General Chemistry: Ionic & Covalent Bonds, Ions, Periodic Table, Solutions, and Electrolytes

Ionic and Covalent Bonds

Bond Types and Electronegativity

Chemical bonds are the forces that hold atoms together in compounds. The two main types are ionic bonds and covalent bonds, distinguished by how electrons are distributed between atoms.

• Ionic Bonds: Formed by the transfer of electrons from a metal to a nonmetal, resulting in oppositely charged ions that attract each other.

• Covalent Bonds: Formed by the sharing of electrons between nonmetals.

• Electronegativity Difference: Determines bond type:

  • Nonpolar covalent: $< 0.4$

  • Polar covalent: $0.4 - 1.8$

  • Ionic: $> 1.8$

Example: The bond in NaCl is ionic, while the bond in H2O is polar covalent.

Bond Type	Electronegativity Difference
Nonpolar covalent	< 0.4
Polar covalent	0.4 - 1.8
Ionic	> 1.8

Monatomic and Polyatomic Ions

Monatomic Ions

Monatomic ions are ions formed from single atoms by gaining or losing electrons. They are classified as cations (positive) or anions (negative).

• Cations: Formed by metals losing electrons (e.g., Na+, Ca2+).

• Anions: Formed by nonmetals gaining electrons (e.g., Cl-, O2-).

Example: Na+ is a cation, Cl- is an anion.

Common Monatomic Ions

• 12 Cations: Barium (Ba2+), Calcium (Ca2+), Sodium (Na+), Potassium (K+), Magnesium (Mg2+), Aluminum (Al3+), Silver (Ag+), Copper I (Cu+), Copper II (Cu2+), Iron II (Fe2+), Iron III (Fe3+), Zinc (Zn2+)

• 6 Anions: Fluoride (F-), Chloride (Cl-), Bromide (Br-), Iodide (I-), Oxide (O2-), Sulfide (S2-)

Polyatomic Ions

Polyatomic ions are ions composed of two or more atoms covalently bonded, carrying a net charge.

Example: Nitrate (NO3-), Sulfate (SO42-), Ammonium (NH4+)

The Periodic Table: Metals, Metalloids, and Nonmetals

Classification of Elements

The periodic table organizes elements by increasing atomic number and groups them by similar chemical properties.

• Metals: Good conductors, malleable, ductile, tend to form cations.

• Metalloids: Have properties intermediate between metals and nonmetals.

• Nonmetals: Poor conductors, brittle, tend to form anions.

Example: Iron (Fe) is a metal, Silicon (Si) is a metalloid, Chlorine (Cl) is a nonmetal.

Important Elements and Their Functions

Trace Elements in the Human Body

Trace elements are required in small amounts for biological functions.

Element	Function
Iron (Fe)	Essential component of hemoglobin
Zinc (Zn)	Enzyme function, immune system
Copper (Cu)	Enzyme cofactor, iron metabolism
Manganese (Mn)	Bone formation, enzyme activation

Solutions: Molarity, Osmolarity, and Dilution

Molarity and Osmolarity

Molarity (M) is the concentration of a solution expressed as moles of solute per liter of solution. Osmolarity measures the total concentration of solute particles.

• Molarity: $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Osmolarity: $\text{Osmolarity} = M \times \text{number of particles per formula unit}$

Example: 1 M NaCl dissociates into 2 osmoles (Na+ and Cl-).

Dilution Equations

Dilution involves adding solvent to decrease the concentration of a solution.

• Equation: $C_1V_1 = C_2V_2$

  • $C_1$ = initial concentration

  • $V_1$ = initial volume

  • $C_2$ = final concentration

  • $V_2$ = final volume

Example: To dilute 1 M solution to 0.5 M, double the volume with solvent.

Electrolytes and Non-Electrolytes

Electrolytes

Electrolytes are substances that conduct electricity when dissolved in water due to the presence of ions.

• Strong Electrolytes: Completely dissociate in water (e.g., NaCl).

• Weak Electrolytes: Partially dissociate (e.g., acetic acid).

Non-Electrolytes

Non-electrolytes do not conduct electricity because they do not produce ions in solution (e.g., sugar).

• Usually polar compounds that dissolve but do not dissociate.

Example: Glucose is a non-electrolyte; NaCl is a strong electrolyte.

Other Key Concepts

Diatomic Elements

Seven elements exist naturally as diatomic molecules:

• Hydrogen (H2), Nitrogen (N2), Oxygen (O2), Fluorine (F2), Chlorine (Cl2), Bromine (Br2), Iodine (I2)

Common Covalent Compounds

Six covalent compounds to know:

• Water (H2O)

• Carbon dioxide (CO2)

• Methane (CH4)

• Ammonia (NH3)

• Oxygen (O2)

• Nitrogen (N2)

Solubility and Solvents

Water is a good solvent for polar and ionic compounds but not for non-polar solutes.

• Polar solutes: Dissolve in water due to interaction with water molecules.

• Non-polar solutes: Do not dissolve in water.

Example: Salt (NaCl) dissolves in water; oil does not.

Balancing Chemical Equations

Balancing equations ensures the same number of atoms of each element on both sides.

• Example: $\text{NaCl} \rightarrow \text{Na}^+ + \text{Cl}^-$

Milliequivalents (mEq)

Milliequivalents measure the chemical combining power of ions in solution.

• Used in medical and chemical calculations.

Example: 1 mEq Na+ = 1 mmol Na+ (for monovalent ions).

Osmolarity and Blood Cells

Osmolarity of IV solutions affects blood cells:

• Hypertonic: Higher osmolarity than blood; cells shrink (crenation).

• Hypotonic: Lower osmolarity; cells swell and may burst.

• Isotonic: Same osmolarity; cells remain unchanged.

Example: 0.9% NaCl is isotonic to blood.

Acids and Bases

Acids release H+ ions in solution; bases release OH- ions.

• Strong acids/bases: Completely dissociate.

• Weak acids/bases: Partially dissociate.

Example: HCl is a strong acid; acetic acid is a weak acid.

Additional info: Some content was inferred and expanded for clarity and completeness, including definitions, examples, and context for ions, electrolytes, and solution chemistry.