Grade 7
  1. Introduction to Chemistry  1.1. What is Chemistry?  1.2. Importance of chemistry in daily life  1.3. Branches of chemistry  1.4. Scientific Method in Chemistry  1.5. Laboratory Safety Rules and Precautions  1.6. Common Laboratory Equipment and Their Uses  1.7. Units of measurement in chemistry
  2. Matter and its properties  2.1. Definition of Matter  2.2. Classification of matter  2.3. Properties of matter  2.3.1. Physical properties  2.3.2. Chemical properties  2.4. States of matter  2.4.1. Solid state  2.4.2. Fluid state  2.4.3. Gaseous state  2.4.4. Plasma and Bose–Einstein condensates  2.5. Changes in the states of matter  2.5.1. Melting and freezing  2.5.2. Evaporation and Condensation  2.5.3. Sublimation and deposition  2.6. Density and its applications  2.7. Diffusion and its importance
  3. Elements, Compounds, and Mixtures  3.1. Definition of the element  3.2. Classification of elements  3.3. Metals, Nonmetals and Metalloids  3.4. Noble gases and their properties  3.5. Introduction to the Periodic Table  3.6. Definition of Compound  3.7. Properties of Compounds  3.8. Introduction to Chemical Formulas  3.9. Definition of Mixture  3.10. Types of mixtures  3.10.1. Homogeneous mixture  3.10.2. Heterogeneous mixtures  3.11. Difference Between Compounds and Mixtures  3.12. Solutions, Colloids and Suspensions
  4. Separation of mixtures  4.1. Need for separation of mixtures  4.2. Separation methods  4.2.1. Filtration  4.2.2. Sedimentation and decantation  4.2.3. Evaporation  4.2.4. Crystallization  4.2.5. Distillation  4.2.6. Chromatography  4.2.7. Magnetic Separation  4.2.8. Centrifugation
  5. Atomic Structure  5.1. Introduction to Atoms  5.2. Structure of the atom  5.2.1. Nucleus and electron cloud  5.3. Subatomic particles  5.3.1. Proton  5.3.2. Neutron  5.3.3. Electrons  5.4. Atomic number and mass number  5.5. Isotopes and their uses  5.6. Ions and ion formation
  6. Periodic table  6.1. Introduction to the Periodic Table  6.2. Groups and periods  6.3. Trends in the Periodic Table  6.3.1. Atomic Size  6.3.2. Metallic and non-metallic character  6.3.3. Electronegativity  6.3.4. Reactivity of the elements  6.4. Alkali metals and their properties
  7. Chemical bond  7.1. Why do atoms form bonds?  7.2. Types of chemical bonds  7.2.1. Ionic bond  7.2.2. Covalent bond  7.2.3. Metal Bond  7.3. Properties of Ionic and Covalent Compounds  7.4. Intermolecular forces
  8. Chemical reactions  8.1. Introduction to Chemical Reactions  8.2. Signs of a chemical reaction  8.3. Types of Chemical Reactions  8.3.1. Combination Reactions  8.3.2. Decomposition reactions  8.3.3. Displacement reactions  8.3.4. Double displacement reactions  8.3.5. Combustion reactions  8.4. Law of conservation of mass  8.5. Balancing simple chemical equations
  9. Acids, Bases and Salts  9.1. Definition of Acid and Base  9.2. Properties of Acids  9.3. Properties of bases  9.4. pH Scale and Indicators  9.5. Neutralization reactions  9.6. Common Acids and Bases in Everyday Life  9.7. Preparation and use of salts  9.8. Strong and weak acids and bases
  10. Solutions and Solubility  10.1. Definition of Solution  10.2. Components of the solution  10.2.1. Solvent  10.2.2. solute  10.3. Factors Affecting Solubility  10.4. Concentration of solutions  10.5. Saturated and unsaturated solutions  10.6. Colloids and suspensions  10.7. Solubility curve and its interpretation
  11. Air and atmosphere  11.1. Composition of air  11.2. Properties of gases in the air  11.3. Importance of Oxygen and Carbon Dioxide  11.4. Air pollution and its effects  11.5. Greenhouse effect and global warming  11.6. Ways to reduce air pollution  11.7. Ozone layer and its importance
  12. Water and its importance  12.1. Properties of Water  12.2. Water as the universal solvent  12.3. Importance of water for life  12.4. Water pollution and its effects  12.5. Water Purification  12.5.1. Filtration  12.5.2. boil  12.5.3. Chlorination in water purification  12.5.4. reverse osmosis  12.6. Hard and soft water  12.7. Water cycle and its importance
  13. Fuel and energy  13.1. Types of fuel  13.1.1. Fossil fuels  13.1.2. Renewable energy sources  13.2. Combustion and energy release  13.3. Global warming and its effects  13.4. Alternative sources of energy  13.5. Ways to conserve energy
  14. Metals and Nonmetals  14.1. Physical and chemical properties of metals  14.2. Physical and Chemical Properties of Non-Metals  14.3. Difference Between Metals and Nonmetals  14.4. Uses of metals and nonmetals  14.5. Corrosion of metals and its prevention  14.6. Alloys and their importance
  15. Plastics and polymers  15.1. Natural and synthetic polymers  15.2. Properties of plastics  15.3. Biodegradable and Non-Biodegradable Plastics  15.4. Environmental impact of plastic  15.5. Recycling and waste management in plastics and polymers  15.6. Daily Uses of Polymers
  16. Introduction to Organic Chemistry  16.1. Basic definitions of organic chemistry  16.2. Carbon Compounds in Daily Life  16.3. Hydrocarbons  16.4. Simple functional groups (alcohols and carboxylic acids)  16.5. Importance of organic compounds in industry

Grade 7 → Chemical bond → Types of chemical bonds ↓


Ionic bond


The chemical bond is an essential concept in chemistry that describes how atoms come together to form molecules and compounds. In simple terms, it is the force that holds atoms together. One of the fundamental types of chemical bonds is the ionic bond.

Introduction to ionic bonding

Ionic bonding occurs when atoms transfer electrons to obtain a full outer shell of electrons, which is typically eight for most elements. This transfer of electrons results in the formation of ions. Ions are atoms that have gained or lost electrons and, as a result, have a positive or negative charge.

How ionic bonds are formed

Ionic bonds form between metals and nonmetals. Metals, found on the left side of the periodic table, lose electrons and form positive ions, also called cations. Nonmetals, found on the right side of the periodic table, gain electrons and form negative ions, also called anions. This gain or loss of electrons allows both atoms to achieve electronic configurations such as the noble gases, which are very stable.

        Na → Na⁺ + e⁻ (Sodium loses an electron) Cl + e⁻ → Cl⁻ (Chlorine gains an electron)

When these oppositely charged ions come together, they attract each other due to electrostatic forces, forming an ionic bond.

Example of ionic bond: sodium chloride

A classic example of ionic bonding is the formation of sodium chloride (NaCl). Sodium (Na) has one electron in its outer shell, while chlorine (Cl) has seven electrons in its outer shell. Sodium can obtain a full outer shell by losing an electron, turning into a sodium ion (Na⁺). Chlorine can obtain a full shell by gaining an electron, turning into a chloride ion (Cl⁻).

        Na (2, 8, 1) + Cl (2, 8, 7) → Na⁺ (2, 8) + Cl⁻ (2, 8, 8) → NaCl

Visualization of ionic bonds

Consider the following example of sodium chloride formation as a visual representation.

No Chlorine

In this view, the sodium atom donates an electron to the chlorine atom, forming an ionic bond, resulting in NaCl.

Properties of ionic compounds

Ionic compounds have distinctive properties due to the nature of the ionic bonds:

  • High melting and boiling points: Ionic bonds are strong, so ionic compounds usually have high melting and boiling points.
  • Solubility in water: Many ionic compounds dissolve in water because the polar water molecules can separate the ions from each other.
  • Electrical conductivity: Ionic compounds do not conduct electricity in the solid state. However, when they are melted or dissolved in water, they conduct electricity because the ions move freely.
  • Brittleness: Ionic compounds are generally brittle, breaking along the plane of the ions due to the arrangement of positive and negative ions.

Boiling point display

Consider table salt (NaCl). Its melting point is about 801°C, which is much higher than many other substances. This high melting point is due to the strong electrostatic forces holding the ions together.

More examples of ionic compounds

Magnesium oxide (MgO)

Magnesium oxide is another example of an ionic compound. Magnesium (Mg) has two electrons in its outer shell and loses both to achieve a stable electronic arrangement like neon. Oxygen (O) needs two electrons to complete its outer shell, just like neon.

        Mg (2, 8, 2) + O (2, 6) → Mg²⁺ (2, 8) + O²⁻ (2, 8) → MgO

Aluminum oxide (Al2O3)

Aluminum oxide is made of aluminum and oxygen. Aluminum (Al) loses three electrons to form Al³⁺ ions, while each oxygen atom needs two electrons to form O²⁻ ions. Therefore, two aluminum ions combine with three oxygen ions to form Al2O3.

        2 Al (2, 8, 3) + 3 O (2, 6) → 2 Al³⁺ (2, 8) + 3 O²⁻ (2, 8) → Al₂O₃

Conclusion

Ionic bonding is a basic and fascinating type of chemical bond that is essential to understanding physical properties and reactions. Ionic compounds exist all around us, from everyday table salt to various minerals. The transfer of electrons between metals and nonmetals and the resulting attraction between the charged ions is an important part of how elements form stable compounds.


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Ionic bond

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