Electron Configuration
To determine electron configuration, the basic sequence is: 1s², 2s², and so on.
There are two types of arrangements for electron configuration:
Ordering Based on Energy: Follows the pattern from the periodic table.
Ordering Based on Shell: Both arrangements are acceptable and students should be prepared to recognize either.
Shorthand Configurations
Shorthand electron configurations utilize the nearest noble gas prior to the element in question. This is represented by placing the noble gas in brackets before adding any additional electrons.
Common exceptions to the usual configurations include elements like Chromium (Cr), Molybdenum (Mo), Copper (Cu), and Silver (Ag), where an electron may be transferred from the s subshell to the d subshell in order to achieve a more stable electron configuration (e.g., a half-filled or fully filled d shell).
Shielding Effect
The concept of shielding refers to the number of core electrons that shield the effect of the nuclear charge (protons) on the valence electrons:
Calculation: Shielding = Protons - Core Electrons.
The shielding effect allows predictions of atomic size (atomic radius), as effective nuclear charge influences the attraction between the nucleus and the electrons.
Energy Levels and Orbitals
According to periodic trends, the energy levels can be understood as:
3d is lower in energy than 4p.
3d is higher in energy than 4s.
Orbital Diagrams: These visual representations can display either all electrons (core and valence) or just valence electrons depending on context.
When drawing orbital diagrams, it is important to abide by:
Pauli Exclusion Principle: No two electrons can have the same four quantum numbers. Thus, if two electrons occupy the same orbital, they must have opposing spins (denoted as one spin-up and one spin-down).
Hund's Rule: When filling degenerate orbitals (orbitals of the same energy), electrons should be placed singularly in each orbital before any pairing occurs. This specifically applies to p, d, and f orbitals.
Magnetic Properties
Understanding whether a substance is paramagnetic or diamagnetic is essential. Paramagnetic substances have unpaired electrons, while diamagnetic substances have all electrons paired.
Periodic Trends
Key trends to be aware of include:
Ionization Energy: Energy required to remove an electron from an atom.
Atomic Radius: The size of an atom, which tends to decrease across a period and increase down a group.
Metallic Character: Generally decreases across a period and increases down a group. It is intuitive that metals, with their tendency to lose electrons, have greater metallic character than nonmetals.
Specific examples of trend behavior include:
Fluorine: Small atomic radius, high ionization energy, high electronegativity, and high electron affinity. Positioned as a 'small, feisty' atom that wants to gain electrons.
Ions and Their Properties
Cations vs. Anions:
Cations (positively charged ions) are smaller than their neutral atoms due to the loss of electrons, while Anions (negatively charged ions) are larger due to the gain of electrons.
For isoelectric species (species with the same electron configuration), shielding can be utilized to predict the relative sizes of ions.
Electron Affinity: The energy change that occurs when an electron is added to a neutral atom.
Understanding when ions form typically involves the loss or gain of electrons, as protons do not generally change outside of nuclear reactions.
Bonding and Compound Naming
Basic naming conventions for compounds include:
For ionic compounds, it follows charge patterns ranging from plus one to minus three depending on the elements involved.
Polyatomic ions require familiarity with specific names and charges.
Transition Metals: Naming can involve the use of Roman numerals to indicate the charge of the metal ions followed by a drop and swap where necessary, making sure that the charges balance.
Hydrates and Covalent Compounds: In covalent compound types, prefixes are used (mono-, di-, tri-, tetra-, etc.) to indicate the number of atoms present.
Calculations
Molar Mass and Conversions
Molar mass is the mass of one mole of an element or compound, typically expressed in grams per mole (g/mol). This is essential when converting between moles, atoms, and molecules:
Avogadro's Number (Approx. $6.022 imes 10^{23}$): Used to convert between moles and individual atoms, molecules, or formula units.
Conversion Factors: Include molar mass, Avogadro's number, and the chemical formula itself for computation between moles and other measurements in chemistry.
Empirical and Molecular Formulas
The steps to calculate empirical formulas from mass percent include:
Convert percentages to grams (assuming 100 g total).
Convert grams to moles by dividing by the atomic weight of each element.
Divide each mole quantity by the smallest number of moles calculated.
If necessary, multiply to get to the smallest whole numbers to represent the empirical formula.
Bonding Patterns
In bond formation, basic sequences are:
Count, connect, fill, and check to confirm bonding.
Recognize and predict common bonding patterns for elements based on group number.
Resonance Structures: Some molecules can be represented by multiple valid Lewis structures, contributing to the resonance hybrid.
Formal Charge: This can be calculated using the following method:
Formal Charge = (Valence Electrons) - (Non-Bonding Electrons) - (1/2 Bonding Electrons).
It’s beneficial to aim for a formal charge of zero for stabilities, especially in resonance scenarios.
Molecular Geometry
Understanding geometry is enhanced through Valence Shell Electron Repulsion (VSEPR) theory, which suggests that electron groups (lone pairs and bonds) will position themselves as far apart as possible.
Key Geometries: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral shapes result depending on the number of bonding and lone electron pairs.
Always recognize that double bonds and triple bonds count as one single group toward the geometry arrangement.
Hybridization
Hybridization Process:
Involves blending atomic orbitals to form hybrid orbitals that can better explain the number of bonds an atom can make as per its electron configuration, e.g.,
sp3 Hybridization: Involves combining one s and three p orbitals creating four equivalent orbitals spaced 109.5 degrees apart (tetrahedral).
sp2 Hybridization: Combines one s and two p orbitals with one remaining unhybridized p orbital for double-bonded structure (120 degrees apart).
sp Hybridization: When one s and one p are hybridized, resulting in linear geometry (180 degrees apart).
Sigma and Pi Bonds
Sigma Bonds: Formed by end-to-end overlap of orbitals. It can occur with:
Two s orbitals, two p orbitals, or one s and one p orbital.
Pi Bonds: Formed by lateral overlap of unhybridized p orbitals. Important in multiple bonds (e.g. double and triple bonds). These bonds enhance molecular interactions and stability beyond what sigma bonds provide.
Summary of Chemical Interactions
Understanding electron configurations, hybridization, bonding types, and molecular geometry are vital for predicting chemical behavior.
This comprehensive knowledge allows chemists to rationalize reactions, compound formation, and molecular properties effectively in practical chemistry applications.
Exam Preparation
Highlight important concepts on the constant sheet provided for examinations, especially related to electron configurations, molecular geometry, and common polyatomic ions. Review past materials and practice essential calculations for proficiency.