Introduction to Chemistry: Matter, Energy, and Measurement

Definition and Classification of Matter

• Matter is anything that has mass and occupies space.

• Energy is the capacity to do work.

• Matter can be classified as:

  • Pure Substances: Have a fixed composition and distinct properties. Examples: Elements (H, Na, O) and Compounds (H2O, NaCl).

  • Mixtures: Physical combinations of two or more substances. Can be:

    • Homogeneous Mixtures (solutions): Uniform composition (e.g., salt dissolved in water, air).

    • Heterogeneous Mixtures: Non-uniform composition (e.g., oil and water, vegetable soup).

• Physical Changes do not alter the chemical composition (e.g., melting ice: H2O(s) → H2O(l)).

• Chemical Changes alter the chemical structure by breaking/forming bonds (e.g., cooking, rusting).

Elements, Compounds, and Mixtures

• Element: A substance that cannot be broken down chemically into simpler substances.

• Compound: A chemical combination of elements in fixed ratios (e.g., NaCl, H2O).

• Mixture: Physical blend of substances; can be separated by physical means.

Examples of Classification

• Iron – element

• Water – compound

• Dissolved sugar in water – homogeneous mixture

• Vegetable soup – heterogeneous mixture

• Air – homogeneous mixture

Measurement and Significant Figures

Significant Figures (Sig Figs)

• All nonzero digits are significant.

• Zeros between nonzero digits are significant (e.g., 3101 has 4 sig figs).

• Leading zeros are not significant (e.g., 0.1 has 1 sig fig).

• Trailing zeros after a decimal point are significant.

• Multiplication/Division: Result should have as many sig figs as the value with the fewest sig figs.

• Addition/Subtraction: Result should have as many decimal places as the value with the fewest decimal places.

Unit Conversions and Metric Prefixes

• Common prefixes:

• Unit conversion example:

Temperature Scales

• Celsius (°C), Fahrenheit (°F), Kelvin (K)

• Kelvin is the SI base unit for temperature.

• Conversion:

Density

• Density is mass per unit volume.

• Formula:

• Example: Calculate the mass of 41.0 mL of mercury (density = 13.53 g/mL):

Atoms, Molecules, and Ions

Atomic Structure and Isotopes

• Atomic Mass is the weighted average of the masses of naturally occurring isotopes.

• Formula:

• Example: Copper has two isotopes, (69.17%, 62.9396 amu) and (30.83%, 64.9278 amu):

Chemical Nomenclature

• Chemical Formulas use element symbols and subscripts to indicate composition (e.g., H2O, Ba3(PO4)2).

• Covalent Compounds (nonmetal + nonmetal): Use Greek prefixes (mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-).

• Ionic Compounds (metal + nonmetal): Name cation (full name), then anion (root + -ide). For transition metals, indicate charge with Roman numerals.

• Polyatomic Ions: Common examples include acetate (CH3COO-), carbonate (CO32-), sulfate (SO42-), phosphate (PO43-).

Acids and Bases Nomenclature

• Binary Acids: "Hydro-" + root + "-ic acid" (e.g., HCl: hydrochloric acid).

• Oxyacids: Based on polyatomic ion name (e.g., HNO3: nitric acid, HNO2: nitrous acid).

Organic Chemistry: Hydrocarbons and Functional Groups

• Alkanes: Single bonds (C–C), general formula CnH2n+2.

• Alkenes: Double bonds (C=C), general formula CnH2n.

• Alkynes: Triple bonds (C≡C), general formula CnH2n-2.

• Common functional groups: alcohol (R–OH), ether (R–O–R'), carboxylic acid (COOH), amine (R–NH2).

Chemical Reactions and Stoichiometry

Balancing Chemical Equations

• Coefficients are used to balance the number of atoms of each element on both sides of the equation.

• Example:

The Mole Concept and Avogadro's Number

• 1 mole = particles (Avogadro's number).

• Number of moles:

• Number of particles:

Molar Mass and Percent Composition

• Molar Mass: Mass in grams of one mole of a substance.

• Percent Composition:

Empirical and Molecular Formulas

• Empirical Formula: Simplest whole-number ratio of atoms in a compound.

• Molecular Formula: Actual number of atoms of each element in a molecule.

Limiting Reactant, Theoretical Yield, Actual Yield, Percent Yield

• Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

• Theoretical Yield: Maximum amount of product possible from given reactants.

• Actual Yield: Amount of product actually obtained.

• Percent Yield:

Reactions in Aqueous Solution

Types of Chemical Reactions

• Combination (Synthesis):

• Decomposition:

• Single Displacement:

• Double Displacement:

• Combustion:

• Acid-Base Reaction:

Net Ionic Equations

• Show only the species that actually participate in the reaction.

• Spectator ions are omitted.

• Example: Net ionic:

Redox (Oxidation-Reduction) Reactions

• Oxidation: Loss of electrons (increase in oxidation state).

• Reduction: Gain of electrons (decrease in oxidation state).

• Rules for assigning oxidation states:

  • Elemental form: 0

  • Monoatomic ion: charge of the ion

  • Oxygen: usually -2

  • Hydrogen: +1 (except when bonded to metals: -1)

  • Sum of oxidation states equals overall charge

Lab Techniques and Mathematical Operations

Solution Concentration: Molarity and Dilution

• Molarity (M):

• Dilution Equation:

• Percent by Mass:

Titration Calculations

• Used to determine the concentration of an unknown solution by reacting it with a solution of known concentration.

• Example:

• Use stoichiometry to relate volumes and concentrations.

Summary Table: Common Polyatomic Ions

Practice Problems and Applications

• Calculate the number of moles, mass, or particles given appropriate data.

• Balance chemical equations and identify reaction types.

• Determine limiting reactant, theoretical yield, and percent yield in reactions.

• Write names and formulas for ionic and covalent compounds.

• Assign oxidation states and identify redox processes.

• Calculate solution concentrations and perform dilution/titration calculations.

Additional info: Some examples and tables were expanded for clarity and completeness. All equations are provided in LaTeX format as per instructions.