CHEM 121 Exam 1 Study Guide: Matter, Measurement, Atomic Structure, and Stoichiometry

Chapter 1: Introduction to Chemistry, Matter, Energy, and Measurement

Definition of Chemistry

Chemistry is the scientific study of matter, its properties, the changes it undergoes, and the energy associated with those changes.

• Matter: Anything that has mass and occupies space.

• Energy: The capacity to do work or produce heat.

Types of Matter and Changes in Matter

Matter can be classified based on its physical and chemical properties.

• Pure Substances: Elements and compounds with fixed composition.

• Mixtures: Physical combinations of two or more substances.

• Physical Changes: Changes that do not alter the composition (e.g., melting, boiling).

• Chemical Changes: Changes that alter the composition (e.g., combustion, rusting).

• Example: Ice melting is a physical change; iron rusting is a chemical change.

Types of Energy and Changes in Energy

Energy exists in various forms and can be transformed during chemical and physical processes.

• Kinetic Energy: Energy due to motion. $KE = (\frac{1}{2})mv^2$

• Potential Energy: Energy due to position or composition.

• Work: Force applied over a distance. $Work = F \times d$

• Example: A moving car has kinetic energy; a stretched spring has potential energy.

Measurement, SI Units, and Metric Prefixes

Scientific measurements use the International System of Units (SI) and metric prefixes to express quantities.

• Length: meter (m)

• Mass: kilogram (kg)

• Temperature: Kelvin (K)

• Time: second (s)

• Amount of Substance: mole (mol)

• Energy: joule (J)

• Volume: 1 dm3 = 1 L; 1 cm3 = 1 mL

Metric Prefixes Table

Prefix	Symbol	Exponent	Meaning
peta-	P	1015	1,000,000,000,000,000
tera-	T	1012	1,000,000,000,000
giga-	G	109	1,000,000,000
mega-	M	106	1,000,000
kilo-	K	103	1,000
deci-	d	10-1	1/10
centi-	c	10-2	1/100
milli-	m	10-3	1/1,000
micro-	µ	10-6	1/1,000,000
nano-	n	10-9	1/1,000,000,000
pico-	p	10-12	1/1,000,000,000,000
femto-	f	10-15	1/1,000,000,000,000,000

Precision, Accuracy, and Significant Figures

Measurement quality is described by precision and accuracy, and significant figures reflect the certainty of measurements.

• Precision: How close repeated measurements are to each other.

• Accuracy: How close a measurement is to the true value.

• Standard Deviation: Quantifies the spread of data.

• Significant Figures: Digits that carry meaning in a measurement; rules determine how many to report.

Conversion Factors and Dimensional Analysis

Conversion factors are used to change units in calculations, ensuring consistency and accuracy.

• Example: To convert 125 lb to grams: $125 \text{ lb} \times \frac{453.59 \text{ g}}{1 \text{ lb}}$

Chapter 2: Atoms, Molecules, and Ions

Laws of Chemical Combination

Several fundamental laws govern the behavior of matter in chemical reactions.

• Law of Definite Composition: A compound always contains the same elements in the same proportion by mass.

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

• Law of Multiple Proportions: When two elements form more than one compound, the masses of one element combine with a fixed mass of the other in ratios of small whole numbers.

Dalton’s Atomic Theory

Dalton proposed a model for the atom based on experimental evidence.

• Elements are composed of tiny particles called atoms.

• Atoms of the same element are identical; atoms of different elements are different.

• Atoms combine in simple whole-number ratios to form compounds.

• Atoms are indivisible in chemical reactions.

Subatomic Particles and Atomic Structure

Atoms consist of protons, neutrons, and electrons, each with distinct properties.

• Proton: Positive charge (+1), mass ≈ 1 amu.

• Neutron: Neutral charge (0), mass ≈ 1 amu.

• Electron: Negative charge (-1), mass ≈ 1/1836 amu.

• Discovery: Electrons (Thomson), nucleus (Rutherford).

• Rutherford’s Model: Atom has a dense, positively charged nucleus with electrons orbiting.

• Modern Model: Electrons occupy regions called orbitals.

Atomic Number, Atomic Symbols, and Isotopes

Atoms are identified by their atomic number and symbol; isotopes are atoms of the same element with different numbers of neutrons.

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

• Atomic Mass Unit (amu): Standard unit for atomic mass.

• Isotopes: Atoms with same Z but different mass numbers (A).

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

Atomic Weight and Periodic Table

Atomic weight is the average mass of an element’s isotopes, weighted by abundance.

• Atomic Weight: $\text{Atomic weight} = \sum (\text{isotope mass} \times \text{fractional abundance})$

• Periodic Table: Arranges elements by increasing atomic number; groups elements with similar properties.

Molecules, Molecular Compounds, and Chemical Formulas

Molecules are groups of atoms bonded together; compounds are substances with two or more elements.

• Diatomic Molecules: H2, O2, N2, etc.

• Empirical Formula: Simplest ratio of elements.

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

• Example: Glucose: Empirical formula CH2O, molecular formula C6H12O6.

Ions and Ionic Compounds

Ions are charged particles formed by loss or gain of electrons; ionic compounds are formed from cations and anions.

• Cation: Positively charged ion (e.g., Na+).

• Anion: Negatively charged ion (e.g., Cl-).

• Polyatomic Ions: Groups of atoms with a charge (e.g., SO42-).

• Example: NaCl is an ionic compound.

Chemical Nomenclature

Rules exist for naming inorganic compounds, acids, and binary molecular compounds.

• Binary Molecular Compounds: Use prefixes to indicate number of atoms.

• Prefixes Table:

Number	Prefix
1	Mono-
2	Di-
3	Tri-
4	Tetra-
5	Penta-
6	Hexa-
7	Hepta-
8	Octa-
9	Nona-
10	Deca-

• Acids: Naming depends on the anion present (e.g., HCl is hydrochloric acid).

• Organic Compounds: Alkanes (single bonds) and alcohols (contain -OH group).

• Isomers: Compounds with same formula but different structures.

Chapter 3: Chemical Reactions and Stoichiometry

Chemical Equations and Balancing

Chemical equations represent reactions; balancing ensures conservation of mass.

• Reactants: Substances consumed.

• Products: Substances formed.

• Balancing: Adjust coefficients to equalize atom counts.

• Types of Reactions: Combination, decomposition, combustion, etc.

Formula and Molecular Weights, Percent Composition

Formula weight is the sum of atomic weights in a compound; percent composition shows the mass percentage of each element.

• Formula Weight: $\text{Formula weight} = \sum \text{atomic weights}$

• Percent Composition: $\% \text{Composition} = \left( \frac{\text{number of atoms} \times \text{atomic weight}}{\text{formula weight}} \right) \times 100$

• Example: Sucrose (C12H22O11): Calculate %C, %H, %O by mass.

Avogadro’s Number, Molar Mass, and Mole Relationships

Avogadro’s number links the mole to the number of particles; molar mass is the mass of one mole of a substance.

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

• Molar Mass: Mass (g) of 1 mole of substance.

• Conversions: Use Avogadro’s number and molar mass to convert between grams, moles, and particles.

Empirical and Molecular Formulas

Empirical formula is the simplest ratio; molecular formula is the actual composition.

• Whole Number Multiple: $\text{Whole number multiple} = \frac{\text{molecular weight}}{\text{empirical formula weight}}$

Stoichiometry and Quantitative Relationships

Stoichiometry uses balanced equations to relate quantities of reactants and products.

• Mole Ratio: Coefficients in equations show ratios.

• Limiting Reagent: Reactant that determines maximum product.

• Theoretical Yield: Maximum possible product.

• Actual Yield: Amount actually obtained.

• Percent Yield: $\% \text{Yield} = \left( \frac{\text{actual yield}}{\text{theoretical yield}} \right) \times 100$

Sample Calculations and Practice Problems

• Unit Conversion: Converting mass, length, volume using SI units and conversion factors.

• Atomic Scale: Calculating how many atoms fit across a given distance (e.g., diameter of a dime vs. silver atom).

• Percent Composition: Calculating mass percentages in compounds (e.g., sucrose).

Equations and Equivalencies

• Density: $d = \frac{m}{v}$

• Density of Water: 1 g/cm3

• Work: $Work = F \times d$

• Kinetic Energy: $KE = (\frac{1}{2})mv^2$

• Temperature Conversions:

  • $0^\circ C = 273 \text{ K} = 32^\circ F$ (freezing point of water)

  • $100^\circ C = 373 \text{ K} = 212^\circ F$ (boiling point of water)

  • $^\circ F = (\frac{9}{5})(^\circ C) + 32$

  • $^\circ C = (\frac{5}{9})(^\circ F - 32)$

• Volume Equivalencies: 1 dm3 = 1 L; 1 cm3 = 1 mL

Practice and Exam Preparation

• Review textbook exercises and sample problems for Chapters 1-3.

• Practice balancing equations, calculating percent composition, and converting units.

• Memorize common ions and their charges for nomenclature and formula writing.

Additional info: For full mastery, students should practice with textbook problems and memorize common ions and polyatomic ions, as well as review the rules for significant figures and chemical nomenclature.