Chemistry (US High School)
A comprehensive NGSS-aligned high school chemistry deck covering atomic structure, bonding, stoichiometry, reactions, gas laws, solutions, acids and bases, thermochemistry, kinetics, equilibrium, nuclear and organic chemistry.
Ämne: Kemi · Nivå: Gymnasium (16–19) · 500 kort
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- An atom consists of a dense central nucleus containing protons (positive charge) and neutrons (no charge), surrounded by electrons (negative charge) in regions called orbitals.
- The atomic number (Z) of an element equals the number of protons in the nucleus and uniquely identifies the element.
- The mass number (A) is the total number of protons and neutrons in a nucleus. Number of neutrons = A − Z.
- Isotopes are atoms of the same element (same Z) with different numbers of neutrons (different A). Example: ¹²C and ¹⁴C.
- The atomic mass on the periodic table is the weighted average of an element's naturally occurring isotopes, expressed in atomic mass units (amu or u).
- A proton's mass ≈ 1.007 u; a neutron's mass ≈ 1.009 u; an electron's mass ≈ 0.00055 u (about 1/1836 of a proton).
- An ion is an atom or molecule with a net electrical charge from gaining or losing electrons. Cations are positive (lost e⁻); anions are negative (gained e⁻).
- Electrons occupy quantized energy levels (n = 1, 2, 3, …) called shells. Each shell contains subshells (s, p, d, f) made of orbitals.
- Maximum electron capacities per subshell: s = 2, p = 6, d = 10, f = 14. Each orbital holds at most 2 electrons with opposite spins.
- The Aufbau principle states that electrons fill the lowest-energy orbitals first. Order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, …
- Hund's rule: within a subshell, electrons occupy each orbital singly with parallel spins before pairing up.
- The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers — each orbital holds at most 2 electrons with opposite spins.
- Valence electrons are the electrons in the outermost shell of an atom. They determine the element's chemical behavior and bonding.
- Niels Bohr (1913) proposed that electrons orbit the nucleus in fixed energy levels and emit/absorb light when transitioning between levels.
- Rutherford's gold foil experiment (1909) showed that atoms are mostly empty space with a small, dense, positively charged nucleus.
- J.J. Thomson discovered the electron in 1897 using cathode-ray tube experiments and proposed the 'plum-pudding' model of the atom.
- John Dalton (early 1800s) proposed the modern atomic theory: matter consists of atoms; atoms of one element are identical; compounds form from fixed ratios of atoms.
- Quantum mechanical model: electrons exist in orbitals — three-dimensional probability regions where an electron is likely to be found.
- The electron configuration of carbon (Z = 6) is 1s² 2s² 2p², giving 4 valence electrons.
- Noble gases (Group 18) have full valence shells (ns²np⁶ except He = 1s²), making them extremely unreactive.
- The periodic table organizes elements by increasing atomic number into 18 vertical groups (families) and 7 horizontal periods.
- Dmitri Mendeleev (1869) created the first widely accepted periodic table, ordered by atomic mass and leaving gaps for undiscovered elements.
- Group 1 (alkali metals — Li, Na, K, Rb, Cs, Fr) are soft, highly reactive metals with one valence electron and form +1 cations.
- Group 2 (alkaline earth metals — Be, Mg, Ca, Sr, Ba, Ra) have 2 valence electrons and form +2 cations.
- Group 17 (halogens — F, Cl, Br, I, At) have 7 valence electrons and readily gain one to form −1 anions; they are highly reactive nonmetals.
- Group 18 (noble gases — He, Ne, Ar, Kr, Xe, Rn) have full valence shells and are largely inert, rarely forming compounds.
- Transition metals occupy groups 3–12 of the periodic table. They have variable oxidation states and form colored compounds and complex ions.
- Lanthanides (Z = 57–71) and actinides (Z = 89–103) are the f-block elements placed below the main table for layout. All actinides are radioactive.
- Metalloids (B, Si, Ge, As, Sb, Te) along the 'stair-step' line have mixed metallic and nonmetallic properties; many are semiconductors.
- Atomic radius generally decreases across a period (increasing nuclear charge pulls electrons closer) and increases down a group (more occupied shells).
- Ionization energy is the energy required to remove the outermost electron from a gaseous atom. It generally increases across a period and decreases down a group.
- Electronegativity is an atom's tendency to attract bonding electrons. Fluorine is the most electronegative element (3.98 on the Pauling scale).
- Electron affinity is the energy change when a gaseous atom gains an electron. More negative values indicate a stronger tendency to accept electrons.
- Metallic character increases down a group and decreases across a period — the lower-left of the periodic table is most metallic, the upper-right is most nonmetallic.
- The octet rule states that atoms tend to bond so each has 8 valence electrons (or 2 for H, He), matching a noble gas configuration.
- An ionic bond forms when electrons transfer from a metal to a nonmetal, creating oppositely charged ions held together by electrostatic attraction.
- A covalent bond forms when two nonmetal atoms share one or more pairs of electrons. Single, double, and triple bonds share 2, 4, and 6 electrons respectively.
- A polar covalent bond results when atoms of different electronegativity share electrons unequally, creating partial charges (δ+, δ−).
- A nonpolar covalent bond forms between atoms of equal (or nearly equal) electronegativity, such as in H₂, O₂, or N₂.
- Metallic bonding: metal cations sit in a 'sea' of delocalized valence electrons, which gives metals their conductivity, malleability, and luster.
- Bond classification by electronegativity difference (ΔEN): nonpolar covalent (ΔEN < 0.5), polar covalent (0.5 ≤ ΔEN < 1.7), ionic (ΔEN ≥ 1.7).
- Ionic compounds form crystal lattices (e.g., NaCl), have high melting points, are brittle, and conduct electricity when molten or dissolved — but not as solids.
- Molecular (covalent) compounds typically have lower melting/boiling points, are softer, and most do not conduct electricity in any state.
- VSEPR theory (Valence Shell Electron Pair Repulsion) predicts molecular geometry by minimizing repulsion between bonding and lone electron pairs.
- Common molecular geometries: linear (180°), trigonal planar (120°), tetrahedral (109.5°), trigonal pyramidal, bent.
- Water (H₂O) has a bent geometry with a bond angle of ~104.5°. The two lone pairs on oxygen push the H atoms closer together than tetrahedral.
- Carbon dioxide (CO₂) is linear (O=C=O, 180°). It is nonpolar overall because two equal C=O dipoles cancel.
- Methane (CH₄) is tetrahedral with 109.5° H–C–H bond angles. It is nonpolar because all C–H dipoles are symmetric and cancel.
- Ammonia (NH₃) is trigonal pyramidal because the lone pair on nitrogen pushes the H atoms down, giving bond angles of ~107°.
- Intermolecular forces (IMFs), from weakest to strongest: London dispersion < dipole–dipole < hydrogen bonding < ion–dipole.