Introduction to Chemistry: Comprehensive Midterm Review Study Notes

Classification of Matter

Homogeneous vs. Heterogeneous Mixtures

Mixtures can be classified based on the uniformity of their composition.

• Homogeneous Mixture: A mixture with a uniform composition throughout. Example: Saltwater, air.

• Heterogeneous Mixture: A mixture with a non-uniform composition, where different parts can be distinguished. Example: Salad, sand in water.

Physical and Chemical Changes

Definitions and Examples

Understanding the difference between physical and chemical changes is fundamental in chemistry.

• Physical Change: A change that affects the form of a chemical substance, but not its chemical composition. Examples: Melting ice, dissolving sugar in water.

• Chemical Change: A change that results in the formation of new chemical substances. Examples: Rusting of iron, burning of wood.

Measurement and Significant Figures

Significant Figures

Significant figures reflect the precision of a measured quantity.

• Count all nonzero digits as significant.

• Zeros between nonzero digits are significant.

• Leading zeros are not significant; trailing zeros are significant only if there is a decimal point.

Examples:

• 441.5 kg: 4 significant figures

• 1.002 2500 cm: 7 significant figures

• 0.00309 g: 3 significant figures

Unit Conversions

Unit conversions are essential for solving chemistry problems.

• Use conversion factors to change from one unit to another.

• Example: 12.56 in to yards. Recall 1 in = 2.54 cm, 1 yd = 36 in.

Density and Experimental Error

Density Calculations

Density is the mass per unit volume of a substance.

• Formula: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Example: If 11.0 kg of titanium has a volume of 4.67 cm3, density = $\frac{11.0\ \text{kg}}{4.67\ \text{cm}^3}$

Percent Error

Percent error quantifies the accuracy of an experimental value.

• Formula: $\text{Percent Error} = \left| \frac{\text{Experimental Value} - \text{Accepted Value}}{\text{Accepted Value}} \right| \times 100\%$

Atomic Structure

Atomic Number and Mass Number

Atoms are composed of protons, neutrons, and electrons.

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

• Mass Number (A): Total number of protons and neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons.

Subatomic Particles

• Protons: Positively charged, in the nucleus.

• Neutrons: Neutral, in the nucleus.

• Electrons: Negatively charged, orbit the nucleus.

Atomic Theories

Key contributors to atomic theory include Dalton, Rutherford, and Thomson.

• Dalton: Proposed that atoms are indivisible particles.

• Thomson: Discovered the electron (plum pudding model).

• Rutherford: Discovered the nucleus via gold foil experiment.

Isotopes and Atomic Mass

Isotope Symbols

• Isotope notation: $^{A}_{Z}\text{X}$, where A = mass number, Z = atomic number, X = element symbol.

Average Atomic Mass

• Formula: $\text{Average Atomic Mass} = \sum (\text{Fractional Abundance} \times \text{Isotope Mass})$

• Example: For carbon with isotopes of 12.0 amu (98.89%) and 13.0 amu (1.11%): $\text{Average Atomic Mass} = (0.9889 \times 12.0) + (0.0111 \times 13.0)$

Chemical Nomenclature and Formulas

Writing Formulas and Naming Compounds

• Use the periodic table to determine charges and write correct formulas.

• Examples:

  • Potassium and dichromate: K2Cr2O7

  • Aluminum and fluoride: AlF3

  • Copper(II) sulfate: CuSO4

  • Barium phosphate: Ba3(PO4)2

  • Sodium nitrate: NaNO3

Polyatomic Ions and Charges

• Common polyatomic ions: sulfate (SO42−), nitrate (NO3−), phosphate (PO43−).

• Transition metals may have multiple charges (e.g., Fe2+, Fe3+).

Chemical Reactions

Types of Chemical Reactions

Chemical reactions can be classified into several types:

Type	General Form	Example
Synthesis	A + B → AB	2H2 + O2 → 2H2O
Decomposition	AB → A + B	2H2O → 2H2 + O2
Single Replacement	A + BC → AC + B	Zn + 2HCl → ZnCl2 + H2
Double Replacement	AB + CD → AD + CB	AgNO3 + NaCl → AgCl + NaNO3
Combustion	Hydrocarbon + O2 → CO2 + H2O	CH4 + 2O2 → CO2 + 2H2O

Balancing Chemical Equations

• Law of Conservation of Mass: Atoms are neither created nor destroyed in a chemical reaction.

• Balance equations by adjusting coefficients, not subscripts.

Stoichiometry and Chemical Quantities

Mole Concept

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

• Molar mass: mass of 1 mole of a substance (g/mol).

Percent Composition

• Formula: $\text{Percent by Mass} = \frac{\text{Mass of Element}}{\text{Total Mass of Compound}} \times 100\%$

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.

Gases and Gas Laws

Gas Laws

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant T and n)

• Charles's Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant P and n)

• Ideal Gas Law: $PV = nRT$

Applications

• Calculate changes in gas volume, pressure, or temperature using the appropriate law.

• STP (Standard Temperature and Pressure): 0°C (273.15 K) and 1 atm.

Solutions and Percent Yield

Percent Yield

• Formula: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

Additional Topics

• Writing and balancing chemical equations for reactions such as double replacement and combustion.

• Calculating the mass of reactants and products using stoichiometry.

• Identifying limiting reactants in chemical reactions.

Additional info: Some context and examples were expanded for clarity and completeness based on standard introductory chemistry curricula.