Atomic Structure, Bonding, and Chemical Principles

Posted on Dec 16, 2024 in Chemistry

  • Atomic Structure and Properties

    • Know the basic structure of an atom, in terms of sub-atomic particles (protons, neutrons and electrons), and discuss their basic properties and how they are arranged in the atom.
      • Bohr theory of atomic structure.
      • Four quantum numbers. (principle, angular momentum, magnetic and spin)
    • Interpret the periodic table, and identify, describe and explain particular trends in the periodic table (ionisation energy, effective nuclear charge, atomic and ionic radii, electronegativity, electron affinity).
    1. Atomic Structure and Properties

    Basic Structure of an Atom

    An atom consists of three main subatomic particles:

    • Protons: Positively charged particles found in the nucleus. They determine the atomic number and thus the identity of the element (e.g., hydrogen has 1 proton, oxygen has 8).
    • Neutrons: Neutrally charged particles also found in the nucleus. Neutrons contribute to the atomic mass but not the charge.
    • Electrons: Negatively charged particles that orbit the nucleus in energy levels (or shells). In a neutral atom, the number of electrons equals the number of protons.

    Bohr Theory of Atomic Structure

    Bohr proposed that electrons move in fixed orbits around the nucleus (like planets around the sun) and that they can absorb or emit energy by jumping between these orbits. Electrons in higher orbits have more energy. This model helps explain atomic emission spectra, where elements emit specific wavelengths of light when electrons drop to lower energy levels.

    Quantum Numbers

    • Principal Quantum Number (n): Indicates the energy level or shell (1, 2, 3, etc.). Higher numbers correspond to higher energy and further distance from the nucleus.
    • Angular Momentum Quantum Number (l): Defines the shape of the orbital (s, p, d, f). For example, s orbitals are spherical, p orbitals are dumbbell-shaped.
    • Magnetic Quantum Number (mₗ): Specifies the orientation of the orbital in space (e.g., px, py, pz for p orbitals).
    • Spin Quantum Number (mₛ): Refers to the spin direction of an electron (either +1/2 or -1/2).

    Periodic Table Trends

    • Ionization Energy: Energy required to remove an electron from an atom. Increases across a period (due to increasing nuclear charge) and decreases down a group (due to increased distance between nucleus and outer electrons).
    • Effective Nuclear Charge (Zₑₗₑₐ): The net positive charge experienced by an electron in a multi-electron atom. Increases across a period, as more protons are added but shielding from inner electrons remains roughly constant.
    • Atomic and Ionic Radii: Atomic radius decreases across a period (more protons pull electrons closer) and increases down a group (more electron shells).
    • Electronegativity: Tendency of an atom to attract electrons in a chemical bond. Increases across a period and decreases down a group.
    • Electron Affinity: The energy change when an electron is added to a neutral atom. More negative electron affinity means the atom more readily accepts electrons.

    Isotopes

    Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons, resulting in different mass numbers. For example:

    • Carbon-12 (¹²C) and Carbon-14 (¹⁴C) are isotopes of carbon, both with 6 protons but 6 and 8 neutrons, respectively.

    Properties of Isotopes:

    • Chemically identical: Isotopes of an element behave the same chemically.
    • Physical differences: Due to different masses, isotopes may behave differently in physical processes (e.g., diffusion).
    • Radioactive isotopes: Some isotopes are unstable and decay, releasing radiation. This property is useful in carbon dating (e.g., Carbon-14 dating).

    Mass Spectrometry

    Mass spectrometry is a technique used to determine the mass and abundance of isotopes in a sample. It works by ionizing atoms or molecules and measuring their mass-to-charge ratio:

    1. Ionization: Atoms are ionized to form charged particles (ions).
    2. Acceleration: Ions are accelerated through an electric field.
    3. Deflection: Ions are deflected in a magnetic field based on their mass-to-charge ratio.
    4. Detection: The ions are detected, and the resulting mass spectrum displays the isotopic abundances.

    This technique allows for the determination of relative atomic mass (RAM) by analyzing the relative abundances of isotopes in a sample.

    Atomic Emission Spectra

    When atoms are heated or electrically excited, electrons move to higher energy levels. As they return to lower levels, they release light in the form of photons. This light produces an atomic emission spectrum that is unique to each element:

    • The light emitted corresponds to the difference in energy levels of electrons.
    • Each element has a unique emission spectrum due to its specific energy levels.

    The emission spectrum provides evidence for quantized energy levels in atoms. For example:

    • Sodium produces a yellow light (589 nm) in streetlights.
    • Neon emits a red light in neon signs.

    This principle is used in spectroscopy to identify elements in samples and to study stellar compositions in astronomy.

  • Chemical Formulae, Equations and Calculations

    • Understand molecular and ionic formulae, physical states of matter, chemical equations, molar mass and the mole.
    • Construct balanced chemical equations, and use them to perform calculations based on mass to moles, gas volumes to moles, solutions and concentrations, reaction yields and limiting reagents.
    1. Chemical Formulae, Equations, and Calculations

    Molar Mass and Mole Concept

    1. The Mole

      • A “mole” is a unit used in chemistry to count particles (atoms, molecules, etc.), similar to how “dozen” counts 12 things.
      • One mole contains 6.022 x 10^23 particles (this number is called Avogadro’s number).

    2. Molar Mass

      • Molar mass is the mass of one mole of a substance, measured in grams per mole (g/mol).
      • To find the molar mass, add up the atomic masses of each element in the compound, based on the periodic table.

      Example

      For water (H₂O), molar mass = 2 times the mass of hydrogen + 1 times the mass of oxygen:

      Molar mass of H₂O = (2 x 1.01) + (1 x 16.00) = 18.02 g/mol

    3. Converting Mass to Moles

      • Use this formula: Moles = Mass (g) / Molar Mass (g/mol)

      Practice Question

      How many moles are in 20 grams of NaCl (sodium chloride)?

      Solution

      1. Find the molar mass of NaCl:
        • Na: 22.99 g/mol
        • Cl: 35.45 g/mol
        • Total molar mass of NaCl = 22.99 + 35.45 = 58.44 g/mol
      2. Calculate moles: Moles of NaCl = 20 g / 58.44 g/mol ≈ 0.342 mol

    Balancing Chemical Equations

    1. Why Balance Equations?

      • Chemical equations must be balanced to obey the law of conservation of mass, which states that matter cannot be created or destroyed.
      • This means the same number of each type of atom must appear on both sides of the equation.

    2. Steps to Balance

      • Write the correct formulas for all reactants and products.
      • Use coefficients (numbers before compounds) to make the number of atoms for each element equal on both sides.
      • Check that each atom type balances and that charges balance if it’s an ionic equation.

      Example

      Balance this equation for the reaction of hydrogen and oxygen to form water:

      Unbalanced: H₂ + O₂ → H₂O Solution

      1. Start by balancing O atoms. We need 2 O atoms on each side, so add a coefficient of 2 before H₂O: H₂ + O₂ → 2H₂O
      2. Now balance H by placing a coefficient of 2 before H₂: 2H₂ + O₂ → 2H₂O

      Practice Question

      Balance the equation for the reaction of magnesium with hydrochloric acid to produce magnesium chloride and hydrogen gas:

      Mg + HCl → MgCl₂ + H₂

      Solution

      1. Balance Cl by putting a 2 before HCl: Mg + 2HCl → MgCl₂ + H₂
      2. Check that each element is balanced (1 Mg, 2 Cl, and 2 H on each side).

    Law of Conservation of Mass

    • Definition In any chemical reaction, the total mass of reactants equals the total mass of products.
    • Application When calculating masses of reactants or products, always ensure the mass on both sides of the reaction equation is the same.

    Practice Question

    If 20 g of hydrogen gas reacts completely with 160 g of oxygen gas to produce water, what will be the mass of the water formed?

    Solution

    1. Total mass of reactants = 20 + 160 = 180 g.
    2. Since mass is conserved, the mass of water produced will also be 180 g.

    Using Moles in Chemical Equations

    1. Mole Ratios
      • In a balanced chemical equation, coefficients represent the mole ratio of reactants and products.
      • Example: In the equation 2H₂ + O₂ → 2H₂O, 2 moles of H₂ react with 1 mole of O₂ to form 2 moles of H₂O.
    2. Conversions
      • To calculate amounts of reactants or products, first convert given quantities to moles.
      • Use mole ratios to find moles of the desired substance, then convert back to grams if needed.

    Practice Question

    In the reaction 2NaOH + H₂SO₄ → Na₂SO₄ + 2H₂O, how many grams of Na₂SO₄ are produced when 4 g of NaOH react completely?

    Solution

    1. Find moles of NaOH: Molar mass of NaOH = 40 g/mol Moles of NaOH = 4 g / 40 g/mol = 0.1 mol
    2. Mole ratio: 2 moles of NaOH produce 1 mole of Na₂SO₄, so 0.1 mol NaOH will produce 0.05 mol Na₂SO₄.
    3. Convert moles of Na₂SO₄ to grams: Molar mass of Na₂SO₄ = 142 g/mol Mass = 0.05 x 142 = 7.1 g

    Limiting Reactant and Percent Yield

    1. Limiting Reactant

      • The reactant that is completely used up first, limiting the amount of product that can form.

      Practice Question

      For the reaction N₂ + 3H₂ → 2NH₃, if you start with 1 mole of N₂ and 2 moles of H₂, which is the limiting reactant?

      Solution

      • Mole ratio requires 3 moles of H₂ for every 1 mole of N₂.
      • Since we only have 2 moles of H₂, H₂ is the limiting reactant.

    2. Percent Yield

      • The efficiency of a reaction, comparing actual yield to theoretical yield.
      • Formula: % Yield = (Actual Yield / Theoretical Yield) x 100

      Practice Question

      If a reaction has a theoretical yield of 10 g but only produces 8 g, what is the percent yield?

      Solution

      % Yield = (8 / 10) x 100 = 80%

  • Intra-Molecular Bonding

    • Describe and compare the three types of intra-molecular bonding (ionic, metallic and covalent), and apply this knowledge to predict the physical and/or chemical properties of a substance.
    • Explain valence bond theory in terms of overlap of atomic orbitals.
    1. Intra-Molecular Bonding

    Types of Chemical Bonds: Ionic, Covalent, and Metallic

    What is an Intra-Molecular Bond?

    Intra-molecular bonds are the forces that hold atoms together within a molecule or compound. These bonds determine a substance’s physical and chemical properties, such as melting point, boiling point, hardness, and electrical conductivity.

    The main types of intra-molecular bonds are ionic, covalent, and metallic.

    1. Ionic Bonding

    • Definition: Ionic bonding occurs when electrons are transferred from one atom to another, resulting in the formation of positive and negative ions. This usually happens between a metal and a non-metal.

      • Metal atoms lose electrons and become positively charged ions (cations).
      • Non-metal atoms gain those electrons and become negatively charged ions (anions).

    • Example: Sodium chloride (NaCl), or table salt

      • Sodium (Na) loses one electron to become Na⁺, and chlorine (Cl) gains that electron to become Cl⁻.
      • The positive Na⁺ and negative Cl⁻ ions are attracted to each other, forming a strong bond.

    • Properties of Ionic Compounds:

      • High melting and boiling points: Strong attraction between oppositely charged ions requires a lot of energy to break.
      • Hard but brittle: They’re hard due to strong ionic bonds but break easily because shifting ions cause like charges to repel.
      • Electrical conductivity: Solid ionic compounds do not conduct electricity because ions are fixed in place. However, when melted or dissolved in water, they do conduct electricity because ions are free to move.

      Practice Comparison: Why does NaCl have a high melting point compared to water (H₂O)?

      • Answer: The ionic bonds in NaCl are much stronger than the hydrogen bonds in water, requiring more energy (heat) to break.

    • More – What is Ionic Bonding?

    2. Covalent Bonding

    • Definition: Covalent bonding occurs when two non-metal atoms share one or more pairs of electrons. This way, both atoms achieve a stable electron configuration (usually completing their outer electron shell).

    • Example: Water (H₂O) and carbon dioxide (CO₂)

      • In H₂O, each hydrogen atom shares one electron with oxygen, forming two covalent bonds.
      • In CO₂, each oxygen shares two electrons with carbon, forming double covalent bonds.

    • Properties of Covalent Compounds:

      • Low melting and boiling points (in simple molecules): These molecules are held together by weak intermolecular forces, requiring less energy to break.
      • Soft or brittle in solid form: Many are soft or powdery because the forces holding the molecules together are weak.
      • Non-conductive: Covalent compounds generally do not conduct electricity because there are no free ions or electrons to carry a charge.

      Practice Comparison: Why is oxygen gas (O₂) a poor conductor of electricity?

      • Answer: In O₂, the covalent bond involves electron sharing, not transferring, so there are no free ions or electrons available to conduct electricity.

    • More – What is Covalent Bonding?

    3. Metallic Bonding

    • Definition: Metallic bonding occurs between metal atoms, where electrons are not bound to any one atom and can move freely. This forms a “sea of electrons” that holds the positively charged metal ions together.

    • Example: Copper (Cu) and aluminum (Al)

      • In copper, each atom contributes free electrons, creating a structure of positive ions in a sea of delocalized electrons.

    • Properties of Metallic Compounds:

      • High melting and boiling points: Strong forces between positive ions and free electrons require a lot of energy to break.
      • Malleable and ductile: Metals can be hammered into shapes or drawn into wires because the ions can shift without breaking the metallic bond.
      • Excellent conductors of electricity and heat: Free electrons in metals can move easily, allowing them to carry electric current and transfer heat.

      Practice Comparison: Why is copper used for electrical wiring instead of sodium?

      • Answer: Copper has strong metallic bonds and is a better conductor due to its high melting point, whereas sodium is soft with weaker bonds and a low melting point, making it unsuitable for wiring.

    • More – What is Metallic Bonding?

      Metallic bonding is the type of bond found between atoms of metallic elements. In metallic bonding:

      • Metal atoms lose their outer (valence) electrons, forming positively charged ions, or cations.
      • These cations arrange themselves in a regular, organized structure known as a lattice.
      • The lost valence electrons do not stay attached to any one atom but instead form a “sea of electrons”that moves freely around the fixed cations.

      This free movement of electrons is what gives metals their unique properties.

      Characteristics of Metallic Bonding

      1. Delocalized Electrons
        • In a metallic bond, electrons are delocalized, meaning they are not attached to any single atom and can move freely through the lattice.
        • This movement creates a strong attraction between the positively charged cations and the negatively charged electrons, holding the metal structure together.
      2. Bond Strength
        • The strength of the metallic bond depends on:
          • Number of delocalized electrons: More delocalized electrons result in stronger bonds.
          • Size of the cations: Smaller cations allow electrons to get closer, strengthening the attraction.
        • Example: Aluminum (Al) forms a stronger metallic bond than sodium (Na) because Al has more delocalized electrons per atom and smaller cations than Na.

      Properties of Metals Due to Metallic Bonding

      1. Electrical and Thermal Conductivity

        • Metals are excellent conductors of electricity and heat.
        • The free-moving electrons in the “sea” allow electric current to flow easily through the metal when a voltage is applied.
        • They also enable metals to conduct heat well by transferring kinetic energy (heat) quickly through the structure.

      2. Malleability and Ductility

        • Metals are malleable (can be hammered into shapes) and ductile (can be drawn into wires).
        • The layers of metal cations can slide over each other without breaking the metallic bond because the “sea of electrons” continues to hold them together, preventing repulsion between the positive ions.

      3. Density and Melting Temperature

        • Metals tend to have high density and high melting points, though these vary across the periodic table.
          • Group 1 metals (e.g., sodium) have relatively low melting points because they lose only 1 electron each, resulting in weaker metallic bonds.
          • Transition metals (e.g., iron, copper) have higher melting points due to having more delocalized electrons and smaller cations, creating stronger bonds.
        • Trend Across a Period: Moving from left to right in a period (row) of the periodic table, melting points generally increase because the number of delocalized electrons per cation increases, strengthening the bond.
        • Trend Down a Group: Moving down a group (column), melting points tend to decrease because cations become larger, which weakens the bond.

        Example: In the third row, the melting points of metals increase in the order: Na < Mg < Al.

        • Explanation: Sodium (Na) has 1 delocalized electron per atom, magnesium (Mg) has 2, and aluminum (Al) has 3. The more delocalized electrons per atom, the stronger the metallic bond, leading to higher melting points.

      4. Ionization Energy and Electronegativity

        • Ionization Energy: Metals generally have low ionization energies (energy required to remove an electron) because their outer electrons are loosely held.
        • Electronegativity: Metals have low electronegativity (tendency to attract electrons in a bond). This is why metals readily lose electrons to form cations.

      Practice Question and Example

      Question: Why does aluminum (Al) have a higher melting point than sodium (Na)?

      Answer:

      • Step 1: Consider the number of delocalized electrons.
      • Sodium (Na) loses 1 electron to form Na⁺, while aluminum (Al) loses 3 electrons to form Al³⁺.
      • Step 2: Compare bond strength.
      • More delocalized electrons in Al mean a stronger metallic bond than in Na.
      • Conclusion: Aluminum has a higher melting point than sodium because its metallic bonds are stronger due to more delocalized electrons.

    Comparing and Explaining Differences in Properties

    Understanding why these substances behave differently can help predict their properties:

    1. Strength of Bonds
      • Ionic bonds are strong due to the attraction between oppositely charged ions.
      • Covalent bonds can vary in strength, but the forces between molecules (intermolecular forces) are typically weaker.
      • Metallic bonds involve a strong attraction between positive ions and free electrons.
    2. Structure
      • Ionic compounds form a crystalline structure, leading to hardness and brittleness.
      • Covalent compounds (like water) usually form simple molecules with weaker forces between them, leading to lower melting points.
      • Metallic compounds have a flexible structure of ions and free electrons, leading to malleability and ductility.
    3. Conductivity
      • Ionic compounds conduct electricity only when ions are free to move (melted or in solution).
      • Covalent compounds generally do not conduct electricity because they lack free electrons or ions.
      • Metallic compounds are excellent conductors due to their sea of free electrons.

    Valence Bond Theory (in Terms of Atomic Orbital Overlap)

    Valence Bond Theory (VBT) is a way to explain how covalent bonds form through the overlapping of atomic orbitals. Here’s a breakdown:

    1. Atomic Orbitals

      • Atoms have orbitals where electrons are likely to be found (like s and p orbitals).
      • When two atoms come close, their orbitals can overlap.

    2. Bond Formation Through Overlap

      • A covalent bond forms when orbitals overlap, allowing atoms to share electrons.
      • Types of Overlap:
        • Sigma (σ) bonds: Formed by head-on overlap of orbitals (e.g., the single bond in H₂).
        • Pi (π) bonds: Formed by side-by-side overlap of p orbitals, as seen in double or triple bonds.

    3. Predicting Bond Strength and Length

      • The more extensive the overlap, the stronger the bond.
      • Sigma bonds are stronger than pi bonds because the overlap is greater.
      • Single bonds (1 sigma) are longer and weaker than double bonds (1 sigma + 1 pi), which are shorter and stronger.

      Example Application: Why is a double bond in oxygen (O₂) stronger than the single bond in hydrogen (H₂)?

      • Answer: O₂ has both sigma and pi bonding, providing stronger attraction between atoms than the single sigma bond in H₂.