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

Chapter 1. Matter, Measurement, and Problem Solving

1.1 Atoms and Molecules

Atoms and molecules are the fundamental building blocks of matter. Chemistry studies their properties, interactions, and transformations.

• Atom: The smallest unit of an element that retains its chemical identity.

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

• Atoms are often represented as spheres in models.

1.2 The Scientific Approach to Knowledge

The scientific method is a systematic way to investigate phenomena, acquire new knowledge, or correct and integrate previous knowledge.

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

• Law: A statement based on repeated experimental observations that describes some aspect of the 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 by its physical state and composition.

• States of Matter: Solid, liquid, gas.

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

• Mixture: A combination of two or more substances where each retains its identity.

• 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: distillation, decantation, filtration.

1.4 Physical and Chemical Changes and Properties

Physical and chemical changes alter matter in different ways.

• Physical Change: Alters the state or appearance but not composition (e.g., melting, boiling).

• Chemical Change: Alters the composition, forming new substances (e.g., rusting, combustion).

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

• Chemical Property: Describes the 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 systems.

• 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

• All measurements have some uncertainty; the last digit is estimated.

• Significant Figures: Digits that carry meaning in a measurement.

• Rules for determining significant figures in calculations.

• Accuracy: Closeness to the true value.

• Precision: Reproducibility of measurements.

1.8 Solving Chemical Problems

• Use dimensional analysis and conversion factors.

• 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 atoms as indivisible units.

• Modern science emphasizes experimentation and the scientific method.

2.3 Modern Atomic Theory and the Laws That Led to It

• Law of Conservation of Mass: Mass is conserved in chemical reactions.

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

• Law of Multiple Proportions: Elements can combine in different ratios to form different compounds.

• Dalton's atomic theory: matter is composed of atoms, which combine in simple ratios to form compounds.

2.4 The Discovery of the Electron

• J. J. Thomson's cathode ray experiments showed the existence of electrons.

• 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 nuclear model of the atom.

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.

• Mass number (A): Number of protons plus 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.

• Groups (columns) contain elements with similar properties.

• Metals, nonmetals, metalloids: classified by properties.

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

• Elements form ions with predictable charges based on their group.

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

• Atomic mass is the weighted average of all isotopes of an element.

• 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.

• Relationship 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: Transfer of electrons from one atom to another, forming ions (e.g., NaCl).

• Covalent bond: Sharing of electrons between atoms (e.g., H2O).

3.3 Representing Compounds: Chemical Formulas and Molecular Models

• Empirical formula: Simplest whole-number ratio of elements.

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

• Structural formula: Shows how atoms are bonded.

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

• Characteristic colors for elements in models (e.g., oxygen = red, carbon = black).

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.

3.5 Ionic Compounds: Formulas and Names

• Ionic compounds are formed from metals and nonmetals.

• Rules for writing formulas: balance charges to achieve electrical neutrality.

• Rules for naming: cation (metal) first, then anion (nonmetal with -ide ending).

• Common ions and their charges should be memorized.

3.6 Molecular Compounds: Formulas and Names

• Molecular compounds are formed from nonmetals.

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

• Binary acids and oxyacids have specific naming conventions.

3.7 Summary of Inorganic Nomenclature

• Understand and use the flowchart for naming compounds.

3.8 Formula Mass and the Mole Concept for Compounds

• Formula mass: Sum of atomic masses in a chemical formula.

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

• Conversions:

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

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

3.9 Composition of Compounds

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

• 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.

• Empirical formula: simplest whole-number ratio of atoms.

• Molecular formula: multiple of the empirical formula.

• Combustion analysis: used to determine empirical formulas of organic compounds.

3.11 Organic Compounds

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

• Alkanes: saturated hydrocarbons (single bonds).

• Alkenes: hydrocarbons with double bonds.

• Alkynes: hydrocarbons with triple bonds.

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

• Recognize common organic functional groups.

Material to be Memorized for the Exam

• 1 mol of particles = $6.022 \times 10^{23}$ particles (Avogadro's number)

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

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