Chapters 1 & 2: Introduction to Chemistry and Measurements

Introduction to Chemistry

Chemistry and Matter

Chemistry is the study of matter, its structure, properties, and the transformations it undergoes. Matter is anything that has mass and occupies space. Matter can change from one form to another, but it cannot be destroyed.

• Matter: Anything with mass and volume.

• Chemistry: The study of matter, its structure, properties, and changes.

• Chemical Changes: Transformations that alter the composition of matter (e.g., chemical reactions).

• Physical Changes: Changes in state or appearance without altering composition (e.g., melting, boiling).

• Example: Ice melting to water (physical change), iron rusting (chemical change).

Chemicals

A chemical is a substance that always has the same composition and properties wherever it is found.

• Pure water from a river is chemically identical to pure water made in a laboratory.

• Aspirin extracted from willow bark is the same as aspirin synthesized in a factory.

The Scientific Method

Steps of the Scientific Method

The scientific method is a systematic approach used for scientific investigation and advancement.

• Observation: Gathering information using the senses (e.g., noticing more sweat in summer).

• Hypothesis: A proposed explanation for an observation, without proof (e.g., "High temperatures cause more sweating").

• Experiment: Testing the hypothesis through controlled procedures.

• Conclusion/Theory: A summary of results that supports or refutes the hypothesis (e.g., "High temperatures contribute to more sweating").

Hypotheses, Theories, and Facts

• Hypothesis: A tentative explanation for an observation or relationship.

• Theory: A well-supported explanation of phenomena, based on evidence. Theories can be modified or rejected if new evidence arises.

• Fact: A statement based on direct evidence.

Chemical Measurements

Significant Figures

Significant figures (sig figs) are the digits in a measurement that are known with certainty plus one estimated digit. They reflect the precision of a measurement.

• Not all digits in a value are significant; only those that are certain and one uncertain digit are counted.

• Example: If a scale reads 3.0 g, it is not correct to write 3.00 g unless the scale can measure to that precision.

Rules for Significant Figures

• Nonzero digits are always significant (e.g., 2.31 has three sig figs).

• Leading zeros (before the first nonzero digit) are never significant (e.g., 0.0055 has two sig figs).

• Trailing zeros in a decimal number are significant (e.g., 3.00 has three sig figs).

• Zeros between nonzero digits are always significant (e.g., 2.045 has four sig figs).

Significant Figures in Calculations

• Addition and Subtraction: The result should have the same number of decimal places as the value with the fewest decimal places.

• Multiplication and Division: The result should have the same number of significant figures as the value with the fewest significant figures.

• Example (Addition): $3.147 + 2.52 = 5.667 \rightarrow 5.67$

• Example (Multiplication): $4.626 \times 1.34 = 6.19884 \rightarrow 6.20$

Scientific Notation

Scientific notation expresses very large or very small numbers as a value between 1 and 10 multiplied by a power of 10.

• Format: $\text{Number} \times 10^{\text{Exponent}}$

• The exponent indicates how many places the decimal point is moved.

• Example: $273,000 = 2.73 \times 10^5$

Units and Measurements

Importance of Units

Units give meaning to numerical values in measurements. Always include units to clarify what is being measured.

• Example: 56 lb (pounds) vs. 56 kg (kilograms) represent very different masses.

• In science, the metric system is used, with units and subunits based on powers of 10.

The Metric System

The metric system is the standard system of measurement in science. Key units include:

Quantity	Unit	Symbol
Length	meter	m
Volume	liter	L
Mass	gram	g
Time	second	s
Temperature	Kelvin	K
Energy	joule	J
Amount of substance	mole	mol

Mole: A counting unit in chemistry, similar to "a dozen," but equal to $6.022 \times 10^{23}$ particles.

Metric System Prefixes

Prefixes indicate multiples or fractions of units.

Prefix	Symbol	Value
giga	G	$10^9$
mega	M	$10^6$
kilo	k	$10^3$
deci	d	$10^{-1}$
centi	c	$10^{-2}$
milli	m	$10^{-3}$
micro	\mu	$10^{-6}$
nano	n	$10^{-9}$

Example: 1 kilometer (km) = 1,000 meters (m); 1 milliliter (mL) = 0.001 liters (L).

Interchanging Between Metric and English Units

Conversion factors are used to convert between different units.

Length	Mass	Volume
1 in. = 2.54 cm	1 oz = 28.35 g	1 qt = 0.946 L
1 m = 39.37 in.	1 lb = 453.6 g	1 gal = 3.785 L
1 mile = 1.609 km	1 kg = 2.205 lb	1 fl oz = 29.57 mL
	1 g = 15.43 grains	1 L = 1.057 qt

• Example: To convert 4 liters to quarts, use the conversion factor $1 \text{ L} = 1.057 \text{ qt}$.

Converting Between Units

When converting units, the amount of substance does not change, only the unit in which it is expressed.

• Use conversion factors to set up the calculation.

• Example: Convert 381 grams to pounds. $1 \text{ lb} = 453.6 \text{ g}$.

Units of Temperature

Temperature Scales

• Fahrenheit (°F): Defined by freezing a salty solution at 0°F and approximating body temperature at 100°F.

• Celsius (°C): Defined by setting the freezing point of water at 0°C and boiling point at 100°C.

• Kelvin (K): Absolute temperature scale; 0 K is the lowest possible temperature. The freezing point of water is 273 K, and the boiling point is 373 K. A Kelvin degree is the same size as a Celsius degree.

Density and Related Concepts

Density

Density is the ratio of mass to volume, or the mass of a substance that occupies a certain volume.

• Formula: $d = \frac{m}{v}$

• Units: g/mL for liquids and solids, g/L for gases.

• Example: What is the density of a metal with a mass of 45.75 g and a volume of 6.95 mL? $d = \frac{45.75}{6.95} = 6.58 \text{ g/mL}$

Density and Concentration

Density is often used in medication dosage calculations, as it relates to the concentration of a solution.

• Example: If a drug is provided at 60 mg/mL and a patient needs 600 mg, the volume administered is $\frac{600 \text{ mg}}{60 \text{ mg/mL}} = 10 \text{ mL}$.

Specific Gravity

Specific gravity is the ratio of the density of a substance to the density of water at a standard temperature. It is dimensionless (no units).

• Formula: $\text{Specific gravity} = \frac{\text{density of substance}}{\text{density of water}}$

• Example: The density of copper at 20°C is 8.92 g/mL. The density of water at this temperature is 1.00 g/mL. Specific gravity = $\frac{8.92}{1.00} = 8.92$

The specific gravity of a substance will always be equal to the value of its density, but without any units.

Additional info: Practice problems and worked examples are included throughout to reinforce key concepts. Students should be familiar with all metric prefixes, unit conversions, and the application of significant figures in calculations.