Models of Atomic Structure Quiz

Models of Atomic Structure

• Solid Sphere Model (Dalton, 1803): Atoms are solid, indivisible spheres. Atoms of the same element are identical, and different elements have different atoms. Atoms combine in fixed ratios to form compounds.

• Plum Pudding Model (Thomson, 1904): Atoms consist of a uniform positive sphere with negatively charged electrons embedded within it, like plums in a pudding. The positive and negative charges balance, making the atom neutral.

• Nuclear Model (Rutherford, 1911): The atom has a tiny, dense, positively charged nucleus at its center containing most of the atom's mass. Negatively charged electrons orbit the nucleus like planets orbiting the sun.

• Planetary Model (Bohr, 1913): Electrons orbit the nucleus in specific, fixed energy levels or shells. Electrons can change energy levels by absorbing or emitting energy.

• Quantum Mechanical Model (Schrödinger, 1926): Electrons are described as waves. They exist in orbitals, defined regions of space where there's a high probability of finding an electron. Electrons exist in specific energy levels. Electrons can only move between energy levels by absorbing or emitting energy in the form of photons.

Quantum Numbers

• Principal (n): Describes the energy level and size of the electron's orbital.

• Azimuthal/Angular Momentum (l): Describes the shape of the electron's orbital (s, p, d, f).

• Magnetic (ml): Describes the orientation of the electron's orbital in space.

• Spin (ms): Describes the intrinsic angular momentum (spin) of the electron (up or down).

Electron Configuration

• Atomic orbitals: Electrons fill lower energy levels before higher levels.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.

• Hund's Rule: Electrons fill orbitals in a subshell such that the total electron spin is maximized before pairing up. A different way of stating the same idea - each orbital in a subshell is singly occupied before any one orbital is doubly occupied.

Valence Electrons

• Valence electrons: Electrons in the outermost energy level of an atom. They are responsible for the chemical properties of an element. The number of valence electrons determines how atoms interact with each other.

Molecular Geometry

• Molecular geometry: The 3-D structure of atoms in a molecule. This is dependent on valence shell electron pair repulsion (VSEPR) theory. The valence shell electron pair repulsion (VSEPR) theory predicts that molecules will arrange themselves so that the electron clouds ( bonds and lone pairs of electrons) are as far from each other as possible. Common geometries include linear, trigonal planar, bent, tetrahedral, trigonal pyramidal, and octahedral. The shape of a molecule determines its polarity.

• Octet Rule: The tendency of atoms to prefer to have eight electrons in their valence shell. This rule helps predict the bonding patterns in many molecules.

Bond Types

• Ionic Bonds: Atoms transfer electrons to form ions (positive and negative). Oppositely charged ions are attracted to form an ionic bond.

• Covalent Bonds: Atoms share electrons to achieve a more stable electron configuration. Shared electrons form a covalent bond. The number of shared electron pairs determines the bond type (single, double, triple bond). Single bonds share one pair, double bonds share two pairs and triple bonds share three pairs.

• Polarity: A molecule is polar if the electrons are not shared equally between the atoms in a covalent bond; if there is an uneven distribution of charge.