Atomic Structure & Chemical Bonding

Chemistry provides the foundation for understanding how drugs work, how the body processes nutrients, and why certain substances are toxic. The HESI A2 Chemistry section covers fundamental chemical concepts.

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Structure of an Atom

• Atomic number = number of protons (defines the element)
• Mass number = protons + neutrons
• Isotopes = atoms of the same element with different numbers of neutrons
• In a neutral atom, the number of protons equals the number of electrons

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Electron Configuration

Electrons are arranged in energy levels (shells) around the nucleus:

Valence electrons are the electrons in the outermost shell. They determine an element's chemical properties and how it bonds with other atoms.

• Noble gases (Group 18) have full outer shells and are chemically inert
• Metals tend to lose electrons → form positive ions (cations)
• Nonmetals tend to gain electrons → form negative ions (anions)
• The octet rule: Atoms tend to gain, lose, or share electrons to achieve 8 valence electrons (stable configuration)

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The Periodic Table

The periodic table organizes elements by increasing atomic number:

Key Groups:

• Group 1: Alkali metals (Li, Na, K) — 1 valence electron, highly reactive
• Group 2: Alkaline earth metals (Mg, Ca) — 2 valence electrons
• Group 17: Halogens (F, Cl, Br, I) — 7 valence electrons, very reactive
• Group 18: Noble gases (He, Ne, Ar) — full outer shell, inert

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Chemical Bonding

Ionic Bonds

• Formed when electrons are transferred from one atom to another
• Typically between a metal and a nonmetal
• The metal loses electrons → becomes a cation (+)
• The nonmetal gains electrons → becomes an anion (-)
• Example: NaCl (sodium chloride) — Na loses 1 electron, Cl gains 1 electron

Properties of ionic compounds:

• High melting and boiling points
• Conduct electricity when dissolved in water
• Form crystal lattice structures
• Usually solid at room temperature

Covalent Bonds

• Formed when electrons are shared between atoms
• Typically between two nonmetals
• Nonpolar covalent: Equal sharing (same or similar electronegativity) — e.g., O2, N2
• Polar covalent: Unequal sharing (different electronegativity) — e.g., H2O

Properties of covalent compounds:

• Lower melting and boiling points than ionic compounds
• Usually do not conduct electricity
• Can be gases, liquids, or solids at room temperature

Hydrogen Bonds

• Weak intermolecular forces (not true chemical bonds)
• Occur between a hydrogen atom bonded to F, O, or N and another F, O, or N atom
• Critical in water properties, DNA structure, and protein folding
• Individually weak but collectively strong in large numbers

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Electronegativity and Polarity

Electronegativity is a measure of how strongly an atom attracts shared electrons:

• Fluorine has the highest electronegativity (4.0 on the Pauling scale)
• Electronegativity generally increases across a period (left to right) and decreases down a group
• Polar molecules have an uneven distribution of charge due to differences in electronegativity
• Nonpolar molecules have an even distribution of charge

Water (H2O) is polar:

• Oxygen is more electronegative than hydrogen
• Electrons spend more time near oxygen → partial negative charge (delta-)
• Hydrogen atoms have partial positive charge (delta+)
• This polarity gives water its unique properties: cohesion, adhesion, surface tension, high specific heat, universal solvent

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States of Matter

Phase Changes:

• Melting: Solid → Liquid (absorbs energy)
• Freezing: Liquid → Solid (releases energy)
• Evaporation/Boiling: Liquid → Gas (absorbs energy)
• Condensation: Gas → Liquid (releases energy)
• Sublimation: Solid → Gas (dry ice)
• Deposition: Gas → Solid

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Mixtures vs. Compounds

Homogeneous mixture = uniform throughout (solutions like saline) Heterogeneous mixture = non-uniform composition (blood, salad dressing)

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Acids, Bases & pH in Biological Context

Understanding pH is critical for interpreting lab values and understanding body chemistry:

Acidosis (blood pH < 7.35) and alkalosis (blood pH > 7.45) are both dangerous conditions:

• Respiratory acidosis: Too much CO2 retained (hypoventilation)
• Respiratory alkalosis: Too much CO2 exhaled (hyperventilation)
• Metabolic acidosis: Too many acids or loss of bicarbonate (diabetic ketoacidosis, renal failure)
• Metabolic alkalosis: Loss of acids or excess bicarbonate (prolonged vomiting)

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Moles and Molar Mass

While not heavily tested, basic mole concepts may appear:

• Avogadro's number: 6.022 x 10^23 particles per mole
• Molar mass: Mass of one mole of a substance (in g/mol), equals the atomic/molecular weight from the periodic table
• Example: Water (H2O) has a molar mass of 18.02 g/mol (2 x 1.01 for H + 16.00 for O)
• Molarity (M): Concentration expressed as moles of solute per liter of solution

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Important Elements in Human Biology

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Electrolytes in the Body

Electrolytes are ions that conduct electricity when dissolved in water. They are critical for nursing:

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Solutions and Concentration

Solution = solute (dissolved substance) + solvent (dissolving substance, usually water)

Types of solutions based on tonicity (osmolarity):

Clinical significance: IV fluid selection depends on the patient's hydration status and electrolyte balance. Normal saline (0.9% NaCl) is isotonic and used most commonly because it does not change cell volume.