Understanding Chemical Bonds, Atomic Structure, and Periodic Trends

Posted on Aug 27, 2024 in Chemistry

Chemical Formulas and Compounds

Chemical Formula Parts

  • Elements
  • Number of Atoms
  • Charge

Molecular Compounds

  • Compounds formed of two or more nonmetals.

Formula Unit

  • The basic unit of ionic compounds.
  • The smallest electrically neutral collection of ions.

Diatomic Elements

  • F, Cl, Br, I, N, O, H

Ionic and Covalent Compounds

Why are Bigger Bonds Stronger?

  • Both atoms’ nuclei are attracted to shared electrons.

Ionic Compounds at Room Temperature

  • Phase: Solid
  • Why? With strong bonds, oppositely charged ions pack tightly together and form a crystal lattice structure that is difficult to break.

Defining Ionic Compounds

  • Consist of charged ions with opposite charges.
  • Generally, a metal and a nonmetal.

Electronegativity

  • Measure of the tendency of an atom to attract a bonding pair of electrons.
  • Fluorine has an electronegativity of 4; it decreases as you move away from fluorine.

Ionic Character

  • Grows as electronegativity difference increases.

Intermolecular Attraction

  • Grows as electronegativity difference increases.

Covalent Compounds at Room Temperature

  • State of Matter: Usually liquid or gas.
  • Why? Due to weak intermolecular forces, they have low melting and boiling points. Unlike ionic compounds, they mainly exist as gases.

Best Ionic Solvent

  • Water
  • Why? Each hydrogen atom carries a slight positive charge, while the oxygen atom carries a slight negative charge. Ionic compounds are crystalline and contain equal numbers of positive and negative ions. When dissolved in water, water molecules attract the ions, causing separation between the ionic molecules.

Best Covalent Solvent

  • Alcohol, oil, any organic or non-polar solvent.
  • Why? There is no strong attraction between the solvent and the covalent molecules.

Melting Point

  • Ionic: High melting point due to strong bonds.
  • Covalent: Lower melting point due to weaker bonds.

Atomic Structure and Periodic Trends

Atomic Size Across a Period

  • Why do atoms get smaller across a period?

1) Electric force of attraction grows stronger as the number of protons and electrons increases.

BUT

2) When a new period begins, size increases because the old valence electrons now shield the new level of valence electrons, which are loosely bound, making the atom larger.

Atomic Size Down a Group

  • Why are atoms larger moving down a group? Each atom has a new energy level, which is further away from the nucleus than the atom above.

Electron Attraction and Distance from Nucleus

  • Is electron attraction related to distance from the nucleus? If so, why? Yes; the closer an electron is to the nucleus, the more attracted it is, and the more difficult it is to remove.

Electron Affinity

  • Amount of energy either required or released for 1 mole of atoms to accept an electron.

Stable Groups

  • s2, d5, d10, p3, Group VIII

Negative vs. Positive Electron Affinity

  • Most electron affinities are negative.
  • Negative EA: Easy to gain an electron.
  • Positive EA: Resists gaining an electron.

Ionization Energy

  • The amount of energy required to remove an electron from a mole of atoms.

Mole

  • 6.02 x 1023

Ionization Energy Equation

  • Energy + Atom = Atom+ + e–

Atoms in the Same Periods

  • Have the same number of electron shells/energy levels.

Atoms in the Same Groups

  • Have the same number of valence electrons.

Groups in Increasing EA Negativity

  1. Group II and Group VIII (positive)
  2. Group I
  3. Group III and Group V
  4. Group IV
  5. Group VI
  6. Group VII

Ionization Energy is Always ____thermic

  • Endothermic

Metal Reactivity

  • Metal atoms are most reactive when large in size and low in ionization energy. This is when valence electrons are most easily removed.

Units of Measurement

  • Atomic radius: picometers (pm)
  • Ionization energy: kilojoules per mole (kJ/mol)

Historical Atomic Theories

Democritus

  • Atomic theory: The universe is composed of atoms and the void in which they exist.
  • Infinite number of atoms, always moving, invisible, never destroyed, differing in size, shape, and temperature.

Dalton

  1. Elements are made of atoms.
  2. All atoms of an element are identical.
  3. Different elements have different properties.
  4. Atoms are not created, destroyed, or split.
  5. In chemical reactions, atoms link and separate.
  6. Atoms combine in simple whole-number ratios to form compounds.

Rutherford

  • Discovered and named the atom’s nucleus and nuclear half-lives.
  • Deliberately turned one atom into another.
  • Gold Foil Experiment: Shot positive particles at gold foil; some deflected, some deflected straight back, some went through, showing there were charged particles but mostly empty space.

Thomson

  • Discovered the electron using a cathode ray tube to negatively charge particles.

Bohr

  • Theorized atomic structure and the Bohr model.

Key Concepts

  • Electrons determine elemental behavior.
  • Protons and neutrons have a mass of 1 amu.
  • Nuclear Attraction vs. Electric Force: In the nucleus, nuclear attraction holds it together, outweighing the effect of electrical force.
  • EMR speed: 3.0 x 108 m/s
  • Planck’s constant: 6.6 x 10-34 Js
  • Frequency unit: 1/s
  • Energy unit: J
  • Wavelength (λ) unit: m
  • Energy equation: E = h * f
  • Speed equation: c = f * λ
  • Electrons per energy level: 2n2