Classifying Matter

Pure Substances and Mixtures

Matter can be classified based on its composition and uniformity. Understanding these classifications is fundamental in chemistry.

• Pure Substance: A form of matter with a constant composition and distinct chemical properties. Examples include elements (e.g., O2, Fe) and compounds (e.g., H2O, NaCl).

• Mixture: A physical blend of two or more substances, each retaining its own properties. Mixtures can be separated by physical means.

• Homogeneous Mixture: Uniform composition throughout (e.g., saltwater).

• Heterogeneous Mixture: Non-uniform composition (e.g., salad, sand in water).

Example: Air is a homogeneous mixture; salad is a heterogeneous mixture.

Elements, Compounds, and the Periodic Table

The periodic table organizes elements by increasing atomic number and similar properties.

• Element: A pure substance consisting of only one type of atom.

• Compound: A pure substance composed of two or more elements chemically combined in fixed ratios.

• Groups: Vertical columns; elements in a group have similar chemical properties.

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

• Metals: Located on the left and center; typically shiny, malleable, and good conductors.

• Nonmetals: Located on the right; often gases or brittle solids, poor conductors.

Example: Sodium (Na) is a metal in Group 1; chlorine (Cl) is a nonmetal in Group 17.

How Matter Changes

Physical and Chemical Changes

Matter can undergo changes that are either physical or chemical in nature.

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

• Chemical Change (Reaction): Alters the composition, forming new substances (e.g., burning wood).

• States of Matter: Solid, liquid, gas.

• Chemical Equation: Represents a chemical reaction using symbols and formulas.

Example: (formation of water from hydrogen and oxygen).

Math in Chemistry

Units and Measurements

Accurate measurement and unit conversion are essential in chemistry.

• Base Units: Meter (distance), gram (mass), liter (volume), Kelvin/Celsius (temperature).

• Metric Prefixes: kilo- (103), centi- (10-2), milli- (10-3).

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

• Scientific Notation: Expresses numbers as .

Example: 0.00045 =

The "Stuff" of Chemistry

Mass, Volume, and Density

Understanding the properties of matter involves measuring mass, volume, and density.

• Mass: Amount of matter in an object (measured in grams).

• Volume: Space occupied by matter (measured in liters or milliliters).

• Density: Mass per unit volume.

• Specific Gravity: Ratio of the density of a substance to the density of water.

Example: If a block has a mass of 10 g and a volume of 2 mL, its density is .

Temperature Scales

• Celsius (°C), Kelvin (K), Fahrenheit (°F): Common temperature scales.

• Conversions:

Kinetic and Potential Energy

• Kinetic Energy: Energy of motion.

• Potential Energy: Stored energy due to position or composition.

Example: A rolling ball has kinetic energy; a stretched bow has potential energy.

States of Matter

• Solid: Definite shape and volume.

• Liquid: Definite volume, takes shape of container.

• Gas: No definite shape or volume.

Measuring Matter

Accuracy and Precision

• Accuracy: How close a measurement is to the true value.

• Precision: How close repeated measurements are to each other.

Unit Conversions

• Convert between SI (metric) and U.S. customary units using provided conversion factors.

• Apply conversion factors, drop units, and convert percentages in health-related measurements.

Atoms and Their Components

Subatomic Particles

• Proton: Positively charged particle in the nucleus.

• Neutron: Neutral particle in the nucleus.

• Electron: Negatively charged particle outside the nucleus.

Example: Carbon-12 has 6 protons, 6 neutrons, and 6 electrons.

Atomic Number and Mass Number

• Atomic Number (Z): Number of protons in the nucleus; defines the element.

• Mass Number (A): Total number of protons and neutrons.

Example: (where N = number of neutrons)

Symbolic Notation

• Written as , where X is the element symbol.

Example: for carbon-14.

Isotopes and Atomic Mass

Isotopes

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

• Mass Number: Number of protons + neutrons in a specific isotope.

• Atomic Mass: Weighted average mass of all isotopes of an element.

Example: and are isotopes of carbon.

Radioactivity and Radioisotopes

Radioactivity

• Radioactivity: The spontaneous emission of particles or energy from unstable atomic nuclei.

• Ionizing Radiation: Includes alpha, beta, and gamma radiation.

• Penetrating Power: Gamma > Beta > Alpha.

Example: Alpha particles can be stopped by paper; gamma rays require lead shielding.

Radiation Units and Half-Lives

• Activity Units: Curie (Ci), Becquerel (Bq).

• Half-Life: Time required for half the atoms in a sample to decay.

Half-life formula:

where = remaining amount, = initial amount, = elapsed time, = half-life.

Medical Applications

• Radioisotopes are used in diagnosis (e.g., PET scans) and treatment (e.g., cancer radiotherapy).

Electron Arrangements and the Octet Rule

Electron Shells and Valence Electrons

• Electrons are arranged in shells (energy levels) around the nucleus.

• Valence Electrons: Electrons in the outermost shell; determine chemical reactivity.

• Main-group elements in the first four periods follow predictable patterns for valence electrons.

The Octet Rule

• Atoms tend to gain, lose, or share electrons to achieve eight valence electrons (octet), leading to stability.

• Three ways to reach stability:

  • Give electrons (form cations)

  • Take electrons (form anions)

  • Share electrons (form covalent bonds)

Example: Sodium gives up one electron to achieve an octet; chlorine takes one electron.

Ion Formation and Ionic Compounds

Ion Formation

• Application of the octet rule to form ions.

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

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

• Predict ionic charge using the periodic table (e.g., Group 1: +1, Group 17: -1).

• Write ion symbols: e.g., , .

• Common polyatomic ions: phosphate (), carbonate (), hydroxide (), hydronium ().

Ionic Compounds

• Formed by the combination of cations and anions.

• Formula reflects the ratio of ions needed for charge neutrality.

• Transition metals may have variable charges, indicated by Roman numerals (e.g., Fe2+ is iron(II)).

Example:

Covalent Compounds and Bonding

Covalent Bond Formation

• Atoms share electrons to achieve an octet.

• Nonmetals typically form covalent bonds.

• Number of bonds relates to the number of valence electrons (e.g., oxygen forms two bonds).

• Lewis structures represent bonding and lone pairs.

Example: : oxygen shares electrons with two hydrogens.

Ionic vs. Covalent Compounds

• Ionic: Metal + nonmetal, transfer of electrons, high melting points, conduct electricity when dissolved.

• Covalent: Nonmetal + nonmetal, sharing of electrons, lower melting points, do not conduct electricity.

The Mole: Counting Atoms and Compounds

• Mole (mol): Amount of substance containing particles (Avogadro's number).

• Relates mass, number of particles, and volume (for gases).

Example: 1 mol of H2O contains molecules.

Shapes and Polarity of Covalent Compounds

Molecular Shape (VSEPR Theory)

• Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on electron pair repulsion.

• Common shapes: linear, bent, trigonal planar, tetrahedral.

• Wedge-and-dash notation shows 3D arrangement.

Example: Methane (CH4) is tetrahedral.

Electronegativity and Molecular Polarity

• Electronegativity: Tendency of an atom to attract electrons in a bond.

• Bond Polarity: Difference in electronegativity leads to polar bonds.

• Molecular Polarity: Determined by bond polarities and molecular shape.

Example: Water (H2O) is polar due to its bent shape and polar O-H bonds.

Summary Table: Classification of Matter