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 → Elements, Compounds, and Mixtures ↓


Chemical symbols and formulas


Chemistry is the scientific study of substances and compounds, and how they interact, combine, and change to form new substances. At the core of chemistry are chemical symbols and formulas. These are the fundamental tools that chemists use to express information about atoms, molecules, and chemical reactions in a concise way. In this article, we will explore the concept of chemical symbols and formulas, focusing specifically on elements, compounds, and mixtures.

Chemical symbols

Chemical symbols are short, one- or two-letter symbols used to represent chemical elements. Each element in the periodic table is given a unique symbol based on its English or Latin name. For example:

  • Hydrogen is represented as H
  • Oxygen is represented as O
  • Carbon is represented by C
  • Gold, which derives from the Latin word "aurum," is represented as Au.

This symbol system helps scientists and students around the world communicate chemical information clearly across language barriers.

H - Hydrogen O - Oxygen C - Carbon Au - Gold

Chemical formula

Chemical formulas provide concise expressions that show the structure of compounds. They include symbols for the constituent elements and numerical sub-digits indicating the number of atoms of each element present in the compound. Formulas show how atoms of different elements form chemical bonds to produce various compounds. There are two main types of chemical formulas:

Empirical formula

The empirical formula shows the simplest ratio of elements in a compound. Unlike other formulas, it does not show the total number of atoms, but rather the relative proportion of each element. For example, the empirical formula of glucose is CH2O, which indicates a 1:2:1 ratio of carbon, hydrogen, and oxygen atoms.

Molecular formula

The molecular formula shows the actual number of atoms of each element in a molecule. It provides more details than the empirical formula. For example, the molecular formula of glucose is C6H12O6, which means that a molecule of glucose contains 6 carbon atoms, 12 hydrogen atoms, and 6 oxygen atoms.

Elements

An element is a pure substance that cannot be broken down into simpler substances. Elements are composed of similar atoms that share common chemical properties. Each element is designated by a unique chemical symbol.

Compounds

When two or more elements combine chemically, they form a compound. All compounds have a certain ratio and their unique properties differ from the properties of the individual elements that make them up. For example:

  • Water (H2O) is composed of two hydrogen atoms for every oxygen atom. Its properties differ from those of its constituent elements.
  • Sodium chloride (NaCl), commonly known as table salt, contains equal numbers of sodium and chlorine atoms.
H2O Sodium chloride

Mixture

Unlike compounds, mixtures are formed when two or more substances combine without chemical bonds. Mixtures can be separated by physical methods, and each substance in a mixture retains its own properties. Mixtures are classified into two types:

Homogeneous mixture

In homogeneous mixtures, the components are evenly distributed, and it is difficult to separate the individual substances. An example of a homogeneous mixture is a solution like salt water. Even though it contains salt and water, the two are mixed evenly, forming a uniform solution.

Heterogeneous mixtures

In heterogeneous mixtures, the different components are not evenly distributed and can be easily identified. For example, in a salad that contains lettuce, tomatoes, and cucumbers, you can easily identify the different ingredients.

Salt water (homogeneous) Salad (heterogeneous)

Chemical reactions

Chemical reactions describe processes where substances (reactants) change into different substances (products). Chemical equations use chemical symbols and formulas to represent these reactions. An example of a chemical reaction is the burning of methane:

CH4 + 2O2 → CO2 + 2H2O

In this reaction, methane (CH4) and oxygen (O2) react to form carbon dioxide (CO2) and water (H2O).

Chemists use these equations to understand and predict the outcome of chemical reactions. It is important to balance these equations to match the number of atoms on both sides, following the law of conservation of mass.

Remember, chemical symbols and formulas are the language of chemistry. They allow scientists to communicate complex information efficiently and are essential in the study of chemistry.


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Chemical symbols and formulas

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