General Chemistry: Fundamental Concepts, Calculations, and Chemical Nomenclature

Chapter 1: Introduction to Chemistry

Physical vs. Chemical Properties and Changes

Chemistry distinguishes between physical and chemical properties and changes. Understanding these differences is essential for classifying matter and predicting behavior in chemical reactions.

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

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

• Physical change: Alters the form or appearance but not the chemical identity (e.g., melting ice).

• Chemical change: Results in the formation of new substances (e.g., rusting iron).

• Example: Dissolving salt in water is a physical change; burning wood is a chemical change.

Units of Measurement and Significant Figures

Accurate measurement is fundamental in chemistry. The use of SI units and proper handling of significant figures ensures precision in calculations.

• SI Units: Standard units used in science (meter, kilogram, second, mole, etc.).

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

• Dimensional analysis: A method to convert between units using conversion factors.

• Metric prefixes: Indicate powers of ten (e.g., kilo-, centi-, milli-).

• Example: $1~\text{kg} = 1000~\text{g}$

Mathematical Operations in Chemistry

Chemical calculations often require multiplication, division, addition, and subtraction, with attention to significant figures.

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

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

• Example: $(1.5 \times 10^4 \times 6.3) \div 2.01$

Temperature Scales

Temperature is measured in Celsius, Kelvin, and Fahrenheit. Conversions between these scales are common in chemistry.

• Celsius to Kelvin: $K = ^\circ C + 273.15$

• Celsius to Fahrenheit: $^\circ F = (9/5) \times ^\circ C + 32$

Density and Volume Calculations

Density is a physical property defined as mass per unit volume. It is used to identify substances and solve problems involving mass and volume.

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

• Example: Gold has a density of $19.3~\text{g/cm}^3$; calculate the volume for a given mass.

Chapter 2: Atomic Theory and Chemical Nomenclature

Dalton’s Atomic Theory and Laws of Multiple Proportions

Dalton’s atomic theory laid the foundation for modern chemistry, describing atoms as indivisible particles and explaining how elements combine in fixed ratios.

• Dalton’s postulates: Elements are made of atoms; atoms of the same element are identical; atoms combine in simple whole-number ratios.

• Law of Multiple Proportions: When elements combine, they do so in ratios of small whole numbers.

• Example: CO and CO2 show different ratios of oxygen to carbon.

Subatomic Particles and Atomic Structure

Atoms consist of protons, neutrons, and electrons. Their arrangement determines the element’s identity and properties.

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

• Neutron (n): Neutral particle in the nucleus.

• Electron (e): Negatively charged particle orbiting the nucleus.

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

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: Same number of protons, different number of neutrons.

• Atomic mass: $\text{Atomic mass} = \sum (\text{isotope mass} \times \text{fractional abundance})$

Periodic Table and Element Classification

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

• Groups: Vertical columns with similar chemical properties.

• Periods: Horizontal rows.

• Metals, nonmetals, metalloids: Classified by physical and chemical properties.

Chemical Formulas and Nomenclature

Chemical nomenclature is the system for naming compounds and writing their formulas. It is essential for clear communication in chemistry.

• Ionic compounds: Named using the cation first, then the anion (e.g., NaCl: sodium chloride).

• Covalent compounds: Use prefixes to indicate the number of atoms (e.g., CO2: carbon dioxide).

• Acids: Named based on the anion (e.g., HCl: hydrochloric acid).

• Hydrates: Compounds containing water molecules (e.g., CuSO4·5H2O).

• Example: Cobalt(II) sulfate heptahydrate: CoSO4·7H2O.

Polyatomic Ions and Common Compounds

Polyatomic ions are charged species composed of two or more atoms covalently bonded. Recognizing their names and formulas is crucial for writing chemical equations.

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

• Common acids: Sulfuric acid (H2SO4), phosphoric acid (H3PO4).

States of Matter

Matter exists in three primary states: solid, liquid, and gas. Each state has distinct physical properties.

• Solid: Definite shape and volume.

• Liquid: Definite volume, indefinite shape.

• Gas: Indefinite shape and volume.

Mixtures and Pure Substances

Chemistry distinguishes between mixtures and pure substances. Mixtures can be homogeneous or heterogeneous.

• Homogeneous mixture: Uniform composition throughout (e.g., salt water).

• Heterogeneous mixture: Non-uniform composition (e.g., sand and iron filings).

• Pure substance: Element or compound with a fixed composition.

Law of Constant Composition

The law of constant composition states that a given compound always contains the same proportion of elements by mass.

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

Chemical Nomenclature Practice

Compound Naming and Formula Writing

Practice naming compounds and writing their formulas is essential for mastering chemical nomenclature.

Name	Formula
Sodium hydroxide	NaOH
Barium chromate	BaCrO4
Hydrochloric acid	HCl
Aluminum acetate	Al(C2H3O2)3
Ammonium nitrate	NH4NO3
Sodium dihydrogen phosphate	NaH2PO4
Phosphoric acid	H3PO4
Carbon tetrachloride	CCl4
Silver chloride	AgCl
Sulfuric acid	H2SO4
Potassium permanganate	KMnO4
Dichlorine heptoxide	Cl2O7
Additional info:	Other compounds and acids can be named using similar rules.

Sample Calculations and Problem Solving

Significant Figures in Calculations

When performing calculations, always report the answer with the correct number of significant figures.

• Example: $(1.5 \times 10^4 \times 6.3) \div 2.01$

Density and Volume Problems

Use the density formula to solve for mass or volume as required.

• Example: Gold has a density of $19.3~\text{g/cm}^3$; find the volume for a given mass.

Dosage Calculations

Dosage calculations require unit conversions and careful attention to significant figures.

• Example: Calculate the volume of solution needed for a given drug dosage based on patient weight.

Additional info:

• Some questions involve real-world applications, such as speed conversions and Olympic race analysis.

• Practice problems cover a range of fundamental chemistry skills, including unit conversions, atomic structure, and chemical nomenclature.