Chemical Elements, Atomic Structure, and Molecular Bonding

Chemical Elements, Matter, and Compounds

• Matter is defined as anything that has mass and occupies space.

• Elements represent the smallest units of matter possessing unique physical and chemical properties.

• Chemical symbols for elements typically originate from German or Latin terms.

• A compound is formed by combining two or more distinct elements:   $\text{Compound} = \text{Element} + \text{Element}$

• Emergent properties refer to characteristics that arise when elements combine into compounds, resulting in chemical and physical properties distinct from those of each individual element on its own.

• Humorous introductory exchange on atomic charge:

  • Atom A: "I lost an electron."

  • Atom B: "Are you positive?"

Element Abundance in the Human Body

• Essential elements constitute approximately $96\%$ of human body weight:

  • Oxygen ($\text{O}$): Atomic number $8$, accounting for $65.0\%$ of human body weight.

  • Carbon ($\text{C}$): Atomic number $6$, accounting for $18.5\%$ of human body weight.

  • Hydrogen ($\text{H}$): Atomic number $1$, accounting for $9.5\%$ of human body weight.

  • Nitrogen ($\text{N}$): Atomic number $7$, accounting for $3.3\%$ of human body weight.

• Secondary essential elements constitute approximately $4\%$ of human body weight:

  • Calcium ($\text{Ca}$): Atomic number $20$, accounting for $1.5\%$ of human body weight.

  • Phosphorus ($\text{P}$): Atomic number $15$, accounting for $1.0\%$ of human body weight.

  • Potassium ($\text{K}$): Atomic number $19$, accounting for $0.4\%$ of human body weight.

  • Sulfur ($\text{S}$): Atomic number $16$, accounting for $0.3\%$ of human body weight.

  • Sodium ($\text{Na}$): Atomic number $11$, accounting for $0.2\%$ of human body weight.

  • Chlorine ($\text{Cl}$): Atomic number $17$, accounting for $0.2\%$ of human body weight.

  • Magnesium ($\text{Mg}$): Atomic number $12$, accounting for $0.1\%$ of human body weight.

• Trace elements are required in extremely small quantities, making up less than $0.01\%$ of human body weight:

  • Boron ($\text{B}$)

  • Chromium ($\text{Cr}$ )

  • Cobalt ($\text{Co}$)

  • Copper ($\text{Cu}$)

  • Fluorine ($\text{F}$)

  • Iodine ($\text{I}$)

  • Iron ($\text{Fe}$)

  • Manganese ($\text{Mn}$)

  • Molybdenum ($\text{Mo}$)

  • Selenium ($\text{Se}$)

  • Silicon ($\text{Si}$)

  • Tin ($\text{Sn}$)

  • Vanadium ($\text{V}$)

  • Zinc ($\text{Zn}$)

Atomic Structure and Subatomic Particles

• An atom is composed of three primary subatomic particles:

  • Protons: Positively charged particles ($+$) located inside the central nucleus.

  • Neutrons: Electrically neutral particles ($\text{neutral}$) located inside the central nucleus.

  • Electrons: Negatively charged particles (-$format) orbiting the nucleus within a negative charge cloud.\n\n# Units of Measure, Atomic Number, Mass Number, and Atomic Mass\n\n- The Dalton (\text{Da}$), or atomic mass unit ($\text{amu}), is the standard unit of measurement for atomic mass.\n\n- Mass equivalencies and conversion:\n - 1.7 \times 10^{-24}\,\text{g} = 1\text{ Dalton (Da)}\n - Protons each weigh 1.7 \times 10^{-24}\,\text{g}$or$1\text{ Da}.\n - Neutrons each weigh 1.7 \times 10^{-24}\,\text{g}$or$1\text{ Da}.\n - Electrons are considered negligible in mass, weighing approximately \frac{1}{2000}$of a Dalton ($\frac{1}{2000}\,\text{Da}).\n\n- Subscript and Superscript Notation:\n - Subscript to the left of the chemical symbol denotes the Atomic Number.\n - Superscript to the left of the chemical symbol denotes the Mass Number.\n\n- Definitions:\n - Atomic Number: The total number of protons in an atom's nucleus. In a neutral atom, this also equals the total number of electrons.\n - Mass Number: The total number of protons plus neutrons in the nucleus.\n - Atomic Mass: The total mass of an atom measured in Daltons (\text{Da}).\n\n- Elemental Examples:\n - Helium (^4_2\text{He}):\n - Protons: 2\n - Neutrons: 2\n - Electrons: 2\n - Atomic mass: 4\text{ Da}\n - Carbon (^{12}_6\text{C}):\n - Protons: 6\n - Neutrons: 6\n - Electrons: 6\n - Atomic mass: 12\text{ Da}\n - Nitrogen (^{14}_7\text{N}):\n - Protons: 7\n - Neutrons: 7\n - Electrons: 7\n - Atomic number: 7\n - Mass number: 14\n - Atomic mass: 14\text{ Da}\n - All atoms containing 7 protons are classified as Nitrogen.\n - Oxygen (^{16}_8\text{O}):\n - Protons: 8\n - Neutrons: 8\n - Electrons: 8\n - Atomic mass: 16\text{ Da}\n\n# Stable and Unstable Isotopes\n\n- Atomic Identity: An atom is strictly defined by the number of protons it contains.\n\n- Isotopes: Atoms that have the exact same number of protons but a different number of neutrons relative to one another.\n\n- Stable Isotopes of Carbon:\n - Carbon-12 (^{12}_6\text{C}$): Contains$6$protons,$6$neutrons, and$6 electrons. Stable.\n - Carbon-13 (^{13}_6\text{C}$): Contains$6$protons,$7$neutrons, and$6 electrons. Stable.\n\n- Unstable (Radioactive) Isotopes:\n - Carbon-14 ( ^{14}_6\text{C}$): Contains$6$protons,$8$neutrons, and$6 electrons. Unstable.\n - An atom containing extra neutrons can become unstable. Over time, a neutron in an unstable nucleus splits into a proton (+$) and an electron ($-$format), releasing energy in the process.

  • Transmutation during radioactive decay:

  • When a neutron splits inside Carbon-14, the proton count increases from $6$ to $7$, and the neutron count decreases from $8$ to $7$.

  • Because the atom now contains $7$ protons, $7$ neutrons, and $7$ electrons, it is no longer Carbon; it transforms into Nitrogen ($^{14}_7\text{N}$) alongside released energy:       $\text{Unstable Carbon-14 } (^{14}_6\text{C}) \rightarrow \text{Nitrogen-14 } (^{14}_7\text{N}) + \text{energy}$

Electron Shell Structure and Chemical Valence

• Electron Shell Filling Rules:

  • First shell: Holds a maximum of $2\,e^-$ and is filled first.

  • Second shell: Holds a maximum of $8\,e^-$ and is filled second.

  • Third shell: Holds a maximum of $8\,e^-$ and is filled third.

• Chemical Nature of Atoms:

  • Valence shell: The outermost electron orbital/shell that contains electrons.

  • Valence electrons: Electrons residing in the valence shell that participate in chemical bonding.

  • Inert elements: Elements with completely filled valence shells. Because their valence shells are full, they are unreactive and do not gain or lose electrons.

Electronegativity

• Definition: Electronegativity measures an atom's affinity or drive to gain electrons to fill its outer orbital. The greater an atom's desire to gain electrons, the higher its electronegativity.

• Relative Electronegativity States:

  • Neutral Electronegativity: An atom is indifferent between gaining or losing electrons (e.g., an atom with a half-full valence shell requiring $4$ electrons to fill or lose).

  • High Electronegativity: An atom strongly prefers to gain electrons (e.g., needing $1$ electron to complete its outer shell rather than losing $7$ electrons).

  • Low Electronegativity: An atom strongly prefers to lose electrons (e.g., losing $1$ electron to reveal a full lower shell rather than gaining $7$ electrons).

• Electronegativity of Biological Elements:

  • Hydrogen ($\text{H}$): Contains $1$ valence electron. Neutral electronegativity; equally indifferent to gaining $1$ electron or losing $1$ electron.

  • Carbon ($\text{C}$): Contains $4$ valence electrons. Neutral electronegativity; equally indifferent to gaining $4$ electrons or losing $4$ electrons.

  • Nitrogen ($\text{N}$): Contains $5$ valence electrons. Slightly electronegative; prefers gaining $3$ electrons over losing $5$ electrons.

  • Oxygen ($\text{O}$): Contains $6$ valence electrons. Highly electronegative (the most electronegative of these four); strongly prefers gaining $2$ electrons over losing $6$ electrons.

Types of Chemical Bonds and Interactions

• Chemical bonds are formed using valence electrons, and all chemical interactions rely upon electronegativity.

• Non-Polar Covalent Bonds:

  • Occur when atoms with equal electronegativity share electrons equally to complete their outer shells.

  • Example — Methane ($\text{CH}_4$):

  • Carbon ($\text{C}$) has $4$ valence electrons; each of four Hydrogen ($\text{H}$) atoms has $1$ valence electron.

  • Because both Carbon and Hydrogen possess neutral electronegativity, they share electrons equally.

  • Carbon forms one shared pair with each of the four Hydrogen atoms.

  • Each Hydrogen atom attains $2$ electrons in its outer orbital, and the Carbon atom attains $8$ electrons in its outer orbital.

  • Result: $4$ non-polar covalent bonds are formed in $\text{CH}_4$

• Polar Covalent Bonds:

  • Occur when atoms with unequal electronegativity share electrons unequally.

  • Example — Water ($\text{H}_2\text{O}$):

  • Oxygen ($\text{O}$) is significantly more electronegative than Hydrogen ($\text{H}$).

  • Oxygen shares one electron pair with each of two Hydrogen atoms, allowing each Hydrogen to complete its shell with $2$ electrons and Oxygen to complete its shell with $8$ electrons.

  • Because Oxygen pulls the shared electrons closer to its nucleus, Oxygen acquires $2$ partial negative charges ($\delta^-$).

  • Because shared electrons are pulled away from the Hydrogen nuclei, each Hydrogen acquires a partial positive charge ($\delta^+$).

  • Result: $2$ polar covalent bonds are formed in $\text{H}_2\text{O}$

• Ionic Bonds:

  • Occur when the electronegativity difference between two atoms is so vast that one atom completely steals an electron from the other rather than sharing.

  • Example — Sodium Chloride ($\text{NaCl}$):

  • Sodium ($\text{Na}$) has very low electronegativity and seeks to lose $1$ electron.

  • Chlorine ($\text{Cl}$) has very high electronegativity and seeks to gain $1$ electron.

  • Chlorine steals $1$ electron from Sodium, resulting in a positively charged sodium cation ($\text{Na}^+$) and a negatively charged chloride anion ($\text{Cl}^-$).

  • An Ionic Bond is the electrostatic attraction between the positive ($+$) charge on one atom and the negative (-$format) charge on the other atom that holds them together.\n\n- Van der Waals Attractions / Forces:\n - Transient interactions driven by the random movement and distribution of electrons within orbitals.\n - Electrons are usually distributed evenly, but at any given split second, they may become concentrated in one specific region.\n - These forces last for only an instant and are weak individually, but collectively become significant.\n\n- Hydrogen Bonds:\n - Weak attractions between the partial positive charge (\delta^+$) on a Hydrogen atom of one molecule and the partial negative charge ($\delta^-) on an electronegative atom of a different molecule.\n - Example — Interaction between Water (\text{H}_2\text{O}$) and Ammonia ($\text{NH}_3):\n - In Ammonia (\text{NH}_3$), Nitrogen ($\text{N}$) is more electronegative than Hydrogen ($\text{H}$), as it prefers gaining$3$electrons over losing$5 electrons.\n - The partial positive Hydrogen of a water molecule is electrostatically attracted to the partial negative Nitrogen atom of an ammonia molecule, forming a Hydrogen Bond.\n\n# Chemical Reactions, Balanced Equations, and Equilibrium\n\n- Balanced Chemical Equations:\n - A chemical equation is balanced when the total number and type of atoms on the reactant side equal the total number and type of atoms on the product side.\n - Example — Formation of Water:\n    2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}\n - Reactant side count: 4$Hydrogen ($\text{H}$) atoms and$2$Oxygen ($\text{O}) atoms.\n - Product side count: 4$Hydrogen ($\text{H}$) atoms and$2$Oxygen ($\text{O}) atoms.\n\n- Chemical Equilibrium:\n - Defined as the condition in which the rate of the forward chemical reaction equals the rate of the reverse chemical reaction.\n - Reversible Water Synthesis and Breakdown:\n    2\text{H}_2 + \text{O}_2 \rightleftharpoons 2\text{H}_2\text{O}$$

  • Forward direction: Water is synthesized from hydrogen and oxygen.

  • Reverse direction: Water is broken down into hydrogen and oxygen.

  • At Chemical Equilibrium, water is synthesized and decomposed simultaneously at the exact same rate.