Atoms and Atomic Structure

Definition and Structure of the Atom

• Atom: The smallest particle of an element that retains its chemical properties.

• Nucleus: The dense central core of the atom, containing protons and neutrons.

• Proton: Positively charged particle found in the nucleus.

• Neutron: Neutral particle found in the nucleus.

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

Energy Levels (Electron Shells)

• Electrons occupy specific regions called energy levels or shells around the nucleus.

• Each energy level can hold a maximum number of electrons:

  • First energy level: 2 electrons

  • Second energy level: 8 electrons

  • Third energy level: 18 electrons

• All occupied energy levels must be full for an atom to be stable (noble gas configuration).

• Example: An atom with 8 electrons: 2 in the first level, 6 in the second. It is not stable because the second shell is not full.

Electron Configuration Examples

• 10 electrons: 2 in the first shell, 8 in the second (stable configuration).

• 16 electrons: 2 in the first shell, 8 in the second, 6 in the third (not stable; third shell not full).

Elements and the Periodic Table

Definition and Classification

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

• There are 90 naturally occurring elements; all are listed on the Periodic Table.

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

Atomic Number and Atomic Mass

• Atomic Number: Number of protons in the nucleus; also equals the number of electrons in a neutral atom.

• Atomic Mass: Sum of protons and neutrons in the nucleus.

• Example: Carbon has atomic number 6 (6 protons) and atomic mass 12 (6 protons + 6 neutrons).

Determining Subatomic Particles

• Protons = Atomic number

• Electrons = Atomic number (in a neutral atom)

• Neutrons = Atomic mass – Atomic number

• Example: Carbon: 6 protons, 6 electrons, 6 neutrons

Ions and Isotopes

• Ion: An atom or molecule with a net electric charge due to loss or gain of electrons.

• Positive ions (cations) have more protons than electrons; negative ions (anions) have more electrons than protons.

• Isotope: Atoms of the same element with different numbers of neutrons (e.g., Carbon-12, Carbon-13, Carbon-14).

Chemical Bonds and Compounds

How Elements Combine

• Compound: Substance formed when two or more different elements bond together (e.g., NaCl, H2O).

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

Covalent Bonds

• Formed when two atoms share electrons.

• Common in organic compounds.

• Example: Water (H2O) is formed by covalent bonds between hydrogen and oxygen.

Polar vs. Nonpolar Covalent Bonds

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

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

Ionic Bonds

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

• Example: Na+ + Cl- → NaCl

Hydrogen Bonds

• Weak bonds 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 (e.g., proteins, DNA).

Water: Structure, Properties, and Importance

Importance of Water

• Medium for chemical reactions.

• Makes up 75–90% of living organisms.

Hydrogen Bonding in Water

• Water molecules are attracted to each other via hydrogen bonds.

• Hydrogen bonds give water many of its unique properties.

Polarity of Water

• Water is a polar molecule due to the unequal sharing 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 (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, stabilizing environments.

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

• Versatile Solvent: Water dissolves many substances due to its polarity.

Solutions, Acids, and Bases

Solutions

• Solution: Homogeneous mixture of two or more substances.

• Solute: Substance dissolved (e.g., iced tea mix).

• Solvent: Substance doing the dissolving (e.g., water).

• Water is known as the "universal solvent" due to its ability to dissolve many substances.

pH, Acids, and Bases

• pH: Measure of hydrogen ion (H+) concentration in a solution; scale ranges from 0 (acidic) to 14 (basic).

• Acid: Substance that increases H+ concentration (e.g., HCl in water).

• Base: Substance that increases OH- concentration (e.g., NaOH in water).

• Examples:

  • Pure water: pH 7 (neutral)

  • Soda: pH 3 (acidic)

  • Hair remover: pH 13 (basic)

Chemical Equations

Writing and Interpreting Chemical Equations

• Chemical equations represent the transformation of reactants into products.

• Example:

• Reactants: Substances present before the reaction.

• Products: Substances formed by the reaction.

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

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