Laws of Chemical Combination:
1. Law of Conservation of Mass - Matter cannot be created or destroyed.
2. Law of Constant Proportions - A given compound always contains the same elements in the same proportion by mass.
Definition: In a chemical reaction, the total mass of the reactants is equal to the total mass of the products. Mass is neither created nor destroyed.
Example: Burning of wood (mass of reactants = mass of products).
Definition: A chemical compound always contains the same elements combined together in a fixed proportion by mass.
Example: Water (H₂O) always has a 1:8 mass ratio of hydrogen to oxygen.
Definition: Proposed by John Dalton, this theory states:
1. Matter is made of tiny particles called atoms.
2. Atoms are indivisible and cannot be created or destroyed.
3. Atoms of the same element are identical.
4. Compounds form when atoms combine in fixed ratios.
Limitation: Atoms can be divided into subatomic particles (protons, neutrons, electrons).
Definition: The smallest particle of an element that retains the chemical properties of that element.
Structure: Nucleus (protons + neutrons) + Electrons.
Example: Hydrogen atom (1 proton, 1 electron).
Table of Atomic Names and Their Latin/Greek Names
| Element | Symbol | Latin/Greek Name |
|---|---|---|
| Gold | Au | Aurum |
| Silver | Ag | Argentum |
| Lead | Pb | Plumbum |
| Copper | Cu | Cuprum |
| Iron | Fe | Ferrum |
Definition: The smallest particle of an element or compound that can exist independently and retain the properties of that substance.
Types:
1. Molecules of Elements: O₂, H₂.
2. Molecules of Compounds: H₂O, CO₂.
Definition: A representation of a compound using symbols of the elements and subscripts to show the number of atoms of each element.
Examples:
1. Water: H₂O.
2. Carbon Dioxide: CO₂.
Definition: Charged particles formed by the gain or loss of electrons.
Types:
1. Cations: Positively charged ions (e.g., Na⁺).
2. Anions: Negatively charged ions (e.g., Cl⁻).
Table of Names, Symbols, and Valency of Common Ions:
| Ion Name | Symbol | Valency |
|---|---|---|
| Sodium | Na⁺ | 1 |
| Magnesium | Mg²⁺ | 2 |
| Calcium | Ca²⁺ | 2 |
| Aluminum | Al³⁺ | 3 |
| Chloride | Cl⁻ | 1 |
| Oxide | O²⁻ | 2 |
| Sulfate | SO₄²⁻ | 2 |
| Nitrate | NO₃⁻ | 1 |
Definition: A group of atoms carrying a net charge due to the gain or loss of electrons.
Examples:
1. Sulfate: SO₄²⁻.
2. Ammonium: NH₄⁺.
Definition: The combining capacity of an element, determined by the number of electrons it can lose, gain, or share.
Example: Oxygen has a valency of 2 (forms O₂).
Definition: Compounds formed by the transfer of electrons from one atom to another, resulting in the formation of ions.
Example: Sodium Chloride (NaCl).
Definition: The mass of an atom compared to 1/12th the mass of a carbon-12 atom.
Example: Relative atomic mass of oxygen = 16 u.
Atomic Mass Table:
| Element | Symbol | Atomic Mass (amu) |
|---|---|---|
| Hydrogen | H | 1.008 |
| Oxygen | O | 16.00 |
| Carbon | C | 12.01 |
| Nitrogen | N | 14.01 |
| Calcium | Ca | 40.08 |
Definition: The number of atoms in exactly 12 g of carbon-12, equal to 6.022 × 1023.
Significance: Defines the mole concept.
Definition: The amount of substance that contains 6.022 × 1023 particles (atoms, molecules, ions, etc.).
Example: 1 mole of carbon = 12 g.
Definition: The mass of one mole of a substance, expressed in grams per mole (g/mol).
Example: Molar mass of water (H₂O) = 18 g/mol.
Please don not use wrong word