Grade 8
  1. Introduction to Chemistry  1.1. Nature and scope of chemistry  1.2. Importance of chemistry in daily life  1.3. Chemistry and its relation with other sciences  1.4. The role of chemistry in industry, medicine and the environment  1.5. Scientific investigation and experimental methods in chemistry
  2. Matter and its properties  2.1. Definition and classification of matter  2.2. Physical and chemical properties of matter  2.3. Law of conservation of matter  2.4. Extensive and Intensive Properties of Matter  2.5. Kinetic molecular theory of matter  2.6. States of matter: solids, liquids, gases, plasmas, and Bose-Einstein condensates  2.7. Changes in state and energy are involved  2.8. Diffusion and Brownian motion
  3. Elements, Compounds, and Mixtures  3.1. Definition and differences between elements, compounds, and mixtures  3.2. Atomic Structure and Atomic Number  3.3. Chemical symbols and formulas  3.4. Trends in the Periodic Table (Groups and Periods)  3.5. Classification of Elements: Metals, Nonmetals and Metalloids  3.6. Rare Earth Elements and Transition Metals  3.7. Types of mixtures: homogeneous and heterogeneous  3.8. Separation of Mixtures Using Physical and Chemical Methods
  4. Separation Techniques  4.1. Filtration, Sedimentation and Decantation  4.2. Evaporation and crystallization  4.3. Distillation and Fractional Distillation  4.4. Chromatography and its applications  4.5. Sublimation in the purification of compounds  4.6. The role of centrifugation in biochemical processes
  5. Atomic Structure  5.1. Dalton's atomic theory  5.2. Structure of the atom: protons, neutrons and electrons  5.3. Bohr's atomic model  5.4. Introduction to Quantum Mechanical Model  5.5. Electron Configuration and Energy Levels  5.6. Excitation and emission spectra  5.7. Isotopes and their applications
  6. Periodic Table and Chemical Trends  6.1. History and Development of the Periodic Table  6.2. Modern periodic law  6.3. Periodicity and trends in atomic and ionic sizes  6.4. Ionization Energy, Electron Affinity and Electronegativity  6.5. Introduction to Lanthanides and Actinides  6.6. Application of periodic trends in chemistry
  7. Chemical Bonding and Molecular Structure  7.1. Types of chemical bonds: ionic, covalent and metallic bonds  7.3. Polar and Nonpolar Molecules  7.4. Hydrogen bonding and its applications  7.5. Intermolecular forces: dipole-dipole, London dispersion force
  8. Chemical Reactions and Stoichiometry  8.1. Chemical Equations and Balance  8.2. Types of Chemical Reactions: Combination, Decomposition, Displacement and Redox  8.3. Introduction to stoichiometry and the mole concept  8.4. Limiting reactant in chemical equations  8.5. Reaction rate, catalyst, and factors affecting reaction rate
  9. Acids, Bases and Salts  9.1. Definition and Properties of Acids and Bases  9.2. Strong vs. Weak Acids and Bases  9.3. pH Scale and pH Calculation  9.4. Neutralization reactions  9.5. Buffer solutions and their importance  9.6. Applications of Acids, Bases and Salts in Daily Life
  10. Solutions and Solubility  10.1. Definition of Solution, Solute, and Solvent  10.2. Factors Affecting Solubility  10.3. Concentration units: molarity and molality  10.4. Saturated, unsaturated and supersaturated solutions  10.5. Concept of osmosis and reverse osmosis  10.6. Concentrative properties of solutions
  11. Gases and Gas Laws  11.1. Properties of gases  11.2. Introduction to Ideal and Real Gases  11.3. Boyle's Law, Charles' Law and Avogadro's Law  11.4. Ideal Gas Equation and Its Applications  11.5. Application of Gas Laws in Real Life
  12. Thermochemistry and energy transformation  12.1. Energy in Chemical Reactions  12.2. Exothermic vs. Endothermic Reactions  12.3. Heat and Enthalpy Changes  12.4. Calorimetry and heat transfer in reactions
  13. Metals and Nonmetals  13.1. Physical and Chemical Properties of Metals and Nonmetals  13.2. Reactivity Series of Metals  13.3. Corrosion and its prevention  13.4. Extraction of metals from ores  13.5. Uses of metals and non-metals in daily life
  14. Introduction to Organic Chemistry  14.1. Characteristics of carbon and organic compounds  14.2. Functional groups and their importance  14.3. Simple Hydrocarbons: Alkenes, Alkenes and Alkynes  14.4. Isomerism in organic compounds  14.5. Introduction to polymers and their uses
  15. Environmental Chemistry and Sustainability  15.1. Green Chemistry Principles  15.2. Effects of Chemical Pollution on Ecosystems  15.3. Acid rain and its effects  15.4. Ozone layer depletion and global warming  15.5. Waste Management and Recycling in Chemistry

Grade 8 → Atomic Structure ↓


Electron Configuration and Energy Levels


Atoms are the building blocks of matter. They are made up of protons, neutrons and electrons. The protons and neutrons are located at the center of the atom called the nucleus, while electrons orbit around the nucleus. Understanding how these electrons are arranged around the nucleus helps us understand how different elements interact and how matter behaves.

Introduction to electrons and energy levels

Electrons are tiny particles that carry a negative charge. They orbit the nucleus of an atom in regions called energy levels or shells. Each energy level can hold a certain number of electrons, and these electrons determine the chemical properties of an element.

Energy levels

Electrons are arranged in energy levels, sometimes called electron shells. The simplest atom, hydrogen, has one proton and, in its most basic form, one electron. This single electron is at the first energy level, the level closest to the nucleus.

Here's a simplified view of how the energy levels relate to the nucleus:

        
            
            
            
            Nucleus

In this diagram, the small black circle represents the nucleus of the atom, and the blue circle represents the energy level, with the electron depicted as another small circle at the energy level.

The number of electrons that fit into the first few energy levels are as follows:

  • First energy level: can hold up to 2 electrons.
  • Second energy level: can hold up to 8 electrons.
  • Third energy level: Can hold up to 18 electrons.
  • Fourth energy level: can hold up to 32 electrons, and so on.

Electron configuration

Electron configuration explains the distribution of electrons in the energy levels of an atom. Electrons will fill the lowest available energy levels before moving to higher levels, following a specific order known as the Aufbau principle.

Let's take a closer look at the structure of electrons in some common elements:

Example: hydrogen atom

Hydrogen has one electron. Its electron configuration is:

1s¹

The number "1" represents the first energy level. The letter "s" represents the type of orbital, and the sub-digit "1" indicates that there is one electron in that orbital.

Example: helium atom

Helium has two electrons. Its electron configuration is:

1s²

This means that both electrons completely fill the first energy level "s" orbital.

Example: oxygen atom

Oxygen has eight electrons. Its electron configuration is:

1s² 2s² 2p⁴

Oxygen's electrons fill the first energy level and partially fill the second energy level, including the "p" orbitals, which can hold up to six electrons.

Visualization of electron configuration

Visual illustrations can help us better understand electron configurations. Consider the following illustration of orbitals for a neon atom, which has 10 electrons:

        
            
            
            
            
            
            

            
            
            
            
            
            
            
            
            
            1st energy level (2 electrons)
            2nd energy level (8 electrons)

The concentric circles represent different energy levels (or shells), and the colored dots represent electrons in those levels.

Subshells and orbitals

In addition to energy levels, electrons are also organized into subshells: s, p, d, and f, each of which has a specific size and capacity for the electrons:

  • s subshell can have a maximum of 2 electrons.
  • p subshell can have up to 6 electrons.
  • d subshell can have up to 10 electrons.
  • f subshell can have up to 14 electrons.

Imagine the order of filling the sub-shells as given below:

        
            1s
            

            2s
            
            2p
            

            3s
            
            3p
            

            4s
            
            3d
            
            4p
            

            5s
            
            4d

How to write electron configuration

When writing electron configurations, following these steps can help simplify the process:

  1. Identify the number of electrons in the atom. The atomic number of the element tells you this.
  2. Fill the electrons into the appropriate energy levels and subshells following the Aufbau principle.
  3. Use a notation of numbers and letters to describe the configuration. Numbers represent energy levels, letters represent subshells, and the numbers above represent electron numbers.

For example, sodium (Na) has an atomic number of 11. So its electron configuration is:

1s² 2s² 2p⁶ 3s¹

This shows that sodium has 2 electrons in 1s sub-shell, 2 electrons in 2s sub-shell, 6 electrons in 2p sub-shell and 1 electron in 3s sub-shell.

Pauli exclusion principle and Hund's rule

Pauli Exclusion Principle: In an atom, no two electrons can have the same set of all four quantum numbers. This means that within the same orbital, the spins of the electrons must be opposite.

Hund's rule: Electrons must fill orbitals of the same energy level at a time, only then can they pair. This minimizes the repulsion between electrons and stabilizes the atom.

Understanding valence electrons

The outermost electrons in an atom are called valence electrons. These are the electrons available for bonding, and they determine the chemical properties of the element.

For example, in an oxygen atom (symbol: O) with the electron configuration 1s² 2s² 2p⁴, the valence electrons are the 2s² and 2p⁴ electrons, making a total of 6 valence electrons.

The number of valence electrons is important in predicting how an atom will react with other atoms. Elements with a full outer energy level are usually inert or nonreactive, such as the noble gases (e.g., helium, neon, argon).

Summary

Electron configuration is an essential concept in chemistry because it helps explain the chemical behavior of an element. By understanding the energy levels, subshells, and electron configurations of atoms, you can understand why elements form specific types of chemical bonds and how they interact with each other. Principles such as the Pauli exclusion principle, Hund's rule, and the concept of valence electrons give us information on a microscopic scale to predict and appreciate the macroscopic world of chemistry.


Grade 8 → 5.5


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Electron Configuration and Energy Levels

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