Structure Of the Atom - Class 9 Science - Chapter 4 - Notes, NCERT Solutions & Extra Questions
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The number of dots in the Lewis symbol represents:
Option A) Atomic mass of the element
Option B) The electronic configuration of the atom
Option C) Atomic number of the element
Option D) The number of valence electrons present in the atom
The number of dots in the Lewis symbol represents the number of valence electrons present in the atom. These dots are positioned around the symbol of the element and indicate the electrons available for bonding or lone pairs.
Hence, the correct answer is:
Option D) The number of valence electrons present in the atom
Foils made of __________ were used in alpha-ray scattering experiments by Rutherford.
In the alpha-ray scattering experiments, specifically in the context of Rutherford's experiments, gold foil was employed instead of aluminum or any other metal due to gold’s exceptional malleability. Gold can be hammered into extremely thin sheets, some only around 1000 atoms thick. This thinness is crucial as it allows for the alpha particles to penetrate the foil and scatter efficiently. Using foils from other metals, which would inherently be thicker than those made of gold, might not yield the same precision in scattering results due to variations in thickness.
The atomic number of an element is 20. In the modern periodic table, this element is placed in $\qquad$
A 3rd period
B 2nd period
C 1st period
D 4th period
The correct option is D - 4th period.
The element with the atomic number 20 is Calcium (symbol: Ca), which belongs to Group IIA. It is located in the 4th period of the modern periodic table.
A nucleus ${}_{Z}^{A} X$ going through a $\beta^{+}$ decay will end up taking what configuration?
(A) ${}_{Z+1}^{A} Y$
(B) ${}{Z}^{A-1} X$ (C) ${}{Z+1}^{A+1} Y$ (D) ${}_{-1}^{A} Y$.
The correct option is (A) ${}_{Z-1}^{A} Y$.
In a Beta-plus (β+) decay, a proton within the nucleus transforms into a neutron, a positron ($e^+$), and an electron neutrino ($\nu_e$). This process results in the decrease of the atomic number $Z$ by 1 because the proton, which contributes to the atomic number, converts into a neutron that does not. However, the mass number $A$, which sums the total number of protons and neutrons, remains unchanged because the neutron remains within the nucleus despite the conversion:
$$ {}{Z}^{A} X \rightarrow {}{Z-1}^{A} Y + e^+ + \nu_e $$
Thus, the daughter nucleus has one less proton, accounting for the decrease in the atomic number, but the same overall number of nucleons as the parent nucleus.
If the atom of an element has atomic number 11 and mass number 23, find the number of protons, electrons, and neutrons in the atom.
The given atom of an element has an atomic number of 11 and a mass number of 23.
The number of protons in an atom is equal to its atomic number. Thus, the number of protons is: $$ \text{Number of protons} = 11 $$
In a neutral atom, the number of electrons is equal to the number of protons. Therefore, the number of electrons is also: $$ \text{Number of electrons} = 11 $$
The number of neutrons in an atom can be calculated by subtracting the atomic number from the mass number: $$ \text{Number of neutrons} = \text{Mass number} - \text{Atomic number} = 23 - 11 = 12 $$
In summary:
Protons: 11
Electrons: 11
Neutrons: 12
An electrically neutral atom must have:
A) Fewer electrons than protons.
B) The same number of neutrons and electrons but a different number of protons.
C) The same number of protons and electrons.
D) The same number of protons and neutrons but a different number of electrons.
Correct Answer: C - Same number of protons and electrons.
Explanation:In an atom:
Protons are positively charged.
Electrons are negatively charged.
Neutrons are neutral, having no charge.
An element is considered electrically neutral when the number of protons (positive charges) equals the number of electrons (negative charges). The charges cancel each other out because the magnitude of the charge on a proton is equal and opposite to that of an electron. Therefore, the neutrality of an atom does not depend on neutrons since they do not contribute to electrical charge.
An atom of an element contains 29 electrons and 35 neutrons. Deduce: (i) the number of protons (ii) the electronic configuration of the element.
Solution
(i) In a neutral atom, the number of protons equals the number of electrons. Therefore, the number of protons in this atom is $29$.
(ii) The element with 29 electrons is copper. The electronic configuration for copper can be expressed as:
$$ 1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^{10} $$
This configuration follows the standard order of filling the electron shells and subshells, respecting the general electron configuration principles.
What is meant by subatomic particle?
Solution
A subatomic particle refers to any of the various constituent elements of matter that are smaller than a hydrogen atom. These particles include, but are not limited to, electrons, neutrinos, photons, and alpha particles. They are the fundamental components that interact with each other in the physical world.
Subatomic particles can be further categorized into types such as protons, neutrons, and electrons, as well as all hadrons and leptons. These particles are essential in the composition and functioning of atoms and molecules. There are two broad categories within subatomic particles: composite particles and elementary particles, which describe their complexity and fundamental nature.
Which amongst the following was not concluded from Rutherford's alpha ray scattering experiment?
A) Most of the space inside an atom is empty.
B) The positive charge in an atom is uniformly distributed.
C) The positive charge in an atom is concentrated in a very small volume.
D) Most of the mass of an atom is concentrated in its center.
Solution
The correct answer is B) The positive charge in an atom is uniformly distributed.
The key findings from Rutherford's alpha ray scattering experiment include:
Most of the space inside an atom is empty, indicating that the electrons occupy very little space.
The positive charge in an atom is concentrated in a very small volume within the nucleus.
Most of the mass of an atom is concentrated in its center, where the nucleus resides.
These conclusions refute option B as Rutherford's experiment demonstrated that the positive charge is not uniformly distributed but rather confined to a small central nucleus.
Who discovered the electron?
Solution
Sir J.J. Thomson in 1897 is credited with the discovery of the electron. Thomson conducted an enhancement of the cathode ray experiments, where he determined the velocity of electrons and the ratio of their charge to mass. This was achieved by applying electrical and magnetic fields perpendicular to each other and to the path of the electrons.
In a solid 'AB' having the $\mathrm{NaCl}$ structure, 'A' atoms occupy the corners of the cubic unit cell. If all the face-centered atoms along one of the axes are removed, then the resultant stoichiometry of the solid is [IIT Screening 2001]
(A) $AB_{2}$
B) $A_{2}B$
(C) $A_{4}B_{3}$
(D) $A_{3}B_{4}$
The correct answer is (D) $A_{3}B_{4}$.
Explanation:
In a $\mathrm{NaCl}$ structure, 'A' atoms normally occupy the corner positions of the cubic unit cell. Each corner atom is shared among eight unit cells, meaning only $\frac{1}{8}$ of each corner atom is part of one particular unit cell. Since there are 8 corners, the number of 'A' atoms contributed by all corners to one unit cell is: $$ 8 \times \frac{1}{8} = 1 \text{ atom of } A $$
Each face-centered 'A' atom is shared between two unit cells, so contributes $\frac{1}{2}$ to the unit cell. Thus, initially for a complete $\mathrm{NaCl}$ structure, with three axes each having two face-centered 'A' atoms, the contribution would be: $$ 3 \times 2 \times \frac{1}{2} = 3 \text{ atoms of } A $$ However, removal of 'A' atoms along one of the axes (removing 2 face-centered atoms) decreases the number of 'A' atoms contributed by the face-centers from 3 atoms to: $$ 2 \times \frac{1}{2} + 1 \times \frac{1}{2} = 1.5 \text{ atoms of } A $$ Adding up contributions from the corners and the modified face-centers gives the total 'A' atoms as: $$ 1 + 1.5 = 2.5 $$ This will typically be normalized in context, for actual crystals, to an integer ratio of positions filled as required physically, resulting in an adjustment yielding 3 when considering multiples for the formula unit.
'B' atoms occupy the center and face-centered positions originally (one in the center shared normally with none, contributing fully, and six face-centered similarly shared but halved due to the cell sharing). However, each 'B' in these face-centered positions still contributes entirely to counting here because no 'B' atom removal is mentioned: $$ 1 + 3 \times 1 = 4 \text{ atoms of } B $$ Therefore, the resultant stoichiometry, counted by sorted integers and not fractionally adjusted, would be $A_{3}B_{4}$. This results in option D.
Which of the following statements about Rutherford's model of an atom is not correct?
A. Most of the atom's mass is concentrated in an extremely small region called the nucleus.
B. Electrons move around the nucleus in orbits.
C. Electrons and the nucleus are held together by electrostatic forces of attraction.
D. The model explained the stability of an atom.
The correct answer is D. The model explained the stability of an atom.
Rutherford's model of the atom was groundbreaking, but it failed to address the stability of an atom effectively. It suggested that electrons revolve around the nucleus, similar to planets orbiting the sun. However, according to classical physics, any charged particle, like an electron, that is in circular motion should emit energy. As a result, an electron would lose energy continuously and ultimately collapse into the nucleus, spiraling inward. This scenario contradicts the observed stability of atoms in nature, thus highlighting a fundamental flaw in the model concerning atomic stability.
Discuss how the structure of the atom developed over the years. [5 MARKS]
The conceptual understanding of the atom's structure has evolved significantly over the years through various models proposed by different scientists. Here's the sequence in which these developments took place:
Dalton Model (1 Mark): Initially, John Dalton proposed that atoms were indivisible particles, essentially the smallest unit of matter. This model was foundational but lacked the intricacies of atomic composition and internal structure.
Plum-Pudding Model (1 Mark): Following Dalton, J.J. Thomson introduced the plum-pudding model. In this model, the atom was described as a large positive sphere with electrons embedded within it. This model suggested a sort of uniformity in the distribution of charge within the atom.
Rutherford's Model (2 Marks): The plum-pudding model was challenged and disproved by Ernest Rutherford through his famous gold foil experiment. Rutherford observed that most alpha particles passed through a thin sheet of gold foil without any deflection, which indicated that most of the atom is empty space. This led to the revolutionary idea that an atom consists of a dense nucleus at its center surrounded by electrons in wider orbits.
Bohr's Model (1 Mark): Building on Rutherford's findings, Niels Bohr advanced the atomic model by introducing the concept of discrete orbital paths for electrons. According to Bohr, electrons revolve in specific, quantized orbits and do not radiate energy as they orbit, which explains why atoms are stable.
Each of these models contributed to the comprehensive understanding of atomic structure, refining and correcting the previous theories based on new experimental evidence and scientific analysis.
Which of the following statements is always correct?
A) An atom has an equal number of electrons and protons.
B) An atom has an equal number of electrons and neutrons.
C) An atom has an equal number of protons and neutrons.
D) An atom has an equal number of electrons, protons, and neutrons.
The correct answer is A) An atom has an equal number of electrons and protons.
Explanation: Atoms are the fundamental building blocks of all matter, and they must maintain electrical neutrality. This neutrality is achieved because atoms contain an equal number of protons and electrons. Protons possess a positive charge, and electrons carry a negative charge of equal magnitude but opposite in sign to protons. Thus, the net charge of an atom that has equal numbers of electrons and protons is zero, maintaining its neutrality.
The statements B, C, and D cannot be assumed to be always true because the number of neutrons can vary even among atoms of the same element, leading to different isotopes. Additionally, the ratio of protons to neutrons can also differ across different elements and their isotopes. Thus, Option A is always correct as it reflects a fundamental property of all atoms: electrical neutrality due to equal numbers of protons and electrons.
When alpha particles are sent through a thin metal foil, only one out of ten thousand rebounded. What did this observation lead to?
A. Unit positive charge is only present in an atom.
B. The size of the positively charged nucleus is very small as compared to the atom.
C. More number of electrons are revolving around the nucleus of the atom.
D. No positive charge is present.
The correct answer is B. The size of the positively charged nucleus is very small as compared to the atom.
When alpha particles, which are helium nuclei, were directed through a thin metal foil, the observation that only one out of ten thousand rebounded indicates that most of the space within the atom is empty. This demonstrates that the positive charge, which can repel the positively charged alpha particles, is highly concentrated in a very small region within the atom.
This led to the conclusion that the positively charged nucleus occupies a tiny portion of the total volume of an atom, which was a key finding in the development of the nuclear model of the atom by Rutherford.
The correct electronic configuration for sodium ion is:
(A) 2,8,1
B) 2,8
C) 2,8,8,1
(D) 2,8,7
The correct answer is Option B: 2,8.
The atomic number of sodium (Na) is 11, which means that a neutral sodium atom has 11 electrons. The electronic configuration of the neutral atom is: $$ 2, 8, 1 $$ When sodium forms an ion, it loses one electron to achieve a more stable electron configuration. This leads to the sodium ion (Na+) having an electronic configuration of: $$ 2, 8 $$ This corresponds to the filled first and second energy levels, lacking the extra electron in the third level that was present in the neutral atom. Thus, the most accurate electronic configuration for the sodium ion is given in Option B.
If by Neil's Bohr formula the number of electrons is 2n2, and if the 3rd shell contains 18 electrons, why did he say that the outermost shell can maximum hold 8 electrons?
In Neil Bohr's formula, the number of electrons in a shell is given by $2n^2$. For the 3rd shell ($n = 3$), this would be $2 \cdot 3^2 = 18$ electrons. However, Bohr also stated that the outermost shell can hold a maximum of 8 electrons. Let's understand this.
When $n = 2$, the number of electrons is $2 \cdot 2^2 = 8$. An atom with atomic number 10 has an electronic configuration of 2,8. This means that up to an atomic number of 10 (and considering the second shell $n = 2$), atoms can hold a maximum of 8 electrons in their outermost shell.
However, for $n > 2$, the capacity increases, and the shell can hold more than 8 electrons. This 8-electron limit is specifically applicable to elements with atomic numbers less than 10. As the atomic number increases, the capacity of the outermost shell can exceed 8 electrons.
In summary:
For the second shell ($n = 2$), the maximum of 8 electrons is observed in the outermost shell for elements with atomic numbers up to 10.
For higher shells ($n > 2$), the capacity increases, allowing more than 8 electrons in the outermost shell as the atomic number increases.
Key takeaways:
The rule of maximum 8 electrons in the outermost shell applies to elements with atomic numbers less than 10.
For elements with higher atomic numbers, the outermost shell can hold more than 8 electrons.
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Compare the properties of electrons, protons and neutrons. (IN TABLE) 100 WORDS
Electrons, protons, and neutrons are subatomic particles that are the fundamental components of atoms. Here is a table comparing their properties:
Property | Electron | Proton | Neutron |
---|---|---|---|
Charge | -1 (negative) | +1 (positive) | 0 (neutral) |
Relative Mass | 1/1836 | 1 | Slightly more than 1 |
Actual Mass | (9.109 × 10^{-31}) kg | (1.673 × 10^{-27}) kg | (1.675 × 10^{-27}) kg |
Location | Orbits the nucleus | Nucleus | Nucleus |
Symbol | $e^-$ | (p^+) | (n^0) |
These particles contribute to atomic structure, with electrons in orbitals enveloping the nucleus, where protons and neutrons reside. Charge determines chemical bonding and electricity flow, with electrons often being exchanged or shared in reactions. Mass contributes to atomic weight; protons and neutrons nearly equal, while electrons are much lighter. The combination of protons and neutrons defines isotopes of elements, while electron arrangements determine chemical properties and reactions.
What are the limitations of J.J. Thomson's model of the atom?
J.J. Thomson's model of the atom, also known as the "plum pudding" model, was proposed after his discovery of the electron in 1897. Thomson suggested that the atom consisted of a positively charged sphere within which electrons were embedded like "plums in a pudding."
This model had several limitations that were revealed by later experiments:
Lack of a Nucleus: The Thomson model couldn't accommodate the nuclear structure of the atom, which was discovered later by Ernest Rutherford in 1911. Rutherford's gold foil experiment showed that atoms have a dense, positively charged nucleus where most of the mass is concentrated, which Thomson's model did not account for.
Inaccurate Representation of Atoms: The plum pudding model suggested that the positive charge was spread out evenly with the electrons interspersed. This didn't represent the discreet structure of the atom, where the electrons occupy specific energy levels or orbits.
Inability to Explain Spectral Lines: The model couldn't explain why atoms emitted light at specific frequencies, known as spectral lines. Niels Bohr later addressed this with his own atomic model by introducing quantized orbits for electrons.
Alpha Particle Scattering: Thomson's model predicted that positively charged alpha particles would pass through the "plum pudding" of positive charge with minimal deflection. However, Rutherford's experiment showed that a small fraction of these particles were deflected at very large angles, which the Thomson model could not explain.
Stability of the Atom: The model didn't provide any explanation for the stability of the atom. According to classical electromagnetism, electrons moving through a diffuse sea of positive charge would lose energy and spiral into the center, causing the atom to collapse.
Quantum Mechanics: The model didn't incorporate quantum mechanics, which we now know plays a critical role in understanding atomic and subatomic behavior. Quantum mechanics provided a more accurate and detailed description of electron behavior through the work of scientists such as Werner Heisenberg and Erwin Schrödinger.
In summary, while J.J. Thomson's model was a step forward in understanding the atomic structure, it was eventually superseded by more advanced models that could better explain the observed experimental data.
What are the limitations of Rutherford's model of the atom?
Rutherford's model of the atom, also known as Rutherford's nuclear model, was a major step forward in our understanding of atomic structure when it was proposed by Ernest Rutherford in 1911. However, this model also had several limitations:
Lack of explanation for atom stability: According to classical electromagnetism, electrons moving in circular orbits should emit electromagnetic radiation. Rutherford's model suggests that electrons circle the nucleus like planets around the sun, which would cause them to lose energy, spiral into the nucleus, and collapse the atom—a situation that does not occur since atoms are stable.
Fixed orbit radii: The model doesn't explain why electrons should only exist at certain distances from the nucleus or why each orbit could hold a fixed maximum number of electrons.
Spectral lines: The model could not account for the discrete spectral lines observed in atomic spectra. It only suggested that atoms had a nucleus and electrons but did not provide a framework for understanding how electrons could produce the specific lines observed.
Quantization of energy levels: Rutherford's model did not include the concept of quantized energy levels, which was necessary to explain why electrons orbit at certain discrete energy levels and why the emission of energy from atoms occurred at specific frequencies.
Heisenberg's Uncertainty Principle: Rutherford's model treated electrons as particles that revolve around the nucleus in well-defined paths, while the Heisenberg uncertainty principle shows that the position and momentum of a particle cannot both be precisely determined at the same time, contradicting the idea of electrons following precise orbits.
Quantum mechanical nature of electrons: Quantum mechanics later showed that the behavior of electrons in an atom is better described by probabilistic wave functions rather than deterministic paths.
Niels Bohr built on Rutherford's model and addressed some of its limitations by proposing that electrons exist in quantized orbits, with a defined set of allowed radii and energy levels, resulting in the more advanced Bohr model of the atom. However, even the Bohr model had its limitations and was eventually superseded by the more complete and sophisticated quantum mechanical models.
Describe Bohr's model of the atom.
Bohr's model of the atom is a theory put forward by Niels Bohr in 1913 to describe the structure of atoms. Here are the key points of Bohr's model:
Quantized Orbits: Bohr postulated that electrons orbit the nucleus in certain allowed circular paths or orbits, each with a fixed energy level. These orbits are called "energy levels" or "shells."
Energy Quantization: Electrons can only occupy these orbits and cannot exist in between them. The energy levels are quantized, which means that the energy of an electron in an atom can only take on certain discrete values.
Photon Emission or Absorption: When an electron transitions from a higher energy level to a lower one, it emits energy in the form of a photon. Conversely, when an electron absorbs energy, it can move from a lower energy level to a higher one.
Angular Momentum Quantization: Bohr proposed that the angular momentum of an electron in a given orbit is quantized and is an integer multiple of the reduced Planck constant divided by $2\pi$ i.e., $L = n\hbar$, where $L$ is the angular momentum, (n) is an integer known as the quantum number, and (\hbar) is the reduced Planck constant).
Stability of Orbits: Within these orbits, the electron does not emit radiation; thus, the orbits are stable, and the electron can circle the nucleus indefinitely without losing energy and spiraling into the nucleus, which was a limitation of the classical models.
Bohr's model successfully explained the spectral lines of the hydrogen atom by relating the energy difference between the quantized orbits to the frequency of the light emitted or absorbed. However, Bohr's model had limitations; it could not accurately predict the spectra of larger atoms and did not fully take into account the wave nature of electrons. It was superseded by the more comprehensive quantum mechanical models, such as the Schrödinger wave equation and the Heisenberg uncertainty principle, which form the basis of modern atomic physics.
Compare all the proposed models of an atom given in this chapter.
The proposed atom models described in the provided chapter are as follows:
J.J. Thomson's Model: Thomson proposed that the atom is composed of electrons embedded in a positive sphere.
Rutherford’s Model: Rutherford's alpha-particle scattering experiment led to the discovery of the atomic nucleus. He proposed the model where a very tiny nucleus is present inside the atom, and electrons revolve around this nucleus. However, this model could not explain the stability of the atom.
Neils Bohr’s Model: Bohr proposed a more successful model in which electrons are distributed in different shells with discrete energy around the nucleus. If the atomic shells are complete, then the atom is stable and less reactive.
Chadwick’s Discovery: In addition to these models, J. Chadwick discovered the presence of neutrons in the nucleus of an atom, identifying the sub-atomic particles constituting an atom as electrons, protons, and neutrons.
These models evolved over time as experiments and theoretical understanding advanced, providing increasingly accurate representations of atomic structure.
Summarise the rules for writing of distribution of electrons in various shells for the first eighteen elements.
The rules for writing the distribution of electrons in various shells for the first eighteen elements are based on the Bohr-Bury scheme and are as follows:
The maximum number of electrons present in a shell is given by the formula (2n^2), where (n) is the orbit number or energy level index (1, 2, 3, ...). Consequently, the maximum number of electrons in different shells are:
First orbit or K-shell: $2 \times 1^2 = 2$
Second orbit or L-shell: $2 \times 2^2 = 8$
Third orbit or M-shell: $2 \times 3^2 = 18$
Fourth orbit or N-shell: $2 \times 4^2 = 32$, and so on .
The maximum number of electrons that can be accommodated in the outermost orbit is 8 .
Electrons are not accommodated in a given shell unless the inner shells are filled—shells are filled in a step-wise manner .
These rules ensure the structured filling of electrons in an atom's shells, and explain the distribution of electrons across the different energy levels for the first eighteen elements of the periodic table.
Define valency by taking examples of silicon and oxygen.
Valency is a measure of the combining power of an element, specifically, the number of hydrogen atoms it can combine with or displace to form a compound. It is also considered as the number of electrons an atom can lose, gain, or share to become stable.
Let's consider the examples of silicon and oxygen:
Silicon (Si): Silicon is in Group 14 of the periodic table and has an atomic number of 14. Its electronic configuration is $1s^2 , 2s^2 , 2p^6 , 3s^2 , 3p^2 $. Si has four electrons in its outermost shell (the 3rd shell). To achieve a stable configuration similar to the nearest noble gas (argon), silicon can share those four electrons with other atoms, most commonly with four atoms of oxygen to form silicon dioxide $SiO₂$. Therefore, the valency of silicon is 4, as it can share four electrons with other atoms to form covalent bonds.
Oxygen (O): Oxygen is in Group 16 of the periodic table and has an atomic number of 8. Its electronic configuration is $1s^2 , 2s^2 , 2p^4 $. Oxygen has six electrons in its outermost shell (the 2nd shell). It needs two more electrons to achieve a stable configuration similar to the nearest noble gas (neon). Therefore, oxygen can either gain two electrons from other atoms or share two electrons through covalent bonding. Thus, the valency of oxygen is 2, as it tends to form compounds such as water $H₂O$ by sharing two electrons with two hydrogen atoms.
In summary, the valency of silicon is 4, and the valency of oxygen is 2, representing their ability to bond with other elements by sharing electrons.
Explain with examples (i) Atomic number, (ii) Mass number, (iii) Isotopes and iv) Isobars. Give any two uses of isotopes.
Certainly! Let's discuss atomic number, mass number, isotopes, and isobars with examples. Additionally, I'll provide two uses of isotopes.
(i) Atomic Number
The atomic number of an element is the number of protons in the nucleus of an atom of that element. It is denoted by the symbol $Z$.
Example: The atomic number of hydrogen is 1, which means a hydrogen atom has one proton in its nucleus.
(ii) Mass Number
The mass number, denoted by $A$, is the total number of protons and neutrons in the nucleus of an atom.
Example: The most common isotope of carbon has 6 protons and 6 neutrons, so its mass number is $A = 6 + 6 = 12$.
(iii) Isotopes
Isotopes are atoms of the same element (same atomic number) that have different mass numbers due to a different number of neutrons in their nuclei.
Example: Carbon has several isotopes, including Carbon-12 (6 protons and 6 neutrons) and Carbon-14 (6 protons and 8 neutrons). Both are isotopes of carbon and share the same atomic number (6) but have different mass numbers $12 and 14, respectively$.
(iv) Isobars
Isobars are atoms that have different atomic numbers (hence, they are different elements) but the same mass number. This means they have the same total number of protons and neutrons.
Example: Argon-40 (18 protons and 22 neutrons) and Calcium-40 (20 protons and 20 neutrons) are isobars. Both have a mass number of 40, but they are different elements with different atomic numbers (18 for Argon and 20 for Calcium).
Uses of Isotopes:
Medical Imaging and Treatment: Certain isotopes are used in medicine for diagnostic imaging and treatment. For example, Technetium-99m is a radioactive tracer widely used in nuclear medicine to image the skeleton and heart muscle. Another isotope, Iodine-131, is used for treating thyroid disorders, including cancer.
Scientific Research and Dating: Isotopes are used in scientific research to understand chemical and biological processes. A well-known use is in radiocarbon dating, where the isotope Carbon-14 is used to determine the age of organic archaeological, geological, and hydrogeological samples.
$\mathrm{Na}^{+}$has completely filled $\mathrm{K}$ and $\mathrm{L}$ shells. Explain.
Sodium (Na) is an element with the atomic number 11, which means it has 11 protons and 11 electrons in its neutral state. The electron configuration of a neutral sodium atom in the ground state would be represented as $1s^2 2s^2 2p^6 3s^1$, corresponding to the following arrangement in its atomic orbitals:
$K$ shell (first energy level): $1s^2$ - contains two electrons
$L$ shell (second energy level): $2s^2 2p^6$ - contains eight electrons
$M $ shell (third energy level): $3s^1$ - contains one electron
Now, when sodium forms a positively charged ion $(\mathrm{Na}^{+})$, it loses one electron. This electron is typically removed from the outermost shell to achieve a more stable, lower energy state. For sodium, this means the loss of the single electron in the (3s) orbital. As a result, the electron configuration for $\mathrm{Na}^{+}$ becomes:
$K$ shell (first energy level): $1s^2$ - completely filled with two electrons
$L$ shell (second energy level):$2s^2 2p^6$- completely filled with eight electrons
The $M $shell no longer has any electrons, and the electron configuration represents a state similar to neon, a noble gas known for its stability due to a full valence shell at the $L $level.
The fact that $\mathrm{Na}^{+}$ has completely filled $ K$ and $ L $ shells indicates a stable electron configuration with full inner shells, which is energetically favorable and contributes to the stability of the $\mathrm{Na}^{+}$ ion.
If bromine atom is available in the form of, say, two isotopes ${ }_{35}^{79} \mathrm{Br}(49.7 \%)$ and ${ }_{35}^{81} \mathrm{Br}(50.3 \%)$, calculate the average atomic mass of bromine atom.
The average atomic mass of an element that has multiple isotopes can be calculated by taking into account the relative abundance of each isotope and its atomic mass. The formula to calculate the average atomic mass is:
$$ \text{Average atomic mass} = \sum (\text{Isotope mass} \times \text{Relative abundance}) $$
The relative abundance is expressed as a percentage but needs to be used as a decimal fraction in the calculation. Now, let's apply the formula to the bromine isotopes given in the question:
For Bromine-79 $Br-79$, the isotope mass is $79$ atomic mass units $u$, and its relative abundance is $49.7%$.
For Bromine-81 $Br-81$, the isotope mass is $81$ atomic mass units$u$ and its relative abundance is (50.3%).
Convert the relative abundance percentages into decimal fractions:
The relative abundance of Br-79: $49.7% = 0.497$
The relative abundance of Br-81: $50.3% = 0.503$
Now, let's calculate the weighted average:
$$ \begin{align*} \text{Average atomic mass} &= (79 \times 0.497) + (81 \times 0.503) \ &= (39.263 , \text{u}) + (40.743 , \text{u}) \ &= 80.006 , \text{u} \end{align*} $$
Therefore, the average atomic mass of bromine with the given isotopic composition is approximately $80.006$atomic mass units $u$.
The average atomic mass of a sample of an element $\mathrm{X}$ is $16.2 \mathrm{u}$. What are the percentages of isotopes ${ }_{8}^{16} \mathrm{X}$ and ${ }_{8}^{18} \mathrm{X}$ in the sample?
The average atomic mass of an element that has multiple isotopes is calculated based on the relative abundances (percentages) and the atomic masses of the isotopes.
Given:
The average atomic mass of element $ \mathrm{X} $ is$ 16.2 \mathrm{u} $.
There are two isotopes: $ {}{8}^{16}\mathrm{X}$, which has an atomic mass of $16 \mathrm{u}$, and $ {}{8}^{18}\mathrm{X} $, which has an atomic mass of $ 18 \mathrm{u}$.
Let's denote the percentage of $ {}{8}^{16}\mathrm{X} $ as $ p $, and therefore the percentage of $ {}{8}^{18}\mathrm{X} $ would be $100 - p$ because the total must be 100%.
The average atomic mass of element X can be computed as follows:
$$ \text{Average atomic mass} = \frac{\text{Mass of } {}{8}^{16}\mathrm{X} \times \text{Percentage of } {}{8}^{16}\mathrm{X} + \text{Mass of } {}{8}^{18}\mathrm{X} \times \text{Percentage of } {}{8}^{18}\mathrm{X}}{100} $$
Plugging in the values, we get:
$$ 16.2 = \frac{(16 \times p) + (18 \times (100 - p))}{100} $$
Solving for ( p ):
$$ 16.2 = \frac{16p + 1800 - 18p}{100} $$
$$ 1620 = 16p + 1800 - 18p $$
$$ 1620 = 1800 - 2p $$
$$ 2p = 1800 - 1620 $$
$$ 2p = 180 $$
$$ p = \frac{180}{2} $$
$$ p = 90 $$
So the percentage of $ {}{8}^{16}\mathrm{X}$ is $ 90%$ and the percentage of $ {}{8}^{18}\mathrm{X} $ is $100 - 90 = 10% $.
Therefore, the sample contains $ 90% $ of $ {}{8}^{16}\mathrm{X} $ and $ 10% $ of $ {}{8}^{18}\mathrm{X} $.
If $Z=3$, what would be the valency of the element? Also, name the element.
The element with atomic number $Z = 3$ is lithium, represented by the atomic symbol Li. Its valency is 1. This means lithium tends to lose one electron to form positive ions with a charge of +1.
Composition of the nuclei of two atomic species $\mathrm{X}$ and $\mathrm{Y}$ are given as under
$\mathrm{X}$ | $\mathrm{Y}$ | |
---|---|---|
Protons | 6 | 6 |
Neutrons | 6 | 8 |
Give the mass numbers of $\mathrm{X}$ and $\mathrm{Y}$. What is the relation between the two species?
The mass number of an atomic species (or nuclide) is the total number of protons and neutrons in its nucleus.
For atomic species ( \mathrm{X} ):
Protons = 6
Neutrons = 6 Hence, the mass number of ( \mathrm{X} ) is ( 6 + 6 = 12 ).
For atomic species ( \mathrm{Y} ):
Protons = 6
Neutrons = 8 Hence, the mass number of ( \mathrm{Y} ) is ( 6 + 8 = 14 ).
Since both ( \mathrm{X} ) and ( \mathrm{Y} ) have the same number of protons, they are isotopes of the same element. Isotopes are atoms of an element that have the same number of protons but different numbers of neutrons. They have different mass numbers but are chemically similar because they have the same number of protons which determines the element's chemical properties.
Therefore, ( \mathrm{X} ) and ( \mathrm{Y} ) have mass numbers of 12 and 14, respectively, and they are isotopes of each other.
For the following statements, write $\mathrm{T}$ for True and $\mathrm{F}$ for False.
(a) J.J. Thomson proposed that the nucleus of an atom contains only nucleons.
(b) A neutron is formed by an electron and a proton combining together. Therefore, it is neutral.
(c) The mass of an electron is about $\frac{1}{2000}$ times that of proton.
(d) An isotope of iodine is used for making tincture iodine, which is used as a medicine.
(a) J.J. Thomson proposed that the nucleus of an atom contains only nucleons. F
J.J. Thomson proposed the "plum pudding" model of the atom in which the atom was a sphere of positive charge with electrons embedded within it, like plums in a pudding. The concept of the nucleus containing neutrons and protons, known as nucleons, was not part of his model.
(b) A neutron is formed by an electron and a proton combining together. Therefore, it is neutral. F
Although a neutron is electrically neutral, it is not formed by a simple combination of an electron and a proton. Neutrons and protons are made up of quarks and are held together by strong nuclear forces in the nucleus. The idea that the neutron is a compound of an electron and proton is an outdated concept that does not reflect current understanding of particle physics.
(c) The mass of an electron is about (\frac{1}{2000}) times that of proton. T
The mass of an electron is indeed approximately (\frac{1}{2000}) times the mass of a proton. More precisely, the mass of a proton is about 1,836 times that of an electron.
(d) An isotope of iodine is used for making tincture iodine, which is used as a medicine. F
Tincture of iodine, used as a disinfectant and antiseptic, is generally made from elemental iodine, not a specific isotope of iodine. However, radioactive isotopes of iodine, such as iodine-131, are used in medical applications, including treatment of thyroid conditions.
Rutherford's alpha-particle scattering experiment was responsible for the discovery of
(a) Atomic Nucleus
(b) Electron
(c) Proton
(d) Neutron
Rutherford's alpha-particle scattering experiment was responsible for the discovery of:
(a) Atomic Nucleus
Isotopes of an element have
(a) the same physical properties
(b) different chemical properties
(c) different number of neutrons
(d) different atomic numbers.
Isotopes of an element have:
(c) different number of neutrons
Isotopes are variants of a particular chemical element that differ in neutron number, while their proton number (atomic number) is the same. They have nearly identical chemical properties because they have the same number of electrons in their outer shell, but they often have different physical properties due to the difference in mass resulting from the different number of neutrons.
Number of valence electrons in $\mathrm{Cl}^{-}$ion are:
(a) 16
(b) 8
(c) 17
(d) 18
The chlorine atom (( \mathrm{Cl} )) normally has 17 electrons, with 7 of them being valence electrons. However, when chlorine gains an extra electron to form the chloride ion (( \mathrm{Cl}^{-} )), it has a total of 18 electrons with the additional one also being a valence electron. Therefore, the number of valence electrons in the ( \mathrm{Cl}^{-} ) ion is 8, which makes the electronic configuration stable as it resembles a noble gas configuration.
The correct option is: (b) 8
Which one of the following is a correct electronic configuration of sodium?
(a) 2,8
(b) $8,2,1$
(c) $2,1,8$
(d) $2,8,1$
The electronic configuration of sodium (chemical symbol: Na), which has an atomic number of 11, follows the pattern of filling electron shells in the order of increasing energy levels. Here is how the electrons would be distributed:
The first energy level (K shell) can hold up to 2 electrons.
The second energy level (L shell) can hold up to 8 electrons.
The third energy level (M shell) can hold more, but for sodium, only one electron will be in this outermost shell, as it is the 11th electron.
Therefore, the correct electronic configuration of sodium is ( 2, 8, 1 ), which corresponds to option (d).
Complete the following table.
Atomic Number | Mass Number | Number of Neutrons | Number of Protons | Number of Electrons | Name of the Atomic Species |
---|---|---|---|---|---|
9 | - | 10 | - | - | - |
16 | 32 | - | - | - | Sulphur |
- | 24 | - | 12 | - | - |
To complete the table, we will use the known relationships between the atomic number, mass number, and the numbers of protons, neutrons, and electrons.
For a neutral atom:
The number of protons is equal to the atomic number.
The number of electrons is equal to the number of protons (since the charges must balance).
The mass number is the sum of the number of protons and neutrons.
Here's the filled table:
For the atomic number 9:
Since the number of neutrons is given as 10, we can deduce the mass number by adding the atomic number (which is equal to the number of protons) to the number of neutrons: (9 + 10 = 19).
The number of protons is the same as the atomic number: 9.
For a neutral atom, the number of electrons is the same as the number of protons: 9.
The atomic species with atomic number 9 is Fluorine.
For the atomic number 16 (Sulfur):
The mass number is given as 32, so the number of neutrons is the mass number minus the number of protons (which equals the atomic number): (32 - 16 = 16).
The number of protons is the same as the atomic number: 16.
For a neutral atom, the number of electrons is the same as the number of protons: 16.
For the mass number 24 and number of protons 12:
The atomic number is the same as the number of protons: 12.
The number of neutrons is the mass number minus the atomic number: (24 - 12 = 12).
For a neutral atom, the number of electrons is the same as the number of protons: 12.
The atomic species with atomic number 12 is Magnesium.
The filled table is as follows:
Atomic Number | Mass Number | Number of Neutrons | Number of Protons | Number of Electrons | Name of the Atomic Species |
---|---|---|---|---|---|
9 | 19 | 10 | 9 | 9 | Fluorine |
16 | 32 | 16 | 16 | 16 | Sulphur |
12 | 24 | 12 | 12 | 12 | Magnesium |
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Comprehensive Structure of the Atom Class 9 Notes: Key Concepts Explained
Understanding the structure of the atom is crucial for students diving into the world of chemistry. Below, we provide comprehensive notes tailored for Class 9 science students, explaining key concepts related to atomic structure.
Fundamental Building Blocks of Matter
In earlier studies, we learned that atoms and molecules are the basic building blocks of matter. The diversity in matter arises from the different types of atoms constituting them. This chapter delves deeper into what differentiates one element's atom from another and whether atoms are truly indivisible, as once proposed by Dalton.
Discovery of Sub-Atomic Particles
Evidence of Charged Particles in Matter
Charged particles in matter were first indicated by experiments involving static electricity. Upon further experimentation, it was discovered that atoms are not indivisible and consist of smaller charged particles.
The Contribution of J.J. Thomson: Discovery of the Electron
J.J. Thomson identified the electron, a negatively charged sub-atomic particle. This marked the first step towards understanding the atomic structure.
The Contribution of E. Goldstein: Discovery of Canal Rays and Protons
E. Goldstein discovered canal rays, positively charged radiations that led to the identification of protons, positively charged sub-atomic particles that are around 2000 times heavier than electrons.
Thomson's Model of the Atom
J.J. Thomson proposed one of the first atomic models, suggesting that:
An atom is a positively charged sphere with electrons embedded within it.
The positive and negative charges within the atom are balanced, making the atom electrically neutral.
Thomson's model is often compared to a Christmas pudding, where the electrons are like raisins within a positively charged sphere.
Rutherford’s Model of the Atom
Ernest Rutherford designed an experiment to investigate the arrangement of electrons in an atom using alpha particles and a thin gold foil. The unexpected results of his experiment led to the discovery of the nucleus.
Rutherford’s Gold Foil Experiment
Rutherford observed:
Most alpha particles passed through the gold foil without deflection.
Some particles were deflected at small angles.
A few particles rebounded at large angles.
These observations suggested that:
Most of the atom is empty space.
Positive charge and most of the mass of the atom are concentrated in a small volume (nucleus).
Key Features of Rutherford’s Nuclear Model
Atoms have a small, dense, positively charged nucleus.
Electrons revolve around the nucleus in circular orbits.
The nucleus is much smaller than the atom itself.
Drawbacks of Rutherford’s Model
The main drawback of Rutherford's model was its inability to explain the stability of atoms. According to classical mechanics, electrons revolving in circular orbits would radiate energy and spiral into the nucleus, making atoms unstable, which contradicts real-world observations.
Bohr’s Model of the Atom
To address defects in Rutherford's model, Neils Bohr proposed that:
Electrons only occupy certain special orbits known as discrete orbits.
While in these discrete orbits, electrons do not radiate energy.
Bohr’s model successfully explained the stability of atoms by introducing the concept of energy levels or shells.
Electron Distribution in Orbits
Bohr and Bury outlined rules for electron distribution in different orbits:
The maximum number of electrons in a shell is determined by (2n^2), where (n) is the orbit number.
The outermost shell can hold a maximum of 8 electrons.
Electrons fill inner shells before occupying outer shells.
Understanding Valency
Valency refers to an atom's ability to combine with other atoms, determined by the number of electrons in the outermost shell.
Elements with complete outer shells are chemically inactive.
Valency is often calculated based on the electrons an atom needs to gain, lose, or share to achieve a full outer shell.
Atomic Number and Mass Number
Atomic Number ((Z))
The number of protons in an atom's nucleus defines its atomic number. This characteristic differentiates elements.
Mass Number ((A))
The total number of protons and neutrons in an atom’s nucleus constitutes its mass number.
Isotopes and Isobars
Isotopes
Isotopes are atoms of the same element with different mass numbers due to varying numbers of neutrons.
Isobars
Isobars are atoms of different elements with the same mass number but different atomic numbers.
Conclusion
Understanding the structure of the atom provides the foundation for comprehending more complex chemical reactions and properties of matter. By recognizing the contributions of early scientists and the evolution of atomic models, students gain a thorough insight into this fundamental topic.
This comprehensive guide ensures students are well-prepared and gives them a strong foundational understanding of atomic structure, essential for future studies in chemistry.
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