Chapter 1: Atomic Structure, Chemical Bonding, and Solutions Notes

Hey, Welcome to Rajasthan Polytechnic (BTER). This blog post is provide you notes of Polytechnic 1st Semester Chemistry Chapter 1: Atomic Structure, Chemical Bonding, and Solutions. \

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1. Introduction to Atomic Structure: Atoms are the basic units of matter and the building blocks of everything around us. Understanding their structure helps us explain how substances interact and bond.

• Atoms: The smallest units of matter that retain the chemical properties of an element.
• Basic Concept: Atoms are composed of a central nucleus surrounded by electrons.
• Importance: Understanding atomic structure helps explain how and why substances interact, bond, and form new materials.

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2. Constituents of the Atom: Atoms consist of three main types of subatomic particles:

• Protons: Positively charged particles found in the nucleus (center) of the atom.

  • Charge: Positive (+1).
  • Location: In the nucleus.
  • Mass: Approximately 1 atomic mass unit (amu).
  • Role: Defines the element (e.g., Carbon has 6 protons).

• Neutrons: Neutral particles (no charge) also located in the nucleus.

  • Charge: Neutral (0).
  • Location: In the nucleus.
  • Mass: Approximately 1 amu.
  • Role: Contributes to the atomic mass and stability of the nucleus.

• Electrons: Negatively charged particles that orbit the nucleus in specific energy levels or shells.

• Charge: Negative (-1).
• Location: Orbiting the nucleus in energy levels or shells.
• Mass: Much smaller than protons and neutrons.
• Role: Participates in chemical bonding and reactions.

The number of protons in the nucleus defines the element (e.g., hydrogen has 1 proton, carbon has 6 protons).

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3. Drawbacks of Rutherford’s Model

• Rutherford’s Model: Depicts electrons orbiting a dense nucleus, similar to planets orbiting the sun. or Rutherford’s model described the atom as having a tiny, dense nucleus with electrons orbiting around it, like planets around the sun. However, this model had some issues:
• Drawbacks:
• Instability Issue: According to classical physics, electrons should lose energy and spiral into the nucleus, causing the atom to collapse. This doesn’t happen in reality.
• Spectral Lines: Rutherford’s model couldn’t explain why atoms emit light in specific colors (line spectra).

Rutherford's Model And It's Drawback Video - Watch Now

4. Bohr’s Theory: Niels Bohr proposed a new model to address Rutherford’s issues:

• Key Concepts:

  • Fixed Orbits: Electrons orbit the nucleus in fixed paths or energy levels, called shells.
  • Quantized Energy: Electrons can only gain or lose energy in specific amounts, moving between these fixed orbits.
  • Spectral Lines: When an electron jumps from a higher energy level to a lower one, it emits energy as light, which is seen as lines in the spectrum.

  This model explained the discrete lines observed in the hydrogen spectrum.

• Bohr’s Model:

• Energy Levels: Electrons occupy specific energy levels (orbits) around the nucleus.
• Energy Absorption/Emission: Electrons absorb energy to move to higher levels or emit energy to drop to lower levels.

Bohr's Model Video - Watch Now

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5. Hydrogen Spectrum and Bohr’s Explanation

• Hydrogen Spectrum: When hydrogen atoms are energized, they emit light of specific wavelengths. This light appears as distinct lines in the spectrum.
• Bohr’s Explanation:
• Electron Transitions: Electrons in a hydrogen atom move between quantized energy levels.
• Emission Lines: When an electron falls from a higher to a lower energy level, it emits a photon (light) with a specific wavelength, creating the spectral lines.

When hydrogen atoms are excited (e.g., by heating), they emit light of specific colors. This emission produces distinct lines known as the hydrogen spectrum. Bohr explained this by:

• Electrons jumping between fixed orbits in the hydrogen atom.
• Each jump releases energy as light, corresponding to specific wavelengths, creating the observed spectral lines.

Hydrogen Spectrum Explanation Based on Bohr's Model Video - Watch Now

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6. Heisenberg’s Uncertainty Principle

• Concept:

  • Principle: It is impossible to precisely measure both the position and momentum (speed and direction) of an electron at the same time.
  • Implication: Electrons are described by probabilistic regions (orbitals) rather than fixed paths.

• Impact on Atomic Model:

• Orbitals: Instead of fixed orbits, electrons are found in regions with a probability distribution around the nucleus.

Werner Heisenberg stated that it is impossible to know both the exact position and the exact speed (momentum) of an electron simultaneously:

• The more precisely you know the position of an electron, the less precisely you can know its speed, and vice versa.
• This principle shows that electrons do not move in fixed orbits but rather exist in regions of space called orbitals.

Heisenberg Uncertainty Principle Video - Watch Now

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7. Quantum Numbers: Quantum numbers describe the properties and positions of electrons in an atom:

• Principal Quantum Number (n):

  • Description: Indicates the main energy level or shell (e.g., n=1, 2, 3).
  • Energy: Higher n values correspond to higher energy levels.

• Azimuthal Quantum Number (l):

  • Description: Describes the shape of the orbital.
  • Types:
    • s-Orbitals: Spherical (l=0)
    • p-Orbitals: Dumbbell-shaped (l=1)
    • d-Orbitals: Cloverleaf-shaped (l=2)
    • f-Orbitals: More complex shapes (l=3)

• Magnetic Quantum Number (m):

  • Description: Specifies the orientation of the orbital in space.
  • Range: For each l value, m can range from -l to +l.

• Spin Quantum Number (s):

• Description: Indicates the direction of electron spin (either +½ or -½).
• Significance: Helps in determining the magnetic properties of the atom.

Quantum Numbers Video - Watch Now

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8. Aufbau Principle

• Concept:
• Principle: Electrons fill orbitals starting from the lowest energy level first before moving to higher levels.
• Order: The general order is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.
• Application: Helps in writing the electronic configuration of elements.

The Aufbau principle states that electrons fill the lowest energy orbitals first before moving to higher energy levels:

• Electrons occupy orbitals in the order of increasing energy (1s, 2s, 2p, 3s, 3p, etc.).
• This helps in determining the electronic configuration of an atom.

Aufbau Principle Video - Watch Now

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9. Orbital Concepts and Shapes: Orbitals are regions around the nucleus where electrons are likely to be found:

• s-Orbitals:

  • Shape: Spherical.
  • Number: 1 s-orbital per energy level.

• p-Orbitals:

  • Shape: Dumbbell-shaped, with three orientations (px, py, pz).
  • Number: 3 p-orbitals (px, py, pz) per energy level starting from n=2.

• d-Orbitals:

  • Shape: Cloverleaf-shaped and more complex.
  • Number: 5 d-orbitals per energy level starting from n=3.

• f-Orbitals:

• Shape: Even more complex shapes, involved in multi-electron atoms.
• Number: 7 f-orbitals per energy level starting from n=4.

Orbit and Orbital Difference Video - Watch Now

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10. Electronic Configuration

• Definition: The distribution of electrons among orbitals in an atom.

• Rules:

  • Aufbau Principle: Electrons fill the lowest energy orbitals first.
  • Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers.
  • Hund’s Rule: Electrons occupy degenerate orbitals singly before pairing up.

• Examples:

• Hydrogen (H): 1 electron → 1s¹
• Carbon (C): 6 electrons → 1s² 2s² 2p²
• Oxygen (O): 8 electrons → 1s² 2s² 2p⁴

The configuration helps in predicting the chemical behavior and bonding of elements.

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11. Concept of Chemical Bonding

• Chemical Bond: An attractive force that holds atoms together in a molecule.

Chemical bonding occurs when atoms combine to form molecules. There are three main types of bonds:

1. Ionic Bond: Formed when one atom donates electrons to another, creating oppositely charged ions that attract each other (e.g., NaCl).
2. Covalent Bond: Formed when atoms share electrons to achieve a full outer shell (e.g., H₂O).
3. Metallic Bond: Found in metals, where electrons are free to move around a lattice of positive metal ions, allowing conductivity (e.g., copper).

• Types:
  • Ionic Bond:
    • Formation: Transfer of electrons from one atom to another, forming ions.
    • Example: Sodium (Na) donates an electron to Chlorine (Cl) to form Na⁺ and Cl⁻, which attract each other.
  • Covalent Bond:
    • Formation: Sharing of electrons between atoms.
    • Example: Two hydrogen atoms share their electrons to form H₂.
  • Metallic Bond:
    • Formation: A "sea of electrons" surrounds positively charged metal ions, allowing electrons to move freely.
    • Properties: Conductivity and malleability (e.g., in copper).

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12. Solution: Solute, Solvent, and Solution

• Solute:
  • Definition: The substance that is dissolved in a solution.
  • Example: Salt in saltwater.
• Solvent:
  • Definition: The substance that dissolves the solute.
  • Example: Water in saltwater.
• Solution:
  • Definition: A homogeneous mixture of solute and solvent.
  • Characteristics: The solute is evenly distributed throughout the solvent.
  • Example: Saltwater.

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13. Methods to Express Concentration of Solutions

1. Molarity (M):

   • Definition: Number of moles of solute per liter of solution.
   • Example: Dissolving 1 mole of NaCl in 1 liter of water gives a 1 M solution.

2. Normality (N):

   • Definition: Number of equivalents of solute per liter of solution.
   • Application: Useful for reactions involving acids and bases.
   • Formula: $$N = \frac{{\text{{equivalents of solute}}}}{{\text{{volume of solution (in liters)}}}}$$

3. Molality (m):

   • Definition: Number of moles of solute per kilogram of solvent.
   • Formula: $$m = \frac{{\text{{moles of solute}}}}{{\text{{mass of solvent (in kg)}}}}$$
   • Example: 1 mole of NaCl dissolved in 1 kg of water is 1 molal.

4. Parts Per Million (ppm):

   • Definition: Mass of solute per million parts of solution.
   • Formula: $$\text{{ppm}} = \frac{{\text{{mass of solute}}}}{{\text{{mass of solution}}}} \times 10^6$$
   • Application: Used for very dilute solutions.

5. Mass Percentage:

   • Definition: Mass of solute divided by the total mass of the solution, multiplied by 100.
   • Formula: $$\text{{Mass \%}} = \frac{{\text{{mass of solute}}}}{{\text{{mass of solution}}}} \times 100$$

6. Volume Percentage:

   • Definition: Volume of solute divided by the total volume of solution, multiplied by 100.
   • Formula: $$\text{{Volume \%}} = \frac{{\text{{volume of solute}}}}{{\text{{volume of solution}}}} \times 100$$

7. Mole Fraction:

   • Definition: Ratio of the number of moles of a component to the total number of moles in the solution.
   • Formula: $$\text{{Mole fraction of A}} = \frac{{\text{{moles of A}}}}{{\text{{total moles of all components}}}}$$

8. pH:

   • Definition: Measure of the hydrogen ion concentration in a solution. or A measure of the acidity or basicity of a solution.
   • Scale:
     • Acidic Solutions: pH < 7
     • Neutral Solutions: pH = 7
     • Basic Solutions: pH > 7
   • Formula: $$\text{{pH}} = -\log[\text{{H⁺}}]$$
   • Importance: pH helps in determining the acidity or alkalinity of a solution, crucial for many chemical reactions and processes.

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These detailed notes should provide a comprehensive understanding of atomic structure, chemical bonding, and solutions. These Notes are Made by Garima Kanwar according to Syllabus provided by BTER (Board of Technical Education Rajasthan, Jodhpur) for Polytechnic 1st Semester Students.

Hope this post will help you. Write down Your query and Suggestion in Comment Section.

Thankyou

Regards

Garima Kanwar