Concise Chemistry Summary

History of the Atom

• 1803: John Dalton introduced atomic theory, stating that atoms are indivisible spheres and that all atoms of the same element are identical in mass and properties. Dalton's theory marked the beginning of modern chemistry.

• 1897: J.J. Thomson discovered the electron through cathode ray experiments, identifying it as a negatively charged subatomic particle.

• 1904: Plum pudding model: Thomson proposed that an atom is a sphere of positive charge with electrons embedded throughout, like plums in a pudding.

• 1911: Ernest Rutherford conducted the gold foil experiment, leading to the nuclear model. He concluded that an atom has a small, dense, positive nucleus surrounded by mostly empty space with orbiting electrons. Most of the atom's mass is concentrated in the nucleus.

• 1913: Niels Bohr refined Rutherford's model, suggesting that electrons orbit the nucleus in specific, quantized energy levels or shells. Electrons can jump between these energy levels by absorbing or emitting energy in the form of coloured light.

• 1932: James Chadwick discovered the neutron in the nucleus, explaining the previously unaccounted-for mass of atoms.

• Modern understanding: The current atomic model describes electrons as existing in probability clouds (orbitals) rather than fixed paths, reflecting the uncertainty in their exact location and momentum.

Elements, Compounds, and Mixtures

• Classifying matter: Matter is classified based on its uniformity and the ability to separate it by physical or chemical means.

• Pure substance: Substances with definite and constant chemical and physical properties.

  • Element: A substance that cannot be broken down into simpler substances through physical or chemical means; consists of atoms with the same number of protons.

  • Compound: A substance consisting of two or more elements chemically bonded in a fixed ratio, represented by a chemical formula; can only be separated through chemical reactions.

• Mixture: A combination of substances that are physically combined but not chemically bonded; components retain their original properties and can be separated through physical processes.

  • Homogeneous: A mixture that is uniform throughout, with evenly distributed components.

  • Heterogeneous: A mixture with non-uniform composition, where components are not evenly distributed.

Atomic Structure and Isotopes

• Sub-atomic particles:

  • Protons (positive): Located in the nucleus, contribute to the atom's mass and define the element.

  • Neutrons (neutral): Located in the nucleus, contribute to the atom's mass and provide nuclear stability.

  • Electrons (negative): Orbit the nucleus in specific energy levels or orbitals; their interactions determine chemical properties. Protons and neutrons are approximately 2000 times more massive than electrons.

• Atomic number: The number of protons in the nucleus of an atom, defining the element's identity.

• Mass number: The total number of protons and neutrons in an atom's nucleus.

• Neutral atoms: Atoms with an equal number of protons and electrons, resulting in no net charge.

• Isotopes: Atoms of the same element (same number of protons) with different numbers of neutrons; they have the same chemical properties but different physical properties, such as mass.

Electronic Configuration

• Bohr Model rules:

  1. Each electron shell has a maximum capacity for electrons (e.g., the first shell holds up to 2 electrons, the second shell up to 8).

  2. Lower energy shells are filled before higher energy shells (for elements 1-18), following the Aufbau principle.

• Valency: The number of electrons an atom can gain, lose, or share to achieve a full outer electron shell, determining its charge and bonding behavior.

Periodic Table and Group Properties

• Elements ordered by atomic number (number of protons), arranged in increasing order.

• Periods: Horizontal rows in the periodic table; elements in the same period have the same number of electron shells.

• Groups: Vertical columns in the periodic table; elements in the same group have the same number of valence electrons, leading to similar chemical properties.

• Metals: Typically found on the left side of the periodic table; characterized by their lustrous appearance, good electrical and thermal conductivity, high density, malleability, ductility, and solid state at room temperature (except mercury).

• Nonmetals: Located on the right side of the periodic table; often dull in appearance, poor conductors of electricity and heat, brittle, and can exist as solids, liquids, or gases at room temperature.

• Group properties:

  • Group 1 (Alkali Metals): Highly reactive metals with 1 valence electron; they react violently with water to form alkaline solutions.

  • Group 2 (Alkaline Earth Metals): Reactive metals with 2 valence electrons; less reactive than Group 1 metals but still readily form compounds.

  • Groups 3-12 (Transition Metals): Metals with variable valencies, forming colored ions; known for their hardness, high density, and high melting/boiling points; used in many industrial applications and as catalysts.

  • Group 17 (Halogens): Reactive non-metals with 7 valence electrons; they readily form salts when they react with metals.

  • Group 18 (Noble Gases): Inert gases with 8 valence electrons (except helium, which has 2); they are extremely stable and unreactive due to their full outer electron shells.

• Metalloids: Elements with properties intermediate between metals and nonmetals; their conductivity can be modified, making them useful as semiconductors in electronic devices.

Ions and Ionic Bonding

• Ion: An atom or group of atoms that carries an electrical charge, either positive or negative.

• Cations: Positive ions formed when an atom loses one or more electrons.

• Anions: Negative ions formed when an atom gains one or more electrons.

• Metals lose electrons to achieve a stable electron configuration, forming cations.

• Non-metals gain electrons to achieve a stable electron configuration, forming anions.

• Naming ions:

  • Cations: Named after the metal from which they are formed (e.g., Aluminum cation, Iron(II) cation).

  • Anions: Named by changing the suffix of the element to '-ide' (e.g., Chloride anion, Oxide anion).

• Ionic compounds: Compounds formed through the transfer of electrons from a metal to a nonmetal, resulting in the formation of ions that are held together by strong electrostatic forces in a lattice structure.

• Octet Rule: Atoms tend to gain, lose, or share electrons in order to achieve a full outer electron shell with 8 electrons (except for elements like hydrogen and lithium, which aim for 2 electrons).

Rules for Writing Chemical Formulas

• Write the metal element (cation) first, followed by the non-metal element (anion) in the chemical formula.

• If the two ions have the same charge magnitude but opposite signs, they combine in a one-to-one ratio to form a neutral compound.

• If the two ions have different charge magnitudes, swap the numbers (charges) and use them as subscripts for the respective ions in the formula to balance the charges.

• Name cations (metals) before anions when naming ionic compounds. If the metal is a transition metal with multiple possible valencies, indicate the valency using Roman numerals in parentheses after the metal's name.