Chemistry Flashcards
Distinction
Electron configuration
refers to the distribution of electrons among the orbitals of an atom.
Electrons filling orbitals
in order of increasing energy, starting from the lowest energy orbital (1s) and moving upwards. This is based on the Aufbau Principle.
Electron filling orbital sequence
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p
Hund’s Rule
Electrons will occupy empty orbitals of the same energy (degenerate orbitals) singly before pairing up.
Pauli Exclusion Principle
No two electrons can have the same set of quantum numbers, so an orbital can hold a maximum of two electrons with opposite spins.
Electron Configurations for Elements
Use the atomic number to determine how many electrons an atom has, and then distribute them according to the orbital filling order.
Determine the Position of an Element in the Periodic Table from its Electronic Structure - Group Number
The group number corresponds to the number of electrons in the outermost shell.
Determine the Position of an Element in the Periodic Table from its Electronic Structure - Period
The period corresponds to the highest energy level (principal quantum number, n) that contains electrons.
Sodium (Na): 1s² 2s² 2p⁶ 3s¹
Period and Group -
Group 1 (1 electron in the outermost s orbital).
Period 3 (highest principal quantum number is 3).
Blocks of the Periodic Table:
Elements are divided into blocks based on their electron configuration
d-block: Transition metals
S-Block
P-block
Transition Metals
Transition metals have partially filled d orbitals, and their electron configurations typically end in d orbitals
Ion Formation
Ions -
Ions form when atoms lose or gain electrons to achieve a full outer shell (usually achieving the same electron configuration as a noble gas).
ions formation
metal -
lose electrons to form positive ions (cations)
Non-metals
gain electrons to form negative ions (anions).
Examples of Ion formation
Sodium (Na): Loses 1 electron to form Na⁺.
Configuration: Na⁺ = 1s² 2s² 2p⁶ (same as Neon).
Chlorine (Cl): Gains 1 electron to form Cl⁻.
Configuration: Cl⁻ = 1s² 2s² 2p⁶ 3s² 3p⁶ (same as Argon).
Formulae of Ionic Compounds
Combine cations and anions to form neutral compounds.
Sodium Chloride (NaCl):
Na⁺ and Cl⁻ combine to form NaCl (ionic bond).
Calcium Oxide (CaO):
Ca²⁺ and O²⁻ combine to form CaO.
Ionic Bonding
Occurs between metals and non-metals.
Involves the transfer of electrons from a metal to a non-metal, forming positive and negative ions
Ionic Bonding properties
High melting and boiling points.
Conduct electricity when dissolved in water (as ions are free to move).
Brittle.
Covalent Bonding
Occurs between non-metals.
Involves the sharing of electron pairs between atoms.
Covalent Bonding Properties
Lower melting and boiling points compared to ionic compounds.
Do not conduct electricity.
Metallic Bonding
Occurs in metals.
Delocalized electrons (free electrons) move throughout the lattice of metal cations.
Metallic Bonding Properties
Conduct electricity.
Malleable and ductile.
Intermolecular Forces
These are forces between molecules (not within molecules)
Van der Waals Forces
Weak interactions between non-polar molecules.
Example: Between O₂ molecules.
Hydrogen Bonding
Stronger than Van der Waals, occurs when hydrogen is bonded to electronegative atoms (N, O, F).
Example: Between H₂O molecules (responsible for water’s high boiling point).
Prediction of Bonding Types
Ionic: Typically formed between metals and non-metals.
Covalent: Typically formed between two non-metals.
Metallic: Found in pure metals and alloys.
Intermolecular Forces: Found between molecules with covalent bonds.
