Classification of Matter

Elements, Atoms, Compounds, and Mixtures

Understanding the basic types of matter is fundamental in chemistry. Matter can be classified based on its composition and properties.

• Element: A pure substance consisting of only one type of atom. Examples: O2 (oxygen gas), Fe (iron).

• Atom: The smallest unit of an element that retains its chemical properties. Example: A single hydrogen atom (H).

• Compound: A substance formed when two or more elements are chemically bonded in fixed proportions. Example: H2O (water).

• Mixture: A physical combination of two or more substances where each retains its own properties. Mixtures can be homogeneous (uniform, e.g., saltwater) or heterogeneous (non-uniform, e.g., salad).

Example: Air is a mixture, water is a compound, and gold is an element.

Structure of the Atom

Subatomic Particles and Their Charges

Atoms are composed of three main subatomic particles:

• Proton: Positively charged particle (+1), located in the nucleus.

• Neutron: Neutral particle (0 charge), located in the nucleus.

• Electron: Negatively charged particle (–1), found in orbitals around the nucleus.

Example: A carbon atom has 6 protons, 6 neutrons, and 6 electrons.

Physical and Chemical Properties and Changes

Distinguishing Physical and Chemical Properties/Changes

• Physical Property: Can be observed without changing the substance's identity (e.g., melting point, color).

• Chemical Property: Describes a substance's ability to undergo chemical changes (e.g., flammability, reactivity).

• Physical Change: Alters appearance but not composition (e.g., melting ice).

• Chemical Change: Alters the chemical composition (e.g., rusting iron).

Example: Boiling water is a physical change; burning wood is a chemical change.

Phase Transitions

Six Phase Changes and Their Linked Phases

• Melting: Solid to liquid

• Freezing: Liquid to solid

• Vaporization: Liquid to gas

• Condensation: Gas to liquid

• Sublimation: Solid to gas

• Deposition: Gas to solid

Example: Dry ice (solid CO2) sublimes directly to gas.

Measurement and Significant Figures

Significant Digits and Scientific Notation

• Significant Figures: Digits in a measurement that are known with certainty plus one estimated digit.

• Rules for Zeros: Leading zeros are not significant; captive zeros are significant; trailing zeros are significant only if there is a decimal point.

• Scientific Notation: Expresses numbers as a product of a coefficient (1 ≤ coefficient < 10) and a power of ten.

Example: 0.00450 has three significant figures; written in scientific notation:

Calculations with Significant Figures

• Addition/Subtraction: Result has the same number of decimal places as the measurement with the fewest decimal places.

• Multiplication/Division: Result has the same number of significant figures as the measurement with the fewest significant figures.

Example: (rounded to two significant figures)

Temperature and SI Units

Temperature Conversions

• Celsius to Kelvin:

• Kelvin to Celsius:

Example: 25°C = K

SI Measurement Concepts

Atoms, Isotopes, and Ions

Distinguishing Atoms, Isotopes, and Ions

• Atom: Neutral, same number of protons and electrons.

• Isotope: Atoms of the same element with different numbers of neutrons.

• Ion: Atom or molecule with a net charge due to loss or gain of electrons.

Example: Na is an atom, Na+ is an ion, Carbon-14 is an isotope of carbon.

Subatomic Particles in Atoms, Isotopes, and Ions

• Protons (p): Equal to atomic number (Z).

• Neutrons (n): (mass number minus atomic number).

• Electrons (e): For neutral atoms, ; for ions, adjust for charge.

Example: For Cl– (chloride ion), Z = 17, A = 35, electrons = 18.

Isotope Notation

• Written as where X is the element symbol, A is mass number, Z is atomic number.

Example: is carbon-14.

The Periodic Table

Reading the Periodic Table

• Atomic Number (Z): Number of protons.

• Mass Number (A): Number of protons plus neutrons.

• Element Symbol: One- or two-letter abbreviation.

Example: For sodium (Na), Z = 11, A ≈ 23.

Metals and Non-metals

• Metals: Left and center of the periodic table; good conductors, malleable, ductile.

• Non-metals: Right side; poor conductors, brittle in solid form.

Groups on the Periodic Table

Electron Configuration

Shells, Subshells, and Orbitals

• Shell: Principal energy level (n = 1, 2, 3, ...).

• Subshell: s, p, d, f (types of orbitals within a shell).

• Maximum electrons:

  • Shell 1: 2 electrons (1s)

  • Shell 2: 8 electrons (2s, 2p)

  • Shell 3: 18 electrons (3s, 3p, 3d)

Writing Electron Configurations (H–Ne)

• Electrons fill orbitals in order of increasing energy (Aufbau principle).

• Example: Oxygen (O, Z=8):

Ions and Ionic Charges

Formation of Ions

• Atoms gain or lose electrons to achieve a stable electron configuration (often a noble gas configuration).

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

Example: Na loses one electron to form Na+; Cl gains one electron to form Cl–.

Predicting Ionic Charges

• Group 1: +1 charge

• Group 2: +2 charge

• Group 16: –2 charge

• Group 17: –1 charge

• Group 18: 0 charge (noble gases, generally do not form ions)

Example: Magnesium (Group 2) forms Mg2+; Oxygen (Group 16) forms O2–.

Additional info: These notes synthesize and expand upon the exam study guide topics, providing definitions, examples, and key tables for foundational chemistry concepts.