Atomic Theory and Structure

• Democritus first used the term "Atom," meaning non-divisible.
• John Dalton proposed the atom is a tiny, solid, indestructible sphere.
• Joseph Thomson discovered the electron and theorized that an atom is electrons floating around in positively charged material.
• Ernest Rutherford introduced a new atomic theory featuring a tiny central nucleus with orbiting electrons.
• Bohr theorized electrons orbit the nucleus in shells, each with a different energy level, emitting unique colors when transitioning to lower energy shells.
• Atoms are neutral, and the number of protons equals the number of electrons.
• The formula 2(n^2) determines the number of electrons that can occupy a shell, where n represents the shell number.
• A valence electron resides in the outermost shell.
• Lewis dot diagrams represent only the valence electrons around the element symbol.
• DeBroglie theorized all matter exhibits wave properties, with the number of waves equaling electron energy levels.
• Schrodinger developed a quantum mechanical model to find the probability of an electron's location.
• To determine the radius of density clouds, the distance between bonded atoms' nuclei is measured and divided by two.
• As energy levels rise, the likelihood of finding electrons farther from the nucleus increases.
• Effective nuclear charge (Zeff) is experienced by valence electrons.
• Radius increases as Zeff decreases.
• Larger atomic radii appear in the bottom left of the periodic table.
• The number at the bottom of an element on the periodic table represents its average atomic mass, which also equals the average number of protons plus neutrons.
• Isotopes are atoms with the same number of protons but differing numbers of neutrons.
• Isotope notation includes the mass number on the top-left and the atomic number on the bottom-left of the element symbol.
• Electromagnetic force causes the nucleus to want to break apart because positives repel.
• Strong nuclear force is 40 times stronger than electromagnetic force.
• A nucleus with an unstable number of protons and neutrons is called a radioisotope.
• A radioisotope decays until it reaches a stable nucleus.
• Every proton needs a neutron as a "partner."
• Alpha decay is a type of radioactive decay producing an alpha particle (helium-4).
• Beta minus decay occurs in nuclei with too many neutrons, where one neutron decays into a proton, releasing a beta minus particle carried by a high speed electron.
• Atomic mass units remain undefined but are roughly equivalent to the mass of one proton or neutron.
• Average atomic mass represents the weighted average atomic mass of an element's isotopes.
• Isotopic abundance is the percentage of each naturally occurring isotope for each element.
• The periodic table is organized into groups of columns and periods of rows.
• Elements are stable when their valence shells are full; Group 18 is all full.
• Groups 1, 2, and 3 tend to lose valence electrons, while groups 15, 16, and 17 tend to gain valence electrons.
• Cations lose electrons, resulting in a positive charge.
• Anions gain electrons, resulting in a negative charge.
• Cations become smaller when they lose a complete electron shell.
• Ionization energy measures the energy needed to remove an electron from an atom.
• High ionization energy indicates that the atom holds its electrons strongly.
• Low ionization energy indicates that the atom holds its electrons loosely.
• The first ionization energy refers to the energy required to remove the first electron.
• The first ionization energy decreases as radius increases because electrons are farther from the nucleus; they feel a lower force of attraction.
• Ionization energy is higher in the top right of the periodic table.
• Reactivity generally increases as ionization energy decreases.
• Elements in the bottom left of the periodic table exhibit high reactivity.
• The opposite of ionization energy is electron affinity.
• Electron affinity is the energy either released or absorbed when an atom gains an electron.
• Gaining an electron, if it makes an atom stable, releases energy.
• Gaining an electron, if it makes an atom unstable, requires energy, and the ion will quickly lose the new electron.
• Electronegativity measures an atom's ability to attract electrons when bonded to another atom.
• In a bond, the atom with greater electronegativity holds the shared electrons closer.
• Electronegativity increases as radius decreases.
• Elements in the top right of the periodic table have high electronegativity.
• Noble gases lack electronegativity because they do not bond, having full valence shells.
• Full valence shells confer stability; atoms gain, lose, or share electrons to achieve eight electrons in their valence shell.
• Ionic bonds form from atoms giving one or more electrons to another atom.
• The atom that gives away electrons gains a positive charge, while the atom that accepts them gains a negative charge.
• Atoms with opposite charges attract each other.
• Ionic compounds consist of compounds held together through ionic bonds.
• Transition elements may lose different numbers of electrons.
• Binary ionic compounds consist of positive cations and negative anions.
• Type one compounds have only one cation/charge.
• Type two compounds transition metals can form 2 or more cations with different charges.
• For type one compounds, the cation name is written first, followed by the anion with the suffix -ide (example: Sodium Bromide).
• For type two compounds, the charge of the metal must be indicated (example: Iron (II) sulfide).
• An older naming system for type two involves using Latin names with suffixes -ic for the higher charge and -ous for the lower charge (Example: Iron (III) becomes ferric and Iron (II) becomes ferrous).
• Covalent bonds form when multiple atoms share electrons.
• In a covalent bond, an unpaired electron in one atom pairs with an unpaired electron in another atom.
• Together, two atoms share their unpaired electrons, so both have a pair.
• A compound of covalent bonds forms a covalent compound.
• Seven elements (nitrogen to iodine) exist as diatomic molecules in their elemental form.
• The electronegativity difference between two bonding atoms determines the bond type.
• A non-polar covalent bond is formed when the electronegativity difference is less than or equal to 0.4, which means the electrons are shared equally.
• A covalent bond is formed when the electronegativity difference is between 0.4 and 1.7, meaning the shared electrons are held more closely by one atom.
• An ionic bond is formed when the electronegativity difference is greater than 1.7, meaning sharing is so unequal that one atom attains a negative charge, leaving the other with a positive charge.
• Type three molecular compounds are named by starting with the less electronegative element, followed by the more electronegative element with the suffix -ide.
• Prefixes are added to indicate the number of each atom per molecule in front of each element name, with the exception that "mono" is not written for the first element. An example is Carbon dioxide.
• Molecular compounds sometimes gain more electrons to fill their valence shell, the electrons form an ionic bond with a cation, and these are called polyatomic ions
• Polyatomic ions are charged; therefore they can bond with other ions in an ionic bond.
• An ionic compound containing polyatomic ions is named by writing the cation followed by the anion, leaving the endings as is.
• Ionic solids like sodium chloride and magnesium fluoride exhibit predictable properties.
• At room temperature, ionic solids are shiny, hard, brittle crystals because this is the most efficient way to pack opposite charges together, which are attracted to each other.
• The solubility of ionic solids in water is usually high due to ion-dipole forces.
• Water's polar nature means the positive H side of the water dipole attracts and surrounds the negative ions, while the negative O side of the water attracts and surrounds the positive ions.
• Electrical conductivity when dissolved indicates it will conduct.
• Ionic solids have high melting points.
• Molecular compounds also have their properties:
• At room temperature, molecular compounds exist as soft solids, liquids, or gases.
• The solubility in water can be either soluble or non-soluble and it depends on the molecule's polarity; polar molecules dissolve in water because water is also polar.
• Molecular compounds are poor conductors of electricity.
• Molecular compounds have a low melting point.
• All organic compounds are molecular compounds.
• The old naming system for acids that do not contain oxygen is Hydro_ic acid. (Example: hydrochloric acid, hydrofluoric acid)
• The modern name for acids that do not contain oxygen is to add "aqueous" and name it like a type 1 ionic compound.
• For acids that contain oxygen, if the root word ends in -ate, the suffix -ic acid is used. Acetic acid is an example.
• For acids that contain oxygen, if the root word ends in -ite, replace it with -ous acid. Examples: nitrous acid, chlorous acid.

Explore the evolution of atomic theory from Democritus to Schrodinger, covering Dalton, Thomson, Rutherford, and Bohr's models. Understand electron configuration, valence electrons, and Lewis dot diagrams. The significance of quantum mechanics and wave-particle duality are discussed.