chem test
Station 1: Bohr-Rutherford Diagrams, Standard Atomic Notation, and Periodic Trends
Bohr-Rutherford Diagrams
Purpose: These diagrams show the arrangement of electrons around the nucleus of an atom.
Structure: You’ll see a central nucleus with circles (representing electron shells) around it.
Electrons in Shells: Know the maximum number of electrons per shell:
1st shell: 2 electrons
2nd shell: 8 electrons
3rd shell: 8 electrons (for simplicity at this level)
Practice: Draw Bohr-Rutherford diagrams for different elements. Start with elements like Carbon, Oxygen, or Neon.
Standard Atomic Notation
Format: Looks like this: \text{^{A}_{Z}X}, where:
X is the symbol of the element (e.g., Na for sodium).
A is the mass number (protons + neutrons).
Z is the atomic number (number of protons).
Example: For Sodium (Na), the standard atomic notation could look like \text{^{23}_{11}Na}.
Periodic Trends
Atomic Radius: Size of atoms increases as you go down a group (column) and decreases as you move across a period (row).
Ionization Energy: The energy required to remove an electron from an atom. It increases across a period (left to right) and decreases down a group.
Electronegativity: The tendency of an atom to attract electrons. It increases across a period and decreases down a group.
Practice: Compare elements based on their position in the periodic table and predict which will have higher/lower atomic radius, ionization energy, or electronegativity.
Station 2: Classification of Matter
Pure Substances vs. Mixtures
Pure Substances: Only one type of particle. Includes elements and compounds.
Elements: Single type of atom (e.g., O₂ for oxygen gas).
Compounds: Two or more types of atoms chemically bonded (e.g., H₂O for water).
Mixtures: Two or more types of particles mixed together but not chemically bonded.
Homogeneous Mixtures (Solutions): Uniform composition, like salt water.
Heterogeneous Mixtures: Non-uniform composition, like salad or sand in water.
Physical vs. Chemical Properties
Physical Properties: Characteristics observed without changing the substance (e.g., color, melting point, density).
Chemical Properties: Describe a substance’s ability to undergo chemical changes (e.g., flammability, reactivity).
Physical vs. Chemical Changes
Physical Changes: Change in form or state, no new substance formed (e.g., melting ice, dissolving sugar in water).
Chemical Changes: New substances are formed with different properties (e.g., rusting, burning).
Station 3: Gas Tests
Oxygen Test
Procedure: Place a glowing splint into a test tube containing the gas.
Result: If the gas is oxygen, the glowing splint will reignite.
Hydrogen Test
Procedure: Place a burning splint near the opening of a test tube with the gas.
Result: If hydrogen is present, you’ll hear a “pop” sound.
Carbon Dioxide Test
Procedure: Bubble the gas through limewater (a solution of calcium hydroxide).
Result: If carbon dioxide is present, the limewater turns cloudy.
Station 4: Building Molecules
Bond Types and Valency
Covalent Bonds: Atoms share electrons. Typically between non-metals.
Ionic Bonds: Atoms transfer electrons, forming positive and negative ions. Usually between a metal and a non-metal.
Valency: The combining power of an atom (how many bonds it can form). E.g., Oxygen has a valency of 2, meaning it forms two bonds.
Molecular Geometry
Linear: 2 atoms or double bonds, e.g., CO₂.
Bent: Typically seen in molecules with lone pairs, e.g., H₂O.
Tetrahedral: 4 bonds, no lone pairs on the central atom, e.g., CH₄.
Practice Building Models
Use a model kit (if available) or sketch diagrams to understand how atoms connect based on their valency.
Example: Try building molecules like water (H₂O), carbon dioxide (CO₂), and methane (CH₄) to see the differences in bond types and geometry.