Chem 100 Study Guide: Units 1-3 Core Concepts

Unit 1: Measurement and Scientific Notation

Metric System and SI Units

The metric system is a standardized system of measurement used in science, based on powers of ten. The International System of Units (SI) is the modern form of the metric system.

• SI Units: Fundamental units include meter (m), kilogram (kg), second (s), mole (mol), kelvin (K), ampere (A), and candela (cd).

• Metric Prefixes: Common prefixes include kilo- (103), centi- (10-2), milli- (10-3), micro- (10-6), nano- (10-9).

• Unit Conversion: Convert between units using conversion factors (e.g., 1 m = 100 cm).

Example: Convert 2500 mg to grams: $2500 \text{ mg} \times \frac{1 \text{ g}}{1000 \text{ mg}} = 2.5 \text{ g}$

Scientific Notation

Scientific notation expresses numbers as a product of a coefficient and a power of ten, making it easier to handle very large or small numbers.

• Format: $a \times 10^n$ where $1 \leq a < 10$ and $n$ is an integer.

• Conversion: Move the decimal point to create a coefficient between 1 and 10.

Example: 0.00056 = $5.6 \times 10^{-4}$

Significant Figures

Significant figures reflect the precision of a measured quantity.

• Rules: All nonzero digits are significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros are significant only if there is a decimal point.

• Calculations: For addition/subtraction, the result has the same number of decimal places as the least precise measurement. For multiplication/division, the result has the same number of significant figures as the measurement with the fewest significant figures.

Example: $2.34 \times 1.2 = 2.8$ (rounded to two significant figures)

Unit 2: Chemical Formulas and Stoichiometry

Atoms, Elements, and Compounds

Chemical formulas represent the composition of substances using element symbols and subscripts.

• Atoms: The basic unit of matter, consisting of protons, neutrons, and electrons.

• Elements: Pure substances made of only one type of atom.

• Compounds: Substances composed of two or more elements chemically combined in fixed ratios.

Balancing Chemical Equations

Balancing equations ensures the conservation of mass and atoms in a chemical reaction.

• Steps:

  1. Write the unbalanced equation.

  2. Count the number of atoms of each element on both sides.

  3. Add coefficients to balance each element.

  4. Check your work.

Example: $\ce{2H_2 + O_2 \rightarrow 2H_2O}$

Calculating Moles and Molar Mass

The mole is a counting unit in chemistry, representing $6.022 \times 10^{23}$ particles (Avogadro's number).

• Molar Mass: The mass of one mole of a substance, expressed in grams per mole (g/mol).

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

Example: Calculate moles in 18 g of water ($H_2O$): $\text{Moles} = \frac{18 \text{ g}}{18.02 \text{ g/mol}} = 1.00$ mol

Empirical and Molecular Formulas

Empirical formulas show the simplest whole-number ratio of atoms in a compound; molecular formulas show the actual number of atoms.

• Empirical Formula: Simplest ratio (e.g., $CH_2$ for ethene).

• Molecular Formula: Actual composition (e.g., $C_2H_4$ for ethene).

Example: Glucose: Empirical formula $CH_2O$, molecular formula $C_6H_{12}O_6$

Unit 3: Chemical Bonding and Structure

Types of Chemical Bonds

Chemical bonds are forces that hold atoms together in compounds.

• Ionic Bonds: Formed by the transfer of electrons from one atom to another, resulting in oppositely charged ions.

• Covalent Bonds: Formed by the sharing of electrons between atoms.

• Metallic Bonds: Involve a 'sea' of delocalized electrons among metal atoms.

Example: Sodium chloride ($NaCl$) is ionic; water ($H_2O$) is covalent.

Lewis Structures and Electron Configuration

Lewis structures represent the arrangement of electrons in molecules, showing bonds and lone pairs.

• Steps:

  1. Count total valence electrons.

  2. Draw skeletal structure.

  3. Distribute electrons to satisfy the octet rule.

• Electron Configuration: Describes the arrangement of electrons in atomic orbitals.

Example: Lewis structure for $H_2O$ shows two lone pairs on oxygen.

Polarity and Intermolecular Forces

Molecular polarity depends on the distribution of electrons and molecular geometry.

• Polar Molecules: Have an uneven distribution of charge (e.g., $H_2O$).

• Nonpolar Molecules: Have an even distribution of charge (e.g., $CO_2$).

• Intermolecular Forces: Include hydrogen bonding, dipole-dipole, and London dispersion forces.

Example: Water exhibits hydrogen bonding, leading to high boiling point.

Unit 4: Chemical Reactions and Stoichiometry

Types of Chemical Reactions

Chemical reactions can be classified based on the changes that occur.

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

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

• Single Replacement: One element replaces another in a compound.

• Double Replacement: Exchange of ions between two compounds.

• Combustion: Reaction with oxygen producing heat and light.

Example: $\ce{2H_2 + O_2 \rightarrow 2H_2O}$ is a synthesis reaction.

Stoichiometry and Limiting Reactant

Stoichiometry involves quantitative relationships between reactants and products in a chemical reaction.

• Balanced Equation: Used to determine mole ratios.

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

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

Example: If 5.0 g of $H_2$ reacts with excess $O_2$, calculate the mass of $H_2O$ produced.

Unit 5: Solutions and Concentration

Types of Solutions and Concentration Units

Solutions are homogeneous mixtures of two or more substances.

• Solvent: The substance present in the greatest amount.

• Solute: The substance dissolved in the solvent.

• Concentration Units: Molarity ($M$), molality ($m$), percent by mass, and parts per million (ppm).

• Molarity Formula: $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

Example: A 0.5 M NaCl solution contains 0.5 moles of NaCl per liter.

Dimensional Analysis

Dimensional analysis is a method for converting between units using conversion factors.

• Steps:

  1. Identify the starting unit and desired unit.

  2. Set up conversion factors so units cancel appropriately.

  3. Multiply through to obtain the answer.

Example: Convert 5.0 cm to meters: $5.0 \text{ cm} \times \frac{1 \text{ m}}{100 \text{ cm}} = 0.050 \text{ m}$

Additional info:

• Some topics (e.g., polyatomic ions, electron configuration, and periodic table trends) were inferred from context and standard general chemistry curricula.

• Specific chemical formulas and equations were expanded for clarity.