Chapter 1: Introduction to Matter, Measurement, and the Periodic Table

Classifying Matter

Pure Substances and Mixtures

Matter can be classified into pure substances and mixtures based on its composition and properties. Understanding these classifications is fundamental in chemistry.

• Pure Substances: Matter with a fixed composition. Includes elements and compounds.

• Elements: Substances made of only one type of atom (e.g., gold (Au), diamond (C)).

• Compounds: Substances composed of two or more elements chemically combined in fixed ratios (e.g., water (H2O), caffeine (C8H10N4O2)).

• Mixtures: Physical combinations of two or more substances. Can be separated by physical means.

• Homogeneous Mixtures (Solutions): Uniform composition throughout (e.g., air, batter).

• Heterogeneous Mixtures: Non-uniform composition (e.g., salad, sand and water).

• Aqueous Mixture: A solution where water is the solvent.

Example: Salt dissolved in water forms a homogeneous aqueous mixture.

Elements, Compounds, and the Periodic Table

The Periodic Table

The Periodic Table is an organized chart of all known elements, arranged by increasing atomic number and grouped by similar chemical properties.

• Chemical Symbol: One- or two-letter abbreviation representing each element (e.g., H for hydrogen, O for oxygen).

• Some symbols do not match their English names (e.g., Ag for silver, Na for sodium).

• Elements are grouped into categories: metals, nonmetals, and metalloids.

Example: The symbol for gold is Au, derived from its Latin name aurum.

Classification of Elements

• Metals: Good conductors of heat and electricity, malleable, ductile, and shiny.

• Nonmetals: Poor conductors, often brittle, and not shiny.

• Metalloids: Have properties intermediate between metals and nonmetals.

Example: Silicon (Si) is a metalloid used in electronics.

Compounds and Chemical Formulas

Compounds

Compounds are substances that contain more than one element, combined in specific ratios. Their composition is represented by chemical formulas.

• Chemical Formula: Indicates the types and numbers of atoms in a compound (e.g., H2O for water).

• Subscripts in formulas show the ratio of elements (e.g., in H2O, there are 2 hydrogen atoms for every 1 oxygen atom).

Example: The formula for ethanol is C2H6O.

Elements in Nutrition

Major and Trace Elements

Certain elements are essential for life and are classified based on the quantities required by the body.

• Major Elements: Required in quantities greater than 100 mg per day (e.g., Na, K, Ca).

• Trace Elements: Required in quantities less than 100 mg per day (e.g., Fe, Zn).

Example: Iron (Fe) is a trace element important for oxygen transport in blood.

Physical and Chemical Changes

Types of Changes

Matter can undergo physical or chemical changes, which affect its properties and composition.

• Physical Change: Alters the state or appearance of matter without changing its chemical composition (e.g., melting ice).

• Chemical Change: Alters the chemical composition, resulting in new substances (e.g., burning wood).

Example: Dissolving sugar in water is a physical change; rusting of iron is a chemical change.

Chemical Equations and Reactions

Writing and Balancing Chemical Equations

Chemical equations represent chemical reactions, showing reactants and products.

• Reactants: Substances present before the reaction.

• Products: Substances formed by the reaction.

• Arrow (→): Separates reactants from products.

• Balancing Equations: Ensures the same number of each atom on both sides of the equation.

Example Equations:

• $\mathrm{CH_4(g) + 2O_2(g) \rightarrow 2H_2O(g) + CO_2(g)}$

• $\mathrm{2H_2(g) + O_2(g) \rightarrow 2H_2O(g)}$

• $\mathrm{NO_2(g) + H_2O(l) \rightarrow HNO_3(aq) + NO(g)}$

Measurement and Units

International System of Units (SI)

Scientific measurements use the SI system for consistency.

• Mass: Kilogram (kg)

• Length: Meter (m)

• Volume: Liter (L)

Example: 1 kg ≈ 2.2 lbs

Scientific Notation

Scientific notation is used to express very large or small numbers conveniently.

• Numbers are written as $a \times 10^n$, where $1 \leq a < 10$ and $n$ is an integer.

• Example: $6,500 = 6.5 \times 10^3$

Significant Figures

Significant figures reflect the precision of a measurement.

• All nonzero digits are significant.

• Zeros between nonzero digits are significant.

• Trailing zeros in a decimal number are significant.

• Rules for calculations:

  • Addition/Subtraction: Result has the same number of decimal places as the measurement with the fewest decimal places.

  • Multiplication/Division: Result has the same number of significant figures as the measurement with the fewest significant figures.

Example: $9.00 \times 4 / 21 = 2$ (rounded to the correct number of significant figures)

Exact and Inexact Numbers

Definitions

• Exact Numbers: Known with complete certainty (e.g., counting objects).

• Inexact Numbers: Obtained by measurement and subject to uncertainty.

Example: There are exactly 12 eggs in a dozen (exact); a measured length of 2.54 cm (inexact).

Dimensional Analysis

Unit Conversions

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

• Conversion Factor: A ratio that expresses how many of one unit are equal to another unit.

• Example: $1 \text{ inch} = 2.54 \text{ cm}$

• Set up calculations so units cancel appropriately.

Example: To convert 10 inches to centimeters: $10 \text{ in} \times \frac{2.54 \text{ cm}}{1 \text{ in}} = 25.4 \text{ cm}$

Properties of Matter

Mass, Volume, and Density

These are fundamental properties used to describe matter.

• Mass: Amount of material in an object (measured in kg or g).

• Volume: Amount of space occupied (measured in L or m3).

• Density: Ratio of mass to volume.

Formula:

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

Example: If a block has a mass of 200 g and a volume of 50 mL, its density is $\frac{200\,g}{50\,mL} = 4\,g/mL$.

Specific Gravity

Specific gravity is the ratio of the density of a substance to the density of water.

• $\text{Specific Gravity} = \frac{\text{Density of substance}}{\text{Density of water}}$

Example: If a liquid has a density of 1.2 g/mL, its specific gravity is $\frac{1.2}{1.0} = 1.2$.

Temperature and Energy

Temperature Scales

Temperature measures the hotness of an object. Common scales include Celsius, Fahrenheit, and Kelvin.

• Celsius (°C): Water freezes at 0°C and boils at 100°C.

• Fahrenheit (°F): Water freezes at 32°F and boils at 212°F.

• Kelvin (K): Absolute temperature scale. $K = °C + 273.15$

Example: 25°C = 298.15 K

Energy and Specific Heat

Energy is the capacity to do work or supply heat. Specific heat is the amount of heat required to raise the temperature of 1 g of a substance by 1°C.

• Kinetic Energy: Energy of motion.

• Potential Energy: Stored energy.

• Law of Conservation of Energy: Energy cannot be created or destroyed.

• Specific Heat Formula: $q = m \times c \times \Delta T$

• Where $q$ = heat (J), $m$ = mass (g), $c$ = specific heat (J/g·°C), $\Delta T$ = change in temperature (°C).

Example: To heat 10 g of water by 5°C ($c = 4.18$ J/g·°C): $q = 10 \times 4.18 \times 5 = 209$ J

States of Matter

Solid, Liquid, and Gas

Matter exists in three primary states, each with distinct properties.

• Solid: Definite shape and volume.

• Liquid: Definite volume, takes the shape of its container.

• Gas: No definite shape or volume, fills its container.

Example: Ice (solid), water (liquid), steam (gas).

Precision and Accuracy in Measurement

Definitions

• Accuracy: How close a measurement is to the true value.

• Precision: How reproducible measurements are.

Example: If repeated measurements of a length are all close to each other, they are precise; if they are close to the actual length, they are accurate.

Unit Conversions and Common Prefixes

Metric and US Customary Units

Conversions between metric and US customary units are often required in chemistry.

Metric Unit	Equivalent US Unit
1 kg	2.2 lbs
1 L	1.06 qt
1 m	39.37 in
1 tsp	4.93 mL
1 fl oz	29.57 mL

Example: To convert 3 L to quarts: $3 \times 1.06 = 3.18$ qt

Additional info: Some content was inferred and expanded for clarity and completeness, including definitions, formulas, and examples relevant to GOB Chemistry students.