Basic Chemistry Fundamental Concepts

Matter, Mass, and Weight

• Matter: Defined as anything that occupies space and has mass. All living and nonliving things are composed of matter.

• Mass: The specific amount of matter contained within an object.

• Weight: The gravitational force acting upon an object of a given mass. For example, the weight of an apple is the result of the force of gravity pulling on its mass.

• Comparison of Mass and Weight in Space: While an astronaut's mass remains constant regardless of location because the amount of matter in their body does not change, their weight would be significantly lower in outer space due to the decreased gravitational pull compared to the Earth's surface.

• Measurement Standards:

  • Kilogram ($\text{kg}$): The base unit for mass in the International System of Units ($\text{SI}$).

  • The Planck Constant ($h$): The modern definition of the unit of mass is based on the fixed value of the Planck constant, defined as $6.62607015 \times 10^{-34}\,\text{kg} \cdot \text{m}^2/\text{s}$. This standard uses "invariants of nature" to ensure consistency, as physical objects with mass constantly emit energy and change over time.

  • Common Equivalents: A 2.2-pound lead weight and 1 liter ($1\,\text{L}$ or $1.06\,\text{qt}$) of water each possess a mass of approximately $1\,\text{kg}$.

  • Gram ($\text{g}$): An object with $1/1000$ the mass of a kilogram has a mass of $1\,\text{g}$.

Elements and Atoms

• Element: The simplest type of matter, possessing unique chemical properties. An element is composed of only one kind of atom.

• Major Elements in the Human Body: Approximately $96\%$ of the body's weight is derived from four elements:

  • Oxygen ($O$): The most abundant element in both the human body (by weight) and the Earth's crust.

  • Carbon ($C$): Crucial to body chemistry because of its propensity to form covalent bonds with itself and other molecules.

  • Hydrogen ($H$).

  • Nitrogen ($N$).

• Trace Elements: Many elements are present only in minute amounts but perform essential roles.

• Variable States of Elements: Elements can exist in different forms and functional roles within the body. For example, mineralized calcium provides structure to the solid matrix of bones, whereas dissolved calcium regulates nervous system signaling and enzyme activities.

• Atomic Symbols: Elements are represented by symbols, typically the first letter or letters of their name (e.g., $C$ for Carbon, $H$ for Hydrogen, $Ca$ for Calcium). Some symbols are derived from Latin, Greek, or Arabic names, such as Sodium ($Na$ from the Latin natrium).

• Atom: Derived from the word atomos (meaning indivisible), it is the smallest particle of an element that retains that element's chemical characteristics.

Detailed Breakdown of Common Elements in the Human Body

• Hydrogen ($H$):

  • Atomic Number: $1$

  • Mass Number: $1$

  • Atomic Mass: $1.008\,\text{Da}$

  • Percent by Weight: $9.5\%$

  • Percent by Number of Atoms: $63.0\%$

• Carbon ($C$):

  • Atomic Number: $6$

  • Mass Number: $12$

  • Atomic Mass: $12.01\,\text{Da}$

  • Percent by Weight: $18.5\%$

  • Percent by Number of Atoms: $9.5\%$

• Nitrogen ($N$):

  • Atomic Number: $7$

  • Mass Number: $14$

  • Atomic Mass: $14.01\,\text{Da}$

  • Percent by Weight: $3.3\%$

  • Percent by Number of Atoms: $1.4\%$

• Oxygen ($O$):

  • Atomic Number: $8$

  • Mass Number: $16$

  • Atomic Mass: $16.00\,\text{Da}$

  • Percent by Weight: $65.0\%$

  • Percent by Number of Atoms: $25.5\%$

• Fluorine ($F$): Atomic Number $9$, Mass Number $19$, Atomic Mass $19.00\,\text{Da}$; present in Trace amounts.

• Sodium ($Na$): Atomic Number $11$, Mass Number $23$, Atomic Mass $22.99\,\text{Da}$; Weight: $0.2\%$, Atoms: $0.3\%$.

• Magnesium ($Mg$): Atomic Number $12$, Mass Number $24$, Atomic Mass $24.31\,\text{Da}$; Weight: $0.1\%$, Atoms: $0.1\%$.

• Phosphorus ($P$): Atomic Number $15$, Mass Number $31$, Atomic Mass $30.97\,\text{Da}$; Weight: $1.0\%$, Atoms: $0.22\%$.

• Sulfur ($S$): Atomic Number $16$, Mass Number $32$, Atomic Mass $32.07\,\text{Da}$; Weight: $0.3\%$, Atoms: $0.05\%$.

• Chlorine ($Cl$): Atomic Number $17$, Mass Number $35$, Atomic Mass $35.45\,\text{Da}$; Weight: $0.2\%$, Atoms: $0.03\%$.

• Potassium ($K$): Atomic Number $19$, Mass Number $39$, Atomic Mass $39.10\,\text{Da}$; Weight: $0.4\%$, Atoms: $0.06\%$.

• Calcium ($Ca$): Atomic Number $20$, Mass Number $40$, Atomic Mass $40.08\,\text{Da}$; Weight: $1.5\%$, Atoms: $0.31\%$.

• Iron ($Fe$): Atomic Number $26$, Mass Number $56$, Atomic Mass $55.85\,\text{Da}$; present in Trace amounts.

• Iodine ($I$): Atomic Number $53$, Mass Number $127$, Atomic Mass $126.9\,\text{Da}$; present in Trace amounts.

Atomic Structure

• Major Subatomic Particles:

  • Neutron: Located in the nucleus; carries no electrical charge; mass is approximately equal to that of a proton.

  • Proton: Located in the nucleus; carries one positive ($+1$) charge.

  • Electron: Located in the electron cloud orbiting the nucleus; carries one negative ($-1$) charge.

• Electrical Neutrality: Each atom has an equal number of protons and electrons, causing the positive and negative charges to cancel each other out.

• Spatial Characteristics:

  • Nucleus: The center of the atom containing protons and neutrons. It accounts for $99.97\%$      of the atom's mass but only $1/10,000,000,000,000$ ($1$ ten-trillionth) of its volume.

  • Electron Cloud: The region where electrons orbit the nucleus; it occupies the vast majority of the atom's volume.

Atomic and Mass Numbers

• Atomic Number: Equal to the number of protons in each atom. Since protons and electrons are equal in a neutral atom, the atomic number also indicates the number of electrons. Each element is uniquely defined by this number.

• Mass Number: The total number of protons plus the number of neutrons in each atom. Electrons are excluded because they have very little mass.

• Sample Calculation for Potassium ($K$):

  • Atomic Number = $19$

  • Mass Number = $39$

  • Number of Protons: $19$

  • Number of Electrons: $19$

  • Number of Neutrons: $39 - 19 = 20$

• Elements: There are $90$ naturally occurring elements; others have been synthesized through the alteration of atomic nuclei.

Isotopes and Atomic Mass

• Isotopes: Forms of the same element that have the same number of protons and electrons but different numbers of neutrons. They share the same atomic number but differ in mass number.

• Hydrogen Isotopes:

  • Hydrogen ($^{1}H$): $1$ proton, $0$ neutrons, $1$ electron.

  • Deuterium ($^{2}H$): $1$ proton, $1$ neutron, $1$ electron.

  • Tritium ($^{3}H$): $1$ proton, $2$ neutrons, $1$ electron.

• Relative Atomic Mass Units:

  • Dalton ($Da$) or Unified Atomic Mass Unit ($u$): Defined as $1/12$ the mass of a Carbon-12 ($^{12}C$) atom.

  • $^{12}C$ has an atomic mass of exactly $12\,\text{Da}$.

• Atomic Mass of an Element: The average mass of all naturally occurring isotopes of that element, weighted by their relative abundance. Because Carbon exists mostly as $^{12}C$ but includes small amounts of $^{13}C$ (6 protons, 7 neutrons), its average atomic mass is $12.01\,\text{Da}$.

Clinical Uses of Atomic Particles

• X-rays: Radiation formed when electrons move from a high energy state to a lower one. These are used to detect bone breaks and dental caries (cavities).

• Mammograms: Low-energy radiographs used to reveal tumors in breast tissue, which are slightly denser than normal tissue.

• Computed Tomography (CT) Scan: Uses computer analysis of several radiographic "slices" taken at different body locations to assemble a three-dimensional image. These are used to detect tumors and other internal abnormalities.

The Mole and Molar Mass

• Avogadro's Number: The specific quantity $6.022 \times 10^{23}$, which defines one mole ($\text{mol}$).

• Mole: A unit representing Avogadro's number of entities (atoms, ions, or molecules).

• Molar Mass: The mass of one mole of a substance expressed in grams ($\text{g}$).

• Equivalency: The atomic mass of an entity in unified atomic mass units is numerically equal to its molar mass in grams.

  • Example: Carbon atoms have an atomic mass of $12.01\,\text{Da}$, so $12.01\,\text{g}$ of carbon contains one mole of carbon atoms.

  • Example: $1.008\,\text{g}$ of hydrogen contains the same number of atoms as $12.01\,\text{g}$ of carbon.

Electrons and Chemical Bonding

• Chemical Level of Organization: Interaction of atoms to form molecules through the sharing or transferring of valence electrons.

• Energy Levels (Electron Shells): Organized as concentric rings around the nucleus.

  • Innermost Shell: Holds a maximum of $2$ electrons.

  • Remaining Shells: Hold a maximum of $8$ electrons.

  • Valence Shell: The outermost shell; its electron count determines the atom's chemical nature.

• Inert Atoms: Atoms with full valence shells (e.g., Helium with 2 electrons) do not naturally form bonds.

• Octet Rule: The tendency of atoms to combine until they have 8 electrons in their valence shell (or 2 for hydrogen to reach a full shell).

• Electronegativity: The ability of an atom's nucleus to attract electrons.

  • Strong Electronegativity: Atoms lacking only 1 or 2 electrons (e.g., Oxygen).

  • Weak Electronegativity: Atoms lacking 6 or 7 electrons.

  • Hydrogen Exception: Hydrogen lacks only 1 electron but has a lower electronegativity than other atoms with more protons.

Bond Types Based on Electronegativity

• Nonpolar Covalent Bond: Formed when atoms have the same electronegativity and share electrons equally (e.g., $H_2$).

• Polar Covalent Bond: Formed when there is a small difference in electronegativities, leading to unequal sharing (e.g., water, $H_2O$).

• Ionic Bond: Formed when atoms have very different electronegativities, leading to a complete transfer of electrons (e.g., $NaCl$).

Ionic Bonds and Ions

• Ion: A charged particle formed when an atom loses or gains electrons.

  • Cation: A positively charged ion formed when a weakly electronegative atom loses an electron (e.g., Sodium ion, $Na^+$).

  • Anion: A negatively charged ion formed when an atom gains an electron (e.g., Chloride ion, $Cl^-$).

• Ionic Bonding: The attraction between oppositely charged cations and anions. Sodium ($Na^+$) and Chloride ($Cl^-$) form an array called sodium chloride ($NaCl$), or table salt.

• Important Ions in the Human Body:

  • Calcium ($Ca^{2+}$): Bone/teeth part, blood clotting, muscle contraction, neurotransmitter release.

  • Sodium ($Na^+$): Membrane potentials, water balance.

  • Potassium ($K^+$): Membrane potentials.

  • Hydrogen ($H^+$), Hydroxide ($OH^-$), Bicarbonate ($HCO_3^-$), Ammonium ($NH_4^+$): Acid-base balance.

  • Chloride ($Cl^-$): Water balance.

  • Phosphate ($PO_4^{3-}$): Bone/teeth, energy exchange ($ATP$), acid-base balance.

  • Iron ($Fe^{2+}$): Red blood cell formation.

  • Magnesium ($Mg^{2+}$): Necessary for enzymes.

  • Iodide ($I^-$): Present in thyroid hormones.

Covalent Bonds

• Formation: Atoms share one or more pairs of electrons due to similar electronegativities.

• Molecule: The structure resulting from a covalent bond that behaves as an independent unit.

• Single Covalent Bond: Sharing of one pair of electrons; represented by a single line (e.g., $H-H$).

• Double Covalent Bond: Sharing of two pairs (4 electrons); represented by a double line (e.g., $O=C=O$ for Carbon Dioxide, $CO_2$).

• Polar Molecules: Molecules like water ($H_2O$) that are electrically asymmetric due to polar covalent bonds. The oxygen side is slightly more negative ($\delta^-$) because it pulls the electron cloud closer, while the hydrogen side is slightly more positive ($\delta^+$).

Molecules and Compounds

• Molecule: Two or more atoms chemically combined. Can be the same type ($H_2$) or different ($H_2O$).

• Compound: A substance resulting from the combination of two or more different types of atoms.

  • $H_2O$ is both a molecule and a compound.

  • $H_2$ is a molecule but not a compound.

  • $NaCl$ is a compound but not a molecule because it forms an organized array rather than distinct independent units.

• Properties Transformation: Properties change when elements combine. Flammable Hydrogen and fire-promoting Oxygen form nonflammable Water. Toxic Chlorine and explosive Sodium form safe Table Salt.

• Molecular Mass: Determined by adding the atomic masses of all atoms or ions in the formula.

  • Example for $NaCl$: $22.99\,\text{(Na)} + 35.45\,\text{(Cl)} = 58.44\,\text{Da}$.

• Glucose ($C_6H_{12}O_6$): Indicates 6 carbon, 12 hydrogen, and 6 oxygen atoms.

Intermolecular Forces and Hydrogen Bonds

• Intermolecular Forces: Weak electrostatic attractions between oppositely charged parts of molecules or between ions and molecules. No electron exchange occurring.

• Hydrogen Bond: Formed when a positively charged hydrogen atom (already covalently bonded to $O$, $N$, or $F$) is attracted to the negative end of another polar molecule.

• Significance: Essential for the unique properties of water, stabilizing the three-dimensional shapes of proteins and nucleic acids (intramolecular hydrogen bonds).

Solubility and Dissociation

• Solubility: The ability of one substance to dissolve in another. Charged/polar substances (sugar, salt) dissolve in water; nonpolar substances (oil) do not.

• Dissociation: The process where ions in an ionic compound separate when dissolved in water because they are attracted to the positive or negative ends of water molecules.

• Electrolytes: Cations and anions that dissociate in water and can conduct an electric current (e.g., those detected by an $ECG$ for heart activity).

• Nonelectrolytes: Molecules that dissolve without dissociating and do not conduct electricity (e.g., pure water or glucose).

• Homeostasis: Electrolyte balance is critical for hydration, blood $pH$, and muscle/nerve function. Sport drinks or intravenous solutions may be used to restore balance after prolonged exercise or in emergencies.