Measurement, Units, and Dimensional Analysis

Physical Quantities and Units

In chemistry, measurements are fundamental for describing matter and its changes. Common physical quantities include mass, volume, density, and temperature. The SI (International System of Units) is used for standardization.

• Density: Defined as mass per unit volume. (units: g/cm3 or kg/m3)

• Temperature: Measured in Kelvin (K), Celsius (°C), or Fahrenheit (°F).

Prefixes and Conversion Factors

SI prefixes are used to express multiples or fractions of units. Conversion factors allow transformation between units.

Dimensional Analysis

Dimensional analysis is a systematic method for converting between units using conversion factors.

• Example: Convert 3.409 mi/hr to km/min. Step 1: Convert miles to kilometers using or . Step 2: Convert hours to minutes (). Calculation:

Significant Figures

Significant figures reflect the precision of a measurement. The number of significant digits in the result should match the least precise measurement used in the calculation.

• Example: (5 significant figures)

Numbers in Science: Exact vs. Measured

Types of Numbers

• Exact numbers: Counted or defined values (e.g., 12 eggs in a dozen).

• Measured numbers: Obtained using instruments, subject to equipment and human errors.

Atomic Theory and Laws of Chemistry

Law of Constant Composition (Law of Definite Proportions)

All samples of a given compound have the same proportions of their constituent elements, regardless of source or preparation method.

• Example: Water (H2O) always contains hydrogen and oxygen in a mass ratio of approximately 1:8.

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.

• Example: Carbon and oxygen form CO and CO2. The mass of oxygen per gram of carbon in CO2 is twice that in CO.

Law of Conservation of Mass

Mass is neither created nor destroyed in a chemical reaction. The total mass of reactants equals the total mass of products.

Dalton's Atomic Theory

• Elements are composed of tiny, indestructible particles called atoms.

• All atoms of a given element are identical in mass and properties.

• Atoms of different elements differ in mass and properties.

• Atoms are not created or destroyed in chemical reactions.

• Compounds are formed by the combination of atoms in fixed ratios.

Structure of the Atom

Subatomic Particles

• Protons and electrons have equal but opposite charges.

• Atoms are electrically neutral when the number of protons equals the number of electrons.

Atomic Number and Mass Number

• Atomic number (Z): Number of protons in the nucleus; defines the element.

• Mass number (A): Total number of protons and neutrons.

• Number of neutrons:

Isotopes

Isotopes are atoms of the same element with identical chemical properties but different masses due to varying numbers of neutrons.

• Example: Carbon-12 (6 protons, 6 neutrons), Carbon-13 (6 protons, 7 neutrons), Carbon-14 (6 protons, 8 neutrons)

Atomic Mass Unit (amu) and Weighted Average

• 1 amu = g

• Atomic mass listed in the periodic table is a weighted average of the masses of naturally occurring isotopes.

• Example: Neon has three isotopes with masses and abundances: (19.9924 amu, 90.48%), (20.9938 amu, 0.27%), (21.9914 amu, 9.25%)

The Periodic Table

Organization and Classification

• Elements are arranged by increasing atomic number (number of protons).

• Rows are called periods; columns are called groups or families.

• Elements in the same group have similar chemical and physical properties.

Element Names and Symbols

It is essential to memorize the names and symbols of the first 36 elements and selected others (Sr, Ba, Sn, Pb, I, Xe, Rn).

Ions and Ionic Compounds

Monatomic and Polyatomic Ions

• Monatomic ions: Ions formed from a single atom (e.g., Na+, Cl-).

• Polyatomic ions: Ions composed of two or more atoms covalently bonded, carrying a net charge (e.g., NO3-, SO42-).

Formation of Ions

• Metals lose electrons to form cations (positively charged ions).

• Nonmetals gain electrons to form anions (negatively charged ions).

• Group number often predicts the charge for main-group elements.

Classification of Matter

Types of Compounds

• Molecular (Covalent) Compounds: Composed of nonmetals bonded by sharing electrons (e.g., H2O, CO2).

• Ionic Compounds: Composed of metals and nonmetals, or polyatomic ions, held together by electrostatic attraction (e.g., NaCl, CaCO3).

Bonding and Formula Units

• Molecular compounds consist of molecules (discrete units).

• Ionic compounds consist of formula units (smallest electrically neutral collection of ions).

Empirical, Molecular, and Structural Formulas

• Empirical formula: Simplest whole-number ratio of atoms (e.g., HO for hydrogen peroxide).

• Molecular formula: Actual number of atoms in a molecule (e.g., H2O2).

• Structural formula: Shows how atoms are connected (e.g., H–O–O–H).

Naming Compounds

Naming Ionic Compounds

• For cations with variable charge (transition metals), indicate charge with Roman numerals (e.g., Fe2+ is iron(II)).

• For polyatomic ions, memorize common names and formulas.

Naming Oxoanions

• Ion with one more O atom than the '-ate' ion: per...ate (e.g., perchlorate, ClO4-).

• Ion with one less O atom than the '-ate' ion: ...ite (e.g., nitrite, NO2-).

• Ion with two fewer O atoms: hypo...ite (e.g., hypochlorite, ClO-).

Practice Example

• Na3PO4: sodium phosphate

• CaBr2: calcium bromide

• Cu2O: copper(I) oxide

• Fe2(SO4)3: iron(III) sulfate

Summary Table: Common Polyatomic Ions

Additional info: Some context and examples have been expanded for clarity and completeness, including formula writing and naming conventions for compounds.