Ghasham International Chemistry SAAT Study Guide

Introduction to Chemistry, Matter, Properties, and Changes

• Branches of Chemistry:     * Analytical Chemistry: Studies the types and composition of substances.     * Atomic Chemistry: Examines theories of matter composition, including bonding and orbital shapes.     * Physical Chemistry: Explores the behavior and changes of matter.     * Nuclear Chemistry: Studies theories of matter composition and nuclear processes.     * Organic Chemistry: Focuses on compounds containing carbon.     * Inorganic Chemistry: Studies compounds that do not contain carbon.     * Biochemistry: Investigates chemical reactions within living organisms.     * Environmental Chemistry: Examines the impact of materials, such as packaging, on the environment.     * Industrial Chemistry: Focuses on chemical processes in industry.

• Matter and its Measurement:     * Definition: Matter is anything that has mass and occupies space.     * Substitutes for Matter: Light, heat, and sound are not considered matter.     * Mass: A measure of the amount of matter only. It remains constant regardless of location.     * Weight: A measure of the amount of matter that also depends on the force of gravity.

• Properties of Matter:     * Intensive Properties: Properties that are independent of the amount of substance present. Examples include density, color, speed, boiling point, and melting point.     * Extensive Properties: Properties that depend on the amount of substance present. Examples include mass, volume, and length.     * Chemical Properties: The ability or inability of a substance to combine with or change into one or more other substances. Examples: iron forming rust in humid air, sodium being corrosive, or sugar decomposing into carbon and water.     * Physical Properties: Characteristics that can be observed or measured without changing the sample's composition. Examples: water being colorless, table salt being a solid crystalline substance, and malleability or ductility of metals.

• States of Matter:     * Solid: Definite shape and volume; particles are closely packed together.     * Liquid: Indefinite shape (takes the shape of the container) but definite volume.     * Gaseous: Indefinite shape and volume; particles are widely spaced and highly compressible.     * Plasma: The primary components of stars and galaxies; characterized by ionized gas.     * Unique Behavior of Water: Water increase in volume when transitioning from a liquid to a solid state.

• Phase Changes:     * Endothermic (Energy-Absorbing): Melting ($\text{solid} \rightarrow \text{liquid}$), Evaporation ($\text{liquid} \rightarrow \text{gas}$), and Sublimation ($\text{solid} \rightarrow \text{gas}$).     * Exothermic (Energy-Releasing): Condensation ($\text{gas} \rightarrow \text{liquid}$), Freezing ($\text{liquid} \rightarrow \text{solid}$), and Deposition ($\text{gas} \rightarrow \text{solid}$).

• Scientific Method:     * Steps: 1. Observation, 2. Asking Questions (Data Collection), 3. Hypothesis, 4. Experimentation, 5. Conclusion, 6. Publishing Results.     * Observation Types:         * Qualitative Data: Describes physical properties like color, smell, or taste.         * Quantitative Data: Uses numerical values to measure properties like mass, volume ($10\text{\ mL}$), pressure, or concentration ($1\text{\ Molar}$).     * Variables:         * Independent Variable: The factor planned to be changed (e.g., temperature in a dissolution experiment).         * Dependent Variable: The factor that changes in response to the independent variable (e.g., dissolution rate).     * Scientific Theory: An explanation of a natural phenomenon based on many observations and investigations over time.     * Scientific Law: A relationship in nature that is supported by many experiments (e.g., Law of Conservation of Mass).

• Ozone ($O_3$):     * Location: Found in the Stratosphere layer of the atmosphere.     * Function: Absorbs harmful Ultraviolet (UV) radiation.     * Formation: Three oxygen atoms ($3O$) form one ozone molecule ($O_3$). (e.g., 12 oxygen atoms form 4 ozone molecules; 18 atoms form 6 molecules).     * Depletion: Caused by Chlorofluorocarbons (CFCs), which consist of Chlorine ($Cl$), Fluorine ($F$), and Carbon ($C$).     * Measurement: G.M.B. Dobson measured ozone levels. The standard natural concentration is $300\text{\ DU}$ (Dobson Units).

Atomic Structure

• Fundamental Models and Discoveries:     * Democritus: First suggested the existence of atoms.     * Dalton: Proposed that matter is composed of tiny particles called atoms and that they are the basic unit of elements.     * Aristotle: Proposed there is no empty space inside the atom.     * J.J. Thomson: Discovered the electron using cathode rays (which have a negative charge).     * Ernest Rutherford: Discovered the nucleus (dense, positively charged center) using the gold foil experiment and alpha rays. He concluded most of the atom is empty space.     * Niels Bohr: Proposed that electrons move around the nucleus in quantized energy levels.

• Atomic Components:     * Atom: The smallest part of an element that retains the properties of that element.     * Nucleus: Contains Protons (positive) and Neutrons (neutral). Holds most of the atom's mass.     * Electrons: Negatively charged particles that revolve around the nucleus. Their mass is much smaller than protons or neutrons.     * Atomic Number: The number of protons in the nucleus (also equals the number of electrons in a neutral atom).     * Mass Number/Atomic Mass: The sum of protons and neutrons in the nucleus.         * $\text{Neutrons} = \text{Mass Number} - \text{Atomic Number}$     * Atomic Mass Unit (amu): Approximately equal to the mass of a single proton or neutron.

• Isotopes:     * Atoms of the same element that have the same number of protons (atomic number) but different numbers of neutrons (mass number).     * Average atomic mass is the weighted average mass of all isotopes of an element found in nature.

• Radioactivity and Nuclear Decay:     * Alpha Decay (̑): Emits an alpha particle ($^4_2He$). Decreases atomic number by 2 and mass number by 4.     * Beta Decay (̒): Emits a high-speed electron ($^0_{-1}e$). Increases atomic number by 1; mass number remains unchanged.     * Gamma Decay (̓): Emits high-energy photons. No change in atomic number or mass number.     * Nuclear Stability: Determined primarily by the ratio of neutrons to protons.

Chemical Reactions and Stoichiometry

• Reaction Types:     * Synthesis (Combination): Multiple reactants form a single product ($A + B \rightarrow AB$).     * Decomposition: A single reactant breaks down into multiple products ($AB \rightarrow A + B$).     * Single Displacement: One element replaces another in a compound ($A + BX \rightarrow AX + B$). Depends on the activity series (e.g., Bromine cannot displace Fluorine because Fluorine is more reactive).     * Double Displacement: Exchange of ions between two compounds ($AX + BY \rightarrow AY + BX$). Often occurs in aqueous solutions (e.g., Acid + Base reactions).     * Combustion: Reaction with Oxygen ($O_2$) resulting in heat and often products like $CO_2$ and $H_2O$.

• Chemical Equations:     * Must be balanced to obey the Law of Conservation of Mass.     * Net Ionic Equation: Shows only those particles that participate in the reaction (e.g., $H^+ + OH^- \rightarrow H_2O$).

• Stoichiometry and the Mole:     * Avogadro's Number: $6.02 \times 10^{23}$ particles per mole.     * Molar Mass: The sum of atomic masses of all atoms in a compound.     * Calculations:         * $\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar Mass (g/mol)}}$         * $\text{Mass (g)} = \text{Moles} \times \text{Molar Mass}$         * $\text{Mole Ratio} = n(n-1)$, where $n$ is the number of substances.

• Yields:     * Theoretical Yield: The maximum amount of product that can be produced from a given amount of reactant.     * Actual Yield: The amount of product produced when the chemical reaction is carried out in an experiment.     * Percent Yield: $\frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100$

• Limiting and Excess Reactants:     * Limiting Reactant: The substance that is totally consumed in a reaction and determines the amount of product.     * Excess Reactant: The substance that remains after the reaction stops.

• Formulas:     * Empirical Formula: Shows the smallest whole-number mole ratio of elements in a compound (e.g., Glucose empirical formula is $CH_2O$).     * Molecular Formula: Shows the actual number of atoms of each element in a molecule.

Electrons in Atoms

• Wave Properties:     * Frequency ($f$): Number of waves that pass a specific point in one second. Measured in Hertz ($Hz$).     * Wavelength ($λ$): Shortest distance between successive peaks or troughs.     * Photon: A massless particle that carries a quantum of energy.     * Energy Equation: $E = hf$ or $E = \frac{hC}{λ}$. Energy is directly proportional to frequency and inversely proportional to wavelength.

• Quantum Theory:     * Heisenberg Uncertainty Principle: Impossible to determine both the position and velocity of an electron simultaneously.     * Photoelectric Effect: Emission of electrons from a metal surface when light of a specific frequency shines on it.     * Quantum Numbers: Principal quantum numbers ($n$) take integer values ($1, 2, 3...$) and determine orbital energy and size.

• Atomic Orbitals and Transitions:     * Ground State: The lowest energy level of an atom.     * Excited State: State after an atom absorbs energy.     * Emission Series:         * Lyman Series: Electrons drop to $n=1$ (Ultraviolet emitted).         * Balmer Series: Electrons drop to $n=2$ (Visible light emitted).         * Paschen Series: Electrons drop to $n=3$ (Infrared emitted).

• Orbital Shapes and Capacity:     * s orbital: Spherical shape. Holds max 2 electrons.     * p orbital: Dumbbell-shaped ($p_x, p_y, p_z$). Holds max 6 electrons.     * d orbital: Multi-lobed. Holds max 10 electrons.     * f orbital: Complex shape. Holds max 14 electrons.     * Max Electrons per Shell: $2n^2$.

Electron Configuration and the Periodic Table

• Governing Principles:     * Aufbau Principle: Electrons occupy the lowest energy orbital first.     * Pauli Exclusion Principle: An orbital can hold a maximum of 2 electrons with opposite spins.     * Hund's Rule: Single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbitals.

• Standard and Ion Configurations:     * Chromium ($Cr=24$): Exception to stability; configuration is $[Ar] 4s^1 3d^5$.     * Copper ($Cu=29$): Exception; configuration is $[Ar] 4s^1 3d^{10}$.     * Ions: Positive ions (cations) lose electrons (e.g., $Na^+$ has 10 electrons). Negative ions (anions) gain electrons.

• Periodic Table Trends:     * Periods: Horizontal rows. Represent the highest principal energy level.     * Groups: Vertical columns. Elements in the same group have the same number of valence electrons and similar chemical properties.     * Atomic Radius: Decreases across a period (left to right) due to increased positive charge; increases down a group.     * Ionization Energy / Electronegativity / Electron Affinity: Increase across a period; decrease down a group.     * Noble Gases (Group 18): Chemically inert; have 8 valence electrons (except Helium - 2); reach the Octet Rule.

Chemical Bonds and Intermolecular Forces

• Bond Types:     * Ionic Bond: Electrostatic attraction between positive metal ions and negative non-metal ions. Electronegativity difference $> 1.7$.     * Covalent Bond: Sharing of electrons between non-metals.         * Polar Covalent: Unequal sharing (diff $0.4 - 1.7$). Examples: $H_2O, NH_3, HCl$.         * Non-polar Covalent: Equal sharing (diff $< 0.4$). Examples: $H_2, Cl_2, CH_4$.     * Metallic Bond: Attraction between positive metal ions and the "sea of free electrons."     * Coordinate Covalent: One atom donates both electrons to the bond (e.g., $BF_4^-$).

• Intermolecular Forces:     * Dispersion Forces (London Forces): Weak forces between non-polar molecules; strength increases with size (e.g., $I_2 > F_2$).     * Dipole-Dipole Forces: Attraction between polar molecules.     * Hydrogen Bonding: Strongest intermolecular force; occurs when Hydrogen is bonded to Oxygen, Nitrogen, or Fluorine.

• Molecular Geometry and Hybridization:     * sp: Linear (180^̲). Example: $BeCl_2, CO_2$.     * sp^2: Trigonal Planar (120^̲). Example: $BF_3, AlCl_3$.     * sp^3: Tetrahedral (109.5^̲). Example: $CH_4, CCl_4$.     * Bent: $sp^3$ with lone pairs (104.5^̲). Example: $H_2O$.

States of Matter and Gas Laws

• Kinetic Molecular Theory: Particles of matter are in constant motion. Temperature is a measure of average kinetic energy ($KE = \frac{1}{2} mv^2$).

• Liquid Properties:     * Viscosity: Resistance to flow. Decreases as temperature increases.     * Surface Tension: Energy required to increase the surface area. Allows spiders to walk on water.     * Capillary Action: Rise of liquids in thin tubes (e.g., water in plant roots).

• Gas Laws:     * Boyle’s Law: Pressure and Volume are inversely proportional ($P_1V_1 = P_2V_2$).     * Charles’s Law: Volume and Temperature are directly proportional (̑\frac{V_1}{T_1} = \frac{V_2}{T_2}̑).     * Gay-Lussac’s Law: Pressure and Temperature are directly proportional (̑\frac{P_1}{T_1} = \frac{P_2}{T_2}̑).     * Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$     * Ideal Gas Law: $PV = nRT$     * Dalton’s Law of Partial Pressures: Total pressure is the sum of partial pressures ($P_{total} = P_1 + P_2...$).     * Graham’s Law: Diffusion rate relates to molar mass. Smaller mass = faster diffusion.

• Phase Diagrams:     * Triple Point: Temperature and pressure where all three phases coexist.     * Critical Point: Point beyond which the substance cannot exist as a liquid regardless of pressure.

Energy and Chemical Changes

• Specific Heat: Amount of heat required to raise the temperature of $1\text{\ g}$ of substance by 1^̲C.     * Formula: $q = m \times c \times ΔT$.     * Water has high specific heat compared to metals like lead.

• Enthalpy ($ΔH$):     * Exothermic: $ΔH$ is negative (e.g., combustion, condensation).     * Endothermic: $ΔH$ is positive (e.g., cold packs, melting).     * Hess's Law: The enthalpy change of a reaction is the same regardless of the number of steps it takes.

Reaction Rates and Chemical Equilibrium

• Collision Theory: Reactants must collide with sufficient energy (Activation Energy) and correct orientation.
• Factors Increasing Rate: Increased surface area (iron filings vs. rod), increased temperature, concentration, and catalysts.
• Catalysts: Proteins called enzymes increase rates by lowering activation energy without being consumed.
• Le Chatelier’s Principle: If a stress is applied to a system at equilibrium, the system shifts to relieve the stress.     * Temperature: Increasing temperature in an exothermic reaction decreases $K_{eq}$.     * Pressure: Increasing pressure shifts equilibrium toward the side with fewer gas moles.

Organic Chemistry and Life Chemistry

• Hydrocarbons:     * Alkanes: Saturated, single bonds only ($C_n H_{2n+2}$).     * Alkenes: Unsaturated, at least one double bond ($C_n H_{2n}$).     * Alkynes: Unsaturated, at least one triple bond ($C_n H_{2n-2}$).

• Functional Groups:     * Alcohols: $-OH$ (Hydroxyl group).     * Halides: $-X$ (Halogen).     * Ethers: $-O-$ (Ether group).     * Aldehydes: $-CHO$ (Carbonyl group at end).     * Ketones: $-CO-$ (Carbonyl group in middle).     * Carboxylic Acids: $-COOH$ (Carboxyl group).     * Esters: $-COO-$ (Fruity odors).     * Amines: $-NH_2$ (Odor of decaying organisms).

• Chemistry of Life:     * Proteins: Polyamides made of amino acids connected by peptide bonds.     * Carbohydrates: Sources of energy (Glucose is monosaccharide, Sucrose is disaccharide, Cellulose is polysaccharide).     * Lipids: Include waxes (ester of long chain alcohol), steroids (4-ring structure), and triglycerides.     * Nucleic Acids: Nitrogen-containing biopolymers; structural unit is the nucleotide.

Mixtures, Solutions, Acids, and Bases

• Solutions:     * Molarity (M): $\text{Moles of solute} / \text{Volume of solution (L)}$.     * Molality (m): $\text{Moles of solute} / \text{Mass of solvent (kg)}$.     * Colligative Properties: Boiling point elevation, freezing point depression, vapor pressure lowering (proportional to solute particles).

• Acids and Bases:     * Arrhenius: Acids produce $H^+$; Bases produce $OH^-$.     * Bronsted-Lowry: Acids donate protons ($H^+$); Bases accept protons.     * Lewis: Acids accept electron pairs; Bases donate electron pairs.     * pH Scale: Acidic ($pH < 7$), Neutral ($pH = 7$), Basic ($pH > 7$). $pH + pOH = 14$.     * Buffer Solution: Resists changes in pH; consists of a weak acid/base and its conjugate salt.

Electrochemistry

• Oxidation and Reduction (Redox):     * Oxidation: Loss of electrons; increase in oxidation number. Occurs at the Anode.     * Reduction: Gain of electrons; decrease in oxidation number. Occurs at the Cathode.     * Oxidizing Agent: The substance that gets reduced.     * Reducing Agent: The substance that gets oxidized.

• Cells and Batteries:     * Galvanic Cell: Converts chemical energy to electrical spontaneously ($E_{cell}$ is positive).         * $E_{cell} = E_{cathode} - E_{anode}$     * Electrolytic Cell: Uses electricity to drive a non-spontaneous reaction.     * Primary Battery: Single-use (e.g., dry cell).     * Secondary Battery: Rechargable (e.g., lead-acid car battery, lithium-ion laptop battery).