week 5 Lecture 3 - Topic 4 (Atomic energy levels)

Periodic Trends

  • Review of previous material and introduction to bonding.
  • Sharing of teaching responsibilities among lecturers.
  • Importance of learning from different teaching styles.

Atomic Radius and Effective Nuclear Charge

  • Atomic size decreases across a period due to increasing effective nuclear charge.
  • Atomic size increases down a group due to additional electron shells.
  • Shielding effect of core electrons.

Electronegativity

  • Electronegativity: An element's ability to attract electrons.
  • Lithium (Li) tends to lose an electron to achieve the electron configuration of helium (He), forming Li^+.
  • Fluorine (F) tends to gain an electron to achieve the electron configuration of neon (Ne), forming F^-.
  • Electronegativity increases from bottom left to top right on the periodic table, excluding noble gases.
  • Noble gases are not electronegative as they have stable electron configurations.
  • Halogens (Group 17) are the most electronegative group.
  • Elements that form +1 cations (Group 1) are the least electronegative.
  • Diagonal line visualization for electronegativity trend.

Electron Affinity

  • Electron affinity: Energy released when an electron is added to an atom.
  • Related to enthalpy change (heat or energy) in thermochemistry.
  • Relevant to bonding and polarity of molecules.

Polarity

  • Uneven distribution of electrons in a molecule.
  • Example: Hydrogen chloride (HCl) where chlorine is more electronegative than hydrogen.
  • Chlorine has a partial negative charge (\delta^-), while hydrogen has a partial positive charge (\delta^+).
  • The electron cloud is denser around the more electronegative element.

Electron Affinity Trends

  • General trend: Electron affinity increases across a period.
  • Fluorine has a high electron affinity because it strongly attracts electrons.

Problem Solving: Applying Periodic Trends

  • Using knowledge of periodic trends to predict properties.

Atomic Size

  • Trend: Atomic size increases down a group and decreases across a period.
  • Example: Comparing the atomic size of sodium, phosphorus, chlorine, potassium, and fluorine.
  • Potassium has the largest atomic size among the given elements.

Exceptions to General Trends

  • Polonium (Po) having a larger atomic radius than bismuth (Bi) despite being to the right of it in the periodic table.
  • Emphasis on general trends and exceptions.

Ionic Size

  • Cations are smaller than their neutral atoms.
  • Anions are larger than their neutral atoms.
  • Example: Comparing the ionic size of sodium, aluminum, fluoride, rubidium, and iodide ions.
  • Aluminum cation is the smallest among the given ions.

Comparing Atoms and Ions

  • Applying knowledge of trends to compare atoms and ions.
  • Example: Comparing potassium cation, chloride anion, aluminum neutral species, bromide anion, and selenium anion.
  • Selenium anion and bromide anion have very similar sizes.

Smallest Ionic Size

  • The smallest between aluminum and potassium ions requires specific data and isn't easily predicted by trends alone.
  • Aluminum cation: 57 picometers.
  • Potassium cation: 33 picometers.
  • Fluoride: 133 picometers
  • Sodium: 98 picometers

Ionization Energy

  • Ionization energy: The energy required to remove an electron from an atom.
  • Largest ionization energy means the element least wants to give up an electron.
  • Ionization energy decreases as atomic size increases (down a group).
  • Ionization energy increases across a period and up a group.

Trends in Ionization Energy

  • Group 1 atoms have low ionization energies.
  • Halogens have high ionization energies.
  • Fluorine has the largest ionization energy among sodium, phosphorus, chlorine, potassium, and fluorine.
  • Potassium has the smallest ionization energy.

Recap of Atomic Energy Levels

  • Atomic structure: Nuclei with protons and neutrons.
  • Atomic number: Number of protons.
  • Electrostatic attraction between protons and electrons.
  • Characteristics of light: Atoms absorb light, leading to different colors.

Orbitals

  • s, p, and d orbitals: Shapes and orientations.
  • s orbitals: One orientation.
  • p orbitals: Three orientations.
  • d orbitals: Five orientations.
  • Principal quantum number (n): Related to atomic size and energy.

Introduction to Bonding

  • Chemical bonding and molecular structure.
  • Qualitative understanding of bonding types.
  • Covalent bonding, ionic bonding, and other interactions.
  • Lewis structures.
  • Geometry of compounds.
  • Properties of bonds.

Polar Covalent Bond

  • Example: Hydrogen chloride (HCl).
  • Unequal sharing of electrons due to electronegativity difference.

Ionic Bond

  • Example: Sodium chloride (NaCl).
  • Transfer of electrons due to large electronegativity difference.

Metallic Bonding

  • Bonding in metals.

Bond Polarity and Electronegativity

  • Bond polarity related to electronegativity differences.
  • Electronegativity increases from bottom left to top right (excluding noble gases).

Fundamentals of Bonding

  • Intramolecular interactions.
  • Repulsion between nuclei.
  • Attraction between electron cloud and nuclei.
  • Energetically favorable distance for bonding.

Bond Strength

  • Relationship between bond length and bond energy.
  • Stronger bonds require more energy to break and release more energy when formed.
  • Breaking bonds requires energy; making bonds releases energy.

Bond Length

  • Distance between neighboring nuclei.

Bond Energy

  • Energy required to break a bond.

Energy Plot

  • Graph of energy versus distance between nuclei.
  • Goldilocks zone: Optimum distance for bond strength.
  • Repulsion at close distances.

Types of Bonds

  • Fluorine gas (F_2): Single bond (sigma bond).
  • Each fluorine atom shares one electron to achieve a full octet.

Uneven Sharing

  • Polar compounds: Uneven sharing of electrons.