General Chemistry 1201: Study Notes for Chapters 1–3 (Matter, Atoms, and Stoichiometry)

Introduction to Matter, Energy, and Measurement

Classification of Matter

Matter is anything that has mass and occupies space. It can be classified based on its physical state and composition.

• Element: A pure substance that cannot be broken down into simpler substances by chemical means (e.g., Helium gas).

• Compound: A substance composed of two or more elements chemically combined in fixed proportions (e.g., Sodium hydrogen carbonate).

• Mixture: A combination of two or more substances in which each retains its own identity (e.g., Sea water).

Example: Mercury is an element, while paint is a mixture.

Physical Quantities and Units

Measurements in chemistry require units. The SI unit for temperature is the kelvin (K).

• Temperature Conversion: To convert Celsius to Kelvin, use $K = ^\circ C + 273.15$.

• Temperature Difference: A difference of 100 °C is equivalent to 100 K.

Example: The temperature difference between 0 °C and 100 °C is 100 K.

Significant Figures

Significant figures reflect the precision of a measured or calculated quantity.

• When performing calculations, the result should have the same number of significant figures as the measurement with the fewest significant figures.

Example: $\frac{25.45 \pm 0.45}{73.55\ \text{mL}}$ should be reported with the correct number of significant figures.

Atoms, Molecules, and Ions

Atomic Structure

Atoms consist of protons, neutrons, and electrons. The number of protons defines the element, while the number of neutrons can vary, resulting in isotopes.

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

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

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

Example: $^{35}\text{Cl}$ and $^{37}\text{Cl}$ are isotopes of chlorine.

Periodic Table

The periodic table organizes elements by increasing atomic number and similar chemical properties.

• Groups: Vertical columns with similar properties.

• Periods: Horizontal rows.

Example: Alkali metals are found in Group 1.

Ions and Ionic Compounds

Ions are atoms or molecules that have gained or lost electrons.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

• Ionic Compounds: Composed of cations and anions in a ratio that results in electrical neutrality.

Example: $\text{Na}^+$ and $\text{Cl}^-$ combine to form $\text{NaCl}$.

Formulas of Compounds

Chemical formulas indicate the elements present and their ratios.

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

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

Example: The empirical formula of hydrogen peroxide is $\text{HO}$, while the molecular formula is $\text{H}_2\text{O}_2$.

Naming Compounds

Systematic naming follows specific rules for ionic and molecular compounds.

• Ionic Compounds: Name the cation first, then the anion (e.g., sodium chloride).

• Molecular Compounds: Use prefixes to indicate the number of atoms (e.g., carbon dioxide).

Example: The systematic name of $\text{ClO}_2$ is chlorine dioxide.

Common Polyatomic Ions

Ion	Formula	Name
Hypochlorite	ClO-	Hypochlorite ion
Carbonate	CO32-	Carbonate ion
Sulfate	SO42-	Sulfate ion
Nitrate	NO3-	Nitrate ion

Chemical Reactions and Reaction Stoichiometry

Types of Chemical Reactions

Chemical reactions involve the transformation of substances into new substances.

• Combination Reaction: Two or more substances combine to form one product.

• Decomposition Reaction: One substance breaks down into two or more products.

• Combustion Reaction: A substance reacts with oxygen to produce energy, carbon dioxide, and water.

Example: $\text{C}_3\text{H}_8 + 5\text{O}_2 \rightarrow 3\text{CO}_2 + 4\text{H}_2\text{O}$ (combustion of propane).

Balancing Chemical Equations

Balanced equations have the same number of each type of atom on both sides.

• Use the smallest whole-number coefficients.

Example: $\text{HCO}_3^- + \text{KOH} \rightarrow \text{K}_2\text{CO}_3 + \text{H}_2\text{O}$

Stoichiometry

Stoichiometry involves calculations based on balanced chemical equations.

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

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

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

Example: If 2.5 mol of $\text{O}_2$ are used in the combustion of propane, calculate the moles of $\text{CO}_2$ produced using the balanced equation.

Percent Composition

Percent composition is the percent by mass of each element in a compound.

• Calculate by dividing the mass of the element by the total molar mass and multiplying by 100%.

Example: Calculate the percent of C in $\text{CH}_3\text{OH}$.

Empirical and Molecular Formulas

Empirical formulas show the simplest ratio, while molecular formulas show the actual number of atoms.

• Determine empirical formula from percent composition or combustion data.

Example: A hydrocarbon produces 0.3931 g $\text{CO}_2$ and 0.0807 g $\text{H}_2\text{O}$; determine its empirical formula.

Isomers

Isomers are compounds with the same molecular formula but different structures.

• Structural Isomers: Differ in the connectivity of atoms.

• Stereoisomers: Same connectivity, different spatial arrangement.

Example: Butane and cyclobutane are not isomers because they have different molecular formulas.

Lab Techniques and Mathematical Operations

Density and Volume Calculations

Density is mass per unit volume ($d = \frac{m}{V}$). Comparing the same mass of two liquids, the one with a smaller volume has a higher density.

• Example: 352 g of two different liquids are poured into beakers; the one with a lower volume is denser.

Significant Figures in Calculations

When performing calculations, report the answer with the correct number of significant figures based on the data provided.

• Example: $\frac{25.45 \pm 0.45}{73.55\ \text{mL}}$

Tables and Data Interpretation

Isotope Table Example

Isotopes can be compared based on their symbol, atomic number, mass number, and charge.

Symbol	Protons	Neutrons	Electrons	Net Charge
$^{34}_{16}\text{S}^{2-}$	16	18	18	2-
$^{36}_{16}\text{S}$	16	20	16	0
$^{38}_{16}\text{S}^{2+}$	16	22	14	2+

Additional info: Table entries inferred from standard isotope notation and periodic table data.

Compound Formula Table Example

	Na+	Fe3+	Al3+
NO3-	NaNO3	Fe(NO3)3	Al(NO3)3
SO42-	Na2SO4	Fe2(SO4)3	Al2(SO4)3

Additional info: Table entries inferred from standard ionic compound formulas.

Bonus: Organic Chemistry Introduction

Simple Alcohols

Alcohols are organic compounds containing the hydroxyl group (-OH) attached to a carbon atom.

• Methanol: CH3OH

• Ethanol: C2H5OH

• Propanol: C3H7OH

• Butanol: C4H9OH

Example: The formula for ethanol is C2H5OH.

Additional info: These study notes cover the main topics assessed in the exam, including matter classification, atomic structure, chemical formulas, reactions, stoichiometry, and basic organic chemistry, as well as lab and mathematical skills relevant to General Chemistry Chapters 1–3.