Q1) Why are
covalent compounds generally poor conductors of electricity?
Ans:-
Covalent compounds are generally poor conductors of
electricity due to their molecular structure and the nature of the chemical
bonds that hold their atoms together.
In covalent compounds, atoms share electrons to form stable
molecules. These shared electrons are localized between the bonded atoms, resulting
in the formation of discrete molecules rather than a continuous lattice
structure. This molecular structure gives rise to several factors that
contribute to the poor conductivity of covalent compounds:
1. Lack of Free Electrons: Covalent compounds typically do
not have free electrons that are mobile and can move through the material to
carry an electric charge. In metals, which are good conductors of electricity,
the presence of delocalized electrons allows for the easy flow of electric
current.
2. No Ion Formation: Covalent compounds do not readily form
ions in solution or in the molten state. In ionic compounds, such as salts,
ions are present that can carry electric charges and facilitate the movement of
charge through the material. In covalent compounds, the electrons are tightly
shared between the atoms, and there are no readily available charged particles
to conduct electricity.
3. Absence of Charged Particles: In covalent compounds, the
individual molecules are electrically neutral, and there are no charged
particles (ions) present within the material that can move to conduct an
electric current. In contrast, in ionic compounds, the presence of positive and
negative ions allows for the movement of charge.
4. Insulating Properties of Molecules: Covalent compounds
are often composed of molecules with non-metallic elements, which tend to have
high electronegativities. This results in strong covalent bonds and the absence
of free mobile charge carriers. Consequently, covalent compounds have
insulating properties and do not allow the flow of electric current.
5. Lack of Crystal Lattice: Unlike ionic and metallic
compounds, covalent compounds do not form a regular three-dimensional crystal
lattice structure. In metals, the crystal lattice provides a pathway for
electrons to move freely throughout the material, facilitating electrical
conductivity.
It's important to note that there are exceptions to the
general rule. Some covalent compounds can conduct electricity under specific conditions.
For example, covalent compounds that are composed of mobile ions or have
delocalized electrons (as seen in some organic conductive polymers) can exhibit
conductivity. Additionally, covalent compounds may conduct electricity when
they undergo chemical reactions that result in the formation of ions, such as
in acids.
In summary, the poor conductivity of covalent compounds is
primarily attributed to their molecular structure, which lacks the presence of
free electrons or mobile ions necessary for the efficient movement of electric
charge.
Q2) Explain, by giving a reason, why is it very difficult to
form (i)cation C4+, (ii) anion c4- .
Ans:-
(i) Formation of C4+ Cations:
Cations are formed when atoms lose electrons and become
positively charged. Forming a C4+ cation, where an atom loses four electrons,
is very difficult because of the following reasons:
1. Stability of Atoms: Atoms like to have a stable
arrangement of electrons. When they lose electrons, they become positively
charged. For carbon to lose four electrons, it would have to give up all its
outer electrons, and this makes the atom very unstable.
2. Strong Attraction: The nucleus of an atom is positively
charged, and electrons are negatively charged. When electrons are lost, the
positive nucleus pulls on the remaining electrons very strongly. If we take
away four electrons, the pull becomes even stronger, and it's hard for the atom
to let go of more electrons.
3. Energy Required: Taking away four electrons requires a
lot of energy. It's like trying to remove something that's stuck very tightly.
The atom needs a huge amount of energy to do this, and it's not something that
happens easily.
(ii) Formation of C4- Anions:
Anions are formed when atoms gain electrons and become
negatively charged. Making a C4- anion, where an atom gains four electrons, is
also very challenging:
1. Electron Cloud Repulsion: Electrons are negatively
charged, and they repel each other. Adding four extra electrons around an atom
would make their negative charges push against each other a lot, like trying to
fit too many people in a small room. This makes the atom uncomfortable and
unstable.
2. Size Increase: Adding electrons makes the atom bigger.
It's like blowing up a balloon. When the atom becomes bigger, the positive
nucleus has a harder time holding onto the extra electrons because they are
farther away.
3. Energy Barrier: Just like with cations, gaining four
electrons needs a lot of energy. It's like trying to catch something very fast
and heavy. The atom has to work really hard to catch all these electrons and
hold onto them.
In simple words, forming C4+ or C4- ions is tough because it
goes against how atoms naturally like to arrange their electrons. It needs a
lot of energy and can make the atom very unstable. Similarly in C4- it will be difficult to have 6 no of Protons and 10 nos. of (in valence shell 4+4=8) electrons. Nature prefers more balanced
and stable ways of forming ions. So it will create a covalent bond.
Q3) Explain how a single covalent bond is formed between.
(i) Two atoms of hydrogen (ii) Two atoms of fluorine.
Support your answer by drawing dot diagrams and structural
diagrams.
Ans:-
(i) Formation of a Single Covalent Bond in Hydrogen (H2):
1. Starting Atoms: Imagine two hydrogen atoms, each with one
electron. These electrons are in their own orbits, and the atoms want to be
more stable. Making 2 electrons in the valence shell.
2. Sharing Electrons: The two hydrogen atoms come close to each other. It has 1 ecotones in their valence shell, it need 1 more electrons to fill its octant and become stable. They decide to share their electrons to become stable. Each hydrogen atom contributes one electron to the shared pair.
3. Bond Formation: This sharing creates a bond between the
two hydrogen atoms. Now, both atoms have two electrons around them, which is
the same as helium, a stable gas.
4. Hydrogen Molecule: The two hydrogen atoms are now
connected by a single covalent bond. They form a molecule called hydrogen gas
(H2).
(ii) Formation of a Single Covalent Bond in Fluorine (F2):
1. Starting Atoms: Imagine two fluorine atoms, each with
seven electrons. These electrons are in different orbits, and the atoms want to
achieve stability.
2. Sharing Electrons: The two fluorine atoms come close to each other. It has 7 ecotones in their valence shell, it need 1 more electrons to fill its octant and become stable. They share one electron from each atom's outer orbit, creating a shared pair of electrons.
3. Bond Formation: This sharing of electrons creates a bond
between the two fluorine atoms. Each atom now has eight electrons, which is the
same as a noble gas called neon.
4. Fluorine Molecule: The two fluorine atoms are now
connected by a single covalent bond. They form a molecule called fluorine gas
(F2).
[ In both cases, the atoms come together to share electrons
and become more stable. This sharing of electrons forms a single covalent bond
between the atoms. This bond helps the atoms stick together to create
molecules. Single covalent bonds are important in holding many different
elements and compounds together. ]
Q4) By drawing dot diagrams explain the formation of
(i) a molecule of oxygen (ii) a molecule of nitrogen.
Ans:-
(i) Formation of a Molecule of Oxygen (O2):
1. Oxygen Atoms: We imagine two oxygen atoms (O). Each oxygen
atom has six electrons in its outermost shell.
2. Sharing Electrons: The two oxygen atoms come close to each other. It has 6 ecotones in their valence shell, it need 2 more electrons to fill its octant and become stable. They decide to share two pairs of electrons to become more stable.
3. Formation of Oxygen Molecule (O2): When the electrons are
shared, it creates a bond between the two oxygen atoms. This bond holds the
atoms together in a molecule. Now, each oxygen atom has a total of eight
electrons around it, which is the same as the noble gas neon (Ne) in its valence shell. Hence this becomes stable now.
Importance of Molecules:
The oxygen molecule (O2) is very important for us to breathe. It's the oxygen that our cells need to work properly, and it's also what we breathe in from the air.
(ii) Formation of a Molecule of Nitrogen (N2):
1. Nitrogen Atoms: We imagine two nitrogen atoms (N). Each
nitrogen atom has seven electrons in its outermost shell.
2. Sharing Electrons: The two nitrogen atoms come close to
each other. It has 5 ecotones in their valence shell, it need 3 more electrons to fill its octant and become stable. They decide to share three pairs of electrons to become more
stable.
3. Formation of Nitrogen Molecule (N2): Sharing these
electrons creates a bond between the two nitrogen atoms, making a molecule.
Now, each nitrogen atom has a total of eight electrons around it, which is
again similar to the noble gas neon (Ne).
Common Gas in the Air:
Nitrogen gas (N2) is a big part of the air we breathe. It makes up about 78% of the air. Even though we don't use nitrogen for breathing like oxygen, it's still important for the environment.
[ In both cases, the atoms come together and share electrons
to become more stable. This sharing creates bonds that hold the atoms in a
molecule. These molecules play important roles in the air we breathe and the
environment around us. ]
Q5) In electrons dot structure, the valence shell electrons
are presented by crosses or dots.
(a) The atomic number of chlorine is 17. Write its
electronic configuration.
Ans:-
The atomic number of chlorine is 17, which means it has 17
electrons. The electronic configuration of chlorine can be represented as
follows:
1s² 2s² 2p⁶ 3s² 3p⁵
Also we can write as 2.8.7
Let's break down what this means:
- The first energy level (closest to the nucleus) can hold
up to 2 electrons, and it is filled in this case with 1s², meaning there are
two electrons in the 1s sublevel.
- The second energy level can also hold up to 8 electrons.
It is filled with 2s² (two electrons in the 2s sublevel) and 2p⁶ (six electrons
in the 2p sublevel).
- The third energy level can hold up to 8 electrons as well,
but it will fill with the rest of the 7 elctrons. It is filled with 3s² (two
electrons in the 3s sublevel) and 3p⁵ (five electrons in the 3p sublevel).
So, the electronic configuration of chlorine is 1s² 2s² 2p⁶
3s² 3p⁵.
This arrangement of electrons helps us understand the
structure of the chlorine atom and how its electrons are distributed in
different energy levels and sublevels.
In short, The atomic number of chlorine is 17, which means it has 17 electrons. The electronic configuration of chlorine 2.8.7
(b) Draw the electron dot structure of chlorine molecule. A
Ans:-
Q6) Name a molecule of an element that has:
(i) One covalent bond
(ii) Two covalent bonds
(iii) Three covalent bonds
Ans:-
Certainly, here are examples of molecules of different
elements that have the specified number of covalent bonds:
(i) One Covalent Bond:
1. Hydrogen Molecule (H2): In a hydrogen molecule, two
hydrogen atoms share a single covalent bond, forming a diatomic molecule.
(ii) Two Covalent Bonds:
1. Oxygen Molecule (O2): In an oxygen molecule, two oxygen atoms share
a double covalent bond, resulting in a stable diatomic molecule.
(iii) Three Covalent Bonds:
1. Nitrogen Molecule (N2): In a nitrogen molecule, two nitrogen atoms are
connected by a triple covalent bond, forming a diatomic molecule.
[ These examples illustrate the concept of covalent bonding,
where atoms share electrons to achieve a more stable electron configuration.
Each covalent bond involves the sharing of one, two, or three pairs of
electrons, which corresponds to single, double, and triple covalent bonds,
respectively. ]
Q7)
(i) Why do covalent
compounds occur in the form of gases or liquids?
Ans:-
Covalent compounds tend to occur in the form of gases or
liquids at room temperature and standard pressure due to the nature of their
molecular interactions and the strength of their intermolecular forces. These
factors play a significant role in determining the physical state of covalent
compounds. Here's why covalent compounds often exist as gases or liquids:
1. Intermolecular Forces: In covalent compounds, molecules
are held together by covalent bonds, which are strong intramolecular forces.
However, the intermolecular forces between covalent molecules are generally
weaker. These intermolecular forces, such as van der Waals forces (London
dispersion forces, dipole-dipole forces, and hydrogen bonding), become more
significant as molecules come closer but are relatively weaker than ionic or
metallic bonds.
2. Weak Intermolecular Attraction: The relatively weak
intermolecular forces in covalent compounds result in low forces of attraction
between molecules. As a result, it takes less energy to overcome these forces
and convert the compound into a liquid or gas state.
3. Molecular Size and Shape: The size and shape of covalent
molecules can influence the strength of intermolecular forces. Larger molecules
with more electrons have stronger London dispersion forces. Additionally,
molecules with polar covalent bonds can exhibit dipole-dipole interactions.
4. Electronegativity: The electronegativity difference
between atoms in covalent compounds can influence the polarity of the molecule.
Polar molecules experience stronger dipole-dipole interactions, which can
affect their physical state.
5. Temperature and Pressure: Covalent compounds typically
have lower melting and boiling points compared to ionic compounds due to the
weaker forces between molecules. This makes it easier for covalent compounds to
exist as gases or liquids at room temperature and standard pressure.
6. Examples: Many simple covalent compounds, such as
diatomic gases (e.g., H2, O2, N2) and small molecular compounds (e.g., CH4,
CO2, H2O), exist as gases or liquids at room temperature and standard pressure
due to the factors mentioned above.
It's important to note that while many covalent compounds
exist as gases or liquids, there are exceptions. Some covalent compounds can
have higher molecular weights, complex structures, or strong intermolecular
forces that lead to them being solids at room temperature, and standard pressure
(e.g., wax, polymers). Additionally, if the molecules has less intermolecular force and weights, then it will be liquid and if it has very small intermolecular force and weights then that will be the gasses.
(ii) Why do covalent compounds have low melting and boiling
points?
Ans:-
Covalent compounds generally have low melting and boiling
points compared to ionic compounds or metals. This characteristic can be
attributed to the nature of covalent bonding and the interactions between
molecules in covalent compounds. Here are the main reasons why covalent
compounds tend to have low melting and boiling points:
1. Weak Intermolecular Forces: Covalent compounds are held
together by strong covalent bonds within molecules, where atoms share
electrons. However, the forces of attraction between individual molecules
(intermolecular forces) are weaker than the forces between ions in ionic
compounds or the metallic bonds in metals. These weak intermolecular forces
make it easier to break apart the molecules and convert the compound from solid
to liquid or from liquid to gas.
2. Dispersion Forces: The dominant intermolecular force in
nonpolar covalent compounds is the London dispersion force, also known as van
der Waals forces. These forces arise from temporary fluctuations in electron
density that create temporary dipoles in molecules. Since dispersion forces are
generally weaker for smaller molecules, covalent compounds with smaller
molecular sizes tend to have lower melting and boiling points.
3. Dipole-Dipole Interactions: Polar covalent compounds have
dipole-dipole interactions, where the positive end of one molecule attracts the
negative end of another molecule. While these interactions are stronger than
dispersion forces, they are still relatively weak, contributing to the low
melting and boiling points of covalent compounds.
4. Hydrogen Bonding: Some covalent compounds, particularly
those containing hydrogen bonded to highly electronegative atoms like oxygen,
nitrogen, or fluorine, can exhibit hydrogen bonding. Hydrogen bonding is a
strong type of dipole-dipole interaction. However, even with hydrogen bonding,
the overall intermolecular forces in covalent compounds are weaker than the
strong electrostatic forces in ionic compounds.
5. Molecular Size and Shape: Larger covalent molecules have
more electrons, leading to stronger London dispersion forces. Additionally, the
shape of the molecules can affect the strength of intermolecular forces.
However, in most cases, the intermolecular forces in covalent compounds remain
weaker than those in ionic compounds or metals.
6. Energy Required: Since covalent compounds have weak
intermolecular forces, relatively low amounts of energy are needed to overcome
these forces and convert the compound from one phase to another (solid to
liquid or liquid to gas). This results in low melting and boiling points.
Overall, the combination of weak intermolecular forces,
dispersion forces, dipole-dipole interactions, and the absence of strong
metallic or ionic bonds contributes to the low melting and boiling points of
covalent compounds, because it needs very less energy to break the bonds.
(iii) Why are covalent compounds either bad or poor
conductors of electricity?
Ans:-
Covalent compounds are generally poor conductors of
electricity due to their molecular structure and the nature of the electrons
involved in covalent bonding. The key factors that contribute to this lack of
conductivity include:
1. Lack of Free Electrons: In covalent compounds, atoms
share electrons to form stable molecules. However, these shared electrons are
localized between the bonded atoms and are not free to move throughout the
entire material. Unlike in metals, where there are free electrons that can
easily move and carry electric charge, covalent compounds lack a
"sea" of mobile electrons.
2. Non-Ionization: Covalent compounds do not readily
dissociate into ions when they are dissolved in water or melted. Conductivity
in solutions or melts is often facilitated by the presence of ions that can
move and carry electric charge. In covalent compounds, the strong covalent
bonds between atoms prevent easy ionization.
3. Insulating Nature: Covalent compounds are typically
composed of molecules with tightly held electrons. This makes them insulators
rather than conductors. The electrons within covalent molecules are involved in
strong covalent bonds and are not easily available to participate in the
conduction of electric current.
4. Electronegativity Differences: Covalent compounds often
involve atoms with similar or identical electronegativities. This means that
the shared electrons are held relatively evenly between the atoms, leading to
non-polar covalent bonds. Non-polar covalent compounds do not have charged
regions that can facilitate the movement of electrons.
5. Exception of Ionic-Covalent Mixtures: While most pure
covalent compounds are poor conductors, there are exceptions. Some covalent
compounds can conduct electricity if they are chemically modified to become
ionized or if they are mixed with ionic compounds to create a conducting
solution.
It's important to note that there are instances where covalent compounds can conduct electricity, but these situations usually involve complex factors such as the presence of mobile ions, the formation of charged species through chemical reactions, or the introduction of specific impurities.
Overall, the absence of free-moving electrons and the strong
covalent bonds characteristic of these compounds are the primary reasons for
their poor or non-existent conductivity.
Q8) Choose form the
list below, the substances that have
(i)single covalent bond only (ii)double covalent bond only (iii)single and double covalent bonds (iv) triple covalent bond only (v) single and triple covalent bonds.
Ans:-
Here's the classification of the given substances based on
the types of covalent bonds they have:
(i) Single Covalent Bond Only:
1. H2O (water)
2. C2H6 (ethane)
3. O2 (oxygen)
4. N2 (nitrogen)
5. C2H2 (acetylene)
(ii) Double Covalent Bond Only:
1. C2H4 (ethylene)
(iii) Single and Double Covalent Bonds:
1. Cl2 (chlorine) - Has a single covalent bond in the Cl2
molecule, but it can also form double and triple bonds in other compounds.
2. C2H2 (acetylene) - Contains both single and triple
covalent bonds.
(iv) Triple Covalent Bond Only:
1. C2H2 (acetylene)
(v) Single and Triple Covalent Bonds:
1. C2H2 (acetylene) - Contains both single and triple
covalent bonds.
Q9) Classify the following as ionic or covalent compounds:
(i) Hydrogen chloride,
(ii) Sodium chloride, (iii) Ammonia (iv) Carbon tetrachloride (v) Methane (vi) Water.
Ans:-
|
Ionic bond |
Covalent bond |
|
Sodium
chloride, |
Hydrogen
chloride, Ammonia, Carbon tetrachloride, Methane, Water. |
