Journey Inside the Atom — Notes

Rediscovering the Roots of Atomic Theory

The fundamental question of what makes up everything has intrigued thinkers for millennia. Ancient Indian philosopher Acharya Kanada proposed the idea of parmanus, the smallest, indivisible particles of matter. He theorised that repeatedly dividing matter would eventually lead to these ultimate particles, which could not be further divided. His thoughts, recorded in the Vaisesika Sutras, suggested that these parmanus combine in groups called dyads (two parmanus) and triads (three parmanus) to form all substances in the universe.

Similarly, ancient Greek philosophers Leucippus and Democritus proposed that matter is composed of indivisible particles they called atomos (meaning indivisible in Greek).

It is important to note that these early concepts of atoms were philosophical ideas, not based on experimental evidence.

Centuries later, in 1808, John Dalton put forth the first scientific atomic theory. This theory was grounded in experimental observations and stated that:

• All matter is composed of indivisible particles called atoms.
• Atoms are the fundamental building blocks of matter and cannot be broken down further.

Dalton's atomic theory marked a significant shift, providing the first scientific framework for understanding matter's composition and laying the groundwork for modern atomic structure. His ideas spurred further scientific curiosity, leading to questions about the internal composition of atoms, their appearance, and what differentiates atoms of one element from another.

Remember

Early ideas of atoms by Kanada, Leucippus, and Democritus were philosophical, while John Dalton's theory was the first to be based on scientific experiments.

Exam-smart

Questions may ask you to differentiate between early philosophical ideas of atoms and Dalton's scientific atomic theory, focusing on the basis of their proposals (philosophical vs. experimental).

A Short Historical Journey Through Atomic Models

Until the late 19th century, atoms were considered the smallest, indivisible units of matter. However, the discovery of radioactivity – the emission of invisible energy and particles from certain elements – revealed that atoms are not indivisible and are composed of smaller particles. This discovery spurred scientists to develop models to understand the atom's internal structure. These early models, though eventually superseded, were crucial steps in the scientific process, demonstrating how theories evolve with new experimental evidence.

Thomson’s Model of an Atom

In 1897, J. J. Thomson investigated the passage of electric current through gases at low pressure using a cathode ray tube. He observed rays emanating from the cathode (negative electrode) towards the anode (positive electrode), known as cathode rays. Through experiments in electric and magnetic fields, he deduced that these rays consisted of streams of negatively charged particles with a mass much smaller than an atom. These particles were later named electrons.

Since electrons were found to be emitted from atoms, Thomson proposed that atoms themselves must contain these subatomic particles. However, atoms are electrically neutral. To account for this neutrality, Thomson proposed his model of the atom in 1904.

• Description: Thomson envisioned the atom as a sphere of uniformly distributed positive charge. Electrons, being negatively charged, were embedded within this sphere, much like plums in a pudding or seeds in a watermelon.
• Neutrality: The total negative charge of the electrons was balanced by the total positive charge of the sphere, making the atom electrically neutral.
• Significance: This was the first scientific model to describe the internal structure of an atom, attempting to explain the co-existence of positive and negative charges and the overall neutrality of an atom.

Analogy:

• Plum Pudding Model: Electrons are the plums, and the pudding is the positively charged sphere.
• Watermelon Model: The red pulp represents the positive charge, and the seeds represent the embedded electrons.

Remember

Thomson's model was the first to incorporate subatomic particles (electrons) and explain atomic neutrality.

Exam-smart

Questions might describe a model with positive and negative charges embedded in it and ask to identify it or its limitations. You would recognize this description as Thomson's model and recall its key features.

Testing Thomson’s Model: The Gold Foil Experiment

In 1911, Hans Geiger and Ernest Marsden, under the guidance of Ernest Rutherford, conducted an experiment to test Thomson's model.

• Experimental Setup: They directed a narrow beam of alpha (α) particles at an extremely thin sheet of gold foil. Alpha particles are positively charged particles, essentially helium nuclei.
• Expectation based on Thomson's Model: According to Thomson's model, the positive charge was spread out evenly throughout the atom. Therefore, the fast-moving alpha particles, being positively charged, were expected to pass through the gold foil with minimal or no deflection, as the diffused positive charge would offer little resistance.
• Observed Results (The Surprise):
• Most alpha particles passed straight through the gold foil undeflected.
• A small fraction of alpha particles were deflected by small angles.
• A very few alpha particles (about 1 in 20,000) were deflected by large angles, some even bouncing back towards the source.
• Conclusion: The experimental results contradicted Thomson's model. The large-angle deflections and backward scattering indicated that the positive charge and most of the mass of the atom were concentrated in a very small, dense region, rather than being spread out. This led to the proposal of a new atomic model by Rutherford.

Remember

The gold foil experiment is also known as the alpha-ray scattering experiment. The key observation was that some alpha particles were deflected through large angles.

Exam-smart

A question might provide data from the gold foil experiment (e.g., percentage of particles deflected) and ask you to infer the structure of the atom or identify which model fails to explain these results. You should link the observed deflections to the concentration of mass and positive charge in a small nucleus.

Rutherford’s model of an atom

Rutherford's gold foil experiment, which involved bombarding a thin gold foil with alpha ($\alpha$) particles, provided crucial evidence that led him to propose a new model of the atom. This model aimed to explain the observed scattering of $\alpha$-particles, which Thomson's model could not adequately account for.

Key Conclusions from the Gold Foil Experiment:

• Most $\alpha$-particles passed through undeflected: This indicated that atoms are mostly empty space.
• A few $\alpha$-particles were deflected by small angles: This suggested the presence of a positively charged region within the atom, which repelled the positively charged $\alpha$-particles.
• A very small number of $\alpha$-particles were deflected by large angles, and some even bounced back: This was the most surprising observation and implied that the positive charge and mass of the atom were concentrated in a tiny, dense region.

Rutherford's Proposed Model (Planetary Model):

Based on these conclusions, Rutherford proposed the following:

• The atom contains a central, extremely small, positively charged region called the nucleus.
• The nucleus contains all the positive charge and most of the mass of the atom.
• Most of the atom is empty space.
• Electrons, which are negatively charged, revolve around the nucleus in definite paths, much like planets orbit the Sun. This is why his model is also known as the planetary model.

Size of the Nucleus and Atom:

Rutherford calculated that the nucleus is remarkably small compared to the overall size of the atom.

• Diameter of an atom: Approximately ${10}^{-10}$ meters (m).
• Diameter of the nucleus: Approximately ${10}^{-15}$ meters (m).

This means the nucleus is about ${10}^5$ (one lakh) times smaller than the atom. An analogy is that if an atom were the size of a cricket ground, the nucleus would be like a small pepper grain at its centre.

Significance of Rutherford's Model:

Rutherford's model was a significant advancement over Thomson's model because it could better explain the results of the gold foil experiment, especially the large-angle deflections and back-scattering of $\alpha$-particles. It introduced the concept of a dense, central nucleus.

Remember

The "planetary model" analogy highlights electrons orbiting the nucleus, but it's important to note that unlike planets, electrons are bound by electrostatic forces, not gravity.

Exam-smart

A question might present data from the gold foil experiment (e.g., percentage of deflected $\alpha$-particles) and ask you to infer properties of the atom's structure, such as the presence of empty space or a dense nucleus, directly linking to Rutherford's conclusions.

Limitations of Rutherford’s model

Rutherford's nuclear model of the atom, while a significant improvement over Thomson's model for explaining the gold foil experiment, had a critical flaw: it could not explain the stability of the atom.

The Problem of Accelerating Electrons

As per physics principles learned in "Describing Motion Around Us" (Chapter 4), any object moving in a circular path is continuously changing its direction. This change in direction means the object is accelerating.

• An electron, being a negatively charged particle, orbiting a positively charged nucleus, would therefore be constantly accelerating.
• According to electromagnetic theory, an accelerating charged particle should continuously lose energy in the form of radiation.
• If an electron loses energy, it would gradually spiral inward and eventually collapse into the positively charged nucleus.

The Contradiction with Reality

This theoretical collapse of the atom directly contradicts what we observe in the real world. Atoms are stable, which is why matter around us remains intact and does not disintegrate. Rutherford's model failed to account for this observed stability. It implied that atoms should not exist as they do, leading to the need for a new model that could explain how electrons maintain their motion without falling into the nucleus.

Remember

The core limitation of Rutherford's model was its inability to explain why atoms are stable, despite the electron being an accelerating charged particle.

Exam-smart

A question might describe the theoretical outcome of Rutherford's model (an electron spiraling into the nucleus) and ask you to identify why this contradicts reality and what phenomenon it fails to explain (atomic stability). You would need to connect the loss of energy by an accelerating charge to the collapse of the atom and then contrast this with the observed stability of elements.

Discovery of the Proton

The discovery of the proton was a significant step in understanding the atom's structure, refining earlier models. It explained the positive charge observed within atoms.

Who Discovered It?

• Ernest Rutherford is credited with the discovery and naming of the proton.

What is a Proton?

• A proton is a subatomic particle found in the nucleus of an atom.
• It carries a positive electrical charge.
• The relative charge of a proton is denoted as +1.

Key Properties of Protons:

• Location: Nucleus of the atom.
• Charge: Positive ($+1$).
• Mass: Protons are significantly heavier than electrons. While their exact mass is $1.672\times {10}^{-27}$ kg, for atomic comparisons, their mass is considered approximately 1 atomic mass unit (amu).

Significance in Atomic Neutrality:

• For an atom to be electrically neutral, the number of protons (positive charges) must be equal to the number of electrons (negative charges).
• Example: A helium atom has 2 protons and 2 electrons, resulting in a net charge of zero. Similarly, a sodium atom has 11 protons and 11 electrons, making it neutral.

Remember

The positive charge in the nucleus comes from protons. The number of protons determines the element.

Exam-smart

Questions might describe an atom's charge and ask to relate the number of protons and electrons, or ask about the particle responsible for the positive charge in the nucleus.

Bohr’s Model of the Atom

Niels Bohr proposed a new atomic model in 1913 to address the stability issues raised by Rutherford's model. He explained why electrons do not collapse into the nucleus while revolving around it.

Key Postulates of Bohr's Model:

1. Fixed Orbits (Energy Levels): Electrons do not move randomly but revolve around the nucleus in specific, fixed circular paths called stationary states, orbits, or shells.
2. Quantized Energy: Each shell has a definite amount of energy. These shells are also known as energy levels.
3. Allowed Orbits: Electrons can only revolve in these allowed shells, not in the space between them.
4. Energy Conservation in Orbit: While revolving in a fixed shell, an electron does not lose energy.
5. Energy Hierarchy:

• The shell closest to the nucleus is the K-shell ($n=1$) and has the least energy.
• Energy increases as we move away from the nucleus. The L-shell ($n=2$) has more energy than the K-shell, the M-shell ($n=3$) has more than the L-shell, and so on.

6. Electron Transitions: An electron can move from one shell to another only by absorbing or releasing a fixed amount of energy. This energy difference is equal to the difference in energy between the two shells.
7. Shell Capacity: Each shell can hold only a specific maximum number of electrons.

Why are Shells Named K, L, M, N...?

• This naming convention originated from early X-ray experiments by Charles Barkla. He designated the first observed X-ray line as 'K' and left room for earlier possibilities. Bohr adopted this notation for atomic shells.

Remember

Bohr's model introduced the concept of quantized energy levels for electrons, explaining atomic stability.

Exam-smart

Questions may ask to identify the energy level closest to the nucleus, compare energy levels, or explain why electrons don't spiral into the nucleus according to Bohr's model.

What Components Contribute to the Mass of an Atom?

Earlier models indicated that the nucleus contains most of the atom's mass. Electrons are so light that their mass is negligible when considering the total atomic mass. However, a puzzle arose regarding the mass of elements like helium compared to hydrogen.

• A hydrogen atom has one proton.
• A helium atom has two protons.
• Yet, a helium atom is about four times heavier than a hydrogen atom, not just double. This suggested the presence of another particle in the nucleus contributing to mass.

Discovery of the Neutron

The puzzle of atomic masses was solved by the discovery of the neutron.

Who Discovered It?

• James Chadwick (a student of Ernest Rutherford) discovered the neutron in 1932.

What is a Neutron?

• A neutron is a subatomic particle found in the nucleus of most atoms.
• It has a mass nearly equal to that of a proton.
• Crucially, it has no electrical charge (it is neutral). Its relative charge is 0.

Key Properties of Neutrons:

• Location: Nucleus of the atom.
• Charge: Neutral (0).
• Mass: Approximately 1 atomic mass unit (amu), similar to a proton.

Contribution to Atomic Mass:

• The mass of an atom is primarily contributed by its protons and neutrons, which are tightly packed in the nucleus.
• This explains why atoms are heavier than the sum of their protons alone.

Presence in Atoms:

• Neutrons are found in the nucleus of all atoms except for the most common isotope of hydrogen (protium), which has only one proton and no neutrons.

Role in Nuclear Stability:

• Protons, being positively charged, repel each other in the nucleus.
• Neutrons play a vital role in nuclear stability. They help reduce the repulsion between protons by:
• Increasing the distance between protons.
• Strengthening the nuclear force (the strong force that binds all particles in the nucleus).
• Heavier atoms generally have more neutrons than protons to overcome the increased electrostatic repulsion and hold the nucleus together.

Table 8.1: Symbols and Relative Charges of Subatomic Particles

Remember

Neutrons contribute significantly to an atom's mass but do not affect its overall charge. They are essential for nuclear stability in heavier elements.

Exam-smart

Questions might provide the number of protons and neutrons and ask for the total mass number, or explain why heavier nuclei need more neutrons. You might also be asked about the particle responsible for mass contribution without charge.

Symbols of Elements

To simplify the study of chemistry and ensure clear communication among scientists globally, elements are represented by unique symbols. These symbols are abbreviations that make it easier to write and discuss chemical reactions and compounds.

Historical Development of Symbols

• John Dalton's Pictorial Symbols (1803): Dalton was one of the first to propose a standardized system of symbols. He used simple drawings (pictograms) to represent elements. For instance, a circle with a dot in the center represented hydrogen. While innovative, these symbols were complex and difficult to draw repeatedly.

• Jöns Jacob Berzelius's Alphabetic Symbols (1813): Berzelius suggested a more practical system where symbols were derived from the Latin names of the elements. This system formed the basis of the symbols used today.

• International Union of Pure and Applied Chemistry (IUPAC): Today, IUPAC is the international body responsible for approving the names and symbols of elements.

Rules for Writing Chemical Symbols

The modern system of chemical symbols follows specific rules established by IUPAC:

1. One or Two Letters: The symbol usually consists of the first letter or the first two letters of the element's name.
2. Capitalization: The first letter of a symbol is always capitalized (uppercase). If a second letter is used, it is always lowercase (lowercase).

• Examples: Hydrogen is H, Aluminium is Al (not AL), Cobalt is Co (not CO).

3. Symbols from Other Letters: Sometimes, the symbol is derived from the first letter and a letter that is not the second letter of the name.

• Examples: Chlorine is Cl, Zinc is Zn.

4. Symbols from Non-English Names: Many element symbols are derived from their Latin, Greek, or German names because these were the common scientific languages when the elements were discovered or named.

• Examples:
• Iron (English) $\to$ Fe (from Latin ferrum)
• Mercury (English) $\to$ Hg (from Greek hydrargyros)
• Tungsten (English) $\to$ W (from German wolfram)

Common Elements and Their Symbols

Why Symbols are Important

• International Recognition: Symbols are understood by scientists worldwide, regardless of their native language.
• Conciseness: They provide a shorthand way to represent elements, making chemical formulas and equations much shorter and easier to write and read.
• Clarity: Using standard symbols prevents confusion that could arise if different scientists used their own unique abbreviations for the same element.

Remember

The first letter of an element's symbol is always capitalized, and the second letter, if present, is always lowercase. This capitalization rule is crucial for distinguishing between symbols of different elements (e.g., Co for Cobalt vs. CO for Carbon Monoxide, a compound).

Exam-smart

You might be given a list of chemical formulas (e.g., $H_{2}O$, $C{O}_2$, $NaCl$) and asked to identify the elements present, or you might be given an element's symbol and asked to deduce its origin (e.g., why is Sodium written as Na?). You could also be asked to name the element from a given symbol.

Atomic Number

The atomic number, represented by the symbol Z, is a fundamental property of an element. It is defined as the number of protons present in the nucleus of an atom of that element.

• Why it matters: The atomic number uniquely identifies an element. All atoms of a given element have the same number of protons, and atoms of different elements have different numbers of protons. This number dictates the element's identity and, consequently, its chemical behaviour.

• Neutral Atoms: In a neutral atom, the number of positively charged protons in the nucleus is equal to the number of negatively charged electrons orbiting the nucleus. Therefore, the atomic number also tells us the number of electrons in a neutral atom.

Examples:

• Hydrogen (H): Has 1 proton and 1 electron. Its atomic number (Z) is 1.
• Helium (He): Has 2 protons and 2 electrons. Its atomic number (Z) is 2.
• Lithium (Li): Has 3 protons and 3 electrons. Its atomic number (Z) is 3.

The atomic number is the primary basis for organizing elements in the periodic table. Elements are arranged in increasing order of their atomic numbers.

Remember

The atomic number (Z) is solely determined by the number of protons, which defines the element.

Exam-smart

A question might provide you with the number of electrons or protons in an atom and ask for its atomic number or the element's identity. For example, if given that an atom has 11 protons, you would identify it as Sodium (Na) because its atomic number is 11. If given that a neutral atom has 17 electrons, you would deduce it has 17 protons, meaning its atomic number is 17, identifying it as Chlorine (Cl).

Mass Number

The mass number (denoted by A) of an atom represents the total number of protons and neutrons found within its nucleus. These particles, protons and neutrons, are collectively known as nucleons because they reside in the atom's core.

The mass of a neutron is approximately equal to the mass of a proton. Electrons, on the other hand, have a mass that is considered negligible and can be ignored when calculating the mass of an atom.

The formula for calculating the mass number is:

Examples of Mass Number

The table below illustrates the calculation of the mass number for a few elements:

Notation for Elements

In the standard notation for an element, the atomic number (Z) is usually written as a subscript to the left of the element's symbol, and the mass number (A) is written as a superscript to the left of the symbol.

For example, carbon has an atomic number of 6 and a mass number of 12. Its notation is:

Here:

• A = 12 (mass number)
• Z = 6 (atomic number)
• C is the symbol for carbon

If we know the atomic number (which is the number of protons) and the mass number, we can determine the number of neutrons:

Let's consider the examples from the "Pause and Ponder" questions:

• Example 1: An atom with an atomic number of 26 has 56 nucleons.

• The atomic number (Z) is the number of protons, so Z = 26.

• The number of nucleons is the mass number (A). So, A = 56.

• The element's notation is ${}_{26}^{56}\text{X}$.

• Example 2: The nucleus of an atom contains 20 protons and its mass number is 41.

• Number of protons = 20. This means the atomic number (Z) is 20.

• Mass number (A) = 41.

• Number of neutrons = A - Z = 41 - 20 = 21.

• The element's notation is ${}_{20}^{41}\text{X}$.

• Example 3: An atom has an atomic number of 17 and 18 neutrons.

• Atomic number (Z) = 17. This means it has 17 protons.

• Number of neutrons (n${}^0$) = 18.

• Mass number (A) = Number of protons + Number of neutrons = 17 + 18 = 35.

• The element's notation is ${}_{17}^{35}\text{X}$.

• Example 4: An atom is represented as ${}^{23}\text{A}$. It has 11 electrons.

• We are given the mass number A = 23.

• In a neutral atom, the number of electrons equals the number of protons (which is the atomic number, Z). So, Z = 11.

• Number of neutrons = A - Z = 23 - 11 = 12.

Remember

The mass number tells us about the total mass contribution from protons and neutrons in the nucleus. The atomic number tells us the identity of the element.

Exam-smart

You might be given the atomic number and mass number of an element and asked to find the number of neutrons, or vice-versa. You could also be given the number of protons and neutrons and asked to write the atomic notation of the element. For example, a question might state: "An element has 11 protons and 12 neutrons. Determine its atomic number, mass number, and write its symbol in the standard notation."

How Are Electrons Distributed in Different Energy Levels?

Electrons in an atom are not randomly scattered around the nucleus. Instead, they occupy specific regions of space called energy levels or shells. These energy levels are associated with different amounts of energy. Electrons with lower energy are found closer to the nucleus, while those with higher energy are further away.

The distribution of electrons into these energy levels follows specific rules, often referred to as the Bohr-Bury rules:

1. The $2n^2$ Rule: The maximum number of electrons that can be accommodated in a particular energy shell is determined by the formula $2n^2$, where 'n' is the principal quantum number or the shell number.

• For the K-shell ($n=1$), the maximum number of electrons is $2\times 1^2=2$.
• For the L-shell ($n=2$), the maximum number of electrons is $2\times 2^2=8$.
• For the M-shell ($n=3$), the maximum number of electrons is $2\times 3^2=18$.
• For the N-shell ($n=4$), the maximum number of electrons is $2\times 4^2=32$.

2. Outermost Shell Limit: The outermost shell of an atom cannot accommodate more than 8 electrons, even if its $2n^2$ capacity is higher. This rule is crucial for understanding chemical bonding and the stability of atoms. The only exception is the K-shell, which can hold a maximum of 2 electrons.

3. Stepwise Filling: Electrons fill these energy shells in a sequential manner, starting from the shell closest to the nucleus (K-shell) and moving outwards (L-shell, then M-shell, and so on). A shell is only filled after the preceding shell closer to the nucleus is completely filled. For instance, the L-shell starts filling only after the K-shell has its maximum 2 electrons.

The arrangement of electrons in different shells of an atom is called its electronic configuration.

Building up atoms

The process of constructing atoms of different elements involves adding electrons to these energy levels according to their atomic number and the rules mentioned above. This stepwise addition of electrons helps us understand the electronic configuration and, consequently, the chemical behavior of elements.

Example:

• Hydrogen (H): Atomic number = 1. It has 1 electron, which occupies the K-shell. Electronic configuration: 1.
• Helium (He): Atomic number = 2. It has 2 electrons, both filling the K-shell. Electronic configuration: 2.
• Lithium (Li): Atomic number = 3. The K-shell is filled with 2 electrons, and the remaining 1 electron goes into the L-shell. Electronic configuration: 2, 1.
• Neon (Ne): Atomic number = 10. The K-shell has 2 electrons, and the L-shell is filled with its maximum of 8 electrons. Electronic configuration: 2, 8.
• Sodium (Na): Atomic number = 11. The K-shell has 2 electrons, the L-shell has 8 electrons, and the remaining 1 electron goes into the M-shell. Electronic configuration: 2, 8, 1.

Table 8.4 in the reference material provides the electronic distribution for the first eighteen elements, illustrating how electrons are filled in the K, L, and M shells.

Remember

The outermost shell can hold a maximum of 8 electrons (except for the first shell, which holds 2), a principle known as the octet rule, which explains the stability of noble gases.

Exam-smart

Case studies might present you with an element's atomic number and ask for its electronic configuration. You'll need to apply the $2n^2$ rule and the stepwise filling principle to determine how many electrons go into each shell. For example, given an atomic number of 13, you would fill the K-shell (2), then the L-shell (8), leaving 3 electrons for the M-shell, resulting in a configuration of 2, 8, 3.

Combining Capacity of an Atom: Valency

The combining capacity of an atom, known as valency, tells us how many atoms of hydrogen or chlorine, or how many electrons, an atom will gain, lose, or share to achieve a stable electron configuration. This stability is typically achieved when an atom's outermost electron shell, called the valence shell, has eight electrons (an octet), or two electrons in the case of the first shell (like in Helium).

Atoms strive for this stable configuration because it makes them less reactive. Atoms with incomplete valence shells are more reactive and will interact with other atoms to achieve this stability.

How Valency is Determined by Electronic Configuration

The electronic configuration (the arrangement of electrons in different energy shells) dictates an atom's valency.

• Valence Shell: The outermost shell of an atom containing electrons.
• Valence Electrons: The electrons present in the valence shell.

Rules for Determining Valency:

1. Losing Electrons: If an atom has fewer than four valence electrons, it tends to lose these electrons to achieve an octet in the shell below. The number of electrons lost equals its valency.

• Example: Sodium (Na) has an electronic configuration of 2, 8, 1. It has 1 valence electron. It loses this electron to become stable, so its valency is 1.

2. Gaining Electrons: If an atom has more than four valence electrons, it tends to gain electrons to complete its octet. The number of electrons gained equals its valency.

• Example: Oxygen (O) has an electronic configuration of 2, 6. It has 6 valence electrons. It needs 2 more electrons to complete its octet. Thus, its valency is 2.

3. Sharing Electrons: If an atom has four valence electrons, it often shares these electrons with other atoms to achieve an octet. The number of electrons shared equals its valency.

• Example: Carbon (C) has an electronic configuration of 2, 4. It has 4 valence electrons. It shares these 4 electrons with other atoms, so its valency is 4.

Valency and Combination:

• In $H_{2}O$, oxygen combines with two hydrogen atoms. Oxygen's valency is 2.
• In $N{H}_3$ (ammonia), nitrogen combines with three hydrogen atoms. Nitrogen's valency is 3.
• In $MgC{l}_2$ (magnesium chloride), magnesium combines with two chlorine atoms. Magnesium's valency is 2.

Valency Table

Remember

Atoms with a complete octet (like Neon) or a duet (like Helium, 2) in their outermost shell are chemically inert (unreactive) and have a valency of 0.

Exam-smart

A question might provide the electronic configuration or atomic number of an element and ask for its valency or the formula of a compound it forms. You need to identify the number of valence electrons and apply the rules for gaining, losing, or sharing to find the valency. For example, an element with atomic number 16 has configuration 2, 8, 6. It has 6 valence electrons and needs 2 more for an octet, so its valency is 2.

A Deeper Look into Atomic Structure

While the basic model of atoms having a fixed number of protons and neutrons is useful, scientists discovered that atoms of the same element can vary in their neutron count.

Isotopes

Isotopes are atoms of the same element that have the same atomic number (same number of protons) but different mass numbers (different number of neutrons).

• Atomic Number (Z): Number of protons. Defines the element.
• Mass Number (A): Number of protons + Number of neutrons.

Why are they "twin atoms"?

They are the "same element" because they have the same number of protons, meaning they have the same number of electrons and, therefore, the same electronic configuration. This similarity in electron arrangement means they exhibit similar chemical properties.

How do they differ?

They differ in the number of neutrons, which affects their mass number and hence their physical properties (like melting point, boiling point, and density).

Examples of Isotopes:

1. Hydrogen:

• Protium (¹H): 1 proton, 0 neutrons (mass number 1). Most common form (~99.98%).
• Deuterium (²H): 1 proton, 1 neutron (mass number 2).
• Tritium (³H): 1 proton, 2 neutrons (mass number 3). Radioactive.

2. Carbon:

• ¹²C: 6 protons, 6 neutrons (mass number 12). Most abundant (~99%).
• ¹³C: 6 protons, 7 neutrons (mass number 13).
• ¹⁴C: 6 protons, 8 neutrons (mass number 14). Radioactive, used in carbon dating.

Applications of Isotopes:

Some isotopes, called radioisotopes, are radioactive and have crucial applications:

• Nuclear Reactors: Uranium-235 (${}^{235}$U) is used as fuel to generate electricity.
• Cancer Treatment: Cobalt-60 (${}^{60}$Co) is used in radiation therapy.
• Medicine: Iodine-131 (${}^{131}$I) is used to treat goitre and thyroid cancer.
• Archaeology & Geology: Carbon-14 (${}^{14}$C) is used to determine the age of fossils and ancient artefacts (carbon dating).

Remember

Isotopes of an element have identical chemical properties because they have the same number of electrons. Their physical properties differ due to the different number of neutrons.

Exam-smart

A question might describe an atom by its mass number and number of protons (e.g., "An atom has a mass number of 23 and 11 protons"). You should deduce that it is Sodium (atomic number = 11). To find the number of neutrons, you use the formula: Mass Number = Protons + Neutrons. So, Neutrons = Mass Number - Protons = 23 - 11 = 12. The question might then ask about its properties or its position in the periodic table, linking to its identity as Sodium.

Average Atomic Mass

The atomic mass of an element represents the mass of its atoms. However, many elements exist naturally as a mixture of isotopes. Isotopes are atoms of the same element that have the same atomic number (same number of protons) but different mass numbers (different number of neutrons).

For example, chlorine exists in nature as two isotopes:

• Chlorine-35 (${}^{35}$Cl), with a mass of approximately 35 atomic mass units (u).
• Chlorine-37 (${}^{37}$Cl), with a mass of approximately 37 atomic mass units (u).

If we were to simply take the arithmetic average of these masses, we would get:

However, this simple average does not accurately represent the actual mass of chlorine found in nature because the isotopes do not occur in equal proportions. Chlorine-35 is much more abundant than Chlorine-37.

To find the true atomic mass of an element as it occurs in nature, we must calculate a weighted average atomic mass. This takes into account the relative abundance (or percentage) of each isotope. The weighted average is calculated by multiplying the mass of each isotope by its relative abundance and then summing these values.

For chlorine, where ${}^{35}$Cl occurs with about 75% abundance and ${}^{37}$Cl with about 25% abundance:

It's important to understand that no single atom of chlorine has a mass of 35.5 u. This value represents the average mass if we were to consider a large sample of chlorine atoms. For instance, in a million chlorine atoms, we would expect approximately 750,000 atoms of ${}^{35}$Cl and 250,000 atoms of ${}^{37}$Cl, resulting in an average mass of 35.5 u.

Isobars

Atoms of different elements that have the same mass number but different atomic numbers are called isobars.

Consider the elements calcium (Ca), potassium (K), and argon (Ar):

• Calcium (Ca) has an atomic number of 20 and a mass number of 40. This means it has 20 protons and 20 neutrons ($20+20=40$).
• Potassium (K) has an atomic number of 19 and a mass number of 40. This means it has 19 protons and 21 neutrons ($19+21=40$).
• Argon (Ar) has an atomic number of 18 and a mass number of 40. This means it has 18 protons and 22 neutrons ($18+22=40$).

Although calcium, potassium, and argon are distinct elements with different numbers of protons (and thus different chemical properties), their most common isotopes all have a mass number of 40. Therefore, they are isobars.

Remember

Isotopes are of the same element with different mass numbers, while isobars are of different elements with the same mass number.

Exam-smart

You might be given the atomic numbers and mass numbers of several species and asked to identify which ones are isotopes of each other and which ones are isobars. Remember to check both the element (atomic number) and the total number of nucleons (mass number). For example, a question might list atomic number and mass number for various elements and ask to find an isobar for a given element. You'd look for a different element with the same mass number.