Ch. 2 study guide

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Biology: Atoms

Composition of Matter

All living organisms are composed of matter, which exists in three physical states: solid, liquid, and gas. Understanding the nature of matter is fundamental to biology, as it forms the basis for all biological structures and processes.

Matter is anything that occupies space and has mass.
Elements are pure substances that cannot be broken down into simpler substances by chemical means.
Compounds are substances formed when two or more elements combine in fixed ratios.
Example: Water (H2O) is a compound made of hydrogen and oxygen.

The Elements of Life

Life depends on a small subset of the elements found on Earth. These essential elements are required for the structure and function of living organisms.

About 20-25% of the 92 naturally occurring elements are essential for life.
Essential elements: Carbon, hydrogen, nitrogen, and oxygen make up approximately 96% of living matter.
Other important elements include calcium, potassium, phosphorus, and sulfur.
Trace elements are required in very small amounts but are vital for health (e.g., iron, iodine).

Atoms and Subatomic Particles

An atom is the smallest unit of an element that retains its chemical properties. Atoms are composed of subatomic particles, each with distinct properties.

Atom: Smallest unit of measurement in chemistry.
Proton: Positively charged particle.
Electron: Negatively charged particle.
Neutron: Neutral particle.
Protons and neutrons are located in the nucleus; electrons orbit the nucleus.
Atomic mass is the average mass of an atom, measured in daltons.

Isotopes

Isotopes are variants of a particular chemical element that have the same number of protons but different numbers of neutrons.

Isotopes have identical atomic numbers but different mass numbers.
Some isotopes are stable, while others are radioactive.

Radioactive Isotopes

Radioactive isotopes have unstable nuclei that decay over time, emitting radiation. This property is useful in biological research and medicine.

Unstable nucleus leads to loss of neutrons or other particles.
There are over 800 known radioactive isotopes.
Radioactive decay occurs naturally and at a stable rate, which can be measured.
Used in metabolic tracing and medical diagnostics.

Energy Levels of Electrons

Electrons in an atom occupy specific energy levels, which influence chemical behavior and bonding.

Energy is the capacity to do work.
Potential energy is energy stored due to an object's position or structure.
Electrons have different potential energies based on their location in electron shells.
Electron shells determine chemical behavior.
When the outer shell is not filled, the atom is unstable and will gain or lose electrons to become stable.
Shell capacities:
- Inner shell: 2 electrons
- Outer shell: 8 electrons
Valence electrons are the electrons in the outermost shell and are involved in chemical bonding.

Periodic Table Organization

The periodic table arranges elements based on their atomic structure and properties.

Period: Horizontal rows (side to side).
Group: Vertical columns (up and down).

Orbital Electrons

Orbitals are three-dimensional spaces around the nucleus where electrons are likely to be found.

Each orbital can hold a maximum of two electrons.
Chemical bonds form when atoms share or transfer valence electrons.

Chemical Bonds

Chemical bonds are forces that hold atoms together in molecules and compounds. The type of bond affects the properties of the substance.

Covalent bonds: Atoms share electrons.
Nonpolar covalent bonds: Electrons are shared equally.
Polar covalent bonds: Electrons are shared unequally; one atom is more electronegative and pulls electrons closer.
Ionic bonds: Electrons are transferred from one atom to another, creating charged ions.
- Cation: Positively charged ion.
- Anion: Negatively charged ion.
- Ionic bonds form between cations and anions.
Hydrogen bonds: Weak attractions between a hydrogen atom covalently bonded to one electronegative atom and another electronegative atom.
Van der Waals interactions: Weak attractions between molecules due to transient partial charges.

Electronegativity

Electronegativity is the ability of an atom to attract electrons in a covalent bond.

Atoms with higher electronegativity pull electrons closer, resulting in polar covalent bonds.

Molecule Shape and Function

The shape and size of molecules are critical for their function in biological systems. Molecular shape is determined by the arrangement of atoms and the type of chemical bonds.

Shape is determined by the arrangement of atoms' orbitals.
Shape affects how molecules interact and recognize each other.
Example: Water (H2O) has a bent shape with a bond angle of 104.5 degrees due to two covalent bonds and two lone pairs.

Chemical Reactions

Chemical reactions involve the making and breaking of chemical bonds, transforming substances into new products.

All chemical reactions are reversible.
Chemical equilibrium is reached when the forward and reverse reactions occur at the same rate, and the relative concentrations of reactants and products remain constant.
Equation for chemical equilibrium:

Oxidation: Removal of electrons from a molecule.
Reduction: Addition of electrons to a molecule.
Oxidation does not always require oxygen.
One molecule is oxidized, and another is reduced in redox reactions.

Summary Table: Types of Chemical Bonds

Bond Type	Description	Example
Covalent	Atoms share electrons	H2O (water)
Ionic	Electrons transferred; ions formed	NaCl (table salt)
Hydrogen	Weak attraction between H and electronegative atom	Between water molecules
Van der Waals	Weak interactions due to transient charges	Between nonpolar molecules

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