Introduction to Chemistry: Scientific Method, Measurement, and Problem Solving

Introduction to Chemistry

Science, Technology, and Chemistry

Chemistry is a branch of science that seeks to understand how matter behaves by studying the properties and interactions of atoms and molecules. Science feeds technology, and technology is the application of scientific knowledge for practical purposes.

• Science: The systematic study of the natural world through observation and experimentation.

• Technology: The application of scientific knowledge to solve problems and create useful products.

• Chemistry: The science that studies the composition, structure, properties, and changes of matter.

Matter and Non-Matter

Matter is anything that has mass and occupies space. Examples include water, air, and rocks. Non-matter includes things that do not have mass or occupy space, such as light and sound.

• Examples of matter: Water, air, wood, metal.

• Examples of non-matter: Heat, light, sound, energy.

The Scientific Method

Steps of the Scientific Method

The scientific method is an organized, logical approach used by scientists to do research and solve problems. It provides a method for scientists to verify the work of others.

1. Observation: Gathering data using the senses.

2. Hypothesis: A tentative, testable statement or prediction about what has been observed.

3. Experiment: A procedure that tests the hypothesis.

4. Conclusion: A judgment based on the information obtained from analyzing the data and checking it against the hypothesis.

• Independent variable: The variable that you plan to change during the experiment (manipulated variable).

• Dependent variable: The variable that responds to or changes because of the independent variable (observed variable).

Scientific Law and Theory

• Scientific Law: A brief statement that summarizes past observations and predicts future ones.

• Scientific Theory: A broader explanation for observations and laws, often predicting behavior that extends beyond the observations and laws on which they are founded.

Measurement and Problem Solving

Measurement Systems

• U.S. Customary System: Used for everyday measurements in the U.S.

• International System of Measurement (SI): Based on multiples of 10, used by most scientists.

• Prefixes: Used with different SI units to make the base unit smaller or larger.

Prefix	Abbreviation	Numerical Value	Power of 10	Size Relative to Base Unit
mega	M	1,000,000	$10^6$	1,000,000 times larger
kilo	k	1,000	$10^3$	1,000 times larger
centi	c	0.01	$10^{-2}$	100 times smaller
milli	m	0.001	$10^{-3}$	1,000 times smaller
micro	μ	0.000001	$10^{-6}$	1,000,000 times smaller

Base Units

• Time: second (s)

• Length: meter (m)

• Mass: kilogram (kg)

• Temperature: kelvin (K)

• Amount of substance: mole (mol)

Weight and Temperature

• Weight: A measure of the pulling force of gravity on an object. It can change with location due to differences in gravitational acceleration.

• Temperature: A measure of the average kinetic energy of particles of matter.

• Celsius scale: Freezing point of water = 0°C, Boiling point of water = 100°C

• Kelvin scale: Absolute zero = 0 K (no kinetic energy)

Temperature conversions:

• $K = ^\circ C + 273.15$

• $^\circ C = K - 273.15$

Volume

• Volume: The amount of space occupied by matter.

• Rectangular solid: $V = \text{length} \times \text{width} \times \text{height}$

• Liquid: Measured in mL using a graduated cylinder (read at the bottom of the meniscus).

• Irregular solid: Use water displacement method.

Density

• Density: The mass per unit volume of a substance.

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

• Units: g/cm³ for solids, g/mL for liquids.

Measurement, Accuracy, and Precision

• Measurement: The determination of a quantitative value using a device such as a ruler or balance.

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

• Precision: How close a series of measurements are to one another.

	Accurate	Not Accurate
Precise	Measurements are close to each other and to the true value.	Measurements are close to each other but not to the true value.
Not Precise	Measurements are not close to each other but average to the true value.	Measurements are neither close to each other nor to the true value.

• Percent error: $\frac{\text{measured value} - \text{accepted value}}{\text{accepted value}} \times 100$

Significant Figures

Rules for Significant Figures

• All nonzero numbers are significant.

• Zeros between nonzero digits are significant.

• Leading zeros are not significant.

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

• Counting numbers and defined constants have infinite significant figures.

When multiplying or dividing, round the answer to the same number of significant figures as the measurement with the fewest significant figures. When adding or subtracting, round to the least precise decimal place.

Scientific Notation

Expressing Numbers in Scientific Notation

• Express numbers as a multiple of two factors: a number (n) between 1 and 10, and ten raised to a power (exponent) (n × 10x).

• Numbers larger than 1 have a positive exponent; numbers smaller than 1 have a negative exponent.

Example: 1,392,000 km = $1.392 \times 10^6$ km

Example: 0.002650 mm = $2.650 \times 10^{-3}$ mm

Dimensional Analysis

Conversion Factors and Problem Solving

Dimensional analysis is a systematic approach to problem solving that uses conversion factors to move, or convert, from one unit to another.

• Conversion factor: A ratio of equivalent values having different units.

Example: 1 hr = 60 min is a conversion factor. Two ratios, both equal to 1, can be written as $\frac{1 \text{ hr}}{60 \text{ min}}$ or $\frac{60 \text{ min}}{1 \text{ hr}}$.

1. Identify what you are given and what you are looking for.

2. Look for one or more conversion factors to convert from what you are given to what you want.

3. Line up conversion factors so that units cancel.

4. Multiply the values above the line, divide by each number below the line.

5. Make sure the answer has the correct number of significant figures and units.

Physical and Chemical Properties and Changes

Properties of Matter

• Physical property: A characteristic that can be observed or measured without changing the substance's identity (e.g., color, mass, melting point).

• Chemical property: A property that describes a substance's ability to undergo changes that transform it into different substances (e.g., flammability, reactivity).

Physical and Chemical Changes

• Physical change: A change that does not alter the chemical composition of a substance (e.g., melting, freezing, dissolving).

• Chemical change: A change that results in the formation of new substances with different properties (e.g., rusting, burning).

Law of Conservation of Mass

• Mass is neither created nor destroyed in a chemical reaction or physical process.

Energy and Heat

Energy and Its Conservation

• Energy: The capacity to do work or produce heat.

• Law of Conservation of Energy: In any chemical reaction or physical process, energy is neither created nor destroyed.

Heat and Temperature

• Heat: Energy that flows from a warmer object to a cooler object.

• Endothermic process: A process in which heat is absorbed.

• Exothermic process: A process in which heat is released.

Heat Capacity and Specific Heat

• Heat capacity: The amount of heat needed to raise the temperature of an object by 1°C.

• Specific heat capacity: The amount of heat required to raise the temperature of 1 gram of a substance by 1°C.

• Formula: $q = m \cdot c \cdot \Delta T$ where $q$ = heat (J), $m$ = mass (g), $c$ = specific heat (J/g°C), $\Delta T$ = change in temperature (°C)

Calorimetry

• Calorimetry: The measurement of heat flow in a chemical or physical process.

• Calorimeter: An insulated device used to measure the amount of heat absorbed or released during a process.

• In calorimetry, the heat lost by the hot object equals the heat gained by the cold object (assuming no heat is lost to the surroundings).

Classification of Matter

Pure Substances and Mixtures

• Element: A pure substance that cannot be broken down into simpler substances by chemical means.

• Compound: A pure substance composed of two or more elements chemically combined in a fixed ratio.

• Mixture: A combination of two or more substances that are not chemically combined.

• Heterogeneous mixture: A mixture that is not uniform in composition (e.g., salad, sand and water).

• Homogeneous mixture (solution): A mixture that is uniform in composition throughout (e.g., salt water, air).

Separation of Mixtures

• Filtration: Separates solids from liquids using a porous barrier.

• Distillation: Separates substances based on differences in boiling points.

• Chromatography: Separates components of a mixture based on their movement through a stationary phase.