Chemistry Basics: Matter, Measurement, and Atomic Structure

Chemistry Basics - Matter and Measurement

Introduction to Chemistry

Chemistry is the scientific study of matter, its properties, composition, and the changes it undergoes. Matter is anything that occupies space and has mass.

• Matter: Anything that takes up space and has mass.

• Chemistry: The study of matter and its transformations.

Classifying Matter

Matter can be classified into pure substances and mixtures. Pure substances include elements and compounds, while mixtures can be homogeneous or heterogeneous.

• Pure Substance: Made up of only one type of substance.

• Element: Contains only one type of atom.

• Compound: Contains two or more elements chemically combined in a fixed ratio.

• Mixture: Combination of two or more pure substances.

• Homogeneous Mixture: Uniform composition throughout.

• Heterogeneous Mixture: Composition varies throughout the sample.

Examples of Matter Classification

Examples help illustrate the classification of matter:

• Blood: Mixture (heterogeneous)

• Diamond: Element (pure carbon)

• Sugar: Compound (C12H22O11)

Visual Classification of Matter

The classification of matter can be visualized as a flowchart, showing the relationships between elements, compounds, and mixtures.

Atoms and Atomic Structure

Atomic Structure

An atom is the smallest unit of matter that retains the properties of an element. Atoms are composed of three basic subatomic particles: protons, neutrons, and electrons.

• Proton: Positive charge (+1), mass = 1 amu, located in the nucleus.

• Neutron: No charge (0), mass = 1 amu, located in the nucleus.

• Electron: Negative charge (-1), negligible mass, located outside the nucleus.

Atoms are electrically neutral, meaning the number of protons equals the number of electrons.

Atomic Number and Isotopes

The atomic number is the number of protons in the nucleus and identifies the element. Atoms of the same element can have different numbers of neutrons, resulting in isotopes.

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

• Mass Number: Sum of protons and neutrons in an atom.

• Atomic Mass: Average mass of an element, accounting for the relative abundance of its isotopes.

The Periodic Table

The Periodic Table organizes elements by increasing atomic number and groups elements with similar properties together. Each element is represented by a unique symbol.

• Element Symbol: One or two letters representing an element.

• Groups: Vertical columns, elements with similar properties.

• Periods: Horizontal rows, elements with increasing atomic number.

Compounds and Chemical Formulas

Compounds

Compounds are substances formed when two or more elements combine in fixed ratios. The chemical formula identifies the type and number of atoms in a compound.

• Chemical Formula: Shows the elements and their ratios in a compound (e.g., H2O).

• Example: Water (H2O) consists of 2 hydrogen atoms and 1 oxygen atom.

Measurement and Units

Scientific Notation and Significant Figures

Scientific notation is used to express very large or small numbers. Significant figures reflect the precision of a measurement.

• Scientific Notation: $C \times 10^i$ where C is a number between 1 and 9, and i is the exponent.

• Significant Figures: Digits with known certainty plus one estimated digit.

Metric System and Prefixes

The metric system uses base units and prefixes to indicate multiples or fractions of units.

• Base Units: Mass (g), Volume (L), Length (m).

• Prefixes: kilo (k, 103), centi (c, 10-2), milli (m, 10-3), micro (μ, 10-6), nano (n, 10-9).

Unit Conversion

Unit conversion uses conversion factors to relate different units. Dimensional analysis is a systematic method for converting units.

• Conversion Factor: Equality between two units (e.g., 1 in. = 2.54 cm).

• Dimensional Analysis: Multiply by conversion factors to reach the desired unit.

The Mole and Molar Mass

The Mole

The mole is a counting unit in chemistry, representing 6.02 × 1023 particles (Avogadro's number).

• 1 mole = 6.02 × 1023 particles

• Molar Mass: Mass (g) of 1 mole of a substance, determined from the periodic table.

Electron Arrangement and Bonding

Electron Shells and Valence Electrons

Electrons are arranged in shells around the nucleus. The outermost electrons are called valence electrons and are involved in chemical bonding.

• Shell 1: Holds max 2 electrons

• Shell 2: Holds max 8 electrons

• Shell 3: Holds max 18 electrons

• Valence Electrons: Given by the group number for main-group elements

Lewis Dot Diagrams and the Octet Rule

Lewis dot diagrams represent valence electrons as dots around the element symbol. Most atoms strive to achieve 8 valence electrons (octet rule).

• Octet Rule: Atoms strive to acquire 8 valence electrons.

• Exceptions: H (2 electrons), B (6 electrons).

Covalent Bonds and Molecules

Covalent bonds are formed when atoms share pairs of electrons. Molecules are neutral entities with covalent bonds and fixed composition.

• Covalent Bond: Shared pair of electrons between atoms.

• Molecule: Neutral group of atoms held together by covalent bonds.

Molecular Geometry and Polarity

VSEPR Theory and Molecular Shape

The shape of a molecule is determined by the repulsion of valence electron pairs (VSEPR theory). Molecular geometry affects physical and chemical properties.

• Linear: 180° bond angle (e.g., CO2)

• Bent: <109.5° bond angle (e.g., H2O)

• Tetrahedral: 109.5° bond angle (e.g., CH4)

Bond Polarity and Molecule Polarity

Bond polarity arises from differences in electronegativity. Molecule polarity depends on both bond polarity and molecular geometry.

• Nonpolar Covalent: Equal sharing of electrons.

• Polar Covalent: Unequal sharing, resulting in dipoles.

• Polar Molecule: Dipoles do not cancel.

• Nonpolar Molecule: Dipoles cancel.

States of Matter and Intermolecular Forces

Intermolecular Forces

Intermolecular forces are attractions between molecules, affecting physical properties like melting and boiling points.

• London Forces: Temporary induced dipoles, weakest force, present in all molecules.

• Dipole-Dipole Attractions: Permanent dipoles, stronger than London forces, present in polar molecules.

• Hydrogen Bonding: Strongest dipole-dipole force, occurs when H is bonded to N, O, or F.

States of Matter

Matter exists in three main states: solid, liquid, and gas. Intermolecular forces keep molecules together in the liquid and solid phases.

• Solid: Fixed shape and volume.

• Liquid: Fixed volume, variable shape.

• Gas: Variable shape and volume, highly compressible.

Changes of State

Phase changes occur when enough energy is supplied to overcome intermolecular forces. Covalent bonds are not broken during phase changes.

• Melting Point: Temperature at which a solid becomes a liquid.

• Boiling Point: Temperature at which a liquid becomes a gas.

Behavior of Gases and Boyle's Law

Gas Properties

Gases are highly compressible and fill their containers. Gas pressure results from collisions between molecules and their surroundings.

• Pressure Units: 1 atm = 760 mm Hg = 760 torr = 29.92 in. Hg = 101,325 Pa.

Boyle's Law

Boyle's Law states that at constant temperature and fixed moles, pressure and volume are inversely proportional.

• Equation: $P_i \times V_i = P_f \times V_f$

• Increasing pressure decreases volume, and vice versa.

Summary Table: Classification of Matter

Type	Definition	Example
Element	Pure substance, one type of atom	Copper, Diamond
Compound	Pure substance, two or more elements in fixed ratio	Water, Sugar
Homogeneous Mixture	Uniform composition	Brass (copper and zinc)
Heterogeneous Mixture	Variable composition	Blood, Water and Copper

Additional info: Academic context and examples were expanded for clarity and completeness. Images were included only when directly relevant to the explanation of the paragraph.