Comprehensive Review Notes for Introduction to Chemistry Final Exam

Review Topics for Introduction to Chemistry

Scientific Notation and Significant Figures

Scientific notation is a method of expressing very large or very small numbers in the form a × 10n, where a is a number between 1 and 10, and n is an integer. Significant figures are the digits in a measurement that are known with certainty plus one estimated digit.

• Converting numbers to scientific notation: Move the decimal point so that only one nonzero digit remains to the left of the decimal.

• Examples: 1200 = $1.2 \times 10^3$; 0.00060 = $6.0 \times 10^{-4}$

Unit Conversions and Dimensional Analysis

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

• Temperature conversions: $T(\degree C) = (T(\degree F) - 32) \times \frac{5}{9}$

• Volume conversions: 1 L = 1000 mL

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

• Example: If a substance has a density of 1.5 g/mL, 7.5 g has a volume of $\frac{7.5\,\text{g}}{1.5\,\text{g/mL}} = 5.0\,\text{mL}$

Classification of Matter

Matter can be classified as elements, compounds, or mixtures. Elements are pure substances consisting of one type of atom. Compounds are substances composed of two or more elements chemically combined. Mixtures are physical combinations of substances.

• Elements: C, Fe, Na

• Compounds: NaCl, H2O

• Mixtures: Air, salt water

Physical and Chemical Properties

Physical properties can be observed without changing the substance's identity, while chemical properties describe a substance's ability to undergo chemical changes.

• Physical properties: Color, melting point, density

• Chemical properties: Reactivity with acid, flammability

Periodic Table and Element Classification

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

• Metals: Good conductors, malleable, shiny

• Nonmetals: Poor conductors, brittle, dull

• Metalloids: Properties intermediate between metals and nonmetals

Atomic Structure and Isotopes

Atoms consist of protons, neutrons, and electrons. Isotopes are atoms of the same element with different numbers of neutrons.

• Atomic number: Number of protons

• Mass number: Number of protons + neutrons

• Example: $^{24}\text{Mg}$ and $^{25}\text{Mg}$ are isotopes of magnesium

Chemical Formulas and Nomenclature

Chemical formulas represent the elements in a compound and their ratios. Nomenclature is the system for naming compounds.

• Empirical formula: Simplest whole-number ratio of elements

• Molecular formula: Actual number of atoms of each element

• Example: H2O (water), NaCl (sodium chloride)

Law of Definite Proportions

The law of definite proportions states that a chemical compound always contains the same proportion of elements by mass.

• Example: Water always contains 2 parts hydrogen to 1 part oxygen by atom count

Chemical Equations and Stoichiometry

Chemical equations represent chemical reactions. Stoichiometry involves calculations based on balanced chemical equations.

• Balancing equations: Ensure the same number of each atom on both sides

• Example: $2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}$

• Mole concept: 1 mole = $6.022 \times 10^{23}$ particles

• Mass-mole conversions: $\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar mass (g/mol)}}$

States of Matter and Intermolecular Forces

Matter exists as solids, liquids, or gases. Intermolecular forces determine physical properties.

• London dispersion forces: Present in all molecules

• Dipole-dipole interactions: Occur in polar molecules

• Hydrogen bonding: Strong dipole interaction involving H and N, O, or F

Solutions and Concentration Calculations

Solutions are homogeneous mixtures. Concentration is often expressed as molarity (M).

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

• Dilution: $M_1V_1 = M_2V_2$

• Example: To prepare 0.100 M KOH from 0.400 M KOH, use $V_1 = \frac{M_2V_2}{M_1}$

Acids, Bases, and pH Calculations

Acids release H+ ions in solution; bases release OH- ions. pH measures the acidity of a solution.

• pH formula: $\text{pH} = -\log[\text{H}^+]$

• Example: If $[\text{H}^+] = 1.0 \times 10^{-4}$, then $\text{pH} = 4$

Lewis Structures and Molecular Geometry

Lewis structures show the arrangement of electrons in molecules. Molecular geometry describes the shape of molecules.

• VSEPR theory: Electron pairs repel and determine molecular shape

• Example: H2O has a bent shape

Physical vs. Chemical Changes

Physical changes do not alter the chemical identity; chemical changes result in new substances.

• Physical change: Dissolving sugar in water

• Chemical change: Burning magnesium

Sample HTML Table: Types of Intermolecular Forces

Type	Description	Example
London Dispersion	Weak, present in all molecules	CH4
Dipole-Dipole	Between polar molecules	HCl
Hydrogen Bonding	Strong, H bonded to N, O, or F	H2O

Additional info:

• Some questions reference calculations and chemical equations; students should practice these using the formulas provided.

• Review includes topics from nearly all major introductory chemistry chapters, including atomic structure, chemical bonding, stoichiometry, solutions, acids/bases, and physical/chemical changes.