Comprehensive Study Guide for Matter, Atomic Structure, and Chemical Bonding

CHEMISTRY: THE STUDY OF MATTER AND ITS PROCESSES

• Definition of Chemistry: The study of the composition, structure, and properties of matter, the processes that matter undergoes, and the energy changes that accompany these processes.

• Scientific Branches:

  • Biological Sciences: Focus on living things.

  • Physical Sciences: Focus on nonliving things.

  • Centrality of Chemistry: Chemistry bridges these categories because all matter, living or nonliving, has a chemical structure.

• Instruments of Observation:

  • Scanning Tunneling Microscope (STM): Beams electrons at materials to reveal the pattern of their microstructure (too small for the unaided eye).

  • X-rays: Patterns created by X-rays reveal the arrangement of atoms and molecules.

• Chemicals: Any substance that has a definite composition (e.g., sucrose, carbon dioxide, water).

BRANCHES OF CHEMISTRY

1. Organic Chemistry: The study of most carbon-containing compounds.

2. Inorganic Chemistry: The study of non-organic substances, including organometallics (organic fragments bonded to metals).

3. Physical Chemistry: The study of the properties and changes of matter and their relation to energy.

4. Analytical Chemistry: The identification of the components and composition of materials.

5. Biochemistry: The study of substances and processes occurring in living things.

6. Theoretical Chemistry: Use of mathematics and computers to understand principles behind chemical behavior and predict properties of new compounds.

RESEARCH AND TECHNOLOGICAL DEVELOPMENT

• Basic Research: Carried out to increase knowledge (how/why reactions occur). Often leads to chance discoveries, such as Roy Plunkett's discovery of Teflon™.

• Applied Research: Driven by a desire to solve a specific problem, such as developing new refrigerants to protect the ozone layer.

• Technological Development: Involves the production and use of products that improve quality of life (e.g., computers, biodegradable materials, fiber optics).

• Overlap: Basic research on crystals/light led to lasers, which led to the applied/technological development of fiber optics.

MATTER AND ITS PROPERTIES

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

• Mass: A measure of the amount of matter (measured with a balance).

• Volume: The amount of three-dimensional space an object occupies.

• Building Blocks of Matter:

  • Atom: The smallest unit of an element that maintains the chemical identity of that element.

  • Element: A pure substance that cannot be broken down into simpler, stable substances and is made of one type of atom.

  • Compound: A substance that can be broken down into simple stable substances; made from atoms of two or more elements chemically bonded.

  • Molecule: The smallest unit of an element or compound that retains all properties of that substance.

CLASSIFICATION OF PROPERTIES

• Extensive Properties: Depend on the amount of matter present (e.g., volume, mass, amount of energy).

• Intensive Properties: Independent of the amount of matter present (e.g., melting point, boiling point, density, conductivity).

• Physical Properties: Characteristics observed/measured without changing identity (e.g., melting point of water at $0^{\circ}C$).

• Chemical Properties: Relate to a substance's ability to undergo changes that transform it into different substances (e.g., charcoal burning in air to form $CO_2$).

STATES OF MATTER AND CHANGES

• Physical Change: Does not involve a change in identity (e.g., grinding, melting).

• Change of State: Physical change from one state to another.

• Solid: Definite volume and shape; particles are packed in relatively fixed positions with strong attractive forces.

• Liquid: Definite volume but indefinite shape; particles are close together but can flow past one another.

• Gas: Neither definite volume nor shape; particles move rapidly and are at great distances.

• Plasma: High-temperature state where atoms lose most of their electrons.

• Chemical Change (Chemical Reaction): One or more substances converted into different substances.

  • Reactants: Substances that react.

  • Products: Substances formed.

• Laws of Conservation:

  • Law of Conservation of Mass: Total amount of matter remains the same before and after a reaction.

  • Law of Conservation of Energy: Energy can be absorbed or released but is neither created nor destroyed.

CLASSIFICATION OF MATTER

• Mixtures: A blend of two or more kinds of matter, each retaining its own identity/properties. Can be separated physically.

  • Homogeneous (Solutions): Uniform in composition (e.g., salt-water, air).

  • Heterogeneous: Not uniform throughout (e.g., granite, wood, blood).

  • Separation Methods: Filtration, decanting, centrifuge, paper chromatography.

• Pure Substances: Fixed composition and always homogeneous. Every sample has identical properties.

  • Compounds: Can be decomposed by chemical means (e.g., electrolysis of water, heating sucrose).

  • Elements: Cannot be decomposed.

• Purity: Laboratory chemicals have varying grades (Primary standard reagent is the highest purity; Technical grade is for industrial use).

THE PERIODIC TABLE

• Groups (Families): Vertical columns (1-18) containing elements with similar chemical properties.

• Periods: Horizontal rows where physical/chemical properties change regularly across the row.

• Categories of Elements:

  • Metals: Located at the left/center. Characterized by metallic luster, high electrical and heat conductivity, malleability (hammered into sheets), ductility (drawn into wire), and high tensile strength. Most are solids (exception: Mercury is liquid).

  • Nonmetals: Poor conductors. Many are gases (Nitrogen, Oxygen). Bromine is liquid. Solids (Carbon, Sulfur) are brittle.

  • Metalloids: Found between metals and nonmetals. Semiconductors; all are solids at room temperature.

  • Noble Gases: Group 18. Generally unreactive and gaseous at room temperature.

MEASUREMENTS AND CALCULATIONS (SI SYSTEM)

• Scientific Method: Logical approach to problem-solving: Observing/Collecting Data, Formulating Hypotheses, Testing, and Theorizing.

• Quantitative vs. Qualitative Data:

  • Quantitative: Numerical (e.g., 25.7 grams).

  • Qualitative: Descriptive (e.g., blue sky).

• System: Specific portion of matter selected for study.

• SI Base Units:

  • Length: Meter ($m$)

  • Mass: Kilogram ($kg$)

  • Time: Second ($s$)

  • Temperature: Kelvin ($K$)

  • Amount: Mole ($mol$)

  • Current: Ampere ($A$)

  • Luminous Intensity: Candela ($cd$)

• Weight vs. Mass: Mass is amount of matter; Weight is gravitational pull ($Weight = Mass \times Gravity$).

• SI Prefixes:

  • Kilo- ($k$): $10^{3}$

  • Centi- ($c$): $10^{-2}$

  • Milli- ($m$): $10^{-3}$

  • Micro- ($\mu$): $10^{-6}$

  • Nano- ($n$): $10^{-9}$

• Derived Units:

  • Area: $m^{2}$

  • Volume: $m^{3}$ ($1 L = 1000 mL = 1000 cm^{3}$)

  • Density: $D = \frac{m}{V}$. Characteristic physical property that decreases as temperature increases for most substances.

SCIENTIFIC MEASUREMENT ACCURACY

• Accuracy: Closeness of measurements to the correct/accepted value.

• Precision: Closeness of a set of measurements of the same quantity made in the same way.

• Percentage Error:
  $\% Error = \frac{Value<em>{experimental} - Value</em>{accepted}}{Value_{accepted}} \times 100$

• Significant Figures: All digits known with certainty plus one final estimated digit.

  • Rules for Zeros:

    1. Zeros between non-zero digits are significant (e.g., 40.7 L has 3 sig figs).

    2. Zeros in front of non-zero digits are NOT significant (e.g., 0.0009 kg has 1 sig fig).

    3. Zeros at the end of a number and to the right of a decimal are significant (e.g., 85.00 g has 4 sig figs).

    4. Zeros at the end of a number but to the left of a decimal point are only significant if a decimal point is present (e.g., 2000. m has 4 sig figs; 2000 m has 1 sig fig).

• Arithmetic Sig Figs:

  • 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 sig figs as the measurement with the fewest sig figs.

• Scientific Notation: $M \times 10^{n}$. Only sig figs are shown in the factor $M$.

PROPORTIONALITY

• Direct Proportion: Ratio of two variables is constant ($\frac{y}{x} = k$). Produces a linear graph passing through the origin.

• Inverse Proportion: Product of two variables is constant ($xy = k$). Produces a hyperbola.

ATOMIC THEORY HISTORY

• Early Ideas: Democritus (400 BCE) proposed the "atomos" (indivisible). Aristotle disagreed, favoring continuous matter.

• Law of Conservation of Mass: Mass is neither created nor destroyed.

• Law of Definite Proportions: A chemical compound contains the same elements in exactly the same proportions by mass regardless of sample size.

• Law of Multiple Proportions: If two elements form more than one compound, the ratios of the masses of the second element that combine with a fixed mass of the first element are ratios of small whole numbers.

• Dalton’s Atomic Theory (1808):

  1. All matter is composed of atoms.

  2. Atoms of an element are identical; atoms of different elements differ.

  3. Atoms cannot be subdivided, created, or destroyed (Modified: atoms are divisible into subatomic particles).

  4. Atoms combine in simple whole-number ratios to form compounds.

  5. In reactions, atoms are combined, separated, or rearranged.

ATOMIC STRUCTURE

• Discovery of Electron: J.J. Thomson (1897) used cathode-ray tubes to discover negatively charged particles (electrons) and their large charge-to-mass ratio. Robert Millikan (1909) measured the electron's charge.

• Plum Pudding Model: Thomson’s early model showing electrons spread through positive charge.

• Discovery of Nucleus: Ernest Rutherford (1911) gold-foil experiment. Alpha particles were deflected at wide angles, proving a small, dense, positively charged nucleus exists.

• Subatomic Particles:

  • Protons ($p^{+}$): Positive charge ($+1$). Mass ≈ $1.673 \times 10^{-27} kg$.

  • Neutrons ($n^{\circ}$): Neutral charge. Mass ≈ $1.675 \times 10^{-27} kg$.

  • Electrons ($e^{-}$): Negative charge ($-1$). Mass ≈ $9.109 \times 10^{-31} kg$.

• Nuclear Forces: Short-range forces (p-n, p-p, n-n) that hold nuclear particles together.

• Atomic Measurements: Atomic radii are expressed in picometers ($pm$). $1 pm = 10^{-12} m$.

QUANTIFYING ATOMS

• Atomic Number ($Z$): Number of protons in the nucleus. Identifies the element.

• Mass Number: Total number of protons and neutrons ($Mass Number = p^{+} + n^{\circ}$).

• Isotopes: Atoms of the same element with different masses (different number of neutrons). Examples: Protium ($1 p^{+}$, $0 n^{\circ}$), Deuterium ($1 p^{+}$, $1 n^{\circ}$), Tritium ($1 p^{+}$, $2 n^{\circ}$).

• Nuclide: General term for a specific isotope.

• Relative Atomic Mass: Standardized against Carbon-12 ($12 u$). 1 unified atomic mass unit ($1 u$) is $1/12$ the mass of a Carbon-12 atom.

• Average Atomic Mass: Weighted average of the atomic masses of naturally occurring isotopes.

• Mole ($mol$): SI unit for amount of substance; contains Avogadro’s number of particles.

• Avogadro’s Number: $6.02214179 \times 10^{23}$.

• Molar Mass: Mass of one mole of a pure substance ($g/mol$).

ELECTRON ARRANGEMENT AND QUANTUM MODEL

• Electromagnetic Radiation: Light exhibits wave behavior. $c = \lambda \nu$ (Speed of light $c = 3.00 \times 10^{8} m/s$).

• Photoelectric Effect: Emission of electrons from metal when light of specific frequencies shines on it. Proves light acts as particles (Photons).

• Planck’s Equation: $E = h\nu$. A quantum is the minimum energy lost/gained by an atom. Planck’s constant $h = 6.626 \times 10^{-34} J \cdot s$.

• Bohr Model (1913): Electrons circle the nucleus in fixed paths called orbits. When falling from an excited state to ground state, an atom emits a photon with energy equal to the difference in energy levels.

• Quantum Mechanics:

  • De Broglie: Electrons have a dual wave-particle nature.

  • Heisenberg Uncertainty Principle: Impossible to determine simultaneously both the position and velocity of an electron.

  • Schrödinger Wave Equation: Foundation of modern quantum theory; treats electrons as waves.

  • Orbital: 3D region around the nucleus indicating the probable location of an electron.

• Quantum Numbers:

  1. Principal ($n$): Main energy level/shell ($n=1, 2, \dots$). Total orbitals in a shell = $n^{2}$.

  2. Angular Momentum ($\ell$): Shape of orbital/sublevel ($s, p, d, f$). $\ell$ can be from $0$ to $n-1$.

  3. Magnetic ($m$): Orientation around nucleus (e.g., $p<em>{x}, p</em>{y}, p_{z}$).

  4. Spin: Indicates the two fundamental spin states ($+\frac{1}{2}, -\frac{1}{2}$).

ELECTRON CONFIGURATION RULES

• Aufbau Principle: Electron occupies the lowest-energy orbital available.

• Pauli Exclusion Principle: No two electrons in the same atom can have the same four quantum numbers; an orbital holds max 2 electrons with opposite spins.

• Hund’s Rule: Orbitals of equal energy are each occupied by one electron before any is occupied by a second, and all singly occupied orbitals must have the same spin.

• Notation Types:

  • Orbital Notation: Use of lines and arrows.

  • Electron-Configuration Notation: Use of superscripts (e.g., $1s^{2} 2s^{2} 2p^{1}$ for Boron).

  • Noble-Gas Notation: Shorthand using the previous noble gas in brackets.

PERIODIC TRENDS

• Atomic Radius: Decreases across a period (increasing nuclear charge pulls electrons closer); Increases down a group (more electron shells).

• Ionization Energy (IE): Energy to remove an electron. Increases across a period; Decreases down a group.

• Electron Affinity: Energy change when an atom acquires an electron. Generally becomes more negative across a period.

• Ionic Radius: Cations (positive) are smaller than neutral atoms; Anions (negative) are larger.

• Electronegativity: Measure of ability to attract electrons in a compound. Increases across a period; Decreases or stays same down a group. Fluorine is the most electronegative ($4.0$).

• Valence Electrons: Electrons in the outermost $s$ and $p$ sublevels. Group 1 has 1; Group 17 has 7.

CHEMICAL BONDING

• Chemical Bond: Mutual electrical attraction between nuclei and valence electrons.

• Ionic Bonding: Electrical attraction between cations and anions; involves large electronegativity differences (>1.7).

• Covalent Bonding: Sharing of electron pairs ($0-1.7$ difference).

  • Nonpolar-Covalent: Equal sharing ($0-0.3$).

  • Polar-Covalent: Unequal attraction ($0.3-1.7$).

• Octet Rule: Compounds form so atoms have 8 electrons in the highest occupied level (Exceptions: $H$, $B$, expanded valence like $PCl_{5}$).

• Lewis Structures: Atomic symbols represent nuclei; dots/dashes represent electron pairs.

• Bond Types:

  • Single Bond: 1 pair of shared electrons.

  • Double Bond: 2 pairs.

  • Triple Bond: 3 pairs (strongest and shortest).

• Resonance: Bonding in molecules that cannot be correctly represented by a single Lewis structure (e.g., Ozone $O_{3}$).

• Lattice Energy: Energy released when one mole of an ionic crystal is formed from gaseous ions. Indicates ionic bond strength.

• Metallic Bonding: Attraction between metal atoms and a "sea of electrons." Explains conductivity, malleability, and luster.

MOLECULAR GEOMETRY

• VSEPR Theory: Valence-shell, electron-pair repulsion. Electron sets orient as far apart as possible.

  • Linear: $AB_{2}$

  • Trigonal-Planar: $AB_{3}$

  • Tetrahedral: $AB_{4}$

  • Trigonal-Pyramidal: $AB_{3}E$ (one lone pair).

  • Bent: $AB<em>{2}E</em>{2}$ (two lone pairs).

• Hybridization: Mixing atomic orbitals (e.g., $sp, sp^{2}, sp^{3}$) to produce new hybrid orbitals for bonding.

• Intermolecular Forces: Weak forces between molecules.

  • Dipole-Dipole: Between polar molecules.

  • Hydrogen Bonding: Strong dipole-dipole force where $H$ is bonded to highly electronegative $F$, $O$, or $N$.

  • London Dispersion Forces: Occur in all atoms/molecules due to constant electron motion and instantaneous dipoles.

NOMENCLATURE AND FORMULAS

• Monatomic Ions: Named by element (cations) or root + -ide (anions).

• Stock System: Uses Roman numerals for elements with multiple oxidation states (e.g., Iron(II) chloride).

• Binary Compounds: Named using prefixes (mono-, di-, tri-, etc.) if molecular.

• Polyatomic Ions: Charged groups of covalently bonded atoms.

• Oxidation Numbers: Assigned to track electron distribution.

  • Pure Element: $0$

  • Fluorine: $-1$

  • Oxygen: $-2$ (usually)

  • Hydrogen: $+1$ (with nonmetals), $-1$ (with metals).

• Empirical Formula: Smallest whole-number mole ratio of elements.

• Molecular Formula: Actual formula determined using molar mass ($x = \frac{Molecular \, Molar \, Mass}{Empirical \, Formula \, Mass}$).

CHEMICAL EQUATIONS AND REACTIONS

• Indicators of Reaction:

  1. Evolution of heat and light.

  2. Production of a gas.

  3. Formation of a precipitate (solid out of solution).

  4. Color change.

• Balancing Equations: Satisfies the Law of Conservation of Mass; add coefficients only, never change subscripts.

• Types of Reactions:

  1. Synthesis: $A + X \rightarrow AX$

  2. Decomposition: $AX \rightarrow A + X$

  3. Single-Displacement: $A + BX \rightarrow AX + B$

  4. Double-Displacement: $AX + BY \rightarrow AY + BX$

  5. Combustion: Substance + $O_{2} \rightarrow energy$.

• Activity Series: List of elements organized by ease of reaction; used to predict if displacement reactions will occur.