O, start subscript, 2, end subscript, and $\text {O}_2$O, start subscript, 2, end subscript. One, two, or three pairs of electrons may be shared between atoms, resulting in single, double, or triple bonds, respectively. The more electrons that are shared between two atoms, the stronger their bond will be.

As an example of covalent bonding, let’s look at water. A single water molecule, $\text H_2\text O$H, start subscript, 2, end subscript, O, consists of two hydrogen atoms bonded to one oxygen atom. Each hydrogen shares an electron with oxygen, and oxygen shares one of its electrons with each hydrogen:

Hydrogen atoms sharing electrons with an oxygen atom to form covalent bonds, creating a water molecule

Image credit: OpenStax Biology.

The shared electrons split their time between the valence shells of the hydrogen and oxygen atoms, giving each atom something resembling a complete valence shell (two electrons for H, eight for O). This makes a water molecule much more stable than its component atoms would have been on their own.

Polar covalent bonds

There are two basic types of covalent bonds: polar and nonpolar. In a polar covalent bond, the electrons are unequally shared by the atoms and spend more time close to one atom than the other. Because of the unequal distribution of electrons between the atoms of different elements, slightly positive (δ+) and slightly negative (δ–) charges develop in different parts of the molecule.

In a water molecule (above), the bond connecting the oxygen to each hydrogen is a polar bond. Oxygen is a much more electronegative atom than hydrogen, meaning that it attracts shared electrons more strongly, so the oxygen of water bears a partial negative charge (has high electron density), while the hydrogens bear partial positive charges (have low electron density).

In general, the relative electronegativities of the two atoms in a bond – that is, their tendencies to "hog" shared electrons – will determine whether a covalent bond is polar or nonpolar. Whenever one element is significantly more electronegative than the other, the bond between them will be polar, meaning that one end of it will have a slight positive charge and the other a slight negative charge.

Nonpolar covalent bonds

Nonpolar covalent bonds form between two atoms of the same element, or between atoms of different elements that share electrons more or less equally. For example, molecular oxygen ($\text {O}_2$O, start subscript, 2, end subscript) is nonpolar because the electrons are equally shared between the two oxygen atoms.

Another example of a nonpolar covalent bond is found in methane ($\text {CH}_4$C, H, start subscript, 4, end subscript). Carbon has four electrons in its outermost shell and needs four more to achieve a stable octet. It gets these by sharing electrons with four hydrogen atoms, each of which provides a single electron. Reciprocally, the hydrogen atoms each need one additional electron to fill their outermost shell, which they receive in the form of shared electrons from carbon. Although carbon and hydrogen do not have exactly the same electronegativity, they are quite similar, so carbon-hydrogen bonds are considered non-polar.

Image modified from OpenStax Biology.

2.ELECTRO SATIC INTERACTIONS-

• Electrostatic interaction due to electro +ve elements like Na,K,Ca,Mg etc donate electrons and become CATION and electronegative elements like Sulphur, Oxygen accept electron and become ANION

3.Van der Waals interactions- 

• TYPE OF INTERACTION WHEN 2 MOLECULES LIE CLOSE TO EACH OTHER.
• WEAKEST
• INTERACTION BETWEEN PERMANENT DIPOLES MUCH WEAKER THAN IONIC INTERACTION & CAN INDUCE DIPOLE MOVEMENT IN A NEIGHBOURING GROUP BY ELECTROSTATIC DE-SHORTING ELECTRON DISTRIBUTION
• DIPOLE-DIPOLE INTERACTION IS MORE STRONGER THAN DIPOLE INDUCED DIPOLE INTERACTION

UMMMMMM WHAT IS DIPOLE??

- A PAIR OF EQUAL & OPPOSITE CHARGED OR MAGNETIZED POLES SEPARATED BY A DISTANCE

                                 fig-1

4.Hydrogen bonding-

In a polar covalent bond containing hydrogen (e.g., an O-H bond in a water molecule), the hydrogen will have a slight positive charge because the bond electrons are pulled more strongly toward the other element. Because of this slight positive charge, the hydrogen will be attracted to any neighboring negative charges. This interaction is called a hydrogen bond.

Hydrogen bonds are common, and water molecules in particular form lots of them. Individual hydrogen bonds are weak and easily broken, but many hydrogen bonds together can be very strong.

hydrogen bond

• bond distance= 1.5-2.6 Amstrong
• weaker bond than covalent

5.London dispersion forces-

Interactions between ions, dipoles, and induced dipoles account for many properties of molecules - deviations from ideal gas behavior in the vapor state, and the condensation of gases to the liquid or solid states. In general, stronger interactions allow the solid and liquid states to persist to higher temperatures. However, non-polar molecules show similar behavior, indicating that there are some types of intermolecular interactions that cannot be attributed to simple electrostatic attractions. These interactions are generally called dispersion forces. The London dispersion force is the weakest intermolecular force. It is a temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles

  fig: Dispersion Interaction with an instantaneous dipole on one He atom inducing a dipole on a nearby He atom.

                                         Figure : Dispersion interaction in the gas phase

COMPARISION BETWEEN DIFFRENT BONDS-

B.Water

and its Properties-

Why should we study water in Biology?
• Living things are made of 70-90% water.
• Water is essential for all organisms.

• Most abundant substance

• Essential for all forms of life

• H and O linked by covalent bonds

• Each water molecule is linked to others

• All the major components in cells (proteins, DNA, RNA

and polysaccharides) can dissolve in water

• It solubilizes and modifies the properties of biomolecules

• Water molecule is not linear and the oxygen atom has a

Water

(H2O)

higher electronegativity than hydrogen atoms, it carries

a slight negative charge, while the hydrogen atoms are

slightly positive.

• As a result, water is a polar molecule with an electrical

dipole moment.

Two very important properties of water

1. Water is a polar molecule. The water molecule is bent, not
linear, and so the distribution of charge is asymmetric.
- The oxygen nucleus draws electrons away from the
hydrogen nuclei, which leaves the region around the
hydrogen nuclei with a net positive charge.
- The water molecule is thus an electrically polar structure.
2. Water is highly cohesive. Water molecules interact strongly with one another through hydrogen bonds.
- These interactions are apparent in the structure of ice .
- Networks of hydrogen bonds hold the structure together;
- similar interactions link molecules in liquid water and
account for the cohesion of liquid water, although, in the liquid
state, some of the hydrogen bonds are broken.
- The highly cohesive nature of water dramatically affects
the interactions between molecules in aqueous solution.

-The polarity and hydrogen-bonding capability of water
make it a highly interacting molecule.
- Water is an excellent solvent for polar molecules.
- The reason is that water greatly weakens electrostatic
forces and hydrogen bonding between polar molecules by
competing for their attractions.
- For example, consider the effect of water on hydrogen
bonding between a carbonyl group and the NH group of an

amide.

- A hydrogen atom of water can replace the amide
hydrogen atom as a hydrogen-bond donor, whereas the
oxygen atom of water can replace the carbonyl oxygen
atom as a hydrogen-bond acceptor.
- A strong hydrogen bond between a CO group and an
NH group forms only if water is excluded.
- The existence of life on Earth depends critically on the
capacity of water to dissolve a remarkable array of polar
molecules that serve as fuels, building blocks, catalysts,
and information carriers.
- High concentrations of these polar molecules can coexist
in water, where they are free to diffuse and interact with
one another.

• Many properties of water are emergent properties due to hydrogen bonding.
• The cohesion of water molecules to each other is exploited by plants and animals.
• Water resists temperature changes by absorbing lots of heat.
• Lower density of ice causes it to float & insulate the 11 water below.
• The polarity of water allows it to dissolve other polar molecules.
• Non-polar compounds are hydrophobic and not easily dissolved in water.

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