Grade 10
  1. Matter and its properties  1.1. States of matter  1.2. Physical and chemical properties of matter  1.3. Classification of substances (elements, compounds, mixtures)  1.4. Changes in Matter (Physical and Chemical Changes)  1.5. Law of conservation of mass and law of definite proportions
  2. Atomic Structure  2.1. Historical development of the atomic model  2.2. Subatomic particles (protons, neutrons, electrons)  2.3. Atomic number, mass number and isotopes  2.4. Electronic Configuration and Energy Levels  2.5. Valency, ion formation and oxidation state
  3. Periodic table  3.1. History and Development of the Periodic Table  3.2. Modern Periodic Law and Periodic Trends  3.3. Groups and Periods in the Periodic Table  3.4. Trends in the Periodic Table  3.5. Metals, non-metals, metalloids and their properties  3.6. Noble gases and their chemical inertness
  4. Chemical bond  4.1. Formation of Chemical Bonds (Octet Rule)  4.2. Ionic Bonding and Properties of Ionic Compounds  4.3. Covalent Bond and Properties of Covalent Compounds  4.4. Metallic bond and its properties  4.5. Molecular geometry and VSEPR theory  4.6. Polarity and dipole moment of molecules  4.7. Intermolecular forces
  5. Chemical Reactions and Equations  5.1. Types of Chemical Reactions  5.2. Writing and balancing chemical equations  5.3. Law of conservation of mass in chemical reactions  5.4. Factors Affecting Chemical Reactions  5.5. Catalysts and their role in chemical reactions  5.6. Exothermic and Endothermic Reactions
  6. Stoichiometry and Chemical Calculations  6.1. Mole concept and Avogadro number  6.2. Molar Mass and Gram-Mole Conversions  6.3. Empirical and molecular formula  6.4. Stoichiometric calculations and chemical yield  6.5. Limiting and Excess Reactants
  7. Acids, Bases and Salts  7.1. Properties and Definitions of Acids and Bases (Arrhenius, Bronsted-Lowry, Lewis)  7.2. pH Scale, pOH, and Indicators  7.3. Neutralization reactions and salt formation  7.4. Strength of Acids and Bases (Strong vs. Weak Acids and Bases)  7.5. Common Acids and Bases in Daily Life  7.6. Salts, their solubility and applications
  8. Metals and Nonmetals  8.1. Physical and chemical properties of metals  8.2. Physical and Chemical Properties of Non-Metals  8.3. Reactivity Series of Metals and Displacement Reactions  8.4. Extraction of metals from ores  8.5. Corrosion, its causes and prevention  8.6. Alloys, their composition and uses
  9. Carbon and its compounds  9.1. Allotropes of Carbon  9.2. Hydrocarbons and their classification (alkanes, alkenes, alkynes)  9.3. Functional groups in organic compounds  9.4. Nomenclature of organic compounds (IUPAC system)  9.5. Isomerism in organic chemistry (structural and geometrical)  9.6. Important organic reactions (combustion, addition, substitution, polymerization)  9.7. Applications of organic compounds in industry and daily life
  10. Environmental Chemistry  10.1. Air Pollution (Sources, Effects and Control)  10.2. Water Pollution (Causes, Effects and Remedies)  10.3. Greenhouse effect and global warming  10.4. Ozone layer depletion and its effects  10.5. Green Chemistry and Its Applications  10.6. sustainable chemistry practice
  11. Thermochemistry  11.1. Heat and Energy Changes in Chemical Reactions  11.2. Exothermic and Endothermic Reactions  11.3. Enthalpy, enthalpy change, and heat of reaction  11.4. Specific heat capacity and calorimetry  11.5. Energy Changes in Combustion Reactions
  12. Electrochemistry  12.1. Electrolytes and non-electrolytes  12.2. Electrolysis and its applications (electroplating, extraction of metals)  12.3. Redox reactions (oxidation and reduction)  12.4. Galvanic cell and electrochemical series  12.5. Corrosion and electrochemical prevention methods
  13. Chemical kinetics and equilibrium  13.1. Rate of chemical reactions and its measurement  13.2. Factors affecting the reaction rate (catalyst, temperature, concentration)  13.3. Reversible and irreversible reactions  13.4. Dynamic equilibrium and equilibrium constant  13.5. Le Chatelier's principle and its applications
  14. Gases and Gas Laws  14.1. Properties of gases and kinetic molecular theory  14.2. Boyle's Law (Pressure-Volume Relation)  14.3. Charles's Law  14.4. Gay-Lussac's Law  14.5. Combined Gas Law and Ideal Gas Law  14.6. Real Gases vs Ideal Gases
  15. Nuclear Chemistry  15.1. Introduction to Radioactivity and Nuclear Reactions  15.2. Types of radioactive decay (alpha, beta, gamma)  15.3. Half-life and applications of radioactive isotopes  15.4. Nuclear fission and fusion reactions  15.5. Nuclear power generation and its environmental impact

Grade 10 ↓


Periodic table


Introduction to the periodic table

The periodic table is a systematic chart of the elements. It shows us all the known chemical elements and arranges them based on their atomic number, which is the number of protons in the nucleus of an atom. The periodic table is like a map for chemists. It helps predict how the elements will behave and react with each other.

History of the periodic table

The history of the periodic table dates back to the early 19th century. One of the most important figures in its development was the Russian chemist Dmitry Mendeleev. In 1869, Mendeleev created a table in which the elements were arranged based on increasing atomic mass and similar chemical properties. Although Mendeleev's table has evolved over time, his contributions laid the groundwork for the modern periodic table.

Structure of the periodic table

The current periodic table is made up of rows and columns. The rows are called periods, and the columns are called groups or families. Each element in the table is represented by its chemical symbol, atomic number, and atomic mass.

Here's a simplified visual example of a section of the periodic table:

H 1 He 2 Took 3 Happen 4 B 5

Understanding atomic number and atomic mass

Atomic number

The atomic number is the number of protons in the nucleus of an atom. Each element has a specific atomic number, which determines its place in the periodic table. For example, hydrogen has an atomic number of 1, so it comes first in the table. Helium has an atomic number of 2, so it comes after hydrogen.

Atomic mass

Atomic mass is the average mass of an element's isotopes, weighted by their natural abundance. Atomic mass is usually expressed in atomic mass units (amu). For example, the atomic mass of carbon is about 12.01 amu.

Groups in the periodic table

The columns in the periodic table are called groups. Elements in the same group have similar properties because they have the same number of electrons in their outer shell. Here are some important groups:

Group 1: Alkali metals

Alkali metals include elements such as lithium (Li), sodium (Na), and potassium (K). These metals are very reactive, especially with water. They have one electron in their outermost shell.

Li, Na, K, Rb, Cs, Fr

Group 2: Alkaline earth metals

The alkaline earth metals include beryllium (Be), magnesium (Mg), and calcium (Ca). These metals are reactive, but not as much as the alkali metals. They have two electrons in their outermost shell.

Be, Mg, Ca, Sr, Ba, Ra

Group 17: The halogens

Halogens include fluorine (F), chlorine (Cl) and bromine (Br). These elements are very reactive nonmetals and have seven electrons in their outermost shell. Halogens are known to react with metals to form salts.

F, Cl, Br, I, At, Ts

Group 18: Noble gases

Noble gases include helium (He), neon (Ne), and argon (Ar). These gases are much less reactive because they have a full set of electrons in their outer shell. Noble gases are often used in lighting and welding because of their stability.

He, Ne, Ar, Kr, Xe, Rn, Og

Periods in the periodic table

The rows in the periodic table are called periods. There are 7 periods in the periodic table. As you move from left to right in a period, the atomic number of the elements increases, which shows a gradual change in properties.

Blocks of the periodic table

The periodic table can also be divided into blocks based on the electron configuration of the elements. These blocks are s-block, p-block, d-block and f-block.

S block

The s-block includes the first two groups: alkali metals and alkaline earth metals. The outermost electrons of the elements in this block are in the s orbital.

P-block

The p-block includes groups 13 to 18. Here the outermost electron of the elements is in the p orbital. This block includes both metals and non-metals.

D-block

The d-block is also known as transition metals. The outermost electrons of these elements are in the d orbital. The d-block includes elements such as iron (Fe), copper (Cu), and gold (Au).

F block

The f-block is made up of the lanthanides and actinides. The outermost electrons of these elements are in the f orbital. They are often displayed separately at the bottom of the periodic table.

Metals, nonmetals and metalloids

Metals

Most of the elements in the periodic table are metals. Metals are good conductors of heat and electricity, are shiny in appearance, and are malleable and ductile. Examples include iron (Fe), gold (Au), and aluminum (Al).

Nonmetals

Nonmetals are found on the right side of the periodic table. They are poor conductors of heat and electricity and their physical properties are more variable. Examples include oxygen (O), carbon (C), and sulfur (S).

Metalloids

Metalloids have properties that are intermediate between metals and nonmetals. They lie between the metals and nonmetals in the periodic table. Examples include silicon (Si) and germanium (Ge).

Trends in the periodic table

The periodic table shows trends or patterns in the properties of elements. Some important trends are as follows:

Atomic radius

The atomic radius is the distance from the nucleus to the outer shell of the atom. As you move down the group, the atomic radius increases because more electron shells are added. As you move across the period, the atomic radius decreases due to the increased attraction between the electron and the nucleus.

Ionization energy

Ionization energy is the energy required to remove an electron from an atom. As you move across a period, ionization energy increases because the attraction between electrons and the nucleus is greater. As you move down a group, ionization energy decreases due to increased distance from the nucleus.

Electronegativity

Electronegativity is a measure of an atom's ability to attract and bond with electrons. As you move across a period, electronegativity increases due to the increased positive charge in the nucleus. As you move down a group, electronegativity decreases due to the increased distance between the nucleus and the outer electrons.

Jet

The reactivity of elements can change both up and down the periodic table. Metals become more reactive as we move down a group and less reactive across a period. Nonmetals show the opposite trend.

Applications of the periodic table

The periodic table is an essential tool in chemistry and many other sciences. Here are some of its applications:

  • Predicting chemical reactions: By understanding an element's position in the periodic table, scientists can predict how it will react with other elements.
  • Creating new materials: Knowledge of the properties of elements enables scientists to create new materials with desired properties, such as stronger metals or more efficient semiconductors.
  • Teaching and learning: The periodic table is a central tool in science education, providing a framework for understanding the elements and chemistry.

Conclusion

The periodic table is a powerful tool that organizes the elements and sheds light on their properties and the way they interact. It is more than a chart; it is a roadmap to understanding the world of atoms and molecules. With each element systematically placed in its place, the periodic table remains an essential resource not only for chemists, but for anyone interested in the physical sciences.


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Periodic table

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