General Chemistry: Measurement, Matter, and Chemical Nomenclature Study Guide

Chapter 1: Introduction to Chemistry

Physical vs. Chemical Properties and Changes

Chemistry distinguishes between physical and chemical properties and changes. Understanding these differences is fundamental to classifying matter and predicting its behavior.

• Physical properties: Characteristics that can be observed or measured without changing the substance's identity (e.g., melting point, density, color).

• Chemical properties: Characteristics that describe a substance's ability to undergo chemical changes (e.g., flammability, reactivity).

• Physical change: A change that does not alter the chemical composition (e.g., melting ice).

• Chemical change: A process that results in the formation of new substances (e.g., rusting iron).

• Example: Boiling water is a physical change; burning wood is a chemical change.

Units of Measurement and Significant Figures

Accurate measurement is essential in chemistry. The International System of Units (SI) is used, and significant figures reflect the precision of measurements.

• SI base units: meter (m), kilogram (kg), second (s), mole (mol), kelvin (K), ampere (A), candela (cd).

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

• Rules: All nonzero digits are significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros are significant if there is a decimal point.

• Example: 0.00450 has three significant figures.

Dimensional Analysis and Metric Prefixes

Dimensional analysis is a method for converting between units using conversion factors. Metric prefixes indicate powers of ten.

• Common prefixes: kilo- (k, $10^3$), centi- (c, $10^{-2}$), milli- (m, $10^{-3}$), micro- (μ, $10^{-6}$), nano- (n, $10^{-9}$).

• Dimensional analysis: Multiply by conversion factors so units cancel appropriately.

• Example: To convert 5.0 cm to meters: $5.0\ \text{cm} \times \frac{1\ \text{m}}{100\ \text{cm}} = 0.050\ \text{m}$

Mathematical Operations with Significant Figures

Rules for 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.

• Example: $2.0 \times 3.00 = 6.0$ (2 significant figures)

Density

Density is a physical property defined as mass per unit volume.

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

• Units: g/cm3 or kg/L

• Example: If a cube of gold has a mass of 1.82 lb and a volume of 45.2 cm3, its density is $\frac{1.82\ \text{lb} \times 453.6\ \text{g/lb}}{45.2\ \text{cm}^3}$

Chapter 2: Atoms, Molecules, and Chemical Nomenclature

Atomic Theory and Subatomic Particles

Atoms are the basic units of matter, composed of protons, neutrons, and electrons.

• Proton: Positively charged particle in the nucleus.

• Neutron: Neutral particle in the nucleus.

• Electron: Negatively charged particle in orbitals around the nucleus.

• Atomic number (Z): Number of protons.

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

• Example: Nickel-62 has 28 protons, 34 neutrons, and 28 electrons.

Isotopes and Atomic Mass

Isotopes are atoms of the same element with different numbers of neutrons. Atomic mass is the weighted average of all isotopes.

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

• Example: $^{14}_{6}\text{C}$ is carbon-14.

Classification of Matter

Matter can be classified as pure substances or mixtures.

• Pure substance: Has a fixed composition (elements and compounds).

• Mixture: Physical combination of two or more substances (homogeneous or heterogeneous).

• Example: Salt water is a homogeneous mixture; sand and iron filings is a heterogeneous mixture.

Law of Constant Composition

This law states that a given compound always contains the same elements in the same proportion by mass.

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

Chemical Nomenclature: Naming Compounds

Chemical nomenclature is the system for naming chemical substances. It includes rules for naming ionic, covalent, and acid compounds.

• Ionic compounds: Name the cation (metal) first, then the anion (nonmetal) with -ide ending.

• Covalent compounds: Use prefixes (mono-, di-, tri-, etc.) to indicate the number of atoms.

• Acids: If the anion ends in -ide, the acid name begins with hydro- and ends with -ic acid. If the anion ends in -ate, the acid name ends with -ic acid; if -ite, ends with -ous acid.

• Hydrates: Compounds with water molecules attached; named with Greek prefixes (e.g., heptahydrate for 7 waters).

• Example: CoCl2·7H2O is cobalt(II) chloride heptahydrate.

Common Polyatomic Ions

Polyatomic ions are ions composed of more than one atom. Memorizing common polyatomic ions is essential for naming compounds.

• Examples: Nitrate (NO3-), sulfate (SO42-), phosphate (PO43-), ammonium (NH4+).

Practice: Naming and Writing Formulas

Practice is essential for mastering chemical nomenclature. Below is a sample table for naming compounds and writing formulas.

Compound Name	Formula
Sodium hydroxide	NaOH
Hydrobromic acid	HBr (aq)
Hydrogen sulfide	H2S (g)
Calcium carbonate	CaCO3
Iron(III) oxide	Fe2O3
Ammonium sulfate	(NH4)2SO4
Phosphoric acid	H3PO4
Aluminum acetate	Al(C2H3O2)3
Chromic acid	H2CrO4
Sodium dichromate	Na2Cr2O7
Additional info:	This table is a partial reconstruction based on the provided list; students should refer to their course materials for a complete list.

Sample Calculations and Problem Solving

Applying concepts through problem solving is crucial in chemistry. Below are examples of typical problems:

• Significant Figures: $(1.5 \times 10^4) \times 6.13 \div 2.01 = 4.6 \times 10^4$ (rounded to two significant figures)

• Density Calculation: $\text{Density} = \frac{\text{mass}}{\text{volume}}$

• Speed Conversion: To convert m/s to km/h, multiply by 3.6.

• Law of Constant Composition: Water from different sources always has the same ratio of hydrogen to oxygen by mass.

Additional info:

• Some content, such as the full list of compounds for naming and formula writing, is inferred from standard general chemistry curricula.

• Students should practice with additional examples for mastery.