Study Notes on Properties & Structure of Matter

Chemistry Module One: Properties & Structure of Matter

1. Learning Outcomes

• 1.3.1: Demonstrate, explain, and predict relationships in observable trends in physical and chemical properties of elements in periods and groups in the periodic table, covering:

  • State of matter at room temperature

  • Electronic configurations and atomic radii

  • First ionisation energy and electronegativity

  • Reactivity with water

2. Periodicity

• Definition: Periodicity indicates that properties of elements occur at regular intervals in the periodic table.

• Importance: Shows how elements share similar properties, enabling predictions about characteristics and behaviors based on position.

3. The Periodic Table of Elements

• Elements arranged by increasing atomic number

• Groups (columns) and periods (rows) provide insights into similarities and patterns.

• Notable aspects of elements presented include atomic numbers, masses, and state of matter at room temperature.

4. States of Matter

• Room Temperature States: Refers to the state (solid, liquid, gas) an element exhibits at 23 degrees Celsius.

  • Solid: High attraction between atoms.

  • Liquid: Moderate attraction.

  • Gas: Minimal attraction.

• Trends:

  • Movement across a period from left (solid) to right (gas).

  • Key Points: Most elements are solids at room temperature; exceptions are two liquids (mercury and bromine) and a few gases (like nitrogen and oxygen).

5. Core Charge

• Definition: Core charge measures the attractive force between valence electrons and the nucleus.

• Calculation:

  • Core Charge = Protons (Nucleus) - Total Electrons in Inner Shells.

  • Example:

  • Sodium (Na): 2, 8, 1

    • Core charge = 11 protons - 10 inner electrons = +1

  • Chlorine (Cl): 2, 8, 7

    • Core charge = 17 protons - 10 inner electrons = +7

• Trends:

  • Increases across a period.

  • Remains the same down a group.

6. Atomic Radius

• Definition: The distance from the center of an atom's nucleus to the outermost electron shell (valence shell).

• Trends:

  • Across a Period: Atomic radius decreases; core charge increases with constant electron shells, leading to greater attraction of valence electrons to nucleus.

  • Down a Group: Atomic radius increases; core charge remains, but additional electron shells reduce nuclear attraction to valence electrons.

7. Electron Configuration

• Definition: Arrangement of electrons in an atom’s electron shells and subshells.

• Significance: Influences chemical reactivity, ionization energy, and atomic radius.

• Patterns: Elements in the same group exhibit similar chemical and physical properties due to the same number of valence electrons.

8. Ionization Energy

• Definition: The energy required to remove an electron from a neutral gaseous atom.

• Trends:

  • Down a Group: Ionization energy decreases; more shells mean valence electrons are farther from the nucleus, and extra shells cause shielding.

  • Across a Period: Ionization energy increases; increased nuclear charge (more protons) means stronger attraction to electrons, while shielding remains constant.

9. Electronegativity

• Definition: A measure of an atom's ability to attract and hold bonding electrons.

• Trends:

  • Down a Group: Electronegativity decreases; valence electrons are shielded by more electron shells.

  • Across a Period: Electronegativity increases; as atoms become smaller, they attract electrons more strongly.

10. Reactivity with Water

• General Trend: Reactivity generally increases down a group.

  • Linked to trends in ionization energy; lower ionization energy allows easier loss of valence electrons in reactions.

  • Notable reactive elements with water are primarily found on the left side of the periodic table.

11. Properties of Matter

• Physical Properties: Characteristics observed without changing the identity of the material (e.g., magnetism).

• Chemical Properties: Reactions of materials with other substances, typically observed during chemical reactions.

  • Examples of Classification:

  • Physical Properties: Intensive (color, melting point, boiling point, density) vs. Extensive (mass, volume, length).

12. Group Classification and Properties

• Metals: Usually hard, shiny, dense, good conductors of heat/electricity, form cations and basic oxides.

• Non-metals: Often dull, low density, poor conductors, form anions and acidic oxides.

• Metalloids: Intermediate properties.

13. Summary of Chemical Properties by Group

• Metals: Tend to donate electrons, displace hydrogen from dilute acids.

  • Example:

  • $( ext{Zn} + 2 ext{HCl} <br>ightarrow ext{ZnCl}<em>2 + ext{H}</em>2)$

• Non-metals: Tend to receive electrons, do not displace hydrogen from dilute acids, form covalent compounds.

  • Reaction examples:

  • $( ext{C} + ext{O}<em>2 ightarrow ext{CO}</em>2)$

  • $( ext{C} + 2 ext{H}<em>2 ightarrow ext{CH}</em>4)$

• Water Chemistry:

  • Metal oxides produce bases (e.g., $( ext{CaO}, ext{ZnO}, ext{BaO})$) and nonmetal oxides produce acids (e.g., $( ext{CO}<em>2, ext{SO}</em>2, ext{NO}_2)$).

  • Acids have pH < 7; bases have pH > 7.

14. Metals and Nonmetals Comparison

15. Conclusion

• Understanding the periodic table is critical. More than a reference, it serves as a tool/map for predicting element properties and behaviors based on their positions.