Key Concept: Energy calculations related to quantum states
Formula: ( E = -2.18 imes 10^{-18} \text{ J} \left( \frac{1}{n_f^2} - \frac{1}{n_i^2} \right) )
Initial state: ( n = 3 )
Final state: ( n = 1 )
Energy change: ( \Delta E = E_f - E_i ) and can be calculated using the above formula.
Important Observations:
Energy values become more negative as you go to lower n levels.
Change in energy is positive when moving to higher energy levels and negative when moving to lower.
Quantum Numbers:
Principal quantum number (n): Represents energy level and size of orbital.
Azimuthal quantum number (l): Defines the shape of the orbital.
Magnetic quantum number (m): Indicates orientation of the orbital.
Spin quantum number (s): Represents the electron's spin.
Pauli Exclusion Principle:
No two electrons in an atom can have the same set of four quantum numbers.
Aufbau Principle:
Electrons fill orbitals in order of increasing energy.
Electron Configurations:
Example of Argon: ( 1s^2 2s^2 2p^6 3s^2 3p^6 )
Includes total number of protons/electrons in neutral atoms.
Cations and Anions:
Cations: Positively charged (loss of electrons).
Anions: Negatively charged (gain of electrons).
Valence Electrons:
Electrons in the outermost shell; dictate chemical properties.
Noble Gases: Inert and stable due to complete valence shells.
Core Electrons: Inner shell electrons, not involved in bonding.
Transition and inner transition metals have unique configurations.
Example for Chromium: ( [Ar] 4s^1 3d^5 )
Example for Copper: ( [Ar] 4s^1 3d^{10} )
Cation (+) and Anion (-) definitions
Effective Nuclear Charge (Z_eff): Calculated as ( Z - S ) (Z = atomic number, S = core electrons).
Trends in the Periodic Table:
Atomic radius increases down a group and decreases across a period.
Ionization energy increases across a period and decreases down a group.
Electron affinity generally more negative across a period.
Ionic vs Covalent Bonds:
Ionic bonds involve electron transfer (metal + non-metal).
Covalent bonds involve electron sharing (non-metal + non-metal).
Examples of molecular formulas vs empirical formulas:
Molecular: C4H8, Empirical: C2H4
Ionic compounds are typically salts, formed from metal cations and non-metal anions.
Formula writing involves balancing charges.
Learn common names for cations/anions and how to name ionic compounds using roman numerals for transition metals.
Example: Iron (III) chloride for ( FeCl_3 )
Important oxyanions: Nitrate (NO3^-), Sulfate (SO4^2-), and Phosphate (PO4^3-).
Ionic compounds act as strong electrolytes, while molecular compounds are typically weak electrolytes.
Hydrated ionic compounds contain water in their crystalline structure ( i.e. ( BaCl_2 \cdot 2H_2O )).
Review how to calculate molecular weights and mass percent composition of compounds.
Use formula: ( , %X = \left( \frac{mass \ of \ X}{mass \ of \ compound} \right) \times 100 % )
Bonds can be classified as ionic, covalent, and metallic based on how electrons are shared or transferred.
Electronegativity (EN) is a key factor in determining bond polarity.
Construct Lewis Dot Structures showing valence electrons and bonding.
Calculate formal charges to assess stability and resonance structures.
Understand that resonance structures denote the same connectivity but differ in electron placement.
Recognize common polyatomic ions and know how to calculate formal charges in compounds.
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