Chem

Introduction to Chemistry

• Matter and its Properties

  • Matter: Anything that has mass and occupies space.

  • Properties:

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

  • Chemical Properties: Describe how a substance reacts to form other substances (e.g., flammability, reactivity).

• Classification of Matter

  • Pure Substances:

  • Elements: Cannot be broken down into simpler substances by chemical means (e.g., H, O).

  • Compounds: Two or more elements chemically combined in fixed proportions (e.g., $H<em>2O$, $CO</em>2$).

  • Mixtures:

  • Homogeneous Mixtures (Solutions): Uniform composition throughout (e.g., saltwater, air).

  • Heterogeneous Mixtures: Non-uniform composition; components are visibly separate (e.g., sand and water, oil and water).

• States of Matter

  • Solid: Definite shape and volume.

  • Liquid: Definite volume but takes the shape of its container.

  • Gas: No definite shape or volume; expands to fill its container.

  • Plasma: Ionized gas.

Atomic Structure

• Atomic Theory

  • Dalton's Atomic Theory: Elements are composed of indivisible atoms; atoms of the same element are identical.

  • Thomson's Plum Pudding Model: Atoms are a sphere of positive charge with electrons embedded.

  • Rutherford's Gold Foil Experiment: Led to the nuclear model (dense, positively charged nucleus with electrons orbiting).

  • Bohr Model: Electrons orbit the nucleus in specific energy levels.

  • Quantum Mechanical Model: Current model, describing electrons in orbitals (regions of probability).

• Subatomic Particles

  • Protons ($p^+$): Positive charge ($+1$), located in the nucleus, mass approx. $1 \text{ amu}$.

  • Neutrons ($n^0$): No charge, located in the nucleus, mass approx. $1 \text{ amu}$.

  • Electrons ($e^-$): Negative charge ($-1$), orbit the nucleus, negligible mass.

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

• Mass Number ($A$): Number of protons + neutrons.

• Isotopes: Atoms of the same element (same $Z$) but different number of neutrons (different $A$).

• Quantum Mechanics and Electron Configuration

  • Light and Wave Properties

    • Electromagnetic Spectrum: Range of all types of EM radiation.

    • Wave-Particle Duality: Light exhibits properties of both waves (wavelength $\lambda$) and particles (photons).

    • Planck's Equation: $E = h\nu$ (where $h$ is Planck's constant, $\nu$ is frequency).

    • Relationship between c, $\lambda$, $\nu$: $c = \lambda\nu$ (where $c$ is the speed of light).

  • Quantum Numbers: Describe the properties of electrons in atomic orbitals.

    • Principal Quantum Number ($n$): Defines the energy level and size of the orbital ($n = 1, 2, 3, \dots$).

    • Angular Momentum (Azimuthal) Quantum Number ($l$): Defines the shape of the orbital ($l = 0, 1, 2, \dots, n-1$ corresponds to s, p, d, f orbitals).

    • Magnetic Quantum Number ($m<em>l$): Defines the orientation of the orbital in space ($m</em>l = -l, \dots, 0, \dots, +l$).

    • Spin Quantum Number ($m_s$): Defines the intrinsic angular momentum of an electron (spin up $+1/2$ or spin down $-1/2$).

    • Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.

    • Hund's Rule: Electrons fill degenerate orbitals singly before pairing up.

    • Aufbau Principle: Electrons fill atomic orbitals of the lowest available energy levels before filling higher ones.

Periodic Table and Periodicity

• Organization: Elements arranged by increasing atomic number.

• Periods (Rows): Indicate the principal energy level.

• Groups (Columns): Elements in the same group have similar chemical properties due to the same number of valence electrons.

• Key Groups:

  • Alkali Metals (Group 1): Highly reactive metals.

  • Alkaline Earth Metals (Group 2): Reactive metals.

  • Halogens (Group 17): Highly reactive nonmetals.

  • Noble Gases (Group 18): Very unreactive.

• Periodic Trends

  • Atomic Radius: Decreases across a period, increases down a group.

  • Ionization Energy: Energy required to remove an electron; increases across a period, decreases down a group.

  • Electronegativity: Ability of an atom to attract electrons in a chemical bond; increases across a period, decreases down a group.

Chemical Bonding

• Ionic Bonding: Transfer of electrons between a metal and a nonmetal, forming ions and electrostatic attraction.

• Covalent Bonding: Sharing of electrons between two nonmetals.

  • Nonpolar Covalent: Equal sharing of electrons (e.g., $O_2$).

  • Polar Covalent: Unequal sharing of electrons, creating partial charges (e.g., $H_2O$).

• Metallic Bonding: Delocalized electrons shared among a lattice of metal atoms.

• Lewis Structures: Diagrams that show bonding and non-bonding electron pairs.

• VSEPR Theory (Valence Shell Electron Pair Repulsion): Predicts molecular geometry based on minimizing electron pair repulsion.

Chemical Reactions and Stoichiometry

• Types of Reactions:

  • Synthesis: $A + B \rightarrow AB$

  • Decomposition: $AB \rightarrow A + B$

  • Single Displacement: $A + BC \rightarrow AC + B$

  • Double Displacement: $AB + CD \rightarrow AD + CB$

  • Combustion: Compound reacts with oxygen, producing heat and light.

• Balancing Chemical Equations: Ensures conservation of mass.

• Mole Concept: $1 \text{ mole} = 6.022 \times 10^{23}$ particles (Avogadro's number).

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

• Stoichiometric Calculations:

  • Relate amounts of reactants and products in a balanced equation.

  • Limiting Reactant: Reactant consumed first, determining maximum product.

  • Percent Yield: (Actual yield / Theoretical yield) $\times 100\%$

Gases

• Ideal Gas Law: $PV = nRT$

  • $P$ = pressure, $V$ = volume, $n$ = moles, $R$ = ideal gas constant ($0.0821 \text{ L} \cdot \text{atm/mol} \cdot \text{K}$), $T$ = temperature (in Kelvin).

• Other Gas Laws:

  • Boyle's Law: Inverse relationship between pressure and volume ($P<em>1V</em>1 = P<em>2V</em>2$).

  • Charles's Law: Direct relationship between volume and temperature ($V<em>1/T</em>1 = V<em>2/T</em>2$).

  • Avogadro's Law: Direct relationship between volume and moles ($V<em>1/n</em>1 = V<em>2/n</em>2$).

  • Dalton's Law of Partial Pressures: Total pressure of a gas mixture is the sum of partial pressures of individual gases ($P<em>{\text{total}} = P</em>1 + P_2 + \dots$).

Solutions

• Solution: Homogeneous mixture of solute (minor component) and solvent (major component).

• Solubility: Maximum amount of solute that can dissolve in a given amount of solvent.

• Concentration Units:

  • Molarity ($M$): Moles of solute / Liters of solution.

  • Molality ($m$): Moles of solute / kg of solvent.

  • Percent by Mass: (Mass of solute / Mass of solution) $\times 100\%$

Acids and Bases

• Definitions:

  • Arrhenius: Acids produce $H^+$ in water; Bases produce $OH^-$ in water.

  • Brønsted-Lowry: Acids are proton ($H^+$) donors; Bases are proton acceptors.

  • Lewis: Acids are electron pair acceptors; Bases are electron pair donors.

• pH Scale: Measures acidity or alkalinity ($pH = -log[H^+]$).

  • pH < 7: Acidic

  • $pH = 7$: Neutral

  • pH > 7: Basic

• Strong vs. Weak:

  • Strong Acids/Bases: Fully dissociate in water.

  • Weak Acids/Bases: Partially dissociate in water.

Thermochemistry

• Thermochemistry: Study of heat changes associated with chemical reactions.

• Energy and Heat:

  • Energy: Capacity to do work or produce heat.

    • Kinetic Energy: Energy of motion.

    • Potential Energy: Stored energy.

  • System and Surroundings:

    • System: The part of the universe being studied (e.g., reactants and products).

    • Surroundings: Everything else.

  • Heat ($q$): Transfer of thermal energy between objects due to a temperature difference.

    • Exothermic Processes: Release heat to the surroundings (q < 0).

    • Endothermic Processes: Absorb heat from the surroundings (q > 0).

  • Temperature: Measure of the average kinetic energy of the particles in a substance.

• Specific Heat Capacity:

  • Specific Heat Capacity ($c$ or $C_s$): Amount of heat required to raise the temperature of $1 \text{ g}$ of a substance by $1^\circ C$ (or $1 \text{ K}$).

  • Heat Calculation: $q = mc\Delta T$ (where $m$ = mass, $c$ = specific heat capacity, $\Delta T$ = change in temperature).

• Enthalpy ($H$):

  • Enthalpy ($H$): Heat content of a system at constant pressure.

  • Enthalpy Change ($\Delta H$): Heat absorbed or released in a chemical reaction at constant pressure.

  • Standard Enthalpy of Formation ($\Delta H_f^\circ$): Enthalpy change when 1 mole of a compound is formed from its elements in their standard states.