General Chemistry I: Exam 1 Study Guide

General Chemistry I: Exam 1 Study Guide

Overview

This study guide outlines the key topics and skills required for success in the first exam of a General Chemistry I course. The material covers foundational concepts in measurement, atomic structure, chemical nomenclature, and stoichiometry, as well as essential problem-solving skills.

Chapter 1: Measurement and Units

Metric System

• Units for Mass, Volume, Length: The metric system uses standardized units such as grams (g) for mass, liters (L) for volume, and meters (m) for length.

• Prefixes and Their Meanings: Common prefixes include kilo- (k, 103), centi- (c, 10-2), milli- (m, 10-3), and micro- (μ, 10-6).

• Conversion: To convert between units, multiply or divide by powers of ten according to the prefix.

• Example: 1 kilometer (km) = 1,000 meters (m); 1 milligram (mg) = 0.001 grams (g).

Conversion Factors

• Basic English to Metric Conversion: Use conversion factors such as 1 inch = 2.54 cm.

• Single-Step and Multi-Step Problems: Set up conversion factors so that units cancel appropriately.

• Dose Calculations: Used in medicine to determine the correct amount of a substance.

• Density Conversions: Density is mass per unit volume ($d = \frac{m}{V}$).

• Example: Convert 5 inches to centimeters: $5 \text{ in} \times \frac{2.54 \text{ cm}}{1 \text{ in}} = 12.7 \text{ cm}$.

Significant Figures

• Definition: Significant figures are the digits in a measurement that are known with certainty plus one estimated digit.

• Rules:

  • Nonzero digits are always significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

  • Trailing zeros are significant only if there is a decimal point.

• Determining Significant Figures in Calculations:

  • Multiplication/Division: The result should have as many significant figures as the measurement with the fewest significant figures.

  • Addition/Subtraction: The result should have as many decimal places as the measurement with the fewest decimal places.

• Example: $2.34 \times 1.2 = 2.8$ (rounded to 2 significant figures).

Scientific Notation

• Definition: A way to express very large or very small numbers using powers of ten.

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

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

Temperature Conversion

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

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

• Example: 25°C = 298.15 K

Chapter 2: Atomic Structure and the Periodic Table

Atomic Theory and Subatomic Particles

• Atoms: The smallest unit of an element that retains its chemical properties.

• Subatomic Particles: Protons (positive charge), neutrons (neutral), electrons (negative charge).

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

• Mass Number (A): Total number of protons and neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons.

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

Reading the Periodic Table

• Groups: Vertical columns; elements in the same group have similar properties.

• Periods: Horizontal rows.

• Key Element Groups: Alkali metals, alkaline earth metals, transition metals, halogens, noble gases, main group elements.

• Example: Sodium (Na) is an alkali metal in Group 1.

Atomic Mass and Isotopic Abundance

• Atomic Mass: Weighted average of the masses of all isotopes of an element.

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

• Example: Chlorine has two main isotopes: $^{35}$Cl and $^{37}$Cl.

Chemical Bonding and Nomenclature

Bonding: Ionic vs. Covalent

• Ionic Bonds: Formed by transfer of electrons from metals to nonmetals.

• Covalent Bonds: Formed by sharing of electrons between nonmetals.

• Predicting Bond Type: Based on the position of elements in the periodic table.

• Example: NaCl is ionic; H2O is covalent.

Ionic Bonding and Nomenclature

• Cation: Positively charged ion (usually a metal).

• Anion: Negatively charged ion (usually a nonmetal).

• Polyatomic Ions: Ions composed of more than one atom (e.g., SO42-).

• Naming: Name the cation first, then the anion. Use Roman numerals for metals with multiple charges.

• Example: FeCl3 is iron(III) chloride.

Covalent Naming

• Prefixes: mono-, di-, tri-, tetra-, penta-, etc.

• Rules: The first element keeps its name; the second gets an -ide suffix.

• Example: CO2 is carbon dioxide.

Molecular Elements

• Diatomic Molecules: H2, N2, O2, F2, Cl2, Br2, I2.

Stoichiometry and Chemical Reactions

The Mole Concept

• Definition: A mole is 6.022 × 1023 particles (Avogadro's number).

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

• Example: 1 mole of H2O = 18.02 g.

Calculations Involving Moles

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

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

• Empirical and Molecular Formulas: Empirical formula is the simplest whole-number ratio; molecular formula is the actual number of atoms.

• Example: Empirical formula of C6H12O6 is CH2O.

Chemical Equations and Stoichiometry

• Balancing Chemical Equations: Ensure the same number of each atom on both sides of the equation.

• Stoichiometric Calculations: Use mole ratios from balanced equations to convert between reactants and products.

• Theoretical Yield: The maximum amount of product that can be formed from given reactants.

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

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

Tables and Reference Data

Common Metric Prefixes

Prefix	Symbol	Factor
kilo-	k	103
centi-	c	10-2
milli-	m	10-3
micro-	μ	10-6

Common Polyatomic Ions

Name	Formula	Charge
Sulfate	SO4	2-
Nitrate	NO3	1-
Carbonate	CO3	2-
Phosphate	PO4	3-
Ammonium	NH4	1+
Hydroxide	OH	1-
Acetate	C2H3O2	1-
Additional info: See course materials for a full list.

Common Diatomic Molecules

Element	Formula
Hydrogen	H2
Nitrogen	N2
Oxygen	O2
Fluorine	F2
Chlorine	Cl2
Bromine	Br2
Iodine	I2

Periodic Table Reference

• Refer to the provided periodic table for element names, symbols, and atomic numbers.

• Key constants:

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

  • Standard Temperature: 273.15 K = 0°C

Skills and Memorization

• Metric and English conversions

• Polyatomic ions (names, formulas, charges)

• Roman numerals for transition metals

• Diatomic molecules

• Scientific notation and calculator use