Levels of Biological Organization

• Life is organized from simple atoms to complex systems.

Levels of Organization

• Atoms are the fundamental units of matter, for instance, carbon, hydrogen, and oxygen.
• Molecules are formed when atoms bond together, examples being water and DNA.
• Organelles are structures within cells, like mitochondria.
• Cells are the basic units of life, such as bacteria and human cells.
• Tissues consist of groups of cells working together, such as muscle tissue.
• Organs are different tissues functioning as a unit, for example, the heart and lungs.
• Organ systems are groups of organs working together, like the digestive system.
• Organisms are living things, including humans and plants.
• Higher levels include: populations, communities, ecosystems, and the biosphere.

Atomic Number

• Atomic number represents the number of protons in an atom's nucleus and defines the element.
• Atoms contain protons (positive), neutrons (neutral), and electrons (negative).

Electron Shells

• Electrons orbit the nucleus in energy levels, also known as shells.
• Atoms achieve greater stability when their outer electron shells are full, influencing bond formation.
• Electrons travel in orbitals around the nucleus.
• The first electron shell, closest to the nucleus, holds up to 2 electrons.
• The second shell holds up to 8 electrons.
• The third shell can hold 8 or more electrons in larger atoms.

Isotope

• Isotopes are atoms of the same element but with varying numbers of neutrons.
• Carbon-12 (6 protons, 6 neutrons) and Carbon-14 (6 protons, 8 neutrons) are isotopes of carbon.
• Some isotopes are radioactive and decay over time.
• Atomic mass is the average mass of an element's isotopes, measured in Daltons or atomic mass units.

Chemical Reactions of Nitrogen

• Nitrogen is stable but reacts in specific ways.
• Nitrogen fixation involves bacteria converting N2 gas into usable forms for plants.
• Ammonia production is when nitrogen reacts with hydrogen to form NH3.
• Nitrate formation occurs when nitrogen reacts with oxygen to form nitrate (NO3-), which plants use.

Types of Bonds

• Bonding involves the interaction of electrons and stabilizes atoms, leading to the formation of molecules.
• Stronger bonds include covalent and ionic bonds.
• Weaker bonds include hydrogen bonds and Van der Waals forces.
• Covalent bonds involve the sharing of electrons between atoms.
• Ionic bonds involve the transfer of electrons between atoms.
• Hydrogen bonds are weak attractions between molecules.
• Van der Waals forces are the weakest intermolecular forces.
• Polar covalent bonds involve the unequal sharing of electrons, such as in H2O.
• Non-polar covalent bonds involve the equal sharing of electrons, such as in O2.
• Ionic bonds occur when one atom donates electrons to another.
• Hydrogen bonds are weak attractions, such as those found in DNA strands.

Hydrophilic & Hydrophobic Substances

• Hydrophilic substances dissolve in water (water-loving), such as salt and sugar.
• Hydrophobic substances do not dissolve in water (water-fearing), such as oils and fats.
• Amphipathic molecules have both hydrophilic and hydrophobic parts, such as phospholipids in cell membranes containing hydrophilic heads and hydrophobic tails.

Properties of Water

• Water makes up 60-90% of cell mass.
• It exists in solid, liquid, and gas states.
• Water has polar covalent bonds.
• Water is a universal solvent.
• Cohesion refers to water molecules sticking together, creating surface tension.
• Adhesion is when water sticks to other surfaces, leading to capillary action.
• Water has a high specific heat, resisting temperature changes.
• Water dissolves many substances due to its solvent ability.
• Ice floats because solid water is less dense than liquid water.

Molarity

• Molarity (M) is defined as moles of solute per liter of solution.
• A 1 M solution of NaCl contains 1 mole of NaCl in 1 liter of water.

Differences between Carbon and Hydrogen

• Carbon forms four covalent bonds, serving as the backbone of organic molecules.
• Hydrogen forms one covalent bond and is often part of larger molecules.

Sodium and Chlorine Interactions

• Sodium (Na) loses an electron and becomes Na+ (a cation).
• Chlorine (Cl) gains an electron and becomes Cl- (an anion).
• Opposite charges attract, forming NaCl (an ionic bond).

Properties of Molecules

• Molecules are two or more atoms bonded together.
• Compounds contain two or more different elements.
• Chemical formulas indicate atom ratios.
• Molecular size and shape determine functions.
• Polarity affects solubility, with polar substances dissolving in water and non-polar substances not.
• Bonding type affects reactivity and stability.

Chemical Bonding

• Chemical bonding is the joining of atoms to form molecules.
• Ionic bonding: Electrons are transferred.
• Covalent bonding: Electrons are shared.
• Hydrogen bonding: Is a weak attraction between molecules.

Creation of an Isotope

• Isotopes are created by adding or removing neutrons from a normal atom while keeping the number of protons constant.
• Example: Oxygen-16 (8 protons, 8 neutrons) becomes Oxygen-18 (8 protons, 10 neutrons).

pH & pH Equation

• pH measures the acidity or basicity of a solution on a scale of 0-14.
• Formula: pH = -log[H+]
• Acidic solutions (pH < 7) have more H+ ions, such as lemon juice.
• Neutral solutions (pH = 7) have equal concentrations of H+ and OH-, such as pure water.
• Basic solutions (pH > 7) have more OH- ions, such as bleach.