General Chemistry I: Study Guide – Matter, Measurement, Atoms, and Compounds

Chapter 1. Matter, Measurement, and Problem Solving

1.1 Atoms and Molecules

Chemistry is the study of matter, its properties, and the changes it undergoes. Atoms and molecules are the fundamental building blocks of matter.

• Atom: The smallest unit of an element that retains the properties of that element.

• Molecule: A group of two or more atoms held together by chemical bonds.

• Atoms are often represented as spheres in models to illustrate their structure.

1.2 The Scientific Approach to Knowledge

The scientific method is a systematic approach to understanding the natural world through observation and experimentation.

• Hypothesis: A tentative explanation for an observation, which can be tested by experiments.

• Law: A statement that summarizes a vast number of experimental observations and describes or predicts some aspect of the natural world (e.g., Law of Conservation of Mass).

• Theory: A well-substantiated explanation of some aspect of the natural world that can incorporate laws, hypotheses, and facts.

• Scientific theories are built from strong experimental evidence.

1.3 The Classification of Matter

Matter can be classified based on its physical state and composition.

• States of Matter: Solid, liquid, and gas.

• Crystalline vs. Amorphous Solids: Crystalline solids have an ordered structure; amorphous solids do not.

• Mixture: A combination of two or more substances that retain their individual properties.

• Pure Substance: Matter with a fixed composition (elements and compounds).

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

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

• Homogeneous Mixture: Uniform composition throughout (solution).

• Heterogeneous Mixture: Non-uniform composition.

• Separation methods include distillation, decantation, and filtration.

1.4 Physical and Chemical Changes and Properties

Physical and chemical changes alter matter in different ways.

• Physical Change: Alters the form or appearance of matter but does not change its composition (e.g., melting, boiling).

• Chemical Change: Alters the composition of matter, resulting in new substances (e.g., rusting, combustion).

• Physical Property: Can be observed without changing the substance (e.g., color, melting point).

• Chemical Property: Describes a substance's ability to undergo chemical changes (e.g., flammability).

1.5 Energy: A Fundamental Part of Physical and Chemical Change

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

• Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed.

1.6 The Units of Measurement

Measurements in chemistry require standardized units.

• Know the differences between the English, metric, and SI (International System) units.

• SI base units: meter (length), kilogram (mass), second (time), kelvin (temperature).

• Temperature scales: Fahrenheit, Celsius, Kelvin.

• Temperature conversions:

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

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

  • $K = ^\circ C + 273.15$

• Use of metric prefixes (e.g., kilo-, centi-, milli-).

1.7 The Reliability of a Measurement

• Measurements have uncertainty; the last digit is estimated.

• Significant Figures: Digits in a measurement that are known with certainty plus one estimated digit.

• Rules for determining significant figures in calculations.

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

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

1.8 Solving Chemical Problems

• Use dimensional analysis and conversion factors to solve problems.

• Problem-solving strategy: sort, strategize, solve, check.

• Convert between units and rearrange equations to solve for unknowns.

Chapter 2. Atoms and Elements

2.1 Brownian Motion: Atoms Confirmed

• Brownian motion: random movement of particles suspended in a fluid, evidence for atoms.

2.2 Early Ideas about the Building Blocks of Matter

• Ancient Greeks proposed that matter is composed of indivisible particles (atoms).

• Development of the scientific method led to modern atomic theory.

2.3 Modern Atomic Theory and the Laws That Led to It

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

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

• 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.

• Dalton's atomic theory: all matter is composed of atoms, atoms of the same element are identical, atoms combine in simple ratios to form compounds, and atoms are rearranged in chemical reactions.

2.4 The Discovery of the Electron

• J. J. Thomson's cathode ray tube experiments provided evidence for the electron.

• Robert Millikan's oil-drop experiment measured the charge of the electron.

2.5 The Structure of the Atom

• Atoms consist of a nucleus (protons and neutrons) and electrons.

• Rutherford's gold-foil experiment demonstrated the existence of a small, dense nucleus.

2.6 Subatomic Particles: Protons, Neutrons, and Electrons in Atoms

• Proton: Positively charged particle in the nucleus.

• Neutron: Neutral particle in the nucleus.

• Electron: Negatively charged particle outside the nucleus.

• 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 charge (cation: positive, anion: negative).

2.7 Finding Patterns: The Periodic Law and the Periodic Table

• The periodic table organizes elements by increasing atomic number.

• Elements with similar properties are grouped in columns (groups or families).

• Metals, nonmetals, and metalloids are distinguished by their properties.

• Special groups: noble gases, alkali metals, alkaline earth metals, halogens.

2.8 Atomic Mass: The Average Mass of an Element's Atoms

• Atomic mass is the weighted average of the masses of an element's isotopes.

• Mass spectrometry is used to determine atomic and molecular masses.

2.9 Molar Mass: Counting Atoms by Weighing Them

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

• Molar mass: mass of one mole of a substance (g/mol).

• Conversions between mass, moles, and number of particles:

$\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

$\text{number of particles} = \text{moles} \times 6.022 \times 10^{23}$

Chapter 3. Molecules and Compounds

3.1 Hydrogen, Oxygen, and Water

• Compounds are substances composed of two or more elements in fixed ratios (e.g., H2O, NaCl).

3.2 Chemical Bonds

• Ionic Bond: Electrostatic attraction between oppositely charged ions (metal + nonmetal).

• Covalent Bond: Sharing of electron pairs between atoms (nonmetal + nonmetal).

• Formation of ionic compounds from elements.

• Formation of covalent bonds in molecules.

3.3 Representing Compounds: Chemical Formulas and Molecular Models

• Empirical Formula: Simplest whole-number ratio of elements in a compound.

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

• Structural Formula: Shows how atoms are bonded in a molecule.

• Ball-and-stick and space-filling models represent molecular structure.

3.4 An Atomic-Level View of Elements and Compounds

• Elements can be atomic (single atoms) or molecular (diatomic or polyatomic molecules).

• Ionic compounds are composed of ions arranged in a lattice; molecular compounds are composed of discrete molecules.

3.5 Ionic Compounds: Formulas and Names

• Ionic compounds are common in Earth's crust and minerals.

• Rules for writing formulas:

  • Combine cations and anions in ratios that yield a neutral compound.

  • Use the charges of ions to determine the formula.

• Rules for naming ionic compounds:

  • Name the cation first, then the anion.

  • For transition metals, indicate the charge with Roman numerals.

3.6 Molecular Compounds: Formulas and Names

• Rules for naming molecular compounds:

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

  • Name the more metallic element first.

• Binary acids and oxyacids have specific naming conventions.

3.7 Summary of Inorganic Nomenclature

• Understand the flowchart for naming inorganic compounds.

3.8 Formula Mass and the Mole Concept for Compounds

• Formula Mass: Sum of the atomic masses of all atoms in a formula unit.

• Molar Mass: Mass of one mole of a compound.

• Conversions between mass, moles, and molecules:

$\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

$\text{number of molecules} = \text{moles} \times 6.022 \times 10^{23}$

3.9 Composition of Compounds

• Mass Percent: $\frac{\text{mass of element in 1 mol of compound}}{\text{mass of 1 mol of compound}} \times 100\%$

• Use mass percent as a conversion factor.

3.10 Determining a Chemical Formula from Experimental Data

• Convert masses to moles and calculate mole ratios to determine empirical formulas.

• Determine empirical and molecular formulas from experimental data.

• Combustion analysis is used to determine the composition of organic compounds.

3.11 Organic Compounds

• Organic compounds contain carbon and hydrogen, often with other elements.

• Alkanes, alkenes, and alkynes are classes of hydrocarbons.

• Know the names and formulas of the first ten alkanes.

• Identify common organic functional groups.

Material to be Memorized for the Exam

• 1 mol of particles = $6.022 \times 10^{23}$ particles

• Prefix multipliers: giga, mega, kilo, deci, centi, milli, micro, nano

• Monoatomic anions (Table 3.2) and polyatomic ions (Table 3.4)