Matter and Measurement – Chapter 1 Study Notes

Matter and Measurement

Introduction to Chemistry

Chemistry is the scientific study of matter and the transformations it undergoes. Matter is defined as anything that has mass and occupies space. Transformations in matter can be either physical or chemical, and these changes often require the transfer of energy.

• Matter: The 'stuff' that makes up the universe; anything with mass and volume.

• Transformation: Any change in matter, classified as physical (no change in composition) or chemical (change in composition).

Classification of Matter

Types of Matter

Matter is categorized into two broad types: pure substances and mixtures.

• Pure Substance: Contains only one type of matter. Examples: table salt (sodium chloride), copper metal.

• Mixture: Contains two or more substances physically combined. Examples: salt and sugar mixture, rocks with various minerals.

Subcategories of Matter

Category	Subcategory	Description	Example
Pure Substance	Element	Cannot be broken down by chemical change	Aluminum (Al)
Pure Substance	Compound	Formed by chemically joining two or more elements	Table salt (NaCl)
Mixture	Homogeneous	Uniform composition throughout	Sugar dissolved in tea
Mixture	Heterogeneous	Non-uniform composition	Oil and water

Elements and Compounds

• Element: Basic building block of matter, organized in the periodic table. About 118 known elements (90 natural, 28 synthetic).

• Compound: Pure substance formed by chemically joining two or more elements. Example: Water (H2O).

Elements are represented by symbols, often derived from English, Latin, or German names (e.g., C for carbon, Fe for ferrum/iron, W for wolfram/tungsten).

Atoms

• The smallest unit of an element is an atom.

• Atoms cannot be seen with normal microscopes; first imaged in the 1980s using a scanning tunneling microscope.

Chemical Formulas

• A chemical formula uses symbols to show the elements present and their proportions. Example: Water (H2O) has 2 hydrogen and 1 oxygen atom.

• Table sugar (C12H22O11) contains 12 carbon, 22 hydrogen, and 11 oxygen atoms.

Practice: Classification Examples

• Iron oxide (Fe2O3): Compound

• Ozone (O3): Elemental substance

• Iron (Fe): Elemental substance

• Carbon monoxide (CO): Compound

• Propane (C3H8): Compound

• Water (H2O): Compound

• Helium (He): Elemental substance

• Graphite (C): Elemental substance

States and Properties of Matter

States of Matter

• Solid: Fixed shape and volume; may be rigid or flexible.

• Liquid: Fixed volume, takes the shape of its container, forms a surface.

• Gas: No fixed shape or volume, expands to fill its container.

The physical state depends on temperature, pressure, and the strength of forces between particles.

Physical and Chemical Properties

• Physical properties: Can be observed or measured without changing the substance's composition (e.g., boiling point, melting point, solubility, color, odor).

• Chemical properties: Describe how a substance can be converted into another substance (e.g., flammability, reactivity).

Physical and Chemical Changes

• Physical change: Alters appearance but not composition (e.g., melting ice, boiling water).

• Chemical change: Alters composition, forming new substances (e.g., burning paper, rusting iron).

Example: Gasoline's odor is a physical property; its flammability is a chemical property.

Phase Changes

• Freezing: Liquid to solid

• Melting: Solid to liquid

• Vaporization: Liquid to gas

• Condensation: Gas to liquid

• Sublimation: Solid to gas

• Deposition: Gas to solid

Numbers in Chemistry

Precision and Accuracy

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

• Precision: How close repeated measurements are to each other.

• Goal: Achieve both precision and accuracy in measurements.

Example: If the true value is 260, and Mike's measurements are 256, 263, 262, 266 (more precise), while Lia's are 250, 242, 270, 278 (less precise), Mike is more precise, but accuracy depends on closeness to 260.

Scientific Notation

• Used to express very large or small numbers conveniently.

• Format: Decimal part (between 1 and 10) × 10exponent.

Example: 2,500 = $2.5 \times 10^3$; 0.036 = $3.6 \times 10^{-2}$

Significant Figures

• All nonzero digits are significant.

• Interior zeros (between nonzero digits) are significant.

• Leading zeros (before the first nonzero digit) are not significant.

• Trailing zeros are significant only if there is a decimal point.

• Exact numbers (counted or defined) have unlimited significant figures.

Examples:

• 0.0035 (2 significant figures)

• 1.080 (4 significant figures)

• 2371 (4 significant figures)

• 2.90 × 105 (3 significant figures)

• 1 dozen = 12 (unlimited significant figures)

Significant Figures in Calculations

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

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

• Rounding: If the digit to be dropped is 5 or greater, round up; if less than 5, round down.

Example: 231.45 rounded to three significant digits is 231.

Units of Measurement

SI Units

• Length: meter (m)

• Mass: kilogram (kg)

• Time: second (s)

• Temperature: kelvin (K)

• Amount of substance: mole (mol)

SI Prefixes

Prefix	Symbol	Multiplier
kilo-	k	1,000
centi-	c	0.01
milli-	m	0.001
micro-	μ	0.000001
nano-	n	0.000000001

Temperature Scales

• Kelvin (K): Absolute temperature scale.

• Celsius (°C): Commonly used in science.

• Fahrenheit (°F): Used in the United States.

Conversion formulas:

• $K = ^\circ C + 273.15$

• $^\circ F = (^\circ C \times \frac{9}{5}) + 32$

• $^\circ C = (^\circ F - 32) \times \frac{5}{9}$

Density

Definition and Calculation

• Density (d): The ratio of mass to volume.

Formula:

• $d = \frac{m}{V}$

Example: If a sample has a mass of 27.2 g and a volume of 25.0 mL, its density is $d = \frac{27.2\ g}{25.0\ mL} = 1.09\ g/mL$.

• Intensive property: Does not depend on the amount of material (e.g., density, temperature).

• Extensive property: Depends on the amount of material (e.g., mass, volume).

Unit Analysis (Dimensional Analysis)

Unit Conversions

• Used to convert a quantity from one unit to another using conversion factors.

• General form: information given × conversion factor(s) = information sought

Example: Convert 44.7 cm to inches (1 in = 2.54 cm):

• $44.7\ cm \times \frac{1\ in}{2.54\ cm} = 17.6\ in$

For units raised to a power, raise the conversion factor to that power. Example: $(2.54\ cm)^2 = (1\ in)^2$ for area conversions.