Chemistry Study Notes: Quantities in Chemical Reactions

  1. Basic Principles

    • All substances are made of tiny particles.

    • There is always space between particles.

    • Particles are always in motion.

    • Particles attract each other.

  2. States of Matter

    • Solids: Definite shape and volume. Particles vibrate but are fixed.

    • Liquids: Take the shape of their container; volume is definite due to strong attractions between particles allowing more freedom of movement.

    • Gases: No definite shape or volume; particles are widely spaced and moving constantly in all directions.

Physical and Chemical Properties

  1. Physical Properties

    • Characteristics that do not change the substance. Examples: density, state, color, size, shape, malleability, ductility, melting point, boiling point, solubility.

    • Qualitative: Described with words (e.g., color of zinc).

    • Quantitative: Described with numbers (e.g., density of zinc is 7.14 g/ml).

  2. Chemical Properties

    • Describes how a substance reacts. Includes toxicity, reactivity, combustibility, stability.

    • Signs of a chemical change: color change, state change, gas production, odor, heat production.

The Atom

  1. Subatomic Particles

    • Particles within an atom: electrons, neutrons, and protons.

    • Electrons: Exist in orbitals and determine reactivity.

    • Nucleus: Contains protons and neutrons, accounting for atomic mass.

  2. Atomic Number and Isotopes

    • Atomic number equals the number of protons. Example: Carbon (C) has 6 protons.

    • Isotopes: Atoms with the same number of protons but different numbers of neutrons (mass number). Example: C-12 has 6 neutrons.

  3. Mass Number

    • Mass number = Number of protons + Number of neutrons.

  4. Charge of an Atom

    • Charge = Number of protons - Number of electrons. Atom is neutral if charges balance.

Periodic Table of Elements

  1. Organization

    • Arranged by increasing atomic number.

    • Groups (columns) indicate valence electrons; periods (rows) indicate energy levels.

  2. Group Names

    • Group 1: Alkali metals.

    • Group 2: Alkaline earth metals.

    • Groups 3-12: Transition metals.

    • Group 17: Halogens.

    • Group 18: Noble gases.

Mass Spectrometry

  1. Process

    • Sample vaporized and ionized.

    • Positively charged ions are accelerated and deflected by magnetic fields.

  2. Law of Inertia

    • Heavier particles deflect less.

  3. Average Atomic Mass Calculation

    • Avg.AtomicMass(AAM)=ext(mass<em>1imesextabundance</em>1+extmass<em>2imesextabundance</em>2+ext…)Avg. Atomic Mass (AAM) = ext{(mass}<em>1 imes ext{abundance}</em>1 + ext{mass}<em>2 imes ext{abundance}</em>2 + ext{…)}

    • Example: For Chlorine, AAM=35(0.80)+37(0.20)=35.4extamuAAM = 35(0.80) + 37(0.20) = 35.4 ext{ amu}.

Chemical Compounds and Bonding

  1. Ionic Bonding

    • Occurs when electrons are transferred from one atom to another to achieve a stable octet.

    • Ionic compounds, e.g. NaCl, are formed between metals and non-metals.

    • High melting points; conduct electricity when dissolved or molten.

  2. Covalent Bonding

    • Occurs when two non-metals share electrons. Example: Methane (CH4).

    • Covalent compounds have lower melting points compared to ionic compounds; generally do not conduct electricity.

  3. Periodic Trends

    • Atomic radius decreases across a period (increased nuclear charge) and increases down a group (more shells).

    • Ionization Energy (IE) increases across a period and decreases down a group.

    • Electron Affinity and Electronegativity have similar trends.

  4. Polarity of Molecules

    • Molecules with asymmetrical charge distribution are polar. Example: Water (H2O) is polar due to its bent structure.

Lewis Structures

  • Method for modeling covalent bonds and molecular shapes. Involves drawing pairs of valence electrons.

  1. Bonding Pairs (BP) and Lone Pairs (LP)

    • BP: Electrons shared between atoms.

    • LP: Non-bonding electrons.

  2. Molecular Shapes

    • Determined by the number of bonding pairs and lone pairs. Examples include linear, tetrahedral, trigonal pyramidal, etc.

Nomenclature

  1. Ionic Compounds

    • Name metal first, followed by the non-metal with an 'ide' suffix (KCl = potassium chloride).

  2. Covalent Compounds

    • Use prefixes to denote the number of atoms (CO2 = carbon dioxide).

  3. Acids

    • Binary acids named as hydro+(element name)+-ic acid. Example: HCl = hydrochloric acid.

    • Oxyacids' names depend on the number of oxygen atoms without 'hydro.'

  4. Common Polyatomic Ions

    • Must be memorized for correct naming of complex compounds.