Chemistry Chapter on Solutions and Drug Delivery

Types of Chemical Equations

• Water is an extremely weak electrolyte, existing mostly as H₂O molecules.
• When sodium hydroxide and hydrochloric acid solutions mix, the ionic and molecular species are: Na⁺ + OH⁻ + H⁺ + Cl⁻ → Na⁺ + Cl⁻ + H₂O
• Sodium (Na⁺) and chloride (Cl⁻) ions do not change in the reaction.
• The net ionic reaction is: H⁺ + OH⁻ → H₂O

How to write Net Ionic Equation?

• Begin with the balanced molecular equation.
• Write the ionic equation, showing the strong electrolytes.
• Identify any precipitates using solubility rules.
• Cancel similar ions on both sides of the ionic equation.

Chemical Bonds

• Compounds are atoms held together by bonds.
• Chemical bonds are forces of attraction between atoms.
• Bonding arises from attractions between protons and electrons in atoms.
• Bonds can exist between atoms of the same or different elements.
• Bonds form because they decrease the potential energy between atoms' charged particles.

Types of Chemical Bonds

• Chemical bonds are classified into three types based on involved atoms: ionic, covalent, and metallic.
• Ionic bonds involve electron transfer between a metal and nonmetal, forming oppositely charged ions.
• Covalent bonds involve electron sharing between nonmetals.
• Metallic bonds involve electrons pooled between metal atoms.

Types of Chemical Bonds: The Ionic Bond

• Ionic bonds form when electrons are transferred between atoms forming oppositely charged ions which attract each other.
• Typically formed when a metal reacts with a nonmetal.
• Electron transfer forms a positively charged cation and a negatively charged anion.

Types of Chemical Bonds: The Covalent Bond

• Covalent bonds result when atoms share electrons.
• Typically form between nonmetals.
• Shared electrons attract both nuclei of the bonded atoms.
• Single, double, or triple covalent bonds exist resulting from sharing one, two, or three electron pairs, respectively.

Types of Chemical Bonds: The Coordinate Bond

• Coordinate bonds (or dative covalent bonds) are covalent bonds in which both electrons in the shared pair originate from the same atom (the donor).
• A donor atom, which has a lone electron pair, donates the electrons to an acceptor atom, lacking a full outer shell.
• Common examples of Lewis acids and bases include H⁺, AlCl₃, BF₃ and NH₃, H₂O, respectively.

Representing Valence Electrons With Dots (Lewis Structures)

• Lewis structures display valence electrons of main-group elements as dots around the symbol of the element.
• The maximum number of dots per side of the element's symbol is two.
• Each dot represents a valence electron.
• The number of valence electrons equals the group number (except for Helium).

Lewis Structures: For Covalent Bonding

• Hydrogen and oxygen have specific Lewis structures, and share valence electrons accordingly.
• Hydrogen forms a duet, oxygen forms an octet, in water molecules.

Lewis Theory Predicts That Hydrogen Should Exist as H2

• Hydrogen atoms have a single valence electron each.
• Two hydrogen atoms share their electrons to form a stable duet configuration.
• Elemental hydrogen exists as a diatomic molecule (H₂).

Lewis Structures: Double and Triple Covalent Bonds

• Oxygen exists as a diatomic molecule (O₂), with a double covalent bond.
• Nitrogen exists as a diatomic molecule (N₂) with a triple covalent bond.

Assessment (Questions)

• Questions are provided to assess student understanding of various chemical concepts.

Intermolecular Forces

• Intermolecular forces are attractive forces between molecules, weaker than intramolecular forces.
• IMF determines physical properties of substances (melting and boiling points, solubility).
• IMF affect the rate and outcome of chemical reactions.

Intermolecular Forces in Covalent Molecules

• London dispersion forces, dipole-dipole interactions, and hydrogen bonds are three types of IMF in covalent molecules, in order of increasing strength.
• London dispersion forces are present in all molecules (weakest); temporary fluctuations in electron distribution creating temporary dipoles,
• Dipole-dipole interactions occur in polar molecules, with attraction between permanent dipoles (moderate strength).
• Hydrogen bonds are strongest IMF, found in molecules with H bonded to N, O, or F; the hydrogen is attracted to another highly electronegative atom in a nearby molecule.

Intermolecular Forces: London Dispersion Forces

• London dispersion forces are weak and arise from temporary fluctuations in electron distribution within molecules.
• Strength is related to molecule size; larger molecules have stronger London dispersion forces.

Intermolecular Forces: Dipole-dipole interactions

• Dipole-dipole interactions involve permanent dipoles in polar molecules.
• The partial positive end of one polar molecule attracts the partial negative end of another.
• Strength is greater than London dispersion forces.

Intermolecular Forces: Hydrogen Bonding

• Hydrogen bonding is an exceptionally strong type of dipole-dipole interaction.
• Occurs in molecules where hydrogen is bonded to highly electronegative atoms (N, O, F).
• The hydrogen is attracted to another highly electronegative atom in a nearby molecule.

Intermolecular Forces: Types of Hydrogen Bonding

• Intermolecular hydrogen bonds are between different molecules.
• Intramolecular hydrogen bonds exist within a single molecule.

Intermolecular Forces: Ion-dipole Forces

• Ion-dipole forces are electrostatic attractions between a charged ion and a polar molecule.
• Strength is directly proportional to the charge of the ion and the magnitude of the dipole.
• Important in dissolving ionic compounds in polar solvents, like water.

Why are Intermolecular forces important in pharmacy?

• IMF play a crucial role in drug solubility, formulation, absorption, bioavailability, stability, delivery systems, and protein-drug interactions.

Intermolecular Forces: Problems

• Analysis for identifying the intermolecular forces present in particular compounds.

Stoichiometry

• Stoichiometry is the calculation of reactant and product amounts in chemical reactions.
• Allows prediction of products based on reactants.
• Allows determination of reactants needed to form a certain amount of product.

Reaction Stoichiometry

• Coefficients in balanced equations represent relative amounts of reactants and products in moles.
• Stoichiometric ratios (conversion factors) relate amounts of substances in a reaction.

Concentration of Solutions

• Solution is a homogeneous mixture of solute and solvent.
• Solvent is the substance present in the largest amount, solute is all other substances.
• Concentration is the amount of solute in a given solution.
• Various concentration units exist (e.g., % by weight, molarity, parts per million).

Concentration of Solutions: Molarity

• Molarity (M) is the number of moles of solute per liter of solution.
• M = moles of solute / liters of solution. Calculation of molarity from given mass and volume.

Types of aqueous solution and Solubility

• Saltwater is a homogeneous mixture of NaCl and water.
• Sugar water is a homogenous mixture of sucrose and water.
• Solids dissolve in water due to attractions between solute and solvent particles.

What Happens When a Solute Dissolves?

• Solute particles are held together by attractions.
• Similar attractions are present between solvent particles.
• Solute dissolves if attractions between solute and solvent are strong enough to overcome the attractions between solute particles.

Dissolving of Sodium Chloride in Water

• Ions of sodium chloride (NaCl) are attracted to water molecules, pulling them apart from the crystal structure.
• Compounds that dissociate into ions when dissolved in water are called electrolytes, allowing the resulting solution to conduct electricity.

Dissolving of Sugar in Water

• Sugar molecules dissolve in water via interactions between sugar and water molecules.
• Sugar does not dissociate into ions.
• Sugar solutions do not conduct electricity.

Electrolytes and Nonelectrolytes

• Electrolytes dissolve in water to produce ions, enabling the solution to conduct electricity.
• Strong electrolytes entirely dissociate into ions.
• Weak electrolytes partially dissociate into ions.
• Nonelectrolytes do not dissociate into ions when dissolved in water, and do not conduct electricity.

Electrolytes Solutions in Pharmacy

• Various electrolytes have specific uses in medicine, including those related to acid-base balance, nutrition, and treatment of deficiencies or imbalances.

Assessment (Questions)

• Questions related to stoichiometry, solutions, chemical reactions, and oxidation-reduction reactions, as well as redox reactions.

Basic Types of Chemical Reactions

• Reactions are categorized for a better understanding of chemical processes.
• Synthesis: forming larger products.
• Decomposition: breaking down larger reactants,
• Replacement: atoms taking place of others.

1- Synthesis Reactions

• Two or more reactants combine to form one product.
• Examples: Fe + S8 → 8FeS, 2H2 + O2 → 2H2O

2- Decomposition Reactions

• A complex substance breaks down into simpler substances.
• Example: 2H2O → 2H2 + O2

3- Replacement Reactions (Single and Double)

• Single Replacement: uncombined element replaces another element from a compound.
• Example: Mg + 2H₂O → Mg(OH)₂ + H₂
• Double Replacement: anions and cations of two compounds switch places to form two new compounds.
• Example: Pb(NO₃)₂ + 2KI → PbI₂ + 2KNO₃

Types of Chemical Reactions: A Summary

• Diagrammatic or graphic representations for synthesis, decomposition, single and double-replacement reactions.

Chemical Reactions between lons

• Formation of products from ion combinations.
• water, weak electrolytes, precipitates, gases or complex ions.

1- Formation of Water

• Formation of water from reactions between metallic hydroxides and acids.
• Example: NaOH + HCl → NaCl + H₂O

2- Formation of Weak Electrolyte

• Formation of weak electrolytes, when strong acid are mixed with solutions of salts.
• Example: HCl + CH3COONa → NaCl + CH3COOH

3- Formation of a Precipitate

• Formation and circumstances around precipitation , which is forming a solid during a chemical reaction.
• Formation of insoluble salts, when ions concentrations exceeds solubility limits.
• Example: AgNO₃ + NaCl → AgCl↓ + NaNO₃

3- Formation of a Precipitate (Another example)

• Formation, giving an example, of ferric hydroxide.
• Example: FeCl₃ + 3NaOH → Fe(OH)₃↓ + 3NaCl

4- Formation of a Gas

• Combination of ions possibly producing a gas.
• Formation of unstable acids, which then decompose producing gases.
• Examples: 2H⁺ + S²⁻ → H₂S↑, or H₂SO₃→ H₂O + SO₂↑

5- Formation of Complex lons

• Formation of a complex ion, from Lewis acid and Lewis base reactions
• Formation of coordination complexes, involving ligands donating electron pairs to a metal atom.
• Example: Cu²⁺ + 4NH₃ → [Cu(NH₃)₄]²⁺

5- Formation of Complex lons (Ligands)

• Ligands, as Lewis bases, may include anions or neutral molecules.
• Examples of anions include Cl⁻, Br⁻, I⁻, F⁻, CN⁻, SCN⁻, S₂O₃²⁻, C₂O₄²⁻.
• Examples of neutral molecules include CO, NH₃, H₂O.
• Various cations act as Lewis acids (metals) such as Ca²⁺, Mg²⁺, Fe²⁺, Fe³⁺, Pb²⁺, Cu²⁺, Zn²⁺, Al³⁺.

Reactions involving Oxidation-Reduction (Redox)

• Oxidation is the loss of electrons, and in contrast reduction is the gain of electrons.
• Redox reactions don't necessarily involve oxygen.
• Example: 2Na(s) + Cl₂(g) → 2NaCl(s)

Oxidizing Agent & Reducing Agent

• Oxidizing agents cause oxidation in a reaction and are themselves reduced.
• Reducing agents cause reduction in a reaction and are themselves oxidized.

Assessment (Questions)

• Various questions related to different chemical concepts.