General Chemistry Study Guide: Atomic Structure, Measurement, and Chemical Nomenclature

Chapter 1: Introduction to Chemistry and Measurement

Pure Substances vs. Mixtures

Understanding the distinction between pure substances and mixtures is fundamental in chemistry. Pure substances have uniform and definite composition, while mixtures consist of two or more substances physically combined.

• Pure Substance: Matter with a fixed composition (e.g., water, sodium chloride).

• Mixture: Combination of two or more substances; can be homogeneous (uniform) or heterogeneous (non-uniform).

• Example: Air is a homogeneous mixture; sand and iron filings are a heterogeneous mixture.

Physical vs. Chemical Change

Chemical changes result in new substances, while physical changes do not alter the chemical identity.

• Physical Change: Change in state or appearance (e.g., melting ice).

• Chemical Change: Formation of new substances (e.g., rusting iron).

Measurement Units and Prefixes

Chemistry uses the International System of Units (SI) for measurements. Prefixes denote powers of ten.

• Length: meter (m)

• Mass: kilogram (kg)

• Time: second (s)

• Temperature: kelvin (K)

• Common Prefixes: kilo- ($10^3$), centi- ($10^{-2}$), milli- ($10^{-3}$)

• Example: 1 kilometer = $1 imes 10^3$ meters

Temperature Conversions

Temperature can be measured in Celsius, Fahrenheit, or Kelvin. Conversion formulas are essential.

• Celsius to Kelvin: $K = ^ ext{\circ}C + 273.15$

• Celsius to Fahrenheit: $^ ext{\circ}F = (^\text{\circ}C \times \frac{9}{5}) + 32$

Significant Figures

Significant figures reflect the precision of a measurement. Rules govern their use in calculations.

• Multiplication/Division: Result has the same number of significant figures as the measurement with the fewest significant figures.

• Addition/Subtraction: Result has the same number of decimal places as the measurement with the fewest decimal places.

Dimensional Analysis and Unit Conversions

Dimensional analysis uses conversion factors to change units.

• Example: Converting inches to centimeters: $1 \text{ in} = 2.54 \text{ cm}$

• Method: Multiply by conversion factors to cancel units.

Chapter 2: Atomic Theory and Structure

Dalton's Atomic Theory

Dalton's theory laid the foundation for modern atomic understanding.

• Statement 1: Elements are composed of tiny, indivisible particles called atoms.

• Statement 2: All atoms of a given element are identical in mass and properties.

• Statement 3: Compounds are formed by combinations of atoms of different elements.

• Statement 4: Chemical reactions involve rearrangement of atoms; atoms are not created or destroyed.

Subatomic Particles

Atoms consist of protons, neutrons, and electrons.

• Proton: Positively charged, found in nucleus.

• Neutron: Neutral, found in nucleus.

• Electron: Negatively charged, found in electron cloud.

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

• Mass Number (A): Number of protons plus neutrons.

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons.

• Example: Carbon-12 and Carbon-14 are isotopes of carbon.

• Representation: $^{A}_{Z}\text{X}$, where X is the element symbol.

Periodic Table Organization

The periodic table arranges elements by increasing atomic number and groups elements with similar properties.

• Groups/Families: Vertical columns (e.g., Alkali metals, Alkaline earth metals, Halogens, Noble gases).

• Periods: Horizontal rows.

• Metals, Nonmetals, Metalloids: Classified by properties.

Atomic Mass and Moles

Atomic mass is the weighted average mass of an element's isotopes. The mole is a counting unit in chemistry.

• Avogadro's Number: $6.022 \times 10^{23}$ particles/mol

• Formula: $\text{Number of moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

Chapter 3: Chemical Nomenclature and Formulas

Covalent vs. Ionic Bonding

Chemical bonds form between atoms to create compounds. Covalent bonds share electrons; ionic bonds transfer electrons.

• Covalent Bond: Nonmetal + Nonmetal; electrons are shared.

• Ionic Bond: Metal + Nonmetal; electrons are transferred.

• Example: NaCl is ionic; H2O is covalent.

Empirical and Molecular Formulas

Formulas represent the composition of compounds.

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

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

• Example: Glucose: Empirical = CH2O, Molecular = C6H12O6

Naming Compounds

Systematic rules are used to name chemical compounds.

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

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

• Acids: Naming depends on anion; "-ate" becomes "-ic acid", "-ite" becomes "-ous acid".

Polyatomic Ions

Polyatomic ions are charged groups of covalently bonded atoms.

• Examples: Nitrate ($\text{NO}_3^-$), Sulfate ($\text{SO}_4^{2-}$), Ammonium ($\text{NH}_4^+$)

Hydrates

Hydrates are compounds that include water molecules in their structure.

• Example: Copper(II) sulfate pentahydrate: CuSO4·5H2O

Chapter 4: Chemical Calculations

Mass and Mole Calculations

Calculating the mass, moles, and number of particles is essential for quantitative chemistry.

• Formula: $\text{Moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• Number of particles: $\text{Particles} = \text{moles} \times 6.022 \times 10^{23}$

Empirical Formula Determination

The empirical formula is determined from percent composition or combustion analysis.

• Steps:

  1. Convert percent to grams (assume 100 g sample).

  2. Convert grams to moles for each element.

  3. Divide by smallest number of moles to get ratios.

  4. Multiply to get whole numbers if necessary.

• Example: A compound with 40% C, 6.7% H, and 53.3% O yields CH2O as the empirical formula.

Combustion Analysis

Combustion analysis is used to determine the empirical formula of organic compounds.

• Method: Burn sample, measure CO2 and H2O produced, calculate moles of C and H, and determine O by difference.

Sample Table: Common Polyatomic Ions

The following table lists some common polyatomic ions and their formulas.

Name	Formula	Charge
Nitrate	NO3-	-1
Sulfate	SO42-	-2
Phosphate	PO43-	-3
Ammonium	NH4+	+1
Hydroxide	OH-	-1

Additional info: Some context and explanations have been expanded for clarity and completeness based on standard General Chemistry curriculum.