Atoms and Atomic Structure

Definition of an Atom

An atom is the smallest particle of an element that retains the properties of that element.

• Nucleus: The dense center of the atom containing protons and neutrons.

• Proton (p+): Positively charged particle found inside the nucleus.

• Neutron (n0): Particle with no charge, also found in the nucleus.

• Electron (e-): Negatively charged particle found in regions (energy levels) around the nucleus.

Energy Levels

Electrons occupy specific energy levels (shells) around the nucleus. Each level can hold a certain maximum number of electrons:

• First energy level: holds up to 2 electrons

• Second energy level: holds up to 8 electrons

• Third energy level: holds up to 18 electrons

• All energy levels being used must be full for an atom to be stable.

Example: An atom with 8 electrons will have 2 in the first level and 6 in the second. The second level is not full, so the atom is not stable.

Electron Configuration

• Electron configuration describes the arrangement of electrons in an atom's energy levels.

• Example: For an atom with 10 electrons (Neon), the configuration is 2 in the first level, 8 in the second (stable).

• Example: For an atom with 16 electrons (Sulfur), the configuration is 2 in the first, 8 in the second, 6 in the third (not full, not stable).

Elements and the Periodic Table

Definition of an Element

An element is a substance that cannot be broken down into simpler substances by chemical means.

• There are 90 naturally occurring elements.

• All elements are listed on the Periodic Table.

• Only 25 elements are essential for living things; 96% of the human body mass is composed of carbon (C), hydrogen (H), nitrogen (N), and oxygen (O).

Atomic Number and Atomic Mass

• Atomic Number: The number found above an element symbol on the periodic table; equals the number of protons (and electrons in a neutral atom).

• Atomic Mass: The number found below an element symbol; equals the sum of protons and neutrons.

Determining Subatomic Particles

• Number of protons = atomic number

• Number of electrons = atomic number (unless the atom is an ion)

• Number of neutrons = atomic mass - atomic number

Ions and Isotopes

Ions

Ions are charged particles formed when atoms gain or lose electrons.

• Number of protons never changes.

• Positive ions (cations) have lost electrons.

• Negative ions (anions) have gained electrons.

• Example: Na+ has lost one electron compared to neutral sodium.

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons.

• Example: Carbon-12 (6p/6n), Carbon-13 (6p/7n), Carbon-14 (6p/8n)

Chemical Bonds and Compounds

How Elements Combine

• Compound: A substance made of two or more different elements chemically bonded (e.g., NaCl, H2O).

• Molecule: A group of atoms held together by covalent bonds (e.g., O2).

Covalent Bonds

• Formed when two atoms share electrons.

• Example: H2O (water)

• Found in organic compounds.

• Polar Covalent Bond: Electrons are shared unequally, resulting in partial charges (e.g., H2O).

• Nonpolar Covalent Bond: Electrons are shared equally (e.g., H2).

Ionic Bonds

• Formed when atoms transfer electrons, resulting in oppositely charged ions that attract each other.

• Example: Na+ + Cl- → NaCl

Hydrogen Bonds

• Weak bonds that form between a hydrogen atom (already covalently bonded to a highly electronegative atom) and another electronegative atom.

• Important in holding water molecules together and stabilizing large biological molecules.

Water and Its Properties

Importance of Water

• Provides a medium for chemical reactions.

• Makes up 75-90% of living organisms.

Polarity of Water

• Water is a polar molecule due to the uneven distribution of electrons between oxygen and hydrogen.

• Oxygen has a stronger pull on electrons, making it slightly negative and hydrogen slightly positive.

Properties of Water

• Cohesion: Attraction between water molecules due to hydrogen bonding (responsible for surface tension).

• Adhesion: Attraction between water molecules and other substances (causes meniscus and capillary action).

• High Specific Heat Capacity: Water absorbs large amounts of heat before changing temperature, helping regulate temperature in organisms and environments.

• Evaporative Cooling: As water evaporates, it removes heat (e.g., sweating).

• Versatile Solvent: Water dissolves many substances due to its polarity, making it essential for biological processes.

Solutions

• A solution consists of a solute (substance dissolved, e.g., iced tea mix) and a solvent (substance doing the dissolving, e.g., water).

• Water is often called the "universal solvent" because it dissolves many substances.

pH, Acids, and Bases

pH Scale

• The pH scale measures the concentration of hydrogen ions (H+) versus hydroxide ions (OH-) in a solution.

• Ranges from 0 (most acidic) to 14 (most basic), with 7 being neutral.

Acids and Bases

• Acid: Substance that increases H+ concentration in solution (pH below 7). Example: HCl in water forms H+ and Cl-.

• Base: Substance that increases OH- concentration in solution (pH above 7). Example: NaOH in water forms Na+ and OH-.

pH Examples

• Pure water: pH 7.0 (neutral)

• Hair remover (Nair): pH 13.0 (strongly basic)

• Soda: pH 3.0 (acidic)

Chemical Equations

• Chemical equations represent chemical reactions, showing reactants and products.

• Reactants: Substances that start a reaction.

• Products: Substances formed by the reaction.

• Coefficients: Numbers before formulas indicating the number of molecules or atoms involved (e.g., 6CO2 means 6 molecules of CO2).

• Subscripts: Numbers written below and after element symbols indicating the number of atoms in a molecule (e.g., O2 means 2 oxygen atoms).

Example Equation:

• Reactants: ,

• Products: , , energy

Example: In , the subscript 2 means there are 2 oxygen atoms in each molecule.