General Chemistry: Classification of Matter, Properties, and Measurement

Classification of Matter and Elements

Elemental Properties and the Periodic Table

The periodic table organizes elements based on their chemical and physical properties. Each element is characterized by its symbol, state at room temperature, diatomic nature, metallic character, period, atomic number, and family/group number.

• Element Symbol: A one- or two-letter abbreviation for each element (e.g., P for phosphorus, Cl for chlorine, Kr for krypton).

• State at Room Temperature: Elements can be solid, liquid, or gas at room temperature.

• Diatomic: Some elements naturally exist as molecules composed of two atoms (e.g., O2, Cl2).

• Metal, Nonmetal, Metalloid, or Noble Gas: Elements are classified based on their properties. Metals are typically shiny and conductive, nonmetals are more variable, metalloids have intermediate properties, and noble gases are inert.

• Period and Family Number: The period is the row on the periodic table; the family (or group) is the column, often indicating similar chemical properties.

Example Table:

Element Symbol	Name	State at Room Temp	Diatomic?	Type	Period	Atomic Number	Family Number
P	Phosphorus	Solid	No	Nonmetal	3	15	V (Group 15)
Cl	Chlorine	Gas	Yes	Nonmetal	3	17	VIIA (Group 17)
Kr	Krypton	Gas	No	Noble Gas	4	36	VIIIA (Group 18)

Additional info: Table entries inferred for illustration.

Similar Chemical and Physical Properties

Elements in the same group (family) of the periodic table often have similar chemical and physical properties due to similar valence electron configurations. For example, chlorine and krypton are both in period 3 and 4, respectively, but only elements in the same group (e.g., halogens) will have closely related properties.

Physical and Chemical Changes

Definitions and Examples

Understanding the difference between physical and chemical changes is fundamental in chemistry.

• Physical Change: A change that affects the form of a chemical substance, but not its chemical composition (e.g., melting, dissolving, breaking).

• Chemical Change: A change that results in the formation of new chemical substances (e.g., burning, rusting, reacting with acids).

Examples:

• When Alka Seltzer is added to water and bubbles form: Chemical change (gas is produced).

• Hydrogen and nitrogen react at high temperature to form ammonia: Chemical change (new substance formed).

• Evaporation of water leaving sodium chloride crystals: Physical change (no new substance formed).

• Mixing two clear liquids to form a solid: Chemical change (precipitate forms).

Classification of Matter

Pure Substances and Mixtures

Matter can be classified as elements, compounds, or mixtures based on its composition.

• Element: A pure substance that cannot be broken down into simpler substances by chemical means (e.g., cobalt).

• Compound: A pure substance composed of two or more elements chemically combined in a fixed ratio (e.g., water, NaCl).

• Mixture: A physical combination of two or more substances that retain their individual properties (e.g., air, alloys, Kool-Aid).

Examples:

• Blue Kool-Aid: Mixture

• Substance that will not break down chemically: Element

• Cobalt: Element

• Alloy of zinc and tin: Mixture (specifically, a homogeneous mixture called an alloy)

Properties of Matter

Physical vs. Chemical Properties

Properties of matter are classified as physical or chemical based on whether they can be observed without changing the substance's identity.

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

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

Examples:

• Silver metal turning black after being in the air: Chemical property (tarnishing, reaction with sulfur compounds).

• Sugar dissolving in water: Physical property (dissolution, no new substance).

• Lead is more dense than sulfur: Physical property (density is a physical property).

Classification of Samples

Homogeneous and Heterogeneous Mixtures

Mixtures can be homogeneous (uniform composition) or heterogeneous (non-uniform composition).

• Tap water: Homogeneous mixture (solution).

• Wet sand: Heterogeneous mixture.

• Salad: Heterogeneous mixture.

Energy in Chemistry

Potential and Kinetic Energy

Energy is the capacity to do work. In chemistry, energy is often discussed as potential or kinetic.

• Potential Energy (PE): Stored energy due to position or composition (e.g., gasoline in a tank).

• Kinetic Energy (KE): Energy of motion (e.g., a thrown football).

Endothermic and Exothermic Processes

Chemical reactions can absorb or release energy.

• Endothermic Process: Absorbs energy from surroundings (temperature decreases).

• Exothermic Process: Releases energy to surroundings (temperature increases).

Example: If a solid chemical is mixed with water and the temperature goes up, the process is exothermic.

Measurement and Calculations in Chemistry

Significant Figures

Significant figures reflect the precision of a measured value. Rules for determining significant figures:

• All nonzero digits are significant.

• Zeros between nonzero digits are significant.

• Leading zeros are not significant.

• Trailing zeros after a decimal point are significant.

Example: 0.00404300 has 6 significant figures.

Scientific Notation

Scientific notation expresses numbers as a product of a coefficient and a power of ten.

• Example: $6.002210 \times 10^{-250}$

Mathematical Operations and Unit Conversions

Calculations in chemistry often require correct use of significant figures and scientific notation.

• When multiplying or dividing, the answer should have as many significant figures as the value with the fewest significant figures.

• When adding or subtracting, the answer should have as many decimal places as the value with the fewest decimal places.

Example Calculation:

• Calculate $\frac{1.92 - 1.32411}{1.32411} \times 100$

Temperature Conversions

Temperature can be converted between Celsius and Kelvin using:

• $K = ^\circ C + 273.15$

Metric Conversions

Metric conversions use powers of ten. For example:

• 1 centigram (cg) = $10^{-2}$ grams

• 1 microgram (μg) = $10^{-6}$ grams

Example: To convert 751 centigrams to micrograms:

• $751 \text{ cg} \times \frac{1 \text{ g}}{100 \text{ cg}} \times \frac{1,000,000 \text{ μg}}{1 \text{ g}} = 7,510,000 \text{ μg}$

Density and Volume Calculations

Density is defined as mass per unit volume:

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

To find volume when density and mass are known:

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

Example: Osmium has a density of 22.60 g/cm3. To find the volume of a block with given dimensions, first calculate the volume in cubic centimeters, then use the density formula.

Dimensional Analysis

Dimensional analysis (factor-label method) is used to convert units by multiplying by conversion factors.

• Example: To convert miles/hour to decimeters/second, use the following conversions:

• 1 mile = 1,609.34 meters

• 1 meter = 10 decimeters

• 1 hour = 3,600 seconds

Volume of Water in the Oceans

To convert the volume of water from km3 to liters:

• 1 km3 = $1 \times 10^{12}$ liters

• Multiply the given volume in km3 by $1 \times 10^{12}$ to get liters.

Example: $1.35 \times 10^9$ km3 × $1 \times 10^{12}$ L/km3 = $1.35 \times 10^{21}$ L

Additional info: Some examples and explanations were expanded for clarity and completeness.