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 → Introduction to Chemistry ↓


Units of measurement in chemistry


Chemistry is the study of matter and the changes that occur in it. To accurately describe these changes and the properties of matter, we need a common language. This language is based on units of measurement, which allow us to measure the characteristics of matter. In this lesson, we will explore the different units of measurement used in chemistry and their importance.

Importance of units of measurement

Units of measurement are essential because they provide a standardized way to express quantities. Without units, scientists would have trouble repeating experiments and comparing their results with others. Imagine you are trying to share a recipe that simply says "add some flour." Without a specific measurement like "cups" or "grams," the recipe is open to interpretation and inconsistency.

International System of Units (SI)

The International System of Units, abbreviated as SI from the French "Système International de Unites", is the most widely used system of measurement in the world. It provides a uniform set of units to simplify communication and calculations across different fields and scientific disciplines. In chemistry, we use many specific units from the SI system to measure various properties of substances.

1 cm

This visual example shows a 1 cm x 1 cm square designated by the measurement of 1 centimeter (cm), which is an example of measuring length.

Base units in chemistry

The SI system is made up of seven base units, but in chemistry we mainly use the following:

  • Mole (mol): Used to measure the amount of a substance.
  • Meter (m): Used to measure length or distance.
  • Kilogram (kg): Used to measure mass.
  • Second (s): Used to measure time.
  • Kelvin (K): Used to measure temperature.

Derived units in chemistry

Derived units are combinations of base units. They are used to measure properties that require more than one dimension of measurement. For example:

  • Volume: Measured in cubic metres (m3) or more commonly in litres (L). A litre is the volume of a cube with a side of 10 cm, which is 1 dm3.
  • Density: Measured in kilograms per cubic meter (kg/m3) or grams per cubic centimeter (g/cm3).
  • Pressure: Measured in pascals (Pa), which is equal to one newton per square meter. In chemistry, pressure is often measured in atmospheres (atm) or millimeters of mercury (mmHg).
Liter

This spherical visualization serves as a representation of measuring volume, and highlights the concept of the liter, a common unit used in chemistry.

Typical units of measurement in chemistry

The Mole

A mole is a unit used to describe the amount of a substance. A mole is defined as 6.022 × 1023 particles (atoms, molecules, ions, or electrons), also known as Avogadro's number. For example, one mole of water (H2O) contains 6.022 × 1023 water molecules.

Example:
1 mole of carbon atoms = 6.022 × 10²³ atoms

Mass

Mass is a measure of the amount of matter in something. In chemistry, it is often measured in grams (g) rather than kilograms because most chemical samples are small and grams are a more convenient unit.

Example:
The mass of a single carbon atom is ≈ 12 amu (atomic mass units), which is ≈ 1.99 × 10⁻²³ grams.

Volume

Volume measures the amount of space occupied by an object. In chemistry, it is often measured in liters or milliliters (ml), where 1 liter = 1000 ml. This unit is especially useful when working with liquids or gases.

Density

Density is the value obtained by dividing the mass of a substance by its volume, expressed by the following formula:

Density = mass / volume
Example:
The density of water is about 1 g/cm³.

Pressure

Pressure is a measure of the force exerted over a certain area. It is important in understanding how gases behave under different conditions. Common units are atmospheres (atm), pascals (Pa), and millimeters of mercury (mmHg).

Example:
Standard atmospheric pressure at sea level is 760 mmHg or 1 atm.

Temperature

While the Celsius scale is often used in everyday life, the Kelvin scale is used in scientific calculations because it starts at absolute zero, which is the theoretical lowest temperature possible.

0 K (Kelvin) = -273.15 °C (Celsius)

Conversion between units

In chemistry, converting between units is common because experiments and calculations may require different units. Using conversion factors allows chemists to switch from one unit to another.

Let us consider some examples:

Converting grams to moles (and vice versa)

To convert grams to moles, divide the mass of the substance (in grams) by its molar mass (in grams per mole). The formula is:

Mole = mass (g) / molar mass (g/mol)

Converting liters to milliliters

Since 1 liter equals 1000 milliliters, you multiply by 1000 to convert liters to milliliters, and divide by 1000 to convert milliliters to liters.

Converting Celsius to Kelvin

To convert degrees Celsius to Kelvin, add 273.15 to the Celsius temperature.

Temperature (K) = Temperature (°C) + 273.15

Converting pressure units

To convert pressure units, you need to know the conversion factors. For example, 1 atm = 101325 Pa = 760 mmHg.

Importance and accuracy in measurement

When making measurements in chemistry, it is important to consider both accuracy and precision. Accuracy refers to how close a measurement is to the true value, while precision refers to how repeatable a measurement is.

Significant figures are used to express the precision of a measurement. They include all known digits and one approximate digit. For example, in a measurement of 12.345, all five digits are significant, indicating high precision.

Conclusion

Understanding units of measurement and using them correctly is fundamental in chemistry. It ensures clear communication and allows scientists to share and compare their findings with confidence. Whether measuring the mass of a reactant, the volume of a solution, or the pressure of a gas, selecting the appropriate units and converting between them accurately is critical for scientific accuracy and reliability.

As you continue to explore chemistry, practicing these concepts will strengthen your understanding and allow you to more effectively engage in chemical analysis and experimentation.


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