Unit 1: Chemical Fundamentals – Structure, Properties and Reactions

TL;DR

Chemistry is all about understanding how atoms are built, how they interact, and how these interactions lead to the diverse properties and reactions of substances. Knowing the fundamental particles and their arrangements helps you predict chemical behaviour. Mastering this topic means you can explain why different materials act the way they do at a microscopic level.

1. The Mental Model

Imagine atoms as tiny LEGO bricks. The type of brick, how many studs it has, and how it connects to other bricks determines what you can build and how strong or flexible the final structure is. Chemistry is like the instruction manual for these LEGO bricks: it tells you how they’re made, how they snap together, and what kind of models they can form.

2. The Core Material

In Unit 1, you're diving into the basics that underpin all of chemistry. It starts with the fundamental building blocks and moves into how these blocks interact.

Atomic Structure

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Every atom has a nucleus (containing protons and neutrons) and electrons orbiting this nucleus.
* Protons: Positively charged (+1), define the element (atomic number, Z).
* Neutrons: No charge (0), contribute to mass (isotopes vary by neutron count).
* Electrons: Negatively charged (-1), located in electron shells (or energy levels) around the nucleus. These are crucial for chemical bonding.

The mass number (A) is the total number of protons and neutrons. Isotopes are atoms of the same element (same number of protons) but with different numbers of neutrons, hence different mass numbers. For example, Carbon-12 and Carbon-14 are isotopes of carbon.

Electron Configuration

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Electrons fill energy levels in a specific order. The outermost electrons, called valence electrons, are the ones involved in chemical reactions.
* Shell 1: max 2 electrons
* Shell 2: max 8 electrons
* Shell 3: max 18 electrons (though typically fills 8 before moving to shell 4 for main group elements)

You write electron configurations by listing the number of electrons in each shell, e.g., Sodium (Na) has 11 electrons: 2, 8, 1. The '1' is its valence electron.

Periodic Table Trends

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The periodic table isn't just a chart; it's a powerful tool! Elements are arranged by increasing atomic number.
* Periods (rows): Indicate the number of electron shells.
* Groups (columns): Indicate the number of valence electrons (for main group elements) and thus similar chemical properties.
* Metallic character: Generally decreases across a period and increases down a group.
* Atomic radius: Generally decreases across a period (more protons pull electrons closer) and increases down a group (more electron shells).
* Ionisation energy: Energy required to remove an electron. Generally increases across a period and decreases down a group.
* Electronegativity: Ability of an atom to attract electrons in a bond. Generally increases across a period and decreases down a group.

graph TD
    A["Element Identity (Protons)"] --> B["Atomic Structure (P, N, E)"]
    B --> C["Electron Configuration (Shells, Valence E-)"]
    C --> D{"Periodic Table Position (Group, Period)"}
    D --> E["Properties & Trends (Size, IE, EN, Metallic)"]
    E --> F["Bonding Type (Ionic, Covalent, Metallic)"]
    F --> G["Substance Properties (Melting Point, Conductivity, State)"]

Chemical Bonding

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Atoms bond to achieve a stable electron configuration, usually a full outer shell (like noble gases).

• Ionic Bonding: Occurs between a metal (loses electrons to form positive ions, cations) and a non-metal (gains electrons to form negative ions, anions). The electrostatic attraction between opposite charges forms a strong ionic lattice.
  • Example: NaCl (Na loses 1e-, Cl gains 1e-)
• Covalent Bonding: Occurs between two or more non-metals. Atoms share electrons to achieve stable outer shells. Can be single, double, or triple bonds.
  • Example: H₂O (Oxygen shares with two Hydrogens)
• Metallic Bonding: Occurs in metals. A "sea" of delocalised electrons is attracted to a lattice of positive metal ions. This explains why metals are good conductors and malleable.

Chemical Reactions

Chemical reactions involve the breaking and forming of chemical bonds, resulting in new substances.
* Reactants $\rightarrow$ Products
* Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction. The total mass of reactants equals the total mass of products.
* Balancing Equations: Ensures conservation of mass by having the same number of each type of atom on both sides of the equation.

3. Worked Example

Let's predict the bonding and some properties for a compound formed between Magnesium (Mg) and Oxygen (O).

1. Locate on Periodic Table:
   • Mg is in Group 2, Period 3. It's a metal.
   • O is in Group 16, Period 2. It's a non-metal.
2. Determine Electron Configuration and Valence Electrons:
   • Mg (Atomic number 12): 2, 8, 2. Has 2 valence electrons. Metals tend to lose valence electrons to form positive ions. Mg will lose 2 electrons to become Mg²⁺.
   • O (Atomic number 8): 2, 6. Has 6 valence electrons. Non-metals tend to gain electrons to fill their outer shell. O will gain 2 electrons to become O²⁻.
3. Determine Bonding Type: Since it's a metal reacting with a non-metal, they will form an ionic bond. Mg transfers 2 electrons to O.
4. Formula: To balance the charges (+2 and -2), you need one Mg²⁺ and one O²⁻. The formula is MgO.
5. Predicted Properties: As an ionic compound, MgO will likely:
   • Have a high melting and boiling point (strong electrostatic forces).
   • Be a solid at room temperature.
   • Conduct electricity when molten or dissolved (ions are free to move), but not as a solid (ions fixed in lattice).
   • Be brittle.

4. Key Takeaways

• The number of protons defines an element; the number of electrons (especially valence) determines its chemical reactivity.
• The periodic table organises elements based on their atomic structure, revealing predictable trends in properties like size and electronegativity.
• Atoms bond to achieve stability, forming ionic (metal-nonmetal), covalent (nonmetal-nonmetal), or metallic bonds.
• The type of bonding directly dictates a substance's macroscopic properties, such as melting point, conductivity, and state.
• Chemical reactions involve rearranging atoms and breaking/forming bonds, always conserving mass.

Common Mistakes to Avoid

• Confusing atomic number with mass number, or protons with neutrons.
• Incorrectly assigning valence electrons or drawing electron configurations.
• Not using the periodic table to predict trends or bonding types.
• Forgetting to balance chemical equations; this is crucial for all stoichiometry.

5. Now Try It

Choose two elements: one from Group 1 (alkali metals) and one from Group 17 (halogens). For your chosen pair:
1. Write their electron configurations.
2. Predict the type of chemical bond they would form.
3. Write the chemical formula for the compound they would create.
4. List three physical properties you'd expect this compound to have.

Success looks like: Correct electron configurations, accurate bonding type (e.g., ionic), the right chemical formula (charges balanced), and three appropriate physical properties (e.g., high melting point, conducts electricity when molten, brittle).