Atomic Structure and the Periodic Table

• The objectives of this lesson include understanding atomic structure and the differences between atoms, ions, molecules, compounds, and elements, and to also identify the organization of the periodic table and its relationship to valence and chemical bonds.

Atomic Structure

• The components of atomic structure include:
• Protons which have a positive charge
• Neutrons which have a neutral charge
• Electrons which have a negative charge

Atomic Theory

• John Dalton in 1803 proposed that matter has indivisible and indestructible atoms.
• J.J. Thomson in 1904 discovered the electron with the plum pudding model.
• Ernest Rutherford in 1911 discovered the small, dense nucleus, containing positively charged protons in his planetary model.
• Niels Bohr in 1913 suggested that electrons move in fixed energy levels, orbiting around the nucleus in fixed orbits.
• Erwin Schrödinger in 1926 described electrons as waves in orbitals, rather than fixed orbits, using the quantum mechanical model.
• James Chadwick in 1932 discovered neutrons, which stabilize the nucleus.
• Werner Heisenberg in 1927 formulated the uncertainty principle, which states that the position and velocity of an electron cannot be precisely known simultaneously.

Electronic Configuration

• Electronic configuration describes how electrons are distributed in the different levels and orbitals within an atom.

Quantum Numbers

• Principal quantum number (n) describes the energy level of an electron.
• Uses positive integers like: n=1,2,3,4
• Oxygen's (O) outermost electron: n=2n because it’s in the second energy level.
• Azimuthal quantum number (l) specifies the type of orbital where an electron is found.
• Values: l=0 (s), l=1 (p), l=2 (d), l=3 (f)
• Oxygen's (O) outermost electron: l=1, being in a p orbital.
• Magnetic quantum number (ml) dictates the orientation of an orbital in space.
• Values range from -l to +l
• Oxygen's (O) outermost electron: ml=-1.
• Spin quantum number (ms) indicates the direction of electron spin within the orbital.
• Options: ms=+1/2 or -1/2
• Oxygen's (O) outermost electron: ms = -1/2.

Calculating Quantum Numbers for Sodium Example

• Determining the electronic configuration for sodium, which has 11 electrons:1s² 2s² 2p⁶ 3s¹
• Identifying the quantum numbers for Na:
• n=3
• l=0
• m=0
• s= +1/2

Atoms, Ions, Elements, Molecules and Compounds

• An atom is the smallest unit of matter.
• Example: hydrogen (H), and oxygen (O).
• An ion is an atom with an electrical charge because of lost or gained electrons.
• Example: Sodium cation (Na+), a sodium atom has lost an electron or Chloride anion (Cl-), a chloride atom that has gained an electron.
• An element is composed of the same type of atoms and it can not be broken down.
• Example: Iron (Fe), gold (Au) and carbon (C).
• A molecule is a unit of two or more atoms, that are the same or different.
• Example: Water (H2O), comprised of hydrogen and oxygen atoms.
• A compound is a pure chemical substance made of two or more different elements chemically bonded in fixed proportions.
• Example: Carbon dioxide (CO2), composed of carbon and oxygen.

Periodic Table

• Groups (Vertical Columns 1, 2, and 13-18): elements with similar properties caused by the number of valence electrons that they share.
• Group 1 (Alkali Metals): highly reactive, with 1 valence electron.
• Group 2 (Alkaline Earth Metals): reactive, with 2 valence electrons.
• Group 13 (Boron Group): 3 valence electrons.
• Group 14 (Carbon Group): 4 valence electrons.
• Group 15 (Pnictogens): 5 valence electrons.
• Group 16 (Chalcogens): 6 valence electrons.
• Group 17 (Halogens): highly reactive, with 7 valence electrons.
• Group 18 (Noble Gases): non-reactive, with fully occupied valence shells.
• Periods (Horizontal Rows) indicate the number of energy levels (n) in an atom's electron configuration, however, there are varying numbers of electrons in the outermost shells.

Chemical Bonding

• The force that holds two or more atoms together and make them function as a unit.
• Measured through bond energy and indicates the energy amount needed to break the bond.
• Valence electrons form chemical bonds.

Types of bonds

• Ionic bonds forms when a (Metal) cedes an electron.
• Atoms that have high affinity to electrons.
• Composed of METAL + NON METAL
• Metallic Bonds are produced by the conjunction of linked metals, resulting in a dense cloud of electrons.
• Covalent Bonds forms between Non Metallic atoms that share electrons.
• Composed of NO METAL + NO METAL
• Polar covalent bonds are made of different elements.
• Apolar covalent bonds are made of elements of the same kind.

Electronegativity and Polar Covalent Bonds

• Electronegativity is the capacity for an atom to attract electrons.
• The polarity of the covalent bond is dependent on the differences in electronegativity between the atoms that make up the bond.

Medical Context

• Sodium chloride (NaCl) is a saline solution used to rehydrate patients.
• Glucose (C6H12O6) provides energy.
• Molecular oxygen (O2) is used by mask to improve blood oxygenation.
• Ferrous sulfate (FeSO4) is an iron supplement to treat existing anemia.
• Metallic copper (Cu) is to make medical instruments.