Atoms and Molecules Short Notes | Class 9
Introduction
Philosophical Foundations: Ancient Indian philosopher Maharishi Kanad and Greek philosophers Democritus and Leucippus hypothesized the divisibility of matter around 500 BC, with ideas of indivisible particles—Kanad's Parmanu and Democritus's atoms.
Combined Matter: Pakudha Katyayama, another Indian philosopher, expounded the concept that these indivisible particles occur in combined forms, leading to various forms of matter.
Lack of Experimental Validation: These philosophical concepts were not tested extensively until the 18th century, with no experimental evidence to validate the concepts of indivisible particles.
Differentiation of Elements and Compounds: By the late 18th century, scientists were aware of the differences between elements and compounds, and this generated interest in their combinations and resulting reactions.
Contributions of Antoine Lavoisier: Antoine L. Lavoisier made significant contributions to chemical science by establishing important laws of chemical combination, which formed the basis of modern chemistry.
Foundation for Modern Chemistry: Lavoisier's research established basic concepts that directed subsequent studies of chemical reactions and the behavior of matter, moving from philosophical speculation to scientific investigation.
Law of Chemical Combination
Chemical combination is a basic concept in chemistry that explains the way substances interact and change into new materials by breaking and forming chemical bonds.
→ These are five basic law
1. Law of Conservation of Mass:
Formulated by Antoine Lavoisier, the law states that the combined mass of the reactants (substances which are undergoing reaction) is equal to the combined mass of the products (substances which are formed due to the reaction) in a chemical reaction.
This indicates that matter cannot be created or destroyed during a chemical reaction; it is merely rearranged.
2. Law of Law of Constant or (Composition Definite Proportions )
This law, developed by Joseph Proust, states that a chemical compound always has the same proportion of elements by mass, irrespective of origin or how it is formed.
For instance, water (H₂O) will always be 2 hydrogen atoms and 1 oxygen atom, so the mass ratio of hydrogen to oxygen in water is always the same.
Dalton's Atomic Theory
All matter, according to Dalton's atomic theory, is made up of very small particles known as atoms, whether an element, a compound or a mixture.
Six Postulates of Dalton's atomic theory:
- Matter is made up of minute particles known as atoms.
- Atoms are indivisible and cannot be broken or formed in chemical reactions.
- Atoms of the same element are identical in mass and properties.
- Atoms of different elements vary in mass and properties.
- Atoms unite in small whole number ratios to form compounds.
- The number and types of atoms in a compound are fixed.
Atoms
An atom is the smallest unit of matter that still has the properties of an element. Atoms are the building blocks of all matter and are made up of three types of major subatomic particles:
- Protons: Positively charged particles which occur in the nucleus of the atom.
- Neutrons: Neutral particles (uncharged) which also occur in the nucleus.
- Electrons: Negatively charged particles which occur in orbit around the nucleus in areas called electron shells or energy levels.
→Atomic Mass
Dalton's Atomic Theory and Atomic Mass
- Atomic Mass Concept: Dalton's atomic theory brought in the concept that every element has a characteristic atomic mass, which predicts chemical behavior.
- Law of Constant Proportions: The theory was able to explain that substances combine in fixed ratios, which led to the measurement of atomic masses.
- Relative Atomic Mass: As it is difficult to find the mass of individual atoms, relative atomic masses were determined by applying chemical combination laws and stoichiometry.
- Illustration of Carbon Monoxide: For the production of carbon monoxide (CO), 3 g of carbon reacts with 4 g of oxygen, indicating that carbon reacts with 4/3 times its weight of oxygen.
- Creation of Atomic Mass Units:
- The atomic mass unit (amu) is now abbreviated as "u" (unified mass).
- Originally, 1/16 of naturally occurring oxygen mass was used as a reference but later fixed.
7. Measurement of Relative Atomic Mass: Relative atomic masses of all elements are measured in reference to carbon-12 in order to achieve uniformity of measurement.
8. Fruit Seller Analogy: The analogy is used to show the creation of standard mass (using watermelon) to measure other masses with reference to a fixed reference mass.
Molecules
Molecules are chemically combined two or more atoms. The atoms can be of the same element or of different elements. Molecules are the smallest entity of a chemical compound that retains the chemical characteristics of the compound.
Molecules of Elements
Definition: A molecule of an element is created when two or more atoms of the same element are joined.
Types of Elemental Molecules:
- Diatomic Molecules: These molecules consist of two atoms of the same element. Examples are:
Hydrogen (H₂)
Nitrogen (N₂)
- Polyatomic Molecules: These molecules consist of three or more atoms of the same element. An example is:
Ozone (O₃), consisting of three oxygen atoms.
Atomicity
The number of atoms in one molecule of an element is referred to as its atomicity.
Chemical Formulae
It is symbolic representation of composition of compound
<< Features of chemical formulae
Chemical formulae play a significant role in specifying the composition and structure of chemical compounds, as they give a concise representation of the elements and their ratios involved.
<<Writing rules for chemical formulae
- Identify Elements: Write the symbols of the elements in the compound.
- Determine Ratios: Use valency to determine the correct ratio of atoms for ionic compounds.
- Use Subscripts: Write subscripts to represent the number of atoms of each element (e.g., H₂O).
- Balance Charges: In ionic compounds, balance the positive and negative charges to determine the simplest ratio.
- Correct Symbol Use: Use correct chemical symbols and enclose the complex ions in parentheses (e.g., Ca(NO₃)₂).
Molecular Mass
Molecular mass is a significant concept in chemistry, as it aids in the understanding of mass relationships in chemical reactions and calculations.
Formula Unit Mass
It is sum of atomic mass of ions and atoms in formula of a compound
Example: In NaCl,
Na = 23 a.m.u.
Cl = 35.5 a.m.u.
So, Formula unit mass = 1×23 + 1×35.5 = 58.5 u
Ions
Definition: Ions are charged particles formed due to gain or loss of electrons by an atom or a group of atoms. Ions are positively charged (cations) or negatively charged (anions).
Types of Ions : -
Cations:
- Description: Positively charged ions formed by loss of electrons.
- Examples:
- Sodium ion (Na⁺)
- Calcium ion (Ca²⁺)
Anions:
- Description: Negatively charged ions formed by gain of electrons.
- Examples:
- Chloride ion (Cl⁻)
- Sulfate ion (SO₄²⁻)
Mole Concept
Definition: Mole is a unit for expressing the amount of substance. It connects the mass of a substance to the number of atoms, molecules, or ions it contains.
- A mole (symbol: "mol") is the amount of substance which contains the same number of particles (atoms, molecules, etc.) as is contained in 12 grams of carbon-12 (12C).
- This quantity is called Avogadro's Number, which is 6.022×10236.022×10^23 particles per mole.
- Molar Mass:
- Molar mass is the mass of a mole of substance, in g/mol. It is numerically equal to the atomic or molecular mass but in g/mol.
- For instance, the molar mass of water (H₂O) is 18 g/mol18g/mol (1 for H and 16 for O).
Calculating Moles:
- To calculate the number of moles from mass, you can use the equation:
- Number of moles=Mass of substance (g)Molar mass (g/mol)
Molar mass (g/mol)
Mass of substance (g)
Finding Mass from Moles:
Mass=Number of moles×Molar mass (g/mol)
Mass=Number of moles×Molar mass (g/mol)