General Chemistry I: Exam 1 Study Guide – Fundamental Concepts and Skills

General Chemistry I: Exam 1 Study Guide

Law of Multiple Proportions, Law of Definite Proportions, and Law of Conservation of Mass

These fundamental laws describe how elements combine to form compounds and how mass is conserved in chemical reactions.

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

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

• Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction; the total mass of reactants equals the total mass of products.

• Example: Water (H2O) always contains 2 hydrogen atoms for every 1 oxygen atom by mass.

The Scientific Method

The scientific method is a systematic approach to research and experimentation in science.

• Steps: Observation, Hypothesis, Experiment, Data Collection, Analysis, Conclusion.

• Example: Observing a chemical reaction, forming a hypothesis about the outcome, testing it, and analyzing results.

Significant Figures and Scientific Notation

Significant figures (sig figs) reflect the precision of a measured quantity. Scientific notation expresses numbers as a product of a coefficient and a power of ten.

• Identifying Sig Figs: All nonzero digits are significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros are significant if after a decimal point.

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

• Example: 0.00450 has 3 significant figures; 4.50 × 10-3.

Mathematical Operations with Significant Figures

Rules for handling significant figures differ for addition/subtraction and multiplication/division.

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

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

Units, Abbreviations, and SI Prefixes

Understanding and converting between units is essential in chemistry.

• Base Units: Length (meter, m), Mass (kilogram, kg), Time (second, s), Temperature (kelvin, K), Amount (mole, mol), Volume (liter, L), Density (g/cm3 or kg/m3).

• Common Prefixes: mega (M, 106), kilo (k, 103), deci (d, 10-1), centi (c, 10-2), milli (m, 10-3), micro (μ, 10-6), nano (n, 10-9), pico (p, 10-12).

• Conversion Example: 2.54 cm = 1 inch; 1 mL = 1 cm3.

Dimensional Analysis and Unit Conversions

Dimensional analysis uses conversion factors to solve problems involving unit changes.

• Method: Multiply by conversion factors so units cancel appropriately.

• Example: To convert 10 inches to centimeters: $10 \text{ in} \times \frac{2.54 \text{ cm}}{1 \text{ in}} = 25.4 \text{ cm}$

Density and Its Equation

Density is the mass per unit volume of a substance.

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

• Units: g/cm3, kg/m3

Temperature Conversions

Temperature can be measured in Celsius (°C), Fahrenheit (°F), or Kelvin (K).

• Equations:

  • $K = ^\circ C + 273.15$

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

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

States of Matter and Their Differences

Matter exists in three primary states: solid, liquid, and gas.

• Solids: Definite shape and volume.

• Liquids: Definite volume, indefinite shape.

• Gases: Indefinite shape and volume.

Pure Substances vs. Mixtures

Understanding the classification of matter is fundamental in chemistry.

• Pure Substances: Have a fixed composition (elements or compounds).

• Mixtures: Physical combinations of two or more substances; can be homogeneous (uniform) or heterogeneous (non-uniform).

Elements, Compounds, and Mixtures

Definitions and distinctions among elements, compounds, and mixtures.

• Element: A substance that cannot be broken down into simpler substances by chemical means.

• Compound: A substance composed of two or more elements chemically combined in fixed proportions.

• Mixture: A combination of two or more substances not chemically combined.

Separation Techniques

Methods for separating mixtures into their components.

• Decanting: Pouring off a liquid to separate it from a solid.

• Distillation: Separating substances based on differences in boiling points.

• Filtration: Separating solids from liquids using a filter.

Physical and Chemical Changes/Properties

Distinguishing between physical and chemical changes and properties.

• Physical Change: Does not alter the chemical composition (e.g., melting, boiling).

• Chemical Change: Alters the chemical composition (e.g., burning, rusting).

• Extensive Properties: Depend on the amount of substance (e.g., mass, volume).

• Intensive Properties: Independent of amount (e.g., density, boiling point).

Structure of the Atom and Subatomic Particles

Atoms are composed of protons, neutrons, and electrons.

• Proton (p): Positively charged, located in the nucleus.

• Neutron (n): Neutral, located in the nucleus.

• Electron (e): Negatively charged, located outside the nucleus.

• Discovery: Protons and neutrons discovered by Rutherford and Chadwick, electrons by Thomson.

Atomic Number, Mass Number, Isotopes, Ions

Atoms are identified by their atomic number and mass number; isotopes and ions are variations of atoms.

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

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

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

• Ion: Atom or molecule with a net electric charge due to loss or gain of electrons.

• Cation: Positively charged ion.

• Anion: Negatively charged ion.

• Isotope Notation: $^{A}_{Z}\text{X}$, where X is the element symbol.

Energy: Definitions and Types

Energy is the capacity to do work or transfer heat.

• Thermal Energy: Energy associated with temperature.

• Kinetic Energy: Energy of motion. $KE = \frac{1}{2}mv^2$

• Potential Energy: Stored energy due to position or composition.

Accuracy, Precision, and Types of Error

Understanding measurement quality is crucial in scientific experiments.

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

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

• Random Error: Error that varies unpredictably.

• Systematic Error: Error that is consistent and repeatable.

Periodic Table: Elements, Groups, and Periods

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

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

• Periods: Horizontal rows.

• Group Names: IA (alkali metals), IIA (alkaline earth metals), VIA, VIIA (halogens), VIIIA (noble gases), and B groups (transition metals).

• Metals, Nonmetals, Metalloids: Metals are typically shiny and conductive; nonmetals are varied; metalloids have intermediate properties.

• Representative Elements: Groups 1A–8A (IA–VIIIA).

• Transition Elements: B groups.

Element Names and Symbols

Each element has a unique symbol and name.

• Examples: 1 (H), 2 (He), 20 (Ca), 22 (Ti), 30 (Zn), 33 (As), 36 (Kr), 38 (Sr), 46 (Pd), 47 (Ag), 50 (Sn), 53 (I), 54 (Xe), 56 (Ba), 78 (Pt), 79 (Au), 80 (Hg), 82 (Pb).

Diatomic Molecules

Certain elements exist naturally as molecules composed of two atoms.

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

• States at Room Temperature: Most are gases; Br2 is a liquid, I2 is a solid.

Compounds: Ionic and Molecular

Compounds can be classified as ionic or molecular based on their bonding.

• Ionic Compounds: Formed from metals and nonmetals; consist of ions held together by electrostatic forces.

• Molecular Compounds: Formed from nonmetals; consist of molecules held together by covalent bonds.

• Identifying Compounds: Use element types and bonding patterns.

Naming and Writing Formulas for Compounds

Rules exist for naming and writing formulas for both ionic and molecular compounds.

• Ionic Compounds: Name the cation first, then the anion. For metals with variable charge, indicate charge with Roman numerals.

• Polyatomic Ions: Common examples include ammonium (NH4+), carbonate (CO32-), sulfate (SO42-), nitrate (NO3-), phosphate (PO43-), hydroxide (OH-), acetate (C2H3O2-), cyanide (CN-), chlorite (ClO2-), chlorate (ClO3-), perchlorate (ClO4-).

• Molecular Compounds: Use prefixes (mono-, di-, tri-, etc.) to indicate the number of each atom.

• Acids: Binary acids (e.g., HCl) and oxyacids (e.g., H2SO4).

Common Polyatomic Ions Table

Name	Formula	Charge
Ammonium	NH4	+
Carbonate	CO3	2-
Sulfate	SO4	2-
Nitrate	NO3	-
Phosphate	PO4	3-
Hydroxide	OH	-
Acetate	C2H3O2	-
Cyanide	CN	-
Chlorite	ClO2	-
Chlorate	ClO3	-
Perchlorate	ClO4	-
Bicarbonate	HCO3	-
Nitrite	NO2	-

Mole Conversions and Calculations

The mole is a fundamental unit for counting particles in chemistry.

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

• Conversions:

  • Particles ↔ Moles: $\text{Moles} = \frac{\text{Particles}}{6.022 \times 10^{23}}$

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

  • Moles of element ↔ Moles of compound: Use subscripts in chemical formulas.

• Example: Convert 12.0 g of carbon to moles: $\frac{12.0 \text{ g}}{12.01 \text{ g/mol}} = 1.00 \text{ mol}$

Empirical and Molecular Formulas

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

• Empirical Formula: Simplest ratio (e.g., CH2O for glucose).

• Molecular Formula: Actual composition (e.g., C6H12O6 for glucose).

Additional info: Some content was inferred and expanded for clarity and completeness, such as the full list of polyatomic ions, the structure of the periodic table, and the rules for naming compounds.