General Biology Study Notes: Measurement, Matter, Energy, and Atomic Structure

Chapter 1 – Science and Measurement

Key Components of Scientific Measurement

Understanding scientific measurement is foundational in biology and other sciences. This section covers SI units, prefixes, significant figures, scientific notation, and unit conversions.

• SI Units: The International System of Units (SI) is the standard for scientific measurement. Important base SI units include meter (length), kilogram (mass), second (time), ampere (electric current), kelvin (temperature), mole (amount of substance), and candela (luminous intensity).

• Derived Units: These are combinations of base units, such as volume (cubic meters, liters), density (kg/m3), and temperature in °C or K.

• Prefixes: Common prefixes include kilo- (103), deci- (10-1), centi- (10-2), milli- (10-3), and micro- (10-6).

Significant Figures and Scientific Notation

• Significant Figures: Indicate the precision of a measurement. All nonzero digits are significant; zeros between nonzero digits or after a decimal point are significant. For example, 0.00320 has three significant figures.

• Scientific Notation: Expresses numbers as a product of a coefficient and a power of ten (e.g., 3.2 × 103).

• Calculations: When multiplying/dividing, the result should have as many significant figures as the factor with the fewest significant figures. When adding/subtracting, the result should have the same number of decimal places as the measurement with the fewest decimal places.

Standard Notation and Unit Conversions

• Standard Notation: Ability to convert between scientific and standard notation is essential.

• Unit Conversions: Use conversion factors to change units (e.g., cm to m, g to kg). Dimensional analysis ensures correct unit cancellation.

Density and Volume Displacement

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

• Volume Displacement: Used to measure the volume of irregularly shaped objects by the amount of water they displace.

Chapter 2 – Matter and Energy

States and Properties of Matter

Matter exists in different states: solids, liquids, and gases. Each state has unique physical properties.

• Physical Properties: Characteristics that can be observed without changing the substance (e.g., melting point, density).

• Chemical Properties: Describe a substance's ability to undergo chemical changes (e.g., flammability, reactivity).

• Extensive vs. Intensive Properties: Extensive properties depend on the amount of matter (e.g., mass, volume), while intensive properties do not (e.g., density, melting point).

Mixtures and Pure Substances

• Heterogeneous Mixtures: Composition is not uniform throughout (e.g., salad, soil).

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

• Pure Substances: Elements and compounds with fixed composition and properties.

Energy and Its Forms

• Definition: Energy is the capacity to do work or transfer heat.

• Kinetic Energy: Energy of motion.

• Potential Energy: Stored energy due to position or composition.

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

Units and Measurement of Energy

• Common Units: Joule (J), calorie (cal), and their multiples.

• Temperature Scales: Celsius (°C), Kelvin (K), Fahrenheit (°F).

• Conversion: $K = ^\circ C + 273.15$

Specific Heat and Heat Calculations

• Specific Heat (c): Amount of heat required to raise the temperature of 1 g of a substance by 1°C.

• Heat Equation: $q = mc\Delta T$ Where $q$ = heat, $m$ = mass, $c$ = specific heat, $\Delta T$ = change in temperature.

• Phase Changes: Know terminology for melting, freezing, boiling, condensation, etc.

• Heating/Cooling Curves: Graphs showing temperature changes as heat is added or removed.

Chapter 3 – Atoms and Nuclear Chemistry

Atomic Structure

Atoms are the basic units of matter, composed of protons, neutrons, and electrons.

• Atomic Number (Z): Number of protons in the nucleus; defines the element.

• Mass Number (A): Total number of protons and neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Neutral Atoms: Have equal numbers of protons and electrons.

• Ions: Atoms that have gained or lost electrons, resulting in a net charge.

Periodic Table Organization

• Groups/Families: Vertical columns; elements in the same group have similar properties.

• Periods: Horizontal rows.

• Major Groups: Alkali metals (Group 1), alkaline earth metals (Group 2), halogens (Group 17), noble gases (Group 18).

Isotopic Notation and Atomic Mass

• Isotopic Symbol: $^{A}_{Z}\text{X}$, where X is the element symbol, A is the mass number, and Z is the atomic number.

• Atomic Mass Unit (amu): Defined as 1/12 the mass of a carbon-12 atom.

• Average Atomic Mass: Weighted average of all naturally occurring isotopes of an element.

Nuclear Chemistry and Radioactivity

• Types of Radiation: Alpha (α), beta (β), and gamma (γ) radiation.

• Balancing Nuclear Reactions: The sum of mass numbers and atomic numbers must be equal on both sides of the equation.

• Half-Life: The time required for half of a radioactive sample to decay. Half-life equation: $N = N_0 \left(\frac{1}{2}\right)^{t/t_{1/2}}$ Where $N$ = remaining amount, $N_0$ = initial amount, $t$ = elapsed time, $t_{1/2}$ = half-life.

• Shielding: Different types of radiation require different materials for protection (e.g., paper for alpha, aluminum for beta, lead for gamma).

Sample Table: Types of Radiation

Type	Symbol	Charge	Penetration	Shielding
Alpha	α	+2	Low	Paper, skin
Beta	β	-1	Moderate	Plastic, aluminum
Gamma	γ	0	High	Lead, concrete

Example: Carbon-14 is an isotope used in radiocarbon dating. It has a half-life of about 5730 years and decays by beta emission.