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

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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 not chemically bonded.
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 chemically combined.
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 the substance (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:
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 of matter.
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, atoms of each element are identical, atoms combine in simple ratios to form compounds.

2.4 The Discovery of the Electron

J. J. Thomson's cathode ray experiments discovered the electron.
Robert Millikan's oil-drop experiment measured the electron's charge.

2.5 The Structure of the Atom

Atoms consist of a nucleus (protons and neutrons) and electrons.
Rutherford's gold-foil experiment showed the nucleus is small and dense.

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 + neutrons.
Isotope: Atoms of the same element with different numbers of neutrons.
Ion: Charged atom (cation = positive, anion = negative).

2.7 Finding Patterns: The Periodic Law and the Periodic Table

The periodic table arranges elements by increasing atomic number.
Groups (columns) contain elements with similar properties.
Metals, nonmetals, metalloids distinguished by properties.
Special groups: noble gases, alkali metals, alkaline earth metals, halogens.
Elements form ions with predictable charges based on group.

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

Atomic mass is the weighted average of all isotopes.
Mass spectrometry measures atomic and molecular masses.

2.9 Molar Mass: Counting Atoms by Weighing Them

Avogadro's Number: particles per mole.
Relationship between mass, moles, and number of atoms:

Chapter 3. Molecules and Compounds

3.1 Hydrogen, Oxygen, and Water

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

3.2 Chemical Bonds

Ionic Bond: Transfer of electrons from one atom to another (metal + nonmetal).
Covalent Bond: Sharing of electrons 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.
Molecular Formula: Actual number of atoms of each element.
Structural Formula: Shows how atoms are bonded.
Ball-and-stick and space-filling models represent molecules visually.
Characteristic colors for elements in models (e.g., carbon = black, oxygen = red).

3.4 An Atomic-Level View of Elements and Compounds

Elements can be atomic or molecular (e.g., O2).
Compounds can be molecular or ionic.
Ionic compounds are composed of formula units; molecular compounds are composed of molecules.

3.5 Ionic Compounds: Formulas and Names

Ionic compounds are common in Earth's crust and minerals.
Rules for writing formulas: charges of ions must balance (electrical neutrality).
Rules for naming ionic compounds, including binary and polyatomic ions.

3.6 Molecular Compounds: Formulas and Names

Rules for naming molecular compounds (prefixes: mono-, di-, tri-, etc.).
Names and formulas for binary acids and oxyacids.

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 atomic masses in a chemical formula.
Molar Mass: Mass of one mole of a compound.
Calculate moles, mass, and number of molecules using:

3.9 Composition of Compounds

Mass Percent:
Use mass percent as a conversion factor.
Use chemical formulas as conversion factors in mole calculations.

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.
Understand combustion analysis for determining empirical formulas.

3.11 Organic Compounds

Organic compounds contain carbon and hydrogen, often with other elements.
Alkanes, alkenes, alkynes: hydrocarbons with single, double, or triple bonds, respectively.
Know the names and formulas of the first ten alkanes.
Identify common organic functional groups (e.g., alcohols, carboxylic acids).

Material to Memorize for the Exam

1 mol of particles = particles
Prefix multipliers: giga, mega, kilo, deci, centi, milli, micro, nano
Monoatomic anions (Table 3.2) and polyatomic ions (Table 3.4)