Chem 20 Unit A Part 1-2 CYU Definitions Flashcards
(28 cards)
What is Bonding capacity?
Bonding Capacity refers to the number of chemical bonds an atom can form based on the number of available valence electrons.
• Determined by the number of unpaired electrons in the valence shell.
Hydrogen (H): 1 bond
What is bond Dipole
Bond Dipole: Unequal electron sharing in a polar covalent bond due to electronegativity differences.
• Arrow (→) points to more electronegative atom (δ⁻).
• Example: H → Cl (H is δ⁺, Cl is δ⁻).
• Stronger dipole = bigger electronegativity difference.
What is a binding electron
A bonding electron is a valence electron that is shared or transferred between atoms to form a chemical bond.
• In covalent bonds: Shared between atoms (e.g., H₂O).
• In ionic bonds: Transferred from one atom to another (e.g., NaCl).
• Pairs of bonding electrons create single, double, or triple bonds.
What is a central atom
A central atom is the atom in a molecule that bonds to multiple other atoms and is usually the least electronegative (except hydrogen).
• Example: In H₂O, oxygen (O) is the central atom.
• In CO₂: Carbon (C) is the central atom.
• Usually has the highest bonding capacity.
What is a covalent bond
A covalent bond is a chemical bond where two non-metal atoms share electrons to achieve a full valence shell.
• Single Bond (1 pair shared): H—H (H₂)
• Double Bond (2 pairs shared): O=O (O₂)
• Triple Bond (3 pairs shared): N≡N (N₂)
• Polar Covalent: Unequal sharing (H₂O).
• Non-Polar Covalent: Equal sharing (O₂).
What is a crystal lattice
A crystal lattice is a repeating, three-dimensional structure of ions, atoms, or molecules arranged in a fixed pattern.
• Common in ionic compounds (e.g., NaCl).
• Held together by strong electrostatic forces.
• Results in high melting/boiling points and hardness.
What is a dipole dipole force
A dipole-dipole force is an intermolecular force between polar molecules, where the positive end (δ⁺) of one molecule attracts the negative end (δ⁻) of another.
• Stronger than dispersion forces but weaker than hydrogen bonding.
• Example: HCl molecules attract each other via dipole-dipole forces.
What is electronegativity
Electronegativity is an atom’s ability to attract shared electrons in a chemical bond.
• Trend: Increases across a period (left to right), decreases down a group (top to bottom).
• High electronegativity: Strong attraction (e.g., Fluorine, F = most electronegative).
• Difference in electronegativity determines bond type:
• 0 - 0.4 → Non-polar covalent
• 0.5 - 1.7 → Polar covalent
• > 1.7 → Ionic
What is Empircal Formula
The empirical formula is the simplest whole-number ratio of atoms in a compound.
• Example:
• Molecular formula: C₆H₁₂O₆ (glucose)
• Empirical formula: CH₂O (divided by 6)
• Represents proportions, not actual number of atoms
What is a hydrogen bond
A hydrogen bond is a strong dipole-dipole force that occurs when hydrogen (H) is bonded to fluorine (F), oxygen (O), or nitrogen (N) and is attracted to a nearby F, O, or N in another molecule.
• Stronger than regular dipole-dipole forces, but weaker than covalent bonds.
• Examples:
• H₂O (water) molecules stick together → high boiling point.
• DNA strands held together by hydrogen bonds.
What is Intermolecular Forces
Intermolecular forces (IMFs) are forces of attraction between molecules that determine physical properties like boiling and melting points.
Types of IMFs:
1. London Dispersion Forces (LDFs): Weak, in all molecules, stronger in larger atoms/molecules.
2. Dipole-Dipole Forces: Between polar molecules (e.g., HCl).
3. Hydrogen Bonds: Strong dipole force with H—N, H—O, or H—F (e.g., H₂O, NH₃).
Stronger IMFs = Higher boiling/melting points.
What is intermolecular forces
Intermolecular forces (IMFs) are forces of attraction between molecules, affecting boiling/melting points and solubility.
Types:
1. London Dispersion Forces (LDFs): Weak, in all molecules, stronger in larger ones.
2. Dipole-Dipole Forces: Between polar molecules (e.g., HCl).
3. Hydrogen Bonds: Strongest, in molecules with H—N, H—O, or H—F (e.g., H₂O).
Stronger IMFs = Higher boiling/melting points.
What is an ionic bond
An ionic bond is a chemical bond formed when a metal transfers electrons to a non-metal, creating oppositely charged ions that attract each other.
• Metal loses electrons → Becomes a cation (positive ion).
• Non-metal gains electrons → Becomes an anion (negative ion).
• Example: NaCl (Na⁺ and Cl⁻).
• Strong bond → High melting/boiling points.
What is the Lewis formula
The Lewis formula (Lewis structure) is a diagram that shows the bonding between atoms in a molecule and the valence electrons as dots.
• Bonding electrons → Shown as lines (shared pairs).
• Lone pairs (non-bonding electrons) → Shown as dots.
• Example: Water (H₂O)
Helps predict molecule shape and reactivity.
What is a lone pair what is a lone pair
A lone pair is a pair of valence electrons that are not shared or bonded with another atom in a molecule.
• Affect molecular shape (VSEPR theory).
• Can influence polarity and reactivity.
• Example: In H₂O, oxygen has two lone pairs:
What is a Molecular formula
A molecular formula shows the exact number of atoms of each element in a molecule.
• Unlike the empirical formula, it is not simplified.
• Example:
• Glucose: Molecular formula = C₆H₁₂O₆
• Water: Molecular formula = H₂O
• Represents the actual composition of the molecule.
What is a nonpolar covalent bond
A nonpolar covalent bond is a bond where electrons are shared equally between two atoms with little to no electronegativity difference (0 - 0.4).
• Occurs between identical atoms or similar nonmetals.
• No charge separation (no dipole).
• Examples:
• O₂ (Oxygen gas) → O=O
• CH₄ (Methane) → C—H bonds are nearly nonpolar.
What is a nonpolar molecule
A nonpolar molecule is a molecule with no overall dipole because electrons are evenly distributed or symmetrical in shape.
• Occurs when:
• All bonds are nonpolar (e.g., O₂, CH₄).
• Polar bonds cancel out due to symmetry (e.g., CO₂, CCl₄).
• Nonpolar molecules do not mix well with water (hydrophobic).
What is the octet rule
The octet rule states that atoms tend to gain, lose, or share electrons to achieve a full valence shell of 8 electrons, similar to noble gases.
• Applies to most main-group elements.
• Exceptions:
• H, He, Li, Be (stable with 2 electrons).
• Expanded octet: Elements in period 3+ (e.g., PCl₅, SF₆).
• Example: NaCl → Na⁺ loses 1 electron, Cl⁻ gains 1 to complete octets.
What is orbital
An orbital is a region around the nucleus where there is a high probability of finding an electron.
Each orbital can hold a maximum of 2 electrons with opposite spins.
What is a peripheral atom
A peripheral atom is an atom in a molecule that is bonded to the central atom but does not bond to multiple other atoms.
• Usually less electronegative than the central atom (except hydrogen).
• Example: In H₂O, hydrogen (H) atoms are peripheral, and oxygen (O) is the central atom.
• In CO₂: Oxygen (O) atoms are peripheral, and carbon (C) is the central ato
What is a polar covalent bond
A polar covalent bond is a bond where electrons are shared unequally between two atoms due to a difference in electronegativity (0.5 - 1.7).
• The more electronegative atom pulls electrons closer, creating partial charges (δ⁺ and δ⁻).
• Example:
• H₂O (Water): Oxygen (O) is δ⁻, Hydrogen (H) is δ⁺.
• HCl: Cl is δ⁻, H is δ⁺.
• Leads to dipole forces and affects solubility (polar dissolves in polar).
What is a polar molecule
A polar molecule is a molecule with an uneven distribution of electrons, creating a dipole (partial positive and negative ends).
Characteristics:
• Has polar bonds (electronegativity difference 0.5 - 1.7).
• Asymmetrical shape prevents dipoles from canceling.
Examples:
• H₂O (Water): Oxygen (δ⁻), Hydrogen (δ⁺) → Bent shape makes it polar.
• NH₃ (Ammonia): Nitrogen (δ⁻), Hydrogen (δ⁺) → Trigonal pyramidal shape.
Polar molecules dissolve in polar substances (e.g., water).
What is the stereochemical formula
The stereochemical formula is a 3D representation of a molecule that shows the spatial arrangement of atoms, including bond angles and molecular shape.
Key Features:
• Solid wedge (▲): Bond coming out of the plane (toward the viewer).
• Dashed wedge (▼): Bond going behind the plane (away from the viewer).
• Straight line (—): Bond in the plane of the paper.
Used to show molecular geometry (e.g., tetrahedral, trigonal pyramidal).
