Fundamental Concepts in General Chemistry: Classification, Properties, and Quantitative Relationships

Classification of Matter

Elements, Compounds, and Mixtures

Matter can be classified based on its composition into three main categories: elements, compounds, and mixtures. Understanding these distinctions is fundamental to the study of chemistry.

• Element: A pure substance consisting of only one type of atom. Examples: O2 (oxygen gas), Fe (iron).

• Compound: A pure substance composed of two or more different elements chemically combined in fixed proportions. Example: H2O (water).

• Mixture: A physical combination of two or more substances where each retains its own properties. Mixtures can be homogeneous (uniform composition, e.g., saltwater) or heterogeneous (non-uniform composition, e.g., sand and iron filings).

Separation of Mixtures

Filtration vs. Distillation

Mixtures can be separated by physical means based on differences in their physical properties.

• Filtration: Used to separate solids from liquids in heterogeneous mixtures. Example: Separating sand from water.

• Distillation: Used to separate components of a homogeneous mixture based on differences in boiling points. Example: Separating ethanol from water.

Properties and Changes in Matter

Chemical vs. Physical Properties and Changes

Properties and changes in matter are classified as either physical or chemical:

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

• Chemical Property: Describes a substance's ability to undergo a specific chemical change (e.g., flammability, reactivity).

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

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

Density and Its Applications

Density Relationship

Density is a physical property defined as mass per unit volume. It is useful for identifying substances and solving quantitative problems.

• Formula:

$d = \frac{m}{V}$

• m = mass (g)

• V = volume (mL or cm3)

Example: If a sample has a mass of 10.0 g and a volume of 2.0 mL, its density is $d = \frac{10.0\ g}{2.0\ mL} = 5.0\ g/mL$.

Significant Figures

Determining Significant Figures in Calculations

Significant figures reflect the precision of a measurement. The rules for determining the number of significant figures in a result depend on the operation:

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

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

Example: $2.34 \times 1.2 = 2.8$ (rounded to 2 significant figures).

Dimensional Analysis

Unit Conversions

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

• Example: Convert 25.0 cm to meters.

$25.0\ \text{cm} \times \frac{1\ \text{m}}{100\ \text{cm}} = 0.250\ \text{m}$

Law of Definite Proportions

Calculating Percent Composition

The Law of Definite Proportions states that a chemical compound always contains the same elements in the same proportion by mass.

• Percent Composition Formula:

$\%\ \text{element} = \frac{\text{mass of element in 1 mol of compound}}{\text{molar mass of compound}} \times 100\%$

• Example: In H2O, % H = $\frac{2.02}{18.02} \times 100\% = 11.2\%$

Atomic Structure

Protons, Neutrons, and Electrons

Atoms are composed of three main subatomic particles:

• Proton: Positively charged, found in the nucleus.

• Neutron: Neutral, found in the nucleus.

• Electron: Negatively charged, found outside the nucleus.

Determining Numbers:

• Number of protons = atomic number (Z)

• Number of electrons = number of protons (in a neutral atom)

• Number of neutrons = mass number (A) - atomic number (Z)

Ions and Isotopes

• Ion: An atom or molecule with a net electric charge due to the loss or gain of electrons.

• Isotope: Atoms of the same element with different numbers of neutrons.

Example: 35Cl and 37Cl are isotopes of chlorine.

Classification of Elements

Metals, Nonmetals, and Metalloids

Elements are classified based on their properties and position on the periodic table.

Type	Location	Properties
Metals	Left and center	Good conductors, malleable, ductile, shiny
Nonmetals	Right side	Poor conductors, brittle, dull
Metalloids	Stair-step line	Intermediate properties

Predicting Ion Charges

Periodic Table Trends

The charge of ions formed by elements can often be predicted by their group on the periodic table.

• Group 1: +1

• Group 2: +2

• Group 17 (Halogens): -1

• Group 16: -2

Example: Sodium (Na) forms Na+; Chlorine (Cl) forms Cl-.

Average Atomic Mass

Using Mass Spectrometry Data

The average atomic mass of an element is calculated using the masses and relative abundances of its isotopes.

• Formula:

$\text{Average atomic mass} = \sum (\text{isotope mass} \times \text{fractional abundance})$

• Example: If an element has two isotopes: 10.0 amu (20%) and 11.0 amu (80%), then average atomic mass = $(10.0 \times 0.20) + (11.0 \times 0.80) = 10.8$ amu.

The Molar Road Map

Conversions Between Moles, Grams, and Number of Atoms

The mole is a fundamental unit in chemistry for counting particles. The molar road map helps convert between mass, moles, and number of particles.

• Key Relationships:

From	To	Conversion Factor
Moles	Particles	1 mol = $6.022 \times 10^{23}$ particles (Avogadro's number)
Moles	Grams	1 mol = molar mass (g)

• Example: How many atoms are in 2.00 g of H? First, convert grams to moles, then moles to atoms.

$2.00\ \text{g} \times \frac{1\ \text{mol}}{1.01\ \text{g}} \times 6.022 \times 10^{23}\ \text{atoms/mol} = 1.19 \times 10^{24}\ \text{atoms}$