Chapter 6 of the first-year chemistry course, titled “Chemical Bonding,” is a foundational and crucial topic in the study of chemistry. This chapter delves into the fundamental principles governing how atoms interact and form chemical compounds. It explores the various types of chemical bonds that exist, including ionic, covalent, and metallic bonds, and elucidates the key concepts related to these bonds. Students learn about the sharing and transfer of electrons between atoms to achieve stable electron configurations.
Moreover, the chapter covers topics like Lewis dot structures, electronegativity, and the octet rule, which are essential for understanding the nature of chemical compounds and predicting their properties. Overall, Chapter 6 provides a solid foundation for students to comprehend the intricacies of chemical bonding, paving the way for a deeper understanding of chemical reactions and the behavior of matter in subsequent chapters of their chemistry course.
Short Questions Chemistry 1st Year Chapter 6 Chemical Bonding
What is a chemical bond, and what does it hold together?
A chemical bond is the force that holds together two or more atoms or ions to form various compounds.
Why do noble gases like helium and neon show little tendency to react chemically with other elements?
Noble gases have stable electronic configurations in their valence shells, with a maximum of eight electrons, making them highly stable and unreactive.
What is the octet rule, and how does it relate to the tendency of atoms to form chemical bonds?
The octet rule states that atoms tend to gain, lose, or share electrons to achieve a maximum of eight electrons in their valence shells. It explains the tendency of atoms to form chemical bonds.
Describe the energetics of bond formation between two hydrogen atoms.
When two hydrogen atoms approach each other, attractive forces decrease the potential energy of the system. At a bond distance of 75.4 pm, the potential energy is minimized, and a stable molecule (H2) is formed. The energy released during bond formation is 436.45 kJ/mol.
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What happens when repulsive forces dominate over attractive forces between two atoms?
When repulsive forces are dominant, the energy of the system increases, leading to instability, and a chemical bond is not formed.
What is the bond distance, and why is it important in the formation of stable molecules?
The bond distance is the distance between two atoms at which the potential energy of the system is minimized, resulting in the formation of a stable molecule. It is also known as bond length or compromise distance.
How much energy is released during bond formation, and what is this energy called?
During bond formation of hydrogen atoms, 436.45 kJ/mol of energy is released, and this energy is referred to as bond formation energy. The same amount of energy is required to break the bond.
Why is it important to know the relative sizes of atoms when studying bonding?
Understanding the relative sizes of atoms is crucial for understanding bonding because it influences the balance between attractive and repulsive forces between atoms, ultimately determining whether a bond can form.
What is the atomic radius, and how is it defined?
The atomic radius is defined as the average distance between the nucleus of an atom and its outermost electronic shell.
Why is it difficult to determine the exact atomic radius of an atom?
The atomic radius cannot be determined precisely due to the absence of a sharp boundary for an atom and the influence of neighboring atoms, which can cause variations in atomic size in different compounds.
How are atomic radii determined?
Atomic radii are determined by measuring the distances between the centers of adjacent atoms using methods such as X-ray diffraction or spectroscopic measurements.
How does atomic radius change across a period and down a group in the periodic table?
Atomic radii generally decrease from left to right across a period due to the increase in nuclear charge, which leads to a stronger pull on electrons. In contrast, atomic radii increase from top to bottom in a group because of the addition of more electron shells and increased shielding effects.
What are ionic radii, and how do they compare to atomic radii?
Ionic radii refer to the size of ions and can be smaller or larger than the atomic radii of the corresponding elements. Cations have smaller ionic radii than the atomic radii of their parent atoms, while anions have larger ionic radii.
Why do cations have smaller ionic radii than their parent atoms?
Cations have smaller ionic radii because they have lost one or more electrons, resulting in a greater effective nuclear charge that attracts the remaining electrons more strongly, causing the electron cloud to contract.
Why do anions have larger ionic radii than their parent atoms?
Anions have larger ionic radii because they have gained one or more electrons, leading to increased electron-electron repulsion in the outermost electron shell, causing the electron cloud to expand.
How can the interionic distance in a crystal lattice be calculated?
The interionic distance ‘R’ in a crystal lattice is equal to the sum of the cationic radius (r+) and the anionic radius (r-), expressed as R = r+ + r-.
How did Linus Pauling contribute to our understanding of ionic radii?
Linus Pauling determined the relationship between the interionic distance ‘R’ and the radii of cations and anions in crystal lattices. He found that R could be calculated as the sum of the radii of the two ions and extended this concept to calculate the radii of various ions.
Do the trends in ionic radii follow the same patterns as atomic radii in the periodic table?
Yes, the trends in ionic radii generally follow the same patterns as atomic radii in the periodic table. However, it’s important to note that ionic radii for metals refer to positive ions, while elements of group VA to VIIA represent negative ions.
What is the covalent radius of hydrogen, and how is it determined?
The covalent radius of hydrogen is 37.7 pm, and it is determined as half of the single bond length (75.4 pm) between two H atoms in an H-H molecule.
How can you calculate the covalent radius of an atom using bond length data?
To calculate the covalent radius of an atom, subtract the known bond length of the atom in a molecule from the bond length data. For example, the covalent radius of C-atom can be calculated as (176.7 pm – 99.4 pm) = 77.3 pm for the C-Cl bond in CH3Cl.
How does the variation of covalent radii relate to atomic radii in the periodic table?
The variation of covalent radii in groups and periods is almost the same as that of atomic radii.
What is ionization energy, and how is it defined?
Ionization energy is the minimum energy required to remove an electron from a gaseous atom to form an ion. It is the energy required for ionization, such as Mg → Mg+ + e^-.
What factors influence ionization energies of atoms?
The ionization energies of atoms depend on several factors, including:
(i) Atomic radius of the atom
(ii) Nuclear charge or proton number of the atom
(iii) Shielding effect of inner electrons
(iv) Nature of the orbital
How does ionization energy vary in the periodic table in terms of periods and groups?
In a period, ionization energies increase from left to right due to an increase in proton number until a maximum value is reached at the end of the period. In groups, ionization energies decrease despite an increase in proton number because valence electrons are placed at a larger distance from the nucleus, and there is an increasing shielding effect.
Which groups in the periodic table exhibit an abnormal trend in ionization energies?
Groups III-A and VI-A exhibit an abnormal trend in ionization energies due to the distribution of electrons.
What is the second ionization energy, and how does it relate to the first ionization energy?
The second ionization energy is the energy required to remove a second electron from an atom, following the removal of the first electron. It is higher than the first ionization energy because the second electron is removed from a positively charged ion rather than a neutral atom.
How does ionization energy change with the number of electrons to be removed?
Ionization energy values increase as the number of electrons to be removed increases. This is because electrons become more tightly held by the dominant positive charge in the ion as more electrons are removed.
What does ionization energy indicate about the metallic character of elements?
Elements with low ionization energies are metals, while those with high ionization energies are non-metals. Elements with intermediate values are typically metalloids.
How can the gaps in first and second ionization energies help determine the valency of an element?
If there is a sufficient gap between the first and second ionization energies of an element, it indicates that the element has a valency of one.
What is electron affinity, and how is it related to the attraction of the nucleus for an extra electron?
Electron affinity is the energy released when an electron is added to an isolated gaseous atom in its valence energy level to form an anion with a unit negative charge. It is a measure of the attraction of the nucleus for an extra electron, and it is given a negative sign because energy is released in the process.
What factors influence electron affinity?
Electron affinity is influenced by factors such as atomic radius, nuclear charge, and the shielding effect of inner electrons. Generally, electron affinities decrease as atomic radius increases because the force of attraction between valence electrons and the nucleus decreases with larger atomic size.
Why does the atomic radius decrease from left to right in a period of the periodic table?
The atomic radius decreases from left to right in a period due to an increase in the nuclear charge.
Why do electron affinities of elements increase from left to right in the periodic table?
Electron affinities of elements increase from left to right in the periodic table because of the decreasing atomic radius and the stronger attraction between the nucleus and electrons.
Which groups of elements have the lowest electron affinities and the highest electron affinities, respectively?
The alkali metals have the lowest electron affinities, while the halogens have the highest electron affinities.
Why do atomic radii increase in groups of the periodic table?
Atomic radii increase in groups of the periodic table due to the successive increase of electronic shells and the shielding effect on the force of attraction between the nucleus and valence electrons.
What causes exceptions to the general trend of increasing electron affinity from left to right in the periodic table?
Exceptions to the trend of increasing electron affinity can occur due to factors such as electron cloud thickness and repulsion of incoming electrons. For example, fluorine has lower electron affinity than chlorine despite being to the right of chlorine in the periodic table.
Why do elements of groups IIA, VA, and VIII show abnormally low values in every period of the periodic table?
Elements of groups IIA, VA, and VIII show abnormally low values in every period due to their electronic configurations, which result in unique properties.
What is electronegativity, and how is it related to ionization energy and electron affinity?
Electronegativity is the tendency of an atom to attract a shared electron pair towards itself. It is related to ionization energy and electron affinity, with elements having higher electronegativity values being more likely to attract electrons in a chemical bond.
How is electronegativity measured, and what element serves as the standard with an arbitrary value of 4.0?
Electronegativity is measured using the Pauling scale, with fluorine assigned an arbitrary standard value of 4.0, making it the most electronegative element.
How do electronegativity values change across a period and within a group in the periodic table?
Electronegativity values generally increase across a period (from left to right) and decrease within a group (top to bottom) in the periodic table.
What does a difference in electronegativity values between two bonded atoms indicate about the bond?
The difference in electronegativity values between bonded atoms indicates the polar nature of the covalent bond. A difference of 1.7 units or more suggests roughly equal contributions of ionic and covalent bonds.
How are polar and non-polar bonds related to electronegativity differences?
Polar bonds result from differences in electronegativity values between bonded atoms, while non-polar bonds occur when the electronegativity difference is zero, indicating equal sharing of electrons.
What is the Lewis concept, and how does it help us understand the tendencies of elements?
The Lewis concept is a chemical theory that helps us understand the tendencies of elements to form bonds with each other by focusing on the sharing or transfer of electrons.
How is an ionic bond formed according to the Lewis theory?
An ionic bond is formed by the complete transfer of electrons from an atom with low ionization energy to another atom with high electron affinity.
Can you provide an example of ionic bond formation using Lewis theory?
Sure, consider the formation of potassium chloride (KCl). Potassium (K) loses its outermost electron to form a K+ ion, and chlorine (Cl) gains this electron to form a Cl- ion, resulting in the strong electrostatic attraction between them.
What is lattice energy, and how does it relate to ionic bonds?
Lattice energy is the energy released when oppositely charged ions, such as K+ and Cl-, arrange themselves in a crystal lattice. It is a measure of the strength of ionic bonding in a compound.
How do Lewis structures help us understand covalent bonds?
Lewis structures illustrate covalent bonds by showing the sharing of electrons between two atoms to complete their valence shells and attain the nearest inert gas configuration.
What is the difference between non-polar and polar covalent bonds?
In non-polar covalent bonds, the bonding electron pairs are equally shared, leading to an even distribution of charge between the bonded atoms. In polar covalent bonds, there is an uneven sharing of electrons, resulting in partial charges on the atoms.
Can you provide an example of a non-polar covalent molecule?
Hydrogen gas (H2) and chlorine gas (Cl2) are examples of non-polar covalent molecules because the electrons in the covalent bonds are equally shared.
Why is carbon tetrachloride (CCl4) considered non-polar overall despite having polar C-Cl bonds?
CCl4 is considered non-polar overall because the polar C-Cl bonds cancel each other out due to the molecule’s symmetric tetrahedral structure.
What factors determine whether a bond is ionic or covalent?
The difference in electronegativity between bonded atoms is a key factor. If the electronegativity difference is 1.7 or more, the bond is considered ionic.
Is it possible for a bond to have 100% ionic character according to Lewis theory?
According to Lewis theory, there is no bond with 100% ionic character; even highly ionic compounds like NaCl have some covalent character.
What happens when two different atoms are joined by a covalent bond?
Answer: When two different atoms are joined by a covalent bond, the electron pair is not equally shared between the bonded atoms, and it is displaced toward the more electronegative atom.
Why does the displacement of electrons in a covalent bond make one end of the molecule partially positive and the other partially negative?
The displacement of electrons in a covalent bond results in one end of the molecule becoming partially positive, and the other end becoming partially negative due to differences in electronegativity between the bonded atoms.
Give an example of a polar covalent molecule.
Methanol is an example of a polar covalent molecule because it contains a polar bond.
What are double and triple bonds in covalent compounds?
Double and triple bonds in covalent compounds involve the sharing of two and three pairs of electrons, respectively, between the bonded atoms.
Which inert gas has a strong triple bond?
Nitrogen (N2) is an inert gas that has a strong triple bond.
What is the structure of carbon dioxide, and why is it considered a non-polar covalent compound?
Carbon dioxide (CO2) has a linear structure, which balances the polar character on both sides of the carbon atom, making it a non-polar covalent compound.
What are the examples of large-sized molecules formed by the mutual sharing of electrons between non-metallic atoms?
Diamond, graphite, and SiC are examples of large-sized molecules formed by the mutual sharing of electrons between carbon atoms and other non-metallic atoms.
What are the three types of covalent bonds that carbon can form in alkanes, alkenes, and alkynes?
Carbon can form single, double, and triple covalent bonds in alkanes, alkenes, and alkynes, respectively.
What is a coordinate covalent bond, and how is it formed?
A coordinate covalent bond is formed when one of the bonded atoms donates a pair of electrons to the other atom. It results from the sharing of electrons, with one atom providing both electrons.
Give an example of a molecule that contains coordinate covalent bonds.
NH3 (ammonia) can form coordinate covalent bonds with BF3 (boron trifluoride), where the lone pair of electrons on nitrogen is donated to BF3 to create the bond.
What is the main limitation of the Lewis Model in explaining covalent bond formation?
The main limitation of the Lewis Model is its oversimplification and inability to justify molecular shapes, bond polarities, bond distances, and various energy transitions during chemical reactions.
How does the Valence Shell Electron Pair Repulsion Theory (VSEPR Theory) address the limitations of the Lewis Model?
The VSEPR Theory addresses the limitations of the Lewis Model by explaining the shapes of molecules based on the arrangement of electron pairs in the outer orbit of the central atom. It takes into account both lone pairs and bond pairs and their repulsions to determine molecular geometry.
What are the postulates of the Valence Shell Electron Pair Repulsion Theory (VSEPR Theory)?
The postulates of the VSEPR Theory are:
(i) Both lone pairs and bond pairs participate in determining molecular geometry.
(ii) Electron pairs are arranged around the central atom to maximize their distance and minimize repulsions.
(iii) Lone pairs occupy more space than bond pairs, leading to greater repulsive forces.
(iv) Double and triple bonds are treated as if they occupy the same region as single bonds in determining molecular geometry.
What is the order of repulsions between electron pairs in a molecule according to the VSEPR Theory?
The order of repulsions between electron pairs in a molecule, as per the VSEPR Theory, is as follows: Lone pair – lone pair > Lone pair – bond pair > Bond pair – bond pair. These repulsions are referred to as van der Waals repulsions.
What is the arrangement of two electron pairs around the central atom in AB2 type molecules?
The two electron pairs are arranged at a farther distance apart at an angle of 180°, forming a linear geometry.
Give an example of a linear molecule that contains two electron pairs.
Beryllium chloride (BeCl2) is a typical linear molecule with two electron pairs. Other examples include MgCl2, CaCl2, SrCl2, CdCl2, and HgCl2.
How many electrons are present in the outermost orbitals of central atoms in AB2 type molecules?
Central atoms in AB2 type molecules have two electrons in their outermost orbitals.
What is the mutual angle between three bonding electron pairs in AB3 type molecules with no lone pairs?
The mutual angle between three bonding electron pairs in such molecules is 120°, resulting in a triangular planar geometry.
Give an example of a molecule with a trigonal planar geometry and no lone pairs.
BH3 (boron hydride) is an example of a molecule with a trigonal planar geometry and no lone pairs.
Which group III-A hydrides and halides are expected to have similar geometries to BH3?
Group III-A hydrides (AlH3, GaH3, InH3) and their halides (BF3, AlCl3, etc.) are expected to have similar trigonal planar geometries.
In AB3 type molecules with one lone pair and two bond pairs, what happens to the structure in the vapor phase?
The structure becomes distorted, as one corner of the triangular arrangement is occupied by a lone pair.
Give an example of a molecule with AB3 type geometry that includes multiple bonds.
SO2 has one corner of the triangular arrangement occupied by a lone pair and two corners by S=O double bonds. SO3 has all three regions occupied by S=O bonds, forming a perfectly triangular structure.
What is the ideal bond angle in AB4 type molecules with no lone pairs?
The ideal bond angle in AB4 type molecules with no lone pairs is 109.5°, resulting in a regular tetrahedral shape.
Name a molecule that exhibits a regular tetrahedral geometry with no lone pairs.
Methane (CH4) is an example of a molecule with a regular tetrahedral geometry and no lone pairs.
How does the presence of one lone pair affect the geometry in AB4 type molecules?
The presence of one lone pair causes compression of bond angles in the molecule, leading to a non-ideal bond angle (e.g., 107.5° in ammonia, NH3).
In which molecule does the bond angle reduce to 102° due to strong polarity of N-F bonds and repulsion effects?
In NF3, the bond angle reduces to 102° due to the strong polarity of N-F bonds and the repulsion between electrons.
What types of repulsion are introduced when a molecule has two lone pairs and two bond pairs?
Presence of two lone pairs introduces three types of repulsion: lone pair-lone pair, lone pair-bond pair, and bond pair-bond pair repulsion.
Why does water (H2O) have a bent or angular geometry instead of a linear AB2 type like BeCl2 and CO2?
Water’s bent geometry is due to the participation of lone pairs, in addition to bond pairs, in determining its overall geometry, as explained by VSEPR theory.
What is the bond angle in the water molecule (H2O)?
The bond angle in the water molecule is approximately 104.5 degrees.
What is the essential condition for chemical bonding according to Valence Bond Theory (VBT)?
The essential condition for chemical bonding, according to VBT, is that the orbitals of atoms participating in bond formation must overlap, and the direction of the bond is determined by the direction of the two overlapping orbitals.
What type of bond is formed when two partially filled atomic orbitals overlap with maximum electron density along the line joining the two nuclei?
When two partially filled atomic orbitals overlap with maximum electron density along the line joining the two nuclei, a sigma (σ) bond is formed.
How is a pi (π) bond formed in covalent bonding?
A pi (π) bond is formed by side-to-side overlap of p orbitals, with the greatest electron density lying above and below the internuclear axis.
What is the bond angle in the H2S molecule?
The bond angle between the two H-S bonds in the H2S molecule is approximately 92 degrees.
How is the geometry of the hydrogen sulfide (H2S) molecule explained by Valence Bond Theory (VBT)?
The VBT explains the geometry of the hydrogen sulfide (H2S) molecule by the overlap of two 3p orbitals of sulfur with the 1s orbitals of hydrogen atoms, resulting in a non-linear geometry.
What is the concept of atomic orbital hybridization?
Atomic orbital hybridization is the process where atomic orbitals with slightly different energies intermix to form new hybrid atomic orbitals with unique shapes and geometries, which are different from the parent atomic orbitals.
How does hybridization explain the valency of elements and molecular shapes?
Hybridization provides a satisfactory explanation for the valency of elements and determines the shape and geometry of molecules by rearranging valence orbitals through the mixing of atomic orbitals.
What happens during sp3 hybridization, and what is its significance?
In sp3 hybridization, one s orbital and three p orbitals combine to form four equivalent sp3 hybrid atomic orbitals. This process is significant as it explains the structures of molecules like CH4, NH3, and H2O and helps determine their shapes.
What is the geometry of sp3 hybridized orbitals, and how are they oriented in space?
Sp3 hybridized orbitals have a tetrahedral geometry, and they are oriented in space so that the angle between them is 109.5 degrees.
How does sp3 hybridization explain the structure of methane (CH4)?
In methane, the sp3 hybrid orbitals of carbon overlap with the 1s orbitals of four hydrogen atoms to form four sigma bonds, resulting in a tetrahedral geometry for the molecule.
What is the electronic configuration of nitrogen in its ground state?
The electronic configuration of nitrogen in its ground state is 1s² 2s² 2p³.
How does nitrogen achieve sp3 hybridization in NH3?
Nitrogen achieves sp3 hybridization in NH3 by hybridizing one 2s orbital and three 2p orbitals, resulting in four sp3 hybrid atomic orbitals.
How is the geometry of NH3 (ammonia) explained in terms of its hybridization and electron pairs?
The geometry of NH3 is pyramidal, with three hydrogen atoms forming the base and a lone pair of electrons at the apex. This geometry is explained by the repulsion between the lone pair and the bond pairs of electrons, causing a deviation from the tetrahedral angle.
What is the experimentally determined angle in ammonia (NH3)?
The experimentally determined angle in ammonia (NH3) is 107.5°.
How does the lone pair in ammonia (NH3) affect bond angles?
The lone pair in ammonia (NH3) is closer to the nitrogen nucleus than the bond pairs, leading to repulsion between them. This repulsion causes the bond angles to decrease from the ideal tetrahedral angle of 109.5° to 107.5°.
What is the electronic configuration of oxygen in its ground state?
The electronic configuration of oxygen in its ground state is 1s² 2s² 2p⁴.
How does oxygen achieve sp3 hybridization in water (H2O)?
Oxygen achieves sp3 hybridization in water (H2O) by hybridizing one 2s orbital and three 2p orbitals to form four sp3 hybrid orbitals.
What is the bond angle in water (H2O)?
The bond angle in water (H2O) is 104.5°.
How is the bent or angular structure of water (H2O) explained in terms of its hybridization and electron pairs?
The bent or angular structure of water (H2O) is explained by the repulsion between two lone pairs of electrons and bond pairs. The lone pairs are closer to the oxygen nucleus, leading to a deviation from the ideal tetrahedral angle.
What type of hybridization occurs in boron trifluoride (BF3)?
Boron trifluoride (BF3) undergoes sp2 hybridization, where one s orbital and two p orbitals of boron intermix to form three sp2 hybrid orbitals.
What is the geometry of boron trifluoride (BF3) and the bond angle?
Boron trifluoride (BF3) has a planar geometry, with bond angles of 120° between the fluorine atoms.
How is the structure of boron trifluoride (BF3) explained in terms of hybridization?
Boron trifluoride (BF3) is formed by the overlap of three half-filled sp2 hybrid orbitals of boron with the lobe-shaped p orbitals of three fluorine atoms, resulting in a triangular planar structure.
What is the electronic configuration of a carbon atom in its ground state?
The electronic configuration of a carbon atom in its ground state is 1s² 2s² 2p².
How many hybrid orbitals are formed when carbon undergoes sp2 hybridization?
When carbon undergoes sp2 hybridization, it forms three hybrid orbitals.
What is the bond angle between the hybrid orbitals in sp2 hybridization?
The bond angle between the hybrid orbitals in sp2 hybridization is 120 degrees.
How is a p-bond formed in the ethene (C2H4) molecule?
A p-bond in the ethene (C2H4) molecule is formed by the sideways overlap of two half-filled, coplanar p-orbitals on adjacent carbon atoms.
What is the geometry of a molecule formed by sp hybridization?
Molecules formed by sp hybridization have a linear geometry.
How many sigma and pi bonds are there in an ethyne (CH≡CH) molecule?
An ethyne (CH≡CH) molecule contains one sigma (σ) bond and two pi (π) bonds between the two carbon atoms.
What does the molecular orbital approach consider in chemical bonding?
The molecular orbital approach considers the entire molecule as a single unit, assuming that atomic orbitals of combining atoms overlap to form molecular orbitals that characterize the entire molecule.
What are bonding and anti-bonding molecular orbitals, and how do they differ in energy?
Bonding molecular orbitals have lower energy and are symmetrical about the molecular axis, while anti-bonding molecular orbitals have higher energy. Bonding orbitals promote bond formation, while anti-bonding orbitals discourage it.
How are electrons filled into molecular orbitals, and which principles govern this process?
Electrons are filled into molecular orbitals according to the Aufbau principle, Pauli’s exclusion principle, and Hund’s rule. These principles dictate the order and spin of electron placement in molecular orbitals.
What types of overlaps occur between p and p atomic orbitals to form molecular orbitals?
Two types of overlaps occur between p and p atomic orbitals: head-on approach, where p-orbitals approach along the same axis, and sideways approach, where the axes of p-orbitals are parallel to each other.
How are sigma (σ) bonds and pi (π) bonds formed in molecular orbitals?
Sigma (σ) bonds result from head-on overlap of atomic orbitals, while pi (π) bonds are formed from sideways overlap of p-orbitals. Sigma bonds are symmetrical about the molecular axis, whereas pi bonds have electron density above and below the nodal plane.
How are the relative energies of molecular orbitals determined, and what is the energy order for diatomic molecules like O2 and F2?
The relative energies of molecular orbitals are determined through spectroscopic measurements. In the case of diatomic molecules like O2 and F2, the energy order of molecular orbitals is s < s* < s < s* < p < p < p* < p* < s*.
Why do some diatomic molecules, such as B2, C2, and N2, show a different energy order for their molecular orbitals?
The energy order for diatomic molecules like B2, C2, and N2 differs because of the mixing of 2s and 2px atomic orbitals, leading to the acquisition of sp-character by the molecular orbitals. This mixing alters their energies, causing some orbitals to become more stable while others become less stable.
What is bond order, and how is it calculated?
Bond order is the number of bonds formed between two atoms after the atomic orbitals overlap. It is calculated as half of the difference between the number of bonding electrons and anti-bonding electrons.
How is the bond order calculated for H2 (hydrogen molecule)?
The bond order for H2 is calculated as (2 – 0) / 2, which equals 1.
What is the bond order of He2 (helium molecule)?
The bond order for He2 is zero because no bonding electrons are present.
What is the bond order of N2 (nitrogen molecule)?
The bond order for N2 is 3, indicating a triple bond consisting of one sigma and two pi bonds.
What is the bond order of O2 (oxygen molecule)?
The bond order for O2 is 2, indicating a double bond.
What is bond energy (bond enthalpy), and how is it measured?
Bond energy is the average amount of energy required to break all bonds of a particular type in one mole of a substance. It is measured experimentally by measuring the heat involved in a chemical reaction.
How is bond energy expressed?
Bond energy is typically expressed in kJ/mol, representing the energy required to break or form one mole of bonds.
What does the bond energy represent in a chemical reaction?
Bond energy represents the energy change (enthalpy change) associated with breaking or forming bonds during a chemical reaction.
What factors determine the strength of a chemical bond?
The strength of a chemical bond depends on factors such as electronegativity difference of bonded atoms, sizes of the atoms, and bond length.
How does electronegativity difference affect bond strength?
Electronegativity difference between bonded atoms influences bond strength. Greater electronegativity difference leads to more polarity and stronger attraction between atoms.
What is bond energy, and how is it calculated?
Bond energy is the measure of the strength of a bond. It can be calculated by dividing the energy required to break a certain number of bonds by that number.
How does the ionic character affect bond length?
Ionic character in a bond can lead to a shorter bond length due to the strong attraction between polar ends.
How do hybridization and orbital contributions affect bond length?
Bonds with a higher s-orbital contribution tend to have shorter bond lengths. Hybridization and orbital characteristics can influence bond length.
How does the periodic table arrangement affect bond length?
Bond length generally increases as you move down a group in the periodic table and decreases as you move from left to right in a period.
What causes a molecule to have a dipole moment?
A molecule becomes polar and has a dipole moment when there is an electronegativity difference between the bonded atoms, resulting in partial positive and negative charges separated on the bonded atoms.
How is the dipole moment of a molecule defined?
The dipole moment (m) of a molecule is defined as the product of the electric charge (q) and the distance between the positive and negative centers (r), expressed as m = q × r.
In what units are dipole moments measured?
Dipole moments are measured in Debye (D) units.
How can you calculate the percentage ionic character in a bond using the dipole moment?
The percentage ionic character in a bond can be calculated using the formul
% Ionic Character = (Observed Dipole Moment / Ionic Dipole Moment) × 100
What information can the dipole moment provide about the molecular structure?
The dipole moment provides information about the percentage ionic character of a bond and the angles between the bonds or the geometry of molecules.
Why does a linear molecule like CO2 have a dipole moment of zero?
A linear molecule like CO2 has a dipole moment of zero because the dipoles of its individual bonds are equal and opposite, canceling each other out.
What is the solubility of ionic compounds in water?
Ionic compounds are mostly soluble in water.
Why do ionic compounds dissolve in water?
Ionic compounds dissolve in water because polar water molecules detach the cations and anions from the crystal lattice by their electrostatic attraction, leading to hydration.
In which solvents are ionic compounds generally insoluble?
Ionic compounds are generally insoluble in non-polar solvents like benzene and hexane.
What types of solvents do covalent compounds typically dissolve in?
Covalent compounds typically dissolve easily in non-polar organic solvents like benzene and ether.
Why are covalent compounds mostly insoluble in water?
Covalent compounds are mostly insoluble in water because the attraction between water molecules cannot overcome the intermolecular forces of attraction in covalent compounds.
Do ionic compounds exhibit isomerism?
No, ionic compounds do not exhibit isomerism because ionic bonds are non-rigid and non-directional.
Why do covalent compounds have the potential for isomerism?
Covalent compounds have a directional and rigid nature, which allows for various forms of isomerism.
What is an example of a covalent compound that shows structural isomerism?
C2H6O exhibits structural isomerism, as seen in the case of dimethyl ether and ethanol.
Why do ionic compounds react rapidly in aqueous solutions?
Ionic compounds react rapidly in aqueous solutions because the ions are already dissociated in solution, so no bonds need to be broken before a new bond is formed.
Why are covalent compounds generally slower to react compared to ionic compounds?
Covalent compounds are generally slower to react because they involve bond breaking and making, and there is no strong electrical force to speed up the reaction, unlike in ionic reactions.
MCQ’s Chemistry 1st Year Chapter 6 Chemical Bonding
Long Questions Chemistry 1st Year Chapter 6 Chemical Bonding
Question: What is a chemical bond? Discuss the formation of ionic and covalent bonds. How does the electronegativity differences differentiate between ionic and covalent bond?
Answer: Chemical Bond
A chemical bond is a force of attraction between two or more atoms that holds them together in a stable arrangement. These bonds are formed due to the interaction of electrons between atoms, leading to the creation of chemical compounds.
Formation of Ionic Bonds
Ionic bonds form between atoms with significantly different electronegativities, typically between a metal and a non-metal. The process involves the transfer of electrons from one atom to another:
Electron Transfer: The atom with a lower electronegativity (the metal) loses one or more electrons to become a positively charged ion (cation), while the atom with higher electronegativity (the non-metal) gains these electrons to become a negatively charged ion (anion).
Attraction: The positively charged cation and negatively charged anion are attracted to each other by electrostatic forces. This attraction creates an ionic bond that holds the ions together in a crystal lattice.
Formation of Ionic Compound: The attraction between multiple cations and anions results in the formation of an ionic compound, such as sodium chloride (NaCl) or calcium oxide (CaO).
Formation of Covalent Bonds
Covalent bonds form between atoms with similar electronegativities, typically between non-metals. In this case, atoms share electrons rather than transferring them:
Electron Sharing: Two atoms come close together, and their outermost electrons are shared between them. Each atom contributes one or more electrons to a shared pair or pairs.
Attraction: The shared electrons are attracted to the positively charged nuclei of both atoms. This attraction creates a stable molecule.
Formation of Covalent Compound: The sharing of electrons results in the formation of a covalent compound, such as hydrogen (H2), oxygen (O2), or water (H2O).
Electronegativity Differences
Electronegativity is a measure of an atom’s ability to attract electrons in a chemical bond. The difference in electronegativity between two bonded atoms can help differentiate between ionic and covalent bonds:
Ionic Bonds: Ionic bonds form when there is a large electronegativity difference between two atoms. Typically, one atom has a much higher electronegativity than the other. This significant difference leads to the complete transfer of electrons, resulting in the formation of ions.
Covalent Bonds: Covalent bonds form when there is a small or negligible electronegativity difference between two atoms. In this case, the electrons are shared between the atoms, and neither atom gains or loses electrons to a significant extent.
In summary, ionic bonds result from the transfer of electrons between atoms with a significant electronegativity difference, while covalent bonds form when electrons are shared between atoms with a small electronegativity difference. Electronegativity differences help classify the type of bond that will likely form in a given chemical interaction.
Question: Define ionization energy and electron affinity. How these quantities change in the periodic table. What factors are responsible for their variation?
Answer: Ionization Energy
Ionization energy is the energy required to remove an electron from an atom or ion in its gaseous state. It’s typically measured in units like kilojoules per mole (kJ/mol) or electronvolts (eV). Ionization energy generally follows the following trends on the periodic table:
Periodic Trend (Left to Right): Ionization energy tends to increase as you move from left to right across a period (horizontal row) of the periodic table. This is because the effective nuclear charge (the positive charge felt by the outermost electrons) increases due to more protons in the nucleus. As a result, it becomes harder to remove an electron.
Periodic Trend (Top to Bottom): Ionization energy generally decreases as you move down a group (vertical column) of the periodic table. This is because the outermost electrons are located in higher energy levels (shells) as you go down a group. These electrons are farther from the nucleus and experience weaker attraction, making them easier to remove.
Electron Affinity
Electron affinity is the energy change that occurs when an atom or ion gains an electron to form a negative ion (anion) in its gaseous state. Like ionization energy, electron affinity is measured in units like kJ/mol or eV. Electron affinity trends on the periodic table are as follows:
Periodic Trend (Left to Right): Electron affinity generally becomes more negative (i.e., more exothermic) as you move from left to right across a period. This means that atoms on the right side of the periodic table tend to have a greater affinity for gaining electrons. This trend is due to the increasing effective nuclear charge, which leads to stronger attraction for additional electrons.
Periodic Trend (Top to Bottom): Electron affinity does not show a strong and consistent trend as you move down a group. In some cases, it becomes more negative, while in others, it remains relatively constant or becomes less negative. Factors such as electron repulsions and electron shielding can influence this trend.
Factors Responsible for Variation
Several factors contribute to the variation in ionization energy and electron affinity:
Effective Nuclear Charge: The effective nuclear charge, which is the net positive charge experienced by the outermost electrons, plays a significant role. As it increases across a period, it makes it harder to remove electrons and easier to attract additional electrons.
Atomic Radius: The size of the atom or ion also affects ionization energy and electron affinity. Smaller atoms have higher ionization energies and greater electron affinities because their electrons are closer to the nucleus.
Electron Shielding: The presence of inner electron shells (core electrons) can shield the outermost electrons from the full effect of the nucleus, reducing ionization energy and influencing electron affinity.
Electron Repulsions: Electron-electron repulsions between negatively charged electrons can affect both ionization energy and electron affinity. Repulsions can make it easier or harder to remove or add electrons, depending on the specific electron configuration.
Question: Explain, what do you understand by the term electronegativity? Discuss its variations in the periodic table. How does it affect the bond strengths?
Answer: Electronegativity is a chemical property that describes the ability of an atom to attract electrons in a chemical bond within a molecule. In other words, it quantifies how strongly an atom pulls or attracts electrons towards itself when it is part of a chemical compound. Electronegativity is an important concept because it helps predict the nature of chemical bonds and the polarity of molecules.
Variations in Electronegativity across the Periodic Table
Electronegativity values tend to increase across a period (from left to right) and decrease down a group (from top to bottom) in the periodic table. Several factors contribute to these trends:
Atomic Size: Across a period, atomic size generally decreases, which leads to a stronger pull on electrons by the nucleus. This increase in effective nuclear charge results in higher electronegativity values.
Nuclear Charge: The number of protons in the nucleus of an atom determines its nuclear charge. As you move across a period, the nuclear charge increases, causing greater electronegativity.
Shielding Effect: Down a group, the number of electron shells increases, providing more electron shielding for outer electrons. This reduces the effective nuclear charge experienced by outer electrons, leading to lower electronegativity values.
Electron Configuration: Elements with a nearly full or completely full outer electron shell are less likely to attract additional electrons. Thus, noble gases have very low electronegativity values.
Effect of Electronegativity on Bond Strengths
Electronegativity differences between atoms in a chemical bond play a crucial role in determining the strength and type of bond formed:
Ionic Bonds: When there is a large electronegativity difference between atoms, typically between a metal and a non-metal, one atom gains electrons (becomes negatively charged) while the other loses electrons (becomes positively charged). The resulting electrostatic attraction between oppositely charged ions forms an ionic bond. The greater the electronegativity difference, the stronger the ionic bond.
Covalent Bonds: In covalent bonds, atoms with similar electronegativities share electrons. The strength of covalent bonds depends on the number of shared electrons and the distance between the nuclei. Generally, larger electronegativity differences result in more polar covalent bonds, where electrons are unevenly shared, and the bond has partial ionic character. Smaller electronegativity differences lead to nonpolar covalent bonds, where electrons are more evenly shared.
Polarity: Electronegativity differences also determine the polarity of molecules. If there is a significant electronegativity difference within a molecule, it results in a polar molecule with regions of partial positive and negative charges. In contrast, molecules with minimal electronegativity differences are nonpolar.
Question: Explain qualitatively the valence bond theory. How does it differ from molecular orbital theory?
Answer: Valence Bond Theory (VBT)
Valence Bond Theory is a qualitative model used to explain the formation of chemical bonds in molecules. It was developed by Linus Pauling and others in the early 20th century. VBT focuses on the concept of atomic orbitals overlapping to form bonds. Here’s a qualitative explanation of the key points of Valence Bond Theory:
Atomic Orbitals: In VBT, chemical bonds are formed when the atomic orbitals of two or more atoms overlap in space. These atomic orbitals contain the electrons that participate in bond formation.
Overlap and Bond Formation: When atomic orbitals overlap, the electrons in these orbitals are shared between the participating atoms. This sharing of electrons results in the formation of a chemical bond.
Types of Overlap: VBT distinguishes between different types of orbital overlap, such as sigma (σ) bonds and pi (π) bonds. Sigma bonds are formed by the direct head-on overlap of orbitals, while pi bonds result from the side-to-side overlap of orbitals.
Bond Strength and Directionality: VBT explains bond strength based on the extent of orbital overlap. The greater the overlap, the stronger the bond. It also considers the directionality of bonds, particularly in molecules with multiple bonds or geometric shapes.
Molecular Orbital Theory (MO Theory)
Molecular Orbital Theory is another qualitative and quantitative model used to describe chemical bonding. Unlike VBT, MO Theory focuses on the behavior of electrons in molecular orbitals, which are formed by the combination of atomic orbitals from all the atoms in a molecule. Here are some key differences between MO Theory and VBT:
Formation of Molecular Orbitals: In MO Theory, atomic orbitals from all atoms in a molecule combine to form molecular orbitals. These molecular orbitals can extend over the entire molecule, not just between two atoms as in VBT.
Mathematical Approach: MO Theory employs mathematical methods to calculate molecular orbital energies and electron distributions. It provides a more quantitative description of electronic structure and bonding.
Electron Behavior: MO Theory treats electrons as occupying molecular orbitals, not localized bonds between specific atoms. Electrons are described as delocalized and can be found in various regions of the molecule.
Bonding and Antibonding Orbitals: MO Theory introduces the concept of bonding and antibonding orbitals. When atomic orbitals combine constructively, they form bonding orbitals that stabilize the molecule. Conversely, when they combine destructively, they form antibonding orbitals that weaken the molecule.
Predictive Power: MO Theory is often more accurate than VBT in predicting molecular properties, such as bond energies, bond lengths, and magnetic properties.
In summary, Valence Bond Theory is a qualitative model that focuses on the overlap of atomic orbitals and provides a localized view of chemical bonds, whereas Molecular Orbital Theory is a more comprehensive model that considers the combination of atomic orbitals to form molecular orbitals, providing a more quantitative and delocalized view of electronic structure and bonding in molecules. Each theory has its strengths and limitations and is used depending on the specific context and level of detail required for understanding chemical bonding.
Question: How the bonding in the following molecules can be explained with respect to valence bond theory? Cl2, O2, N2, HF, H2S.
Answer: Valence Bond Theory (VBT) can be used to explain the bonding in the following molecules:
Cl2 (Chlorine Gas)
In Cl2, two chlorine atoms come together to form a diatomic molecule. According to VBT, each chlorine atom has seven valence electrons, with three in the 3p subshell.
When they bond, one 3p orbital from each chlorine atom overlaps head-on to form a sigma (σ) bond. This sigma bond consists of two electrons shared between the two atoms.
Since no other orbitals overlap, Cl2 has a single sigma bond.
O2 (Oxygen Gas)
In O2, two oxygen atoms form a diatomic molecule. According to VBT:
Each oxygen atom has six valence electrons, with two in the 2p subshell.
When they bond, one 2p orbital from each oxygen atom overlaps head-on to form a sigma (σ) bond. This sigma bond consists of two electrons shared between the two atoms.
Additionally, two more electrons from each oxygen atom are involved in a pi (π) bond formed by the side-to-side overlap of unhybridized 2p orbitals. This results in a total of one sigma bond and one pi bond in O2.
N2 (Nitrogen Gas)
In N2, two nitrogen atoms form a diatomic molecule. According to VBT:
Each nitrogen atom has five valence electrons, with one in the 2p subshell.
When they bond, one 2p orbital from each nitrogen atom overlaps head-on to form a sigma (σ) bond. This sigma bond consists of two electrons shared between the two atoms.
Additionally, two more electrons from each nitrogen atom are involved in a pi (π) bond formed by the side-to-side overlap of unhybridized 2p orbitals. This results in a total of one sigma bond and one pi bond in N2.
HF (Hydrogen Fluoride)
In HF, hydrogen and fluorine form a molecule. According to VBT
Hydrogen has one valence electron in a 1s orbital, while fluorine has seven valence electrons, with one in the 2p subshell.
When they bond, the 1s orbital of hydrogen overlaps with one of the 2p orbitals of fluorine to form a sigma (σ) bond. This sigma bond consists of two electrons shared between the two atoms.
H2S (Hydrogen Sulfide)
In H2S, two hydrogen atoms and one sulfur atom form a molecule.
According to VBT, Hydrogen has one valence electron in a 1s orbital, while sulfur has six valence electrons, with two in the 3p subshell.
Each hydrogen atom forms a sigma (σ) bond with one of the 3p orbitals of sulfur. These sigma bonds consist of two electrons each, resulting in two sigma bonds in H2S.
In all of these molecules, Valence Bond Theory explains the formation of sigma and, where applicable, pi bonds based on the overlap of atomic orbitals. This theory helps us understand the nature of the chemical bonds and the arrangement of electrons in these molecules.
Question: Explain VSEPR theory. Discuss the structures of CH4, NH3, H2O, BeCl2, BF3,S02, SO3 with reference to this theory.
Answer: VSEPR Theory (Valence Shell Electron Pair Repulsion Theory):
VSEPR theory is a model used to predict the molecular geometry and shape of molecules and polyatomic ions. It is based on the idea that electron pairs (bonding pairs and lone pairs) around a central atom will arrange themselves in a way that minimizes repulsion between them, leading to a stable, three-dimensional molecular shape. Here’s a summary of the key principles of VSEPR theory:
Electron Pair Arrangement: In VSEPR theory, you start by determining the electron pair arrangement around the central atom. This arrangement is based on the number of bonding pairs and lone pairs of electrons.
Steric Number: The steric number is the sum of the number of bonding pairs and lone pairs around the central atom.
Electron Pair Repulsion: Electron pairs (bonding and lone pairs) repel each other due to their negative charge. The repulsion between these electron pairs determines the molecular geometry.
Molecular Geometry: VSEPR theory predicts various molecular geometries, including linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral, among others.
Now, let’s discuss the structures of the given molecules with reference to VSEPR theory:
CH4 (Methane)
Steric Number: 4 (four bonding pairs around the central carbon atom)
Electron Pair Arrangement: Tetrahedral
Molecular Geometry: Tetrahedral
CH4 has a tetrahedral shape, with four hydrogen atoms surrounding the central carbon atom.
NH3 (Ammonia)
Steric Number: 4 (three bonding pairs and one lone pair around the central nitrogen atom)
Electron Pair Arrangement: Tetrahedral
Molecular Geometry: Trigonal Pyramidal
NH3 has a trigonal pyramidal shape, with three hydrogen atoms and one lone pair around the central nitrogen atom.
H2O (Water)
Steric Number: 4 (two bonding pairs and two lone pairs around the central oxygen atom)
Electron Pair Arrangement: Tetrahedral
Molecular Geometry: Bent (or V-shaped)
H2O has a bent (or V-shaped) molecular geometry due to the presence of two lone pairs on oxygen.
BeCl2 (Beryllium Dichloride)
Steric Number: 2 (two bonding pairs around the central beryllium atom)
Electron Pair Arrangement: Linear
Molecular Geometry: Linear
BeCl2 has a linear molecular geometry, with the two chlorine atoms on opposite sides of the central beryllium atom.
BF3 (Boron Trifluoride)
Steric Number: 3 (three bonding pairs around the central boron atom)
Electron Pair Arrangement: Trigonal Planar
Molecular Geometry: Trigonal Planar
BF3 has a trigonal planar molecular geometry, with the three fluorine atoms arranged in a flat, triangular fashion around the central boron atom.
SO2 (Sulfur Dioxide)
Steric Number: 3 (two bonding pairs and one lone pair around the central sulfur atom)
Electron Pair Arrangement: Trigonal Planar
Molecular Geometry: Bent (or V-shaped)
SO2 has a bent (or V-shaped) molecular geometry due to the presence of one lone pair on sulfur.
SO3 (Sulfur Trioxide)
Steric Number: 3 (three bonding pairs around the central sulfur atom)
Electron Pair Arrangement: Trigonal Planar
Molecular Geometry: Trigonal Planar
SO3 has a trigonal planar molecular geometry, with three oxygen atoms arranged in a flat, triangular fashion around the central sulfur atom.
VSEPR theory provides a systematic way to predict the shapes of molecules based on the arrangement of electron pairs around the central atom and the principles of electron pair repulsion.