Honors Chemistry - Solutions and Kinetics Flashcards

Solutions and Concentration

• Matter: mixture and pure substance

  • Takes up space and has mass.

  • Can have variable composition (Mixture) or definite composition (Pure Substance).

• Mixtures:

  • Heterogeneous: Not uniform throughout (e.g., suspension).

  • Homogeneous: Uniform throughout; also known as a solution.

• Pure Substance:

  • Compound: Combination of two or more atoms (e.g., CO_2).

  • Element: Made up of atoms; cannot be futher separated by cannot be further separated by physical or chemical means.

• Density:

  • Density increases from top to bottom in a substance.

• Solution:

  • A homogeneous mixture made up of two parts: solute and solvent.

  • Can be any combination of two phases of matter.

  • Particles (atoms, ions, compounds) do not scatter light.

• Suspension

  • A heterogeneous and mixture where a solid is not completely dissolved in a liquid.

• Colloid: homo and hetero

  • Particles larger than molecules (solid in liquid, not completely dissolved).

  • Scatters light (e.g., fog).

• Solution:

  • A homogeneous mixture

  • Solution = Solute + Solvent

  • Solute: Gets dissolved; lesser quantity (e.g., sugar).

  • Solvent: Does the dissolving; greater quantity (e.g., water is a universal solvent).

• Solutes Classification:

  • Electrolytes: Yield dissolved ions (e.g., NaCl); conduct electricity.

  • Non-electrolytes: Yield dissolved molecules (e.g., sugar); do not conduct electricity.

• Solvation:

  • Process of surrounding solute particles with solvent particles in aqueous solutions.

  • "Like dissolves like": Polar solvents dissolve polar solutes and ionic solutes; nonpolar solvents dissolve nonpolar solutes.

  • Polar solutes do not dissolve in nonpolar solvents, and vice versa.

  • Solvation in water is called hydration.

Solution Concentration

• Definition: A measure of how much solute is dissolved in a given amount of solvent.

• Qualitative Descriptions:

  • Dilute: Small amount of solute.

  • Concentrated: Large amount of solute.

• Quantitative Measures:

  • Molarity (M)

  • Molality (m)

  • % by mass

  • ppm

• Types of Solutions:

  • Unsaturated: Contains less solute than the solvent can hold; more solute can be added.

  • Saturated: Contains the maximum amount of solute a solvent can hold; precipitate forms at the bottom.

  • Supersaturated: Contains more solute dissolved than a saturated solution would contain under the same conditions.

Solubility

• Definition: The maximum amount of solute that can be dissolved in a solvent at a specific temperature (grams solute/100 g H_2O ).

• Units: Typically grams solute / 100 g solvent.

• Amount of substance required to form a saturated solution.

• Factors Affecting Solubility:

  • Temperature:

    • Solids: Temperature increase = solubility increase.

    • Gases: Temperature increase = solubility decrease.

  • Pressure: Affects only gases; pressure increase = solubility increase.

• Solubility Curves:

  • On the line = saturated solution.

  • Below the line = unsaturated solution.

  • Above the line = supersaturated solution.

Molarity (M)

• Definition: Number of moles of solute per liter of solution (molar concentration).

• Formula: M=\frac{\text{ mol solute}}{\text{L solution}}

Dilution

• Concentrated solutions of standard molarities are called stock solutions.

• Solutions of lower concentrations are prepared by dilution.

• As volume increases (more solvent added), molarity decreases.

• Only the volume of solvent changes. moles of solute before= moles of solute after

• Formulas:

  • M1V1=V1V2

Molality (m)

• Definition: Used when calculating colligative properties.

• Formula: m=\frac{\text{ mol solute}}{\text{kg solvent}}

• Be careful in identifying solute versus solvent (solvent usually has greater quantity).

Colligative Properties

• Definition: Physical properties of solutions that depend ONLY on the number of solute particles present, NOT the kind of solute.

• Ionic compounds have a larger effect than molecular compounds because when dissolved, they dissociate into multiple ions (electrolytes), while molecular compounds do not (non-electrolytes).

• Examples:

  • Covalent; i=1

  • NaCl(s) \rightarrow Na^+(aq) + Cl^-(aq), i = 2 (electrolyte)

• Four Colligative Properties:

  1. Vapor Pressure Reduction:

     • Vapor pressure of a liquid is the equilibrium pressure of gas molecules from that liquid.

     • A solvent in a solution has lower vapor pressure than a pure solvent because it must overcome the attractive forces of both solvent and solute molecules.

     • Adding a solute lowers the solvent's vapor pressure.

  2. Osmotic Pressure:

     • Osmosis: Flow of solvent through a semipermeable membrane from areas of low solute concentration to areas of high solute concentration.

     • Osmotic Pressure: Difference in pressure between pure solvent and solution once equilibrium is re-established.

     • Increases when the concentration of a solution increases

  3. Boiling Point Elevation:

     • The boiling point of a solution is higher than that of the pure solvent.tb=Kbm\left(solution\right)i

     • For water, K_b = 0.512 \frac{°C}{m}.

  4. Freezing Point Depression:

     • The freezing point of a solution is lower than that of the pure solvent.

     • tf=Kfm\left(solution\right)i

     • For water, K_f = 1.86 \frac{°C}{m}.

Steps to Calculate Freezing Point Depression and Boiling Point Elevation

1. If the solution is ionic, write the dissociation equation to determine the i value.

2. Convert mass of solute to moles, and mass of solvent to kg, to calculate the molality (m).

3. Calculate boiling point +100

4. Calculate freezing point 0-x

Rates of Chemical Reactions

• A rate=speed

• Average rate of a rxn= concentration over time

• Chemical kinetics involves manipulating factors to speed up or slow down a chemical reaction.

• Average rate of a reaction: \text{Rate} = \frac{\Delta \text{concentration}}{\Delta \text{time}}.

• For a reaction to occur:

  • Reactant particles need to collide.

  • Some chemical bonds need to break, and new bonds need to form.

  • Not all collisions result in a new product.

Collision Theory

• The rate of a chemical reaction depends on the number of effective collisions between molecules.

• Three parts to the collision theory:

  1. Atoms or particles must collide.

  2. They must collide at the correct orientation.

  3. They must collide with a sufficient amount of energy.

• Activation Energy (E_a):

  • The minimum amount of energy the colliding particles must have for productive collisions.

  • How fast or slow a reaction occurs depends on the magnitude of E_a.

  • Low E_a: Reaction is faster; higher number of productive collisions.

  • High E_a: Reaction is slower; lower number of productive collisions.

  • The more energy needed to bond, the fewer molecules will react.

Exothermic and Endothermic Reactions LOOK AT NOTES

• Exothermic Reaction:

  • Releases energy.

  • Products have lower energy than reactants.

• Endothermic Reaction:

  • Requires energy.

  • Products have higher energy than reactants.

• Reaction rates are determined from experimental data, such as:

  • Change in volume

  • Change in mass

  • Change in concentration

  • Change in color

  • Change in pressure of gas produced

Factors That Affect Reaction Rates

• Nature of Reactants:

  • (Activity series of metals, halogens)

  • Certain chemicals are more reactive than others.

  • Ions, small molecules, and gas molecules react faster than liquids and large molecules.

• Concentration of Reactants:

  • Lower concentration = fewer collisions = slower reaction rate.

  • Higher concentration = more collisions = faster reaction rate.

• Surface Area:

  • Smaller surface area = fewer potential collision sites = slower reaction rate.

  • Larger surface area = more potential collision sites = faster reaction rate.

• Temperature:

  • Higher temperature = faster reaction rate.

  • Higher temperature = more kinetic energy = more collisions (with more energy).

  • The increase of 10°C will cause the reaction rate to double.

• Catalyst:

  • Increases the reaction rate without being consumed.

  • Lowers activation energy.

  • Biological catalysts: enzymes.

• Inhibitors:

  • Slows down a reaction (e.g., food preservative).