Structure Of The Atom

An atom is the basic building block of everything around us. It’s made up of three tiny particles:

1. Protons: These are positively charged particles found in the center of the atom, which is called the nucleus.
2. Neutrons: These are neutral (no charge) particles that are also located in the nucleus, alongside protons.
3. Electrons: These are negatively charged particles that move around the nucleus in different layers or shells.

The number of protons in the nucleus decides what element an atom is. For example, if an atom has one proton, it is hydrogen. The number of electrons usually matches the number of protons, so the atom has no overall charge.

The electrons orbit the nucleus in specific paths or energy levels. The arrangement of these electrons determines how the atom will bond with other atoms.

The discovery of electrons and cathode rays marked significant milestones in the understanding of atomic structure.

Discovery of Cathode Rays:

1. Early Experiments (Late 1800s): Scientists like Michael Faraday and Heinrich Geissler first observed cathode rays while experimenting with electrical discharge in vacuum tubes. They noticed that when a high voltage was applied across a cathode (negative electrode) and an anode (positive electrode) in a vacuum tube, a stream of rays was emitted from the cathode and traveled toward the anode.

2. J.J. Thomson’s Contribution (1897): The most crucial developments came from J.J. Thomson's experiments. Using a cathode ray tube, Thomson discovered that these rays were deflected by magnetic and electric fields, suggesting that they were composed of charged particles. He determined that these particles were much smaller than atoms and carried a negative charge. Thomson concluded that cathode rays were actually streams of electrons, which he later confirmed as fundamental components of atoms.

Discovery of the Electron:

1. Identification and Charge (1897): Thomson’s experiments showed that electrons were the smallest known particles and carried a negative electrical charge. He measured the charge-to-mass ratio of electrons, leading to the understanding that electrons were subatomic particles found within all atoms.

2. Impact on Atomic Theory: The discovery of electrons revolutionized atomic theory. It led to the development of the first models of the atom, such as Thomson's "plum pudding" model, which depicted the atom as a sphere of positive charge with embedded electrons.

In summary, the discovery of cathode rays and electrons helped scientists understand that atoms are made of smaller particles. Thomson's work in the late 19th century was crucial in identifying the electron as a fundamental building block of matter, setting the stage for further developments in atomic theory.

A discharge tube is a special kind of glass tube used in experiments to study electricity and gases. Here’s an easy explanation for Class 9:

What is a Discharge Tube?

• Shape and Material: It’s a long, sealed glass tube with electrodes (metal parts) at each end.
• Inside the Tube: The tube is usually filled with a gas at a low pressure.

How It Works:

1. Applying Voltage: When a high voltage is applied across the electrodes, it creates an electric current inside the tube.
2. Observing Effects: This current makes the gas inside the tube glow or produce light. The color of the glow can change depending on the type of gas used.

What It Shows:

• Cathode Rays: The glowing light is due to cathode rays, which are streams of tiny particles moving from the negative electrode (cathode) to the positive electrode (anode)

J.J. Thomson’s Experiment:

1. Equipment Used:

   • Thomson used a cathode ray tube, which is a long glass tube with a vacuum inside and two metal parts at each end called electrodes (one negative and one positive).

2. What Happened:

   • When Thomson applied a high voltage across the electrodes, a glowing beam of rays appeared inside the tube. This beam was made up of particles.

3. Key Observations:

   • Deflection: Thomson noticed that this glowing beam (called cathode rays) could be bent or deflected by magnets and electric fields.
   • Charge: The way the beam bent showed that the rays were made of tiny particles with a negative charge.

4. Conclusion:

   • Discovery of Electrons: Thomson discovered that these tiny, negatively charged particles were called electrons.
   • Model: This led him to propose a new model of the atom, where electrons were spread out inside a positively charged "soup" or sphere, known as the "plum pudding" model.

In summary, Thomson’s experiment with the cathode ray tube showed that atoms contain small, negatively charged particles called electrons, changing our understanding of atomic structure.

Properties of Cathode Rays:

1. They are Streams of Particles: Cathode rays are made up of tiny particles, which are actually electrons.

2. They Travel in Straight Lines: When there is no interference, cathode rays move in straight lines from the negative electrode (cathode) to the positive electrode (anode) in a vacuum tube.

3. They Can Be Deflected: Cathode rays can be bent or deflected by magnetic and electric fields, which shows that they are charged particles.

4. They Produce Light: When cathode rays hit a fluorescent material inside the tube, it makes the material glow.

5. They Carry Negative Charge: The particles in cathode rays have a negative charge, which means they are attracted to positive charges and repelled by negative charges.

Properties of Electrons:

1. Tiny Size: Electrons are very small particles that are much smaller than atoms.

2. Negative Charge: Electrons carry a negative electric charge.

3. Movement: Electrons move around the nucleus of an atom in specific energy levels or shells.

4. Charge-to-Mass Ratio: Electrons have a specific ratio of charge to mass, which Thomson measured and found to be very high compared to other particles.

5. In All Atoms: Electrons are found in all atoms and are important for chemical reactions and bonding between atoms.

In summary, cathode rays are streams of electrons, and electrons have key properties like being very small, carrying a negative charge, and being crucial to the structure of atoms.

Discovery of Anode Rays:

1. What Are Anode Rays?

   • Anode rays are rays that travel in the opposite direction of cathode rays. They move from the positive electrode (anode) to the negative electrode (cathode) in a vacuum tube.

2. How They Were Discovered:

   • In the late 1800s, scientists like Eugen Goldstein noticed that when they applied a high voltage in a cathode ray tube, rays also traveled from the anode and created a glow on the tube's glass.

3. Key Observations:

   • Deflection: Anode rays could be bent by electric and magnetic fields, showing that they were made of charged particles.
   • Positive Charge: Unlike cathode rays, which are negatively charged, anode rays are positively charged particles.

Discovery of Protons:

1. What Are Protons?

   • Protons are positively charged particles found in the nucleus of an atom.

2. How They Were Discovered:

   • Scientists like Ernest Rutherford discovered protons by studying how anode rays (positive rays) were deflected in experiments. He concluded that these positive rays were made up of positively charged particles.

3. Key Findings:

   • Charge and Mass: Protons have a positive charge and are heavier than electrons.
   • Atomic Identity: The number of protons in the nucleus of an atom determines what element it is. This is known as the atomic number.

In summary, anode rays are streams of positively charged particles moving from the anode to the cathode in a tube, and protons are the positively charged particles in the nucleus of an atom.

Goldstein's Experiment:

1. Setup:

   • Goldstein used a cathode ray tube, similar to the one used by J.J. Thomson, but with a few differences.
   • He placed a perforated (hole-filled) anode inside the tube.

2. What Happened:

   • When a high voltage was applied, rays began to travel from the anode (positive electrode) to the cathode (negative electrode).
   • These rays were different from the cathode rays that Thomson had studied.

3. Observations:

   • Glow: Goldstein noticed that these rays caused a glow on the glass of the tube.
   • Deflection: These rays could be bent by electric and magnetic fields, just like cathode rays, but in the opposite direction.

4. Key Findings:

   • Positive Rays: Goldstein concluded that these rays were made of positively charged particles, unlike the negatively charged cathode rays.
   • Anode Rays: These rays were later named anode rays or positive rays.

5. Importance:

   • Goldstein’s experiment helped identify the existence of positively charged particles in atoms. These particles were later called protons.

In summary, Goldstein’s experiment with the cathode ray tube revealed the existence of positively charged particles, which helped scientists understand more about the structure of atoms.

How Cathode Rays Are Produced:

1. Setup:

   • In a cathode ray tube, there are two electrodes: a negative one (cathode) and a positive one (anode). The tube is filled with a gas at low pressure.

2. Applying Voltage:

   • When a high voltage is applied between the cathode and anode, it creates an electric field inside the tube.

3. Ray Production:

   • Cathode Rays: Electrons (tiny, negatively charged particles) are emitted from the cathode (the negative electrode). These electrons travel through the tube toward the anode (the positive electrode). The stream of these electrons is called cathode rays.

4. Effect:

   • Cathode rays make the glass of the tube glow when they strike it, and they can be bent by magnetic and electric fields.

How Anode Rays Are Produced:

1. Setup:

   • In the same tube, when the cathode rays travel toward the anode, some rays can pass through small holes in the anode.

2. Ray Production:

   • Anode Rays: These rays, which travel from the anode to the cathode, are made of positively charged particles. They are produced when the positive particles, moving through the holes in the anode, travel through the tube in the opposite direction of the cathode rays.

3. Effect:

   • Anode rays can also make the tube glow and can be bent by electric and magnetic fields, but in the opposite direction compared to cathode rays.

Why They Are Produced:

1. Cathode Rays:

   • They are produced because the high voltage causes electrons to be released from the cathode. These electrons move toward the anode, creating a beam of cathode rays.

2. Anode Rays:

   • They are produced as a result of the interaction between the electrons and the gas inside the tube. When electrons hit the gas, they cause the release of positively charged particles that move toward the cathode.

In summary, cathode rays are streams of electrons moving from the negative to the positive electrode, while anode rays are streams of positively charged particles moving in the opposite direction. Both are created by applying a high voltage in a vacuum tube.

Properties of Anode Rays:

1. Positive Charge:

   • Anode rays are made up of particles that have a positive charge. This is why they move towards the negative electrode (cathode) when inside a vacuum tube.

2. Deflection:

   • These rays can be bent by electric and magnetic fields, but they bend in the opposite direction compared to cathode rays. This shows that they have a positive charge.

3. Glow:

   • When anode rays hit the glass of the tube, they can make it glow. The color of the glow depends on the type of gas inside the tube.

4. Mass:

   • The particles in anode rays are heavier than electrons. They are different from the negatively charged cathode rays.

5. Charge-to-Mass Ratio:

   • Anode rays have a different charge-to-mass ratio compared to cathode rays, which helps in identifying different types of positive particles.

Properties of Protons:

1. Positive Charge:

   • Protons have a positive electrical charge. This is the opposite of the negative charge of electrons.

2. Location:

   • Protons are found in the nucleus (the center) of an atom.

3. Mass:

   • Protons are much heavier than electrons. They are about 2000 times heavier than an electron.

4. Stability:

   • Protons are stable particles and do not break down under normal conditions. They are a key part of the atom’s nucleus.

5. Atomic Identity:

   • The number of protons in an atom determines what element it is. For example, if an atom has one proton, it’s hydrogen. If it has six protons, it’s carbon.

In summary, anode rays are positively charged particles that move in the opposite direction to cathode rays and cause the glass of the tube to glow, while protons are positively charged particles in the nucleus of an atom that determine the element’s identity.

Thomson's Atomic Model:

1. Name:

   • Thomson's model is also known as the "Plum Pudding Model".

2. Description:

   • Imagine an atom as a big, positively charged ball, like a pudding. Inside this ball are small, negatively charged particles, like plums in a pudding.

3. Positive Charge:

   • The positive charge is spread throughout the atom, forming a big, positive sphere.

4. Negative Particles:

   • The negative particles (electrons) are scattered randomly inside this positive sphere.

5. Overall Neutrality:

   • The atom as a whole is electrically neutral. This means the total positive charge balances out the total negative charge.

6. What It Explains:

   • This model was the first to suggest that atoms are not just tiny, indivisible spheres but have internal structures. It explained why atoms are neutral and how they might be built from smaller particles.

In summary, Thomson's atomic model described the atom as a positive sphere with negatively charged electrons scattered inside it, similar to a plum pudding. This was an early attempt to explain the structure of atoms.

Rutherford’s Alpha Particles Scattering Experiment:

1. Setup:

   • Rutherford used a thin gold foil, only a few atoms thick.
   • He fired alpha particles (which are positively charged) at the gold foil. These particles come from a radioactive source.

2. Observation:

   • Around the gold foil, Rutherford placed a fluorescent screen coated with a special material that would glow when hit by alpha particles. He also used a microscope to detect where the particles hit the screen.

3. Findings:

   • Most of the alpha particles passed through the gold foil without any deflection. This suggested that most of the atom is empty space.
   • Some alpha particles were deflected at small angles, and a very few were deflected back at large angles.

4. Conclusion:

   • Rutherford concluded that there must be a very small, dense, and positively charged center within the atom that repels the positively charged alpha particles. He named this center the nucleus.

5. Discovery of the Nucleus:

   • The nucleus is a tiny part of the atom but contains most of its mass. It has a positive charge and is surrounded by the negatively charged electrons, which orbit around it.

Summary:

Rutherford’s experiment showed that atoms have a small, dense, positively charged center called the nucleus, surrounded by mostly empty space where the electrons are located. This was a key discovery in understanding atomic structure.

Bohr’s Atomic Model:

1. Basic Idea:

   • Bohr’s model describes the atom as having a small, positively charged nucleus at the center and electrons orbiting around it in specific paths or shells.

2. Nucleus:

   • The nucleus is very tiny compared to the whole atom and contains most of the atom's mass. It has a positive charge due to protons.

3. Electron Orbits:

   • Electrons move in fixed, circular paths or orbits around the nucleus. These orbits are at specific distances from the nucleus and are called energy levels or shells.

4. Energy Levels:

   • Electrons in these orbits have certain energy levels. The closer an orbit is to the nucleus, the lower the energy level. Electrons can jump between these orbits by absorbing or emitting energy.

5. Stable Orbits:

   • Electrons in these orbits do not lose energy and spiral into the nucleus. They stay in stable orbits as long as they don’t gain or lose energy.

6. Energy Changes:

   • When electrons move from one orbit to another, they absorb or emit energy in the form of light. This is why atoms can give off colors of light when heated or excited.

Summary:

Bohr’s model of the atom shows a central nucleus surrounded by electrons moving in fixed orbits or shells. This model explains why electrons stay in specific energy levels and how they can absorb or emit energy when they move between these levels.

Discovery of the Neutron:

1. Background:

   • Before the discovery of the neutron, scientists knew that atoms had a nucleus made up of protons and that electrons orbited around this nucleus. However, the total mass of the nucleus was more than what could be explained by protons alone.

2. James Chadwick’s Experiment (1932):

   • James Chadwick conducted an experiment where he bombarded a beryllium target with alpha particles (positively charged particles).
   • The bombardment caused the target to emit a new type of radiation that was not deflected by electric or magnetic fields.

3. Observations:

   • Chadwick found that this radiation had no electrical charge, but it could knock protons out of other materials. This suggested that the radiation was made of particles with mass.

4. Conclusion:

   • Chadwick concluded that these particles were neutral (having no charge) and had a similar mass to protons. He named them neutrons.

5. Importance:

   • The discovery of neutrons helped explain the extra mass in the nucleus that could not be accounted for by protons alone. Neutrons, along with protons, are found in the nucleus of every atom.

Summary:

James Chadwick discovered neutrons by observing particles with no charge and mass similar to protons. This discovery completed the understanding of the atomic nucleus, explaining the missing mass and providing a full picture of atomic structure.

Atomic Number:

1. Definition:

   • The atomic number of an element is the number of protons found in the nucleus of an atom of that element.

2. Importance:

   • It tells us the identity of an element. Each element has a unique atomic number.
   • For example, hydrogen has an atomic number of 1 because it has 1 proton, and carbon has an atomic number of 6 because it has 6 protons.

3. Periodic Table:

   • Elements are arranged in the Periodic Table in order of increasing atomic number. This helps in understanding the properties of elements and how they relate to each other.

4. Relationship with Electrons:

   • In a neutral atom (one without an overall charge), the number of protons is equal to the number of electrons. So, the atomic number also tells us the number of electrons in the atom.

5. Example:

   • If an element has an atomic number of 8, it means it has 8 protons and, in a neutral atom, also 8 electrons.

Summary:

The atomic number is the number of protons in an atom's nucleus and defines the element's identity. It also tells us the number of electrons in a neutral atom and helps organize elements in the Periodic Table.

Mass Number:

1. Definition:

   • The mass number of an atom is the total number of protons and neutrons in its nucleus.

2. How It’s Calculated:

   • To find the mass number, you add the number of protons and neutrons together.
   • For example, if an atom has 6 protons and 6 neutrons, its mass number is 12 (6 + 6 = 12).

3. Importance:

   • The mass number helps determine the weight of an atom. It is used to identify different isotopes of an element. Isotopes are versions of the same element that have the same number of protons but different numbers of neutrons.

4. Example:

   • Carbon has several isotopes. The most common one is carbon-12, which has 6 protons and 6 neutrons, giving it a mass number of 12. Another isotope, carbon-14, has 6 protons and 8 neutrons, giving it a mass number of 14.

5. Notation:

   • The mass number is usually written as a superscript before the element’s symbol. For example, $_{6}^{12}C$ where 12 is the mass number.

Summary:

The mass number is the sum of protons and neutrons in an atom’s nucleus. It helps in identifying isotopes and understanding the atom’s overall mass.

Mass Number:

1. Definition:

   • The mass number of an atom is the total number of protons and neutrons in its nucleus.

2. How It’s Calculated:

   • Add the number of protons and neutrons together.
   • For example, if an atom has 6 protons and 6 neutrons, its mass number is 12 (6 + 6 = 12).

3. Purpose:

   • The mass number helps identify different isotopes of an element. Isotopes are atoms of the same element with different numbers of neutrons.

Atomic Mass:

1. Definition:

   • Atomic mass is the average mass of an atom of an element, taking into account all of its isotopes and their abundances in nature.

2. How It’s Measured:

   • Atomic mass is usually measured in atomic mass units (amu). It’s an average value because elements often exist as mixtures of different isotopes.

3. Purpose:

   • The atomic mass gives a more precise idea of the mass of an atom compared to just the mass number. It reflects the weighted average of the masses of an element’s isotopes.

4. Example:

   • For carbon, the mass number of the most common isotope is 12, but the atomic mass is about 12.01 amu. This is because carbon has a mix of isotopes like carbon-12 and carbon-14, and their relative abundances affect the average atomic mass.

Summary:

• Mass Number: The total number of protons and neutrons in an atom’s nucleus. It’s used to identify isotopes.
• Atomic Mass: The average mass of an atom of an element, considering all its isotopes and their natural abundance. It’s a more precise measure compared to the mass number.

1. Finding Number of Electrons, Protons, and Neutrons

Example 1:

• Element: Carbon
• Atomic Number: 6
• Mass Number: 12

Calculations:

1. Number of Protons:

   • The atomic number is equal to the number of protons.
   • For carbon, the number of protons = 6.

2. Number of Electrons:

   • In a neutral atom, the number of electrons is equal to the number of protons.
   • For carbon, the number of electrons = 6.

3. Number of Neutrons:

   • To find the number of neutrons, subtract the number of protons from the mass number.
   • Number of neutrons = Mass number - Number of protons
   • For carbon, number of neutrons = 12 - 6 = 6.

Summary:

• Protons: 6
• Electrons: 6
• Neutrons: 6

Example 2:

• Element: Oxygen
• Atomic Number: 8
• Mass Number: 16

Calculations:

1. Number of Protons:

   • The atomic number tells us the number of protons.
   • For oxygen, the number of protons = 8.

2. Number of Electrons:

   • In a neutral atom, the number of electrons equals the number of protons.
   • For oxygen, the number of electrons = 8.

3. Number of Neutrons:

   • Number of neutrons = Mass number - Number of protons
   • For oxygen, number of neutrons = 16 - 8 = 8.

Summary:

• Protons: 8
• Electrons: 8
• Neutrons: 8

Example 3:

• Element: Sodium (Na)
• Atomic Number: 11
• Mass Number: 23

Calculations:

1. Number of Protons:

   • The atomic number = 11
   • For sodium, the number of protons = 11.

2. Number of Electrons:

   • In a neutral atom, the number of electrons = number of protons.
   • For sodium, the number of electrons = 11.

3. Number of Neutrons:

   • Number of neutrons = Mass number - Number of protons
   • For sodium, number of neutrons = 23 - 11 = 12.

Summary:

• Protons: 11
• Electrons: 11
• Neutrons: 12

Key Points to Remember:

• Protons = Atomic Number
• Electrons = Number of Protons (in a neutral atom)
• Neutrons = Mass Number - Number of Protons

These examples should help you understand how to calculate the number of protons, electrons, and neutrons in various atoms!

Electronic Configuration:

1. Definition:

   • Electronic configuration is the arrangement of electrons in the shells or energy levels of an atom.

2. Shells and Energy Levels:

   • Electrons are arranged in different shells around the nucleus of an atom. Each shell can hold a specific number of electrons:
     • First shell: Can hold up to 2 electrons.
     • Second shell: Can hold up to 8 electrons.
     • Third shell: Can hold up to 18 electrons.
     • Fourth shell: Can hold up to 32 electrons.

3. How to Write Electronic Configuration:

   • Start filling the electrons in the lowest energy level (closest to the nucleus) first, then move to the next higher level.
   • Use numbers to represent the shells and write the number of electrons in each shell.

4. Example: Hydrogen (H):

   • Atomic Number: 1
   • Electronic Configuration: 1 (1 electron in the first shell)

5. Example: Neon (Ne):

   • Atomic Number: 10
   • Electronic Configuration: 2, 8 (2 electrons in the first shell and 8 electrons in the second shell)

6. Example: Sodium (Na):

   • Atomic Number: 11
   • Electronic Configuration: 2, 8, 1 (2 electrons in the first shell, 8 in the second shell, and 1 in the third shell)

7. Writing the Configuration:

   • Use a notation where each shell is listed with the number of electrons it holds:
     • 1s² 2s² 2p⁶ represents the configuration for Neon, where 1s² means 2 electrons in the first shell, and 2s² 2p⁶ means 8 electrons in the second shell.

Summary:

• Electronic Configuration is the way electrons are arranged in an atom's shells.
• Electrons fill the shells starting from the closest to the nucleus.
• Each shell has a maximum number of electrons it can hold.
• The configuration helps to understand how atoms bond and interact with other atoms.

Atomic Structures of the First Twenty Elements

1. Hydrogen (H)

   • Atomic Number: 1
   • Electrons: 1
   • Configuration: 1 (1 electron in the first shell)

2. Helium (He)

   • Atomic Number: 2
   • Electrons: 2
   • Configuration: 2 (2 electrons in the first shell)

3. Lithium (Li)

   • Atomic Number: 3
   • Electrons: 3
   • Configuration: 2, 1 (2 electrons in the first shell, 1 in the second shell)

4. Beryllium (Be)

   • Atomic Number: 4
   • Electrons: 4
   • Configuration: 2, 2 (2 electrons in the first shell, 2 in the second shell)

5. Boron (B)

   • Atomic Number: 5
   • Electrons: 5
   • Configuration: 2, 3 (2 electrons in the first shell, 3 in the second shell)

6. Carbon (C)

   • Atomic Number: 6
   • Electrons: 6
   • Configuration: 2, 4 (2 electrons in the first shell, 4 in the second shell)

7. Nitrogen (N)

   • Atomic Number: 7
   • Electrons: 7
   • Configuration: 2, 5 (2 electrons in the first shell, 5 in the second shell)

8. Oxygen (O)

   • Atomic Number: 8
   • Electrons: 8
   • Configuration: 2, 6 (2 electrons in the first shell, 6 in the second shell)

9. Fluorine (F)

   • Atomic Number: 9
   • Electrons: 9
   • Configuration: 2, 7 (2 electrons in the first shell, 7 in the second shell)

10. Neon (Ne)

    • Atomic Number: 10
    • Electrons: 10
    • Configuration: 2, 8 (2 electrons in the first shell, 8 in the second shell)

11. Sodium (Na)

    • Atomic Number: 11
    • Electrons: 11
    • Configuration: 2, 8, 1 (2 electrons in the first shell, 8 in the second shell, 1 in the third shell)

12. Magnesium (Mg)

    • Atomic Number: 12
    • Electrons: 12
    • Configuration: 2, 8, 2 (2 electrons in the first shell, 8 in the second shell, 2 in the third shell)

13. Aluminum (Al)

    • Atomic Number: 13
    • Electrons: 13
    • Configuration: 2, 8, 3 (2 electrons in the first shell, 8 in the second shell, 3 in the third shell)

14. Silicon (Si)

    • Atomic Number: 14
    • Electrons: 14
    • Configuration: 2, 8, 4 (2 electrons in the first shell, 8 in the second shell, 4 in the third shell)

15. Phosphorus (P)

    • Atomic Number: 15
    • Electrons: 15
    • Configuration: 2, 8, 5 (2 electrons in the first shell, 8 in the second shell, 5 in the third shell)

16. Sulfur (S)

    • Atomic Number: 16
    • Electrons: 16
    • Configuration: 2, 8, 6 (2 electrons in the first shell, 8 in the second shell, 6 in the third shell)

17. Chlorine (Cl)

    • Atomic Number: 17
    • Electrons: 17
    • Configuration: 2, 8, 7 (2 electrons in the first shell, 8 in the second shell, 7 in the third shell)

18. Argon (Ar)

    • Atomic Number: 18
    • Electrons: 18
    • Configuration: 2, 8, 8 (2 electrons in the first shell, 8 in the second shell, 8 in the third shell)

19. Potassium (K)

    • Atomic Number: 19
    • Electrons: 19
    • Configuration: 2, 8, 8, 1 (2 electrons in the first shell, 8 in the second shell, 8 in the third shell, 1 in the fourth shell)

20. Calcium (Ca)

    • Atomic Number: 20
    • Electrons: 20
    • Configuration: 2, 8, 8, 2 (2 electrons in the first shell, 8 in the second shell, 8 in the third shell, 2 in the fourth shell)

Summary:

• Electrons are arranged in shells around the nucleus.
• Each shell has a maximum number of electrons it can hold.
• Configuration shows how electrons are distributed among these shells for each element.

Valency:

1. Definition:

   • Valency is the ability of an atom to combine with other atoms. It tells us how many bonds an atom can form with other atoms.

2. Why It Matters:

   • Valency helps determine how elements combine to form compounds. It’s a way to understand the chemical behavior of an element.

3. Determining Valency:

   • For Elements in the Main Groups:
     • Group 1: Elements like hydrogen (H) and sodium (Na) have a valency of 1. They can form one bond.
     • Group 2: Elements like calcium (Ca) have a valency of 2. They can form two bonds.
     • Group 15: Elements like nitrogen (N) have a valency of 3. They can form three bonds.
     • Group 16: Elements like oxygen (O) have a valency of 2. They can form two bonds.
     • Group 17: Elements like chlorine (Cl) have a valency of 1. They can form one bond.
     • Group 18 (Noble Gases): These elements (e.g., neon) usually don’t form bonds easily because they already have full outer shells.

4. Examples:

   • Hydrogen (H): Valency is 1. It can combine with one atom of oxygen (O) to form water (H₂O).
   • Oxygen (O): Valency is 2. It can combine with two atoms of hydrogen to form water (H₂O).

5. Chemical Bonds:

   • Single Bond: One bond formed (e.g., H-Cl).
   • Double Bond: Two bonds formed (e.g., O=O).
   • Triple Bond: Three bonds formed (e.g., N≡N).

Summary:

• Valency is the number of bonds an atom can form with other atoms.
• It helps in understanding how elements combine to form compounds.
• Elements in different groups have different valencies, affecting their bonding behavior.

Types of Valency:

1. Basic Valency (or Positive Valency):

   • Definition: This is the ability of an atom to donate or lose electrons to form bonds.
   • Example:
     • Sodium (Na): Has a valency of 1 because it can lose 1 electron to form a positive ion (Na⁺).
     • Calcium (Ca): Has a valency of 2 because it can lose 2 electrons to form a positive ion (Ca²⁺).

2. Negative Valency (or Negative Valency):

   • Definition: This is the ability of an atom to accept or gain electrons to form bonds.
   • Example:
     • Chlorine (Cl): Has a valency of 1 because it can gain 1 electron to form a negative ion (Cl⁻).
     • Oxygen (O): Has a valency of 2 because it can gain 2 electrons to form a negative ion (O²⁻).

3. **Neutral Valency:

   • Definition: This occurs when an atom does not gain or lose electrons but shares electrons with other atoms to form bonds.
   • Example:
     • Carbon (C): Has a valency of 4. It shares 4 electrons with other atoms to form bonds (e.g., in methane, CH₄).

4. **Variable Valency:

   • Definition: Some elements can have more than one valency depending on the chemical reaction.
   • Example:
     • Iron (Fe): Can have a valency of 2 or 3. It can form Fe²⁺ or Fe³⁺ ions.
     • Copper (Cu): Can have a valency of 1 or 2, forming Cu⁺ or Cu²⁺ ions.

5. **Zero Valency:

   • Definition: This is when an atom does not have a tendency to form bonds because it already has a full outer shell of electrons.
   • Example:
     • Noble Gases (e.g., Neon, Ne): Have a valency of 0 because their outer electron shells are full, so they do not readily form bonds.

Summary:

• Basic Valency: Atom loses electrons (positive valency).
• Negative Valency: Atom gains electrons (negative valency).
• Neutral Valency: Atom shares electrons (no charge).
• Variable Valency: Atom can have different valencies in different compounds.
• Zero Valency: Atom has a full outer shell and does not bond.

These types help us understand how elements interact and combine to form various compounds.

Calculation of Valency:

1. Understand the Basics:

   • Valency is the number of bonds an atom can form with other atoms.
   • It is determined by the number of electrons in the outermost shell (valence shell) of an atom.

2. Determine the Number of Valence Electrons:

   • Find the Atomic Number: The atomic number tells you the number of electrons in an atom.
   • Identify the Number of Valence Electrons: Look at the outermost shell to find out how many electrons are there.

3. Calculate Valency:

   • For Atoms with Less than 4 Electrons in the Outer Shell:
     • Valency = Number of electrons in the outer shell.
     • Example: Carbon (C) has 4 electrons in its outer shell. Its valency is 4.
   • For Atoms with More than 4 Electrons in the Outer Shell:
     • Valency = 8 - Number of electrons in the outer shell.
     • Example: Oxygen (O) has 6 electrons in its outer shell. Its valency is 8 - 6 = 2.

4. Check the Group Number:

   • Groups 1 and 2 (Left side of Periodic Table): Valency is the same as the group number.
     • Example: Sodium (Na) in Group 1 has a valency of 1.
   • Groups 13 to 18 (Right side of Periodic Table): Valency = 8 - Group number (for elements in Groups 15 to 17).
     • Example: Chlorine (Cl) in Group 17 has a valency of 8 - 7 = 1.

5. Examples:

   • Hydrogen (H):

     • Atomic Number: 1
     • Valence Electrons: 1
     • Valency: 1 (needs 1 more electron to complete its outer shell)

   • Oxygen (O):

     • Atomic Number: 8
     • Valence Electrons: 6
     • Valency: 8 - 6 = 2 (needs 2 more electrons to complete its outer shell)

   • Magnesium (Mg):

     • Atomic Number: 12
     • Valence Electrons: 2 (in the outer shell)
     • Valency: 2 (can lose these 2 electrons to form bonds)

Summary:

• Valency is calculated based on the number of electrons in the outermost shell.
• For fewer than 4 electrons: Valency equals the number of valence electrons.
• For more than 4 electrons: Valency equals 8 minus the number of valence electrons.
• Use the Group Number to quickly determine valency for many elements.

Isotopes:

1. Definition:

   • Isotopes are different forms of the same element. They have the same number of protons but different numbers of neutrons.

2. Same Element, Different Neutrons:

   • Protons: All isotopes of an element have the same number of protons.
   • Neutrons: Isotopes have different numbers of neutrons.
   • Example: Carbon has isotopes with different numbers of neutrons. Carbon-12 and Carbon-14 are both isotopes of carbon.

3. How to Identify Isotopes:

   • Atomic Number: The number of protons, which is the same for all isotopes of an element.
   • Mass Number: The sum of protons and neutrons. Different isotopes have different mass numbers.
   • Example:
     • Carbon-12: 6 protons + 6 neutrons = mass number 12
     • Carbon-14: 6 protons + 8 neutrons = mass number 14

4. Notation:

   • Isotopes are often written with the element’s name or symbol followed by the mass number.
   • Example: Carbon-12 can be written as $^{12}C$, and Carbon-14 can be written as $^{14}C$.

5. Uses of Isotopes:

   • Medical: Some isotopes are used in medical treatments and diagnostics. For example, radioactive iodine is used to treat thyroid problems.
   • Dating: Carbon-14 is used to date ancient fossils and artifacts.

Summary:

• Isotopes are different forms of the same element with the same number of protons but different numbers of neutrons.
• They have different mass numbers because of the different numbers of neutrons.
• Isotopes are used in various fields, including medicine and archaeology.