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


Atoms are the building blocks of matter. They are the smallest unit of an element that retains the properties of that element. In this discussion, we will explore the basic principles of atomic structure, including the components of an atom, the arrangement of electrons, and how these aspects affect the behavior of different elements.

Basic components of an atom

An atom is composed of three main particles:

  • Protons - These positively charged particles are found in the nucleus of the atom.
  • Neutrons - Found with protons in the nucleus, neutrons have no charge; they are neutral.
  • Electrons - These negatively charged particles orbit the nucleus at different energy levels.

Center

The nucleus is the core of the atom and contains both protons and neutrons. The number of protons in the nucleus determines an element's atomic number, which is important in identifying the element. For example, the element with 6 protons is carbon, represented by the symbol C and atomic number 6.

NucleusPN

Illustration of a simple atomic structure showing the nucleus containing protons (p) and neutrons (n).

Electrons and energy levels

Electrons revolve around the nucleus in specific paths called orbitals or energy levels. These levels can hold a certain number of electrons:

  • The first energy level can have a maximum of 2 electrons.
  • The second one can hold up to 8 electrons.
  • The third shell can also have 8 electrons, but in some cases it can extend to 18.

A diagram showing circles representing an electron cloud with different energy levels.

The arrangement of electrons in these energy levels determines how atoms interact with each other. For example, if an atom's outermost energy level has fewer electrons than its maximum capacity, the atom can form bonds with other atoms to achieve a full outer shell.

Valence electrons

The electrons in the outermost energy level are known as valence electrons. These electrons play an important role in the chemistry of an atom because they are involved in forming chemical bonds. For example, oxygen has 6 valence electrons and needs 2 more to complete its outer energy level. This is why oxygen often bonds with elements such as hydrogen to form water (H 2 O).

Oxygen (O): valence electrons: 6 tries to gain 2 more electrons -> forms 2 bonds

Isotopes

Isotopes are different forms of the same element that have different numbers of neutrons but the same number of protons. For example, carbon has several isotopes, such as carbon-12 and carbon-14, which are identified by their mass number (number of protons + neutrons).

Carbon Isotopes: - Carbon-12: 6 protons, 6 neutrons - Carbon-14: 6 protons, 8 neutrons

Anions

Atoms can become ions by losing or gaining electrons. When an atom gains an electron, it becomes a negatively charged ion (anion). Conversely, when it loses an electron, it becomes a positively charged ion (cation). Understanding ions is important to understand how compounds are formed. For example, in salt (NaCl), the sodium atom loses an electron to become Na +, while the chlorine atom gains an electron to become Cl-.

Sodium (Na) loses 1 electron: Na -> Na + Chlorine (Cl) gains 1 electron: Cl -> Cl -

Atomic model

Over time, scientists have developed several models to describe the structure of the atom:

Dalton's model

John Dalton proposed that atoms were solid, indivisible spheres. Although simple, this model laid the groundwork for modern atomic theory.

Thomson's model

J.J. Thomson discovered the electron and proposed the "plum pudding" model, in which electrons were scattered within a positively charged "soup". This was later proven wrong, but led to important advances in atomic theory.

Rutherford's model

Ernest Rutherford proposed that the atom consisted of a small nucleus around which electrons revolved, just as planets revolved around the sun. This model introduced the concept of the atomic atom.

Bohr's model

Niels Bohr improved upon Rutherford's model by suggesting that electrons orbited around the nucleus in specific paths or energy levels that arose from quantized energy levels.

Quantum mechanical model

The modern understanding of the atom is based on the quantum mechanical model. This model states that electrons exist in clouds around the nucleus, rather than in fixed orbits defined by regions where there is a high probability of finding the electron.

Conclusion

Understanding atomic structure reveals the fundamental principles of chemistry and the science behind the interactions and bonds that create the diversity of matter in the universe. The study of atoms and their structures continues to be a vital aspect of scientific progress and applications.

By understanding the basic structures, behaviors, and models of atoms, we can better understand the complexities of the elements and compounds around us.


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Atomic Structure

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