Stoichiometry and Chemical Reactions Quiz

General Chemistry for Pharmaceutical Sciences - Pharm-101

• Course presented by Dr. Azza H. Rageh, Associate Professor of Pharmaceutical Analytical Chemistry at Taibah University.

Stoichiometry, Solution Concentration, and Chemical Reactions

• Stoichiometry involves calculating the amounts of reactants and products in a chemical reaction.
• It allows predicting product amounts based on reactant amounts.
• It also determines the necessary reactants for a given amount of product.
• Coefficients in balanced chemical equations represent relative amounts in moles.

Reaction Stoichiometry

• The example reaction 2 C8H18(l) + 25 O2(g) → 16 CO2(g) + 18 H₂O(g) demonstrates the mole ratios between octane, oxygen, carbon dioxide, and water.
• 2 moles of octane react with 25 moles of oxygen to produce 16 moles of carbon dioxide and 18 moles of water.

Stoichiometric Ratio

• The balanced equation 2 C8H18(l) + 25 O2(g) → 16 CO2(g) + 18 H2O(g) relates the amounts of substances in moles.
• The ratio of the number of moles of octane to carbon dioxide is 2:16 or 1:8.

Stoichiometric Ratio Example

• If you want to find out how much CO2 will be formed when 22 moles of C8H18 are burned, use the ratio.
• 22 moles C8H18 x 16 moles CO2 / 2 moles C8H18 = 176 moles of CO2.

Concentration of Solutions

• A solution is a homogeneous mixture of two or more substances.
• The solvent is the substance present in the largest amount.
• The solute is all other materials present.
• Examples include sugar dissolved in water; water is the solvent, sugar is the solute.
• Concentration units include percentage weight, percentage volume, and parts per million (ppm).

Concentration Units

• Concentration units like percentage weight (%, w/w), percentage volume (%, v/v), percentage weight per volume (%, w/v), parts per million (ppm), molarity (M), formality (F), normality (N), and molality (m) are used to describe the amount of solute in a solution.

Molarity

• Molarity (M) is a method to express concentration, showing the relationship between moles of solute and liters of solution
• Molarity = amount of solute (in mol) / volume of solution (in L)
• Unit of molarity is moles of solute/liter of solution (mol/L).

Example Calculating Molarity

• Finding the molarity of a solution with 25.5 g KBr dissolved in 1.75 L of solution.
• Calculate moles of KBr (25.5 g / 119.0 g/mol = 0.21429 mol)
• Molarity (M) = moles of KBr / volume of solution (0.21429 / 1.75 = 0.122 M)

Example Calculating Volume

• Using molarity to find volume: 0.255 moles NaOH in a 0.125 M NaOH solution.
• Volume (L) = moles / molarity =0.255/0.125 = 2.04 L.

Preparing 1 Liter of 1 M NaCl Solution

• Weigh out 1.00 mol (58.44g) of NaCl
• Add water to the solid until dissolved
• Add more water until the solution reaches 1 liter.

Types of Aqueous Solutions and Solubility

• Salt water is a homogeneous mixture of NaCl and H2O.
• Sugar water is a homogeneous mixture of C12H22O11 and H2O.

What Happens When a Solute Dissolves

• Attractive forces exist between solute particles and solvent molecules.
• When mixed, attractive forces arise between solute and solvent particles
• Strong attractions lead to solute dissolving.

Dissolving of Sodium Chloride in Water

• Each ion in a salt is attracted to surrounding water molecules, detaching from the crystal lattice.
• Salts dissociate into ions in water forming electrolytes, enabling the solution to conduct electricity

Dissolving of Sugar in Water

• Sugar molecules mix homogeneously with water.
• Dissolving sugar into water does not result in the formation of ions.
• Solutions of sugar, therefore, are nonelectrolytes and do not conduct electricity.

Electrolytes and Nonelectrolytes

• Electrolytes are substances that dissolve in water to produce solutions that conduct electricity.
• Strong electrolytes completely ionize in water, while weak electrolytes only partially ionize.
• Nonelectrolytes do not ionize in water and do not conduct electricity.

Strong, Weak and Nonelectrolytes

• Strong electrolytes (soluble ionic salts, strong acids, strong bases) completely ionize.
• Weak electrolytes (weak acids, weak bases) partially ionize
• Nonelectrolytes (molecular substances, e.g., sugar and alcohol) do not ionize.

Chemical Kinetics

• Chemical kinetics is the study of the rate of chemical reactions.
• Includes how rapidly food spoils or a drug decomposes, and factors that determine the rusting of steel.

Speed or Rate of Reaction

• Rate of reaction measures the change in concentration of a reactant or product with time.
• For example, if the concentration decreases over time it is the rate of disappearance.

Factors Affecting Reaction Rate

• Temperature: higher temperature increases the rate as particles move faster and collide more.
• Catalyst: alters reaction rate without being consumed in the reaction.
• Surface Area: increased surface area increases the rate for solid/liquid or gas/solid reactions.
• Concentration: higher concentration of reactants increases reaction rate as more particles collide.

Concentration of Reactants

• According to the law of mass action, at constant temperature, the rate of a chemical reaction is directly proportional to the product of the molar concentrations of the reactants, each raised to an appropriate power
• This is captured in the rate law.

Order of Reaction

• The order describes how the rate of a reaction depends on the concentration of each reactant. It is determined experimentally.
• Zero order: rate is independent of concentration.
• First order: doubling the concentration doubles the rate.
• Second order: doubling the concentration increases the rate by a factor of four.

Chemical Equilibrium in Aqueous Solutions

• Dynamic equilibrium exists when the rate of the forward reaction equals the rate of the reverse reaction.
• Equilibrium constant (Keq): represents the relationship between the amounts of products and reactants when the reaction is at equilibrium.
• Keq is expressed using molar concentrations raised to powers equal to the coefficients in the balanced equation.

Equilibrium Constant: Kc or Keq

• Equilibrium constant, Kc or Keq, is expressed using molar concentrations of products and reactants at equilibrium.
• Products are in numerator, reactants in denominator.
• Exponents equal to coefficients in the balanced equation.

The Significance of Equilibrium Constants (Keq) Values

• When Keq is much larger than 1 (>>1), the reaction strongly favors the formation of products.
• When Keq is significantly less than 1 (<<1), the reaction strongly favors the formation of the reactants.
• When Keq is approximately equal to 1 (≈1), the concentrations of reactants and products are nearly equal, and the reaction is closer to a balanced state.

Calculating the Keq from Measured Equilibrium Concentrations

• The equilibrium constant, Keq, is calculated as the ratio of the products' concentrations to the reactants' concentrations, each raised to the power of its stoichiometric coefficients in the balanced chemical equation.

Le Châtelier's Principle

• If a chemical system at equilibrium is disturbed (e.g., by changing concentration, temperature, or volume/pressure), the system will shift to counteract the change and re-establish equilibrium.
• The shift will occur to a side minimizing the change.

The Effect of Concentration Change on Equilibrium

• Adding more reactants shifts equilibrium to the product side.
• Adding more products shifts equilibrium toward the reactants side.
• Adding a solid or liquid does not change equilibrium.

The Effect of Volume/Pressure Change on Equilibrium

• Increased pressure favors reactions making fewer moles of gaseous products.
• Decreased pressure favors reactions making more moles of gaseous products.
• Changes in pressure/volume have no effect on reactions without gaseous species.

The Effect of Temperature Change on Equilibrium

• Exothermic reaction (releases heat): increased temperature favors reactants, decreased temperature favors products.
• Endothermic reaction (absorbs heat): increased temperature favors products, decreased temperature favors reactants.

Acid–Base Neutralization

• Under the Arrhenius definition, acid+base --> water + salt.
• Neutralization reactions involve reacting acids and bases to form water and a salt.

Classifying Acids by The Number of H⁺

• Monoprotic acids: furnish one ionizable hydrogen proton (e.g., HCl, HNO3).
• Polyprotic acids: furnish more than one ionizable hydrogen proton (e.g., H2SO4, H3PO4, and sequential releases).

Strong and Weak Acids and Bases: Common Examples

• A table lists common strong acids and strong bases.
• A table also lists some common weak acids and weak bases.

Definitions of Acids and Bases: Brønsted-Lowry Definition

• Brønsted–Lowry acid is a proton donor.
• Brønsted–Lowry base is a proton acceptor.
• Acid-base reactions always involve a conjugate acid-base pair.

Definitions of Acids and Bases: Lewis Definition

• Lewis acid: an electron pair acceptor.
• Lewis base: an electron pair donor.
• The product obtained from a Lewis acid–base reaction is called an "adduct."

lonization Constant of Weak Acids and Bases

• Ka, for weak acid, describes the extent of ionization.
• Kb for weak base, represents the extent of ionization.

pH Scale

• pH is a logarithmic scale for expressing the concentration of hydronium ions (H3O+) in a solution.
• pOH describes the hydroxide ion concentration [OH⁻].
• pH + pOH = 14.

Base Solutions

• Strong bases completely dissociate in water, releasing hydroxide ions, e.g., NaOH.
• Weak bases react with water to accept a proton; example: ammonia (NH3).

Salts

• Salts formed by strong acids and strong bases are neutral (pH = 7) (example: NaCl).
• Salts formed by strong acids and weak bases are acidic (pH < 7) (e.g., NH4Cl).
• Salts formed by strong bases and weak acids are basic (pH > 7) (e.g., Na2CO3).

Buffer Solutions

• A buffer solution is a mixture of a weak acid and its conjugate base, or weak base and its conjugate acid.
• Buffers resist changes in pH from additions of acids or bases.

Buffer Solutions in Pharmacy

• Buffers are crucial in maintaining the stability, improving patient tolerance, and enhancing compatibility of medications.

Assessment questions (selected examples)

• Given specific solution concentrations, determine pH, pOH, [H₃O⁺], and [OH⁻].
• Identify Brønsted-Lowry acids, bases, their conjugate pairs.
• Write formulas for the conjugate base of different acids.
• Classify chemical species as Lewis acids or bases.