Structure of the Atom, Chemical Reactions, and Acids and Bases
Classified in Chemistry
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Structure of the Atom
Protons, Electrons, and Neutrons
Protons: Equal to the atomic number
Electrons: Equal to the atomic number
Neutrons: Equal to the atomic mass minus the atomic number
Bohr-Rutherford Diagram
Electron shells: 2, 8, 8, 18, 18, 32
Counting Atoms
Example 1: Na2CO3
- Sodium (Na): 2
- Carbon (C): 1
- Oxygen (O): 3
Example 2: 4Al2(CO3)3
- Aluminum (Al): 4 x 2 = 8
- Carbon (C): 3 x 4 = 12
- Oxygen (O): 9 x 4 = 36
- Total: 56
Ions and Ionic Compounds
Ions are atoms that have either lost or gained electrons. While atoms are neutral, ions are charged particles.
Ionic Compounds are a combination of a cation (positive ion) and an anion (negative ion), typically formed between a metal and a nonmetal.
Example: Calcium Chloride (CaCl2)
Calcium (Ca+2) has a charge of +2, and Chloride (Cl-1) has a charge of -1.
To form a neutral compound, we need two chloride ions for every calcium ion: (+2) + 2(-1) = 0
The formula is CaCl2, which is derived using the criss-cross method.
Note: The metal ion (cation) is always written first in the formula.
Molecular Compounds
Molecular Compounds are formed when atoms of two or more different elements share electrons. They are usually formed between two or more nonmetals and are also called covalent compounds.
How Atoms Bond
- Ionic Bond: Transfer of electrons between atoms of opposite charges.
- Covalent Bond: Sharing of electrons to complete the valence shell.
Bonding Combinations
- Ionic: Metal to nonmetal
- Molecular: Nonmetal to nonmetal
Compound Formation
- Ionic: Solids
- Molecular: Smaller compounds with 2-3 atoms linked (e.g., F2)
Example: Fluorine (F2)
Fluorine has 7 valence electrons and needs one more electron to complete its outer shell. It achieves this by sharing an electron with another fluorine atom, forming a covalent bond.
Naming Molecular Compounds
Compound | Formula | Prefix | # Atoms | Example | Formula |
Hydrogen gas Oxygen gas Nitrogen gas Fluorine gas Chlorine gas Bromine (liquid) Iodine (solid) | H2 O2 N2 F2 Cl2 Br2 I2 | Mono Di Tri Tetra Penta Hexa Hepta Octa Nona Deca | 1 2 3 4 5 6 7 8 9 10 | CO = Carbon Monoxide CO2 = Carbon Dioxide SO3 = Sulfur Trioxide | P2Br4 = Diphosphorus Tetrabromide Pb3P2 = Lead(II) Phosphide (because ionic) B2H4 = Diboron Tetrahydride Note: If the compound is ionic, use the -ide ending for the anion. If it's covalent, use prefixes for both elements. |
Chemical Equations
Example: Formation of Water
Hydrogen + Oxygen → Water
Unbalanced: H2 + O2 → 2HO (Incorrect)
Balanced: 2H2 + O2 → 2H2O
Example: Double Displacement Reaction
2FeCl3 + 3Na2S → 6NaCl + Fe2S3
Law of Conservation of Mass
The total mass of the products is always the same as the total mass of the reactants.
Types of Reactions
- Synthesis: Combination of smaller atoms into larger ones (A + B → AB)
- Decomposition: Opposite of synthesis (AB → A + B)
- Single Displacement: One element replaces or displaces a second element (A + XY → AY + X)
- Double Displacement: Both elements in different compounds displace each other (AB + XY → AY + XB)
Properties of Acids and Bases
Acids
- Sour taste
- Dissolve in water, producing hydrogen ions (H+)
- Conduct electricity (electrolytes)
- Corrosive to certain materials
- pH less than 7 (lower pH = stronger acid)
Bases
- Bitter taste
- Dissolve in water, producing hydroxide ions (OH-)
- Conduct electricity (electrolytes)
- Caustic to living tissues
- pH greater than 7 (higher pH = stronger base)