CHEM 151 Exam 1 Study Guide: Essential Concepts and Skills

Unit Conversions & Scientific Notation

Temperature Conversions

Understanding how to convert between different temperature units is fundamental in chemistry. The three main units are Celsius (°C), Kelvin (K), and Fahrenheit (°F).

• Celsius to Kelvin: Add 273.15 to the Celsius value. $K = \text{°C} + 273.15$

• Celsius to Fahrenheit: Multiply by 9/5 and add 32. $\text{°F} = \frac{9}{5} \times \text{°C} + 32$

• Fahrenheit to Celsius: Subtract 32 and multiply by 5/9. $\text{°C} = \frac{5}{9} \times (\text{°F} - 32)$

Scientific Notation

Scientific notation expresses numbers as a coefficient (1 ≤ coefficient < 10) multiplied by a power of ten.

• Convert decimals to scientific notation: Move the decimal point so only one nonzero digit remains to the left.

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

• Example: 0.00045 = $4.5\cdot10^{-4}$

Significant Figures

Rules for Counting Significant Figures

Significant figures reflect the precision of a measurement.

• Leading zeros: Do not count as significant (e.g., 0.0025 has 2 significant figures).

• Trailing zeros after a decimal: Count as significant (e.g., 2.300 has 4 significant figures).

• Trailing zeros without a decimal: Ambiguous; may or may not be significant.

• Exact numbers: Have infinite significant figures (e.g., counted objects).

• Rounding: Round to the specified number of significant figures as required.

Metric Prefixes & Mass Conversions

Common Metric Prefixes

Metric prefixes indicate multiples or fractions of base units.

• kilo- (k): $10^3$ (1,000)

• milli- (m): $10^{-3}$ (0.001)

• micro- (μ): $10^{-6}$ (0.000001)

• nano- (n): $10^{-9}$ (0.000000001)

Dimensional Analysis

Dimensional analysis is used to convert between units using conversion factors.

• Example: To convert 5.0 g to mg: $5.0 \text{ g} \times \frac{1000 \text{ mg}}{1 \text{ g}} = 5000 \text{ mg}$

Density

Definition and Formula

Density is the mass per unit volume of a substance.

• Formula: $\text{density} = \frac{\text{mass}}{\text{volume}}$

• Units: Typically g/cm3 or g/mL

• Solving for mass or volume: Rearranged formulas: $\text{mass} = \text{density} \times \text{volume}$ $\text{volume} = \frac{\text{mass}}{\text{density}}$

Atomic Structure & Isotopes

Basic Atomic Structure

Atoms consist of protons, neutrons, and electrons.

• Protons: Equal to the atomic number (Z)

• Neutrons: $\text{mass number} - \text{atomic number}$

• Electrons: Equal to protons in a neutral atom

Isotopes and Average Atomic Mass

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

• Average atomic mass: Weighted average based on percent abundance. $\text{Average atomic mass} = \sum (\text{isotope mass} \times \text{fractional abundance})$

• Example: If isotope A (10 amu, 60%) and isotope B (12 amu, 40%): $10 \times 0.60 + 12 \times 0.40 = 6 + 4.8 = 10.8$ amu

Periodic Table Basics

Groups and Element Types

The periodic table organizes elements by properties.

• Alkali metals: Group 1A (e.g., Li, Na, K)

• Noble gases: Group 8A (e.g., He, Ne, Ar)

• Hydrogen: Not a metal; unique properties

• Nonmetals: Elements like C, N, O, F, P, S, Cl

Polyatomic Ions

Common Polyatomic Ions

Polyatomic ions are charged groups of atoms.

Ion	Name
CO32−	carbonate
NO3−	nitrate
ClO3−	chlorate
PO43−	phosphate

Naming Compounds & Acids

Ionic Compounds

Ionic compounds consist of metals and nonmetals.

• Naming: Name the metal first, then the anion.

• Transition metals: Use Roman numerals to indicate charge (e.g., FeCl2 is iron(II) chloride).

Molecular Compounds

• Prefixes: mono-, di-, tri-, tetra-, penta-, hexa-, etc. (e.g., CO2 is carbon dioxide)

• No Roman numerals: Charges are not specified.

Acids

• Example: H3PO4 is phosphoric acid

Reaction Types

Classification of Reactions

Chemical reactions can be classified by their patterns.

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

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

• Combustion: A substance reacts with oxygen, often producing CO2 and H2O.

• Double replacement: Two compounds exchange ions to form new compounds.

Balancing Chemical Equations

Principles of Balancing

Balancing equations ensures the conservation of mass.

• Change coefficients only: Never alter subscripts.

• Use coefficients for mole ratios: Essential for stoichiometry.

• Example: $2H_2 + O_2 \rightarrow 2H_2O$

Molar Mass & Counting Particles

Molar Mass Calculation

Molar mass is the mass of one mole of a substance, calculated by summing atomic masses.

• Include parentheses: For compounds with grouped atoms (e.g., Ca(NO3)2)

• Example: Molar mass of H2O: $2 \times 1.01 + 16.00 = 18.02$ g/mol

Counting Particles

• Molecules ↔ moles ↔ grams: Use Avogadro's number ($6.022 \times 10^{23}$ molecules/mol)

• Counting atoms: Multiply the number of atoms per molecule by the number of molecules.

Stoichiometry

Stoichiometric Calculations

Stoichiometry involves quantitative relationships in chemical reactions.

• Mass-to-mass problems: Convert grams of one substance to grams of another using mole ratios.

• Balanced equations: Required for correct mole ratios.

• Limiting reagent: The reactant that determines the maximum amount of product.

• Percent yield: Measures efficiency of a reaction. $\text{Percent yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100$

• General setup: grams → moles → mole ratio → grams

Chemical vs Physical Changes

Types of Changes

Understanding the difference between chemical and physical changes is essential.

• Chemical changes: Result in new substances (e.g., rusting, decomposition reactions)

• Physical changes: Do not alter chemical identity (e.g., freezing, compression)