CHM1046 Mastery Checkpoint 1 Study Guide

• ΔG Equations

  ΔG = ΔH - TΔS: This equation predicts that the solubility of every compound should increase with increasing temperature.

  
  

  ΔG < 0: This process is exergonic and will proceed spontaneously in the forward direction to form more products.

  
  

  ΔG > 0: This process is endergonic and not spontaneous in the forward direction. Instead, it will proceed spontaneously in the reverse direction to make more starting materials.

  
  

  ΔG standard is positive: The reaction is non-spontaneous at standard conditions.

  
  

  ΔG standard is zero: The system is at equilibrium at standard conditions.
• Explain why free energy change is the thermodynamic function that allows us to predict solubility. (Why can't we just use entropy or energy changes?)

  Because it involves the second law of thermodynamics, enthalpy, and entropy of the system in a same equation, considering only the system as the energy from surroundings cannot be quantified.
• Describe the solution process in terms of entropy changes for both solute, and solvent.

  (Note: Entropy INCREASES when bonds or forces are broken, and DECREASES when bonds or forces are formed).

  
  

  1. IMFs in the solute are broken (entropy increases)

  
  

  2. Some IMFs in the solvent are broken (entropy increases)

  
  

  3. IMFs are formed between the solute and the solvent (entropy decreases)
• Relate temperature changes to the interactions that are broken and formed during the solution process.

  Temperature increases (in the system) means bonds or interactions are being broken, temperature decreases (in the system) means bonds or interactions are being formed
• Understand and explain the terms: solute, solvent, solution, solvation, molarity, concentration, dilute, concentrated, micelle, colloid, emulsion.

  Solute: the thing being dissolved (the lesser amount, ex. salt)

  
  

  Solvent: the thing dissolving (the greater amount, ex. water)

  
  

  Solution: the homogeneous mixture of a solute and solvent (ex. salt water)

  
  

  Solvation: the process of making a solute

  
  

  Molarity (M) : moles / liter

  
  

  Concentration: can be shown by M, g/mL, ppm etc.

  
  

  Dilute: to add more solvent to a solution, reducing the concentration of the solute

  
  

  Concentrated: a lot of solute

  
  

  Micelle: ball shaped thing with polar heads on outside and non polar tails on the inside, for amphipathic molecules.
• Describe the structure of ionic compounds and use it to explain the properties of ionic compounds such as high mp and bp, hard, and brittle.

  ionic compounds form a crystalline lattice. A crystalline lattice is a repeating pattern of ions. The cations (+ ion) and the anions (- ions) attract each other. The attraction between the ions is called and ionic bond.

  
  

  High melting and boiling points: Ionic bonds have a very strong attraction between positive and negative ions - a lot of energy, heat, and pressure is needed to break them. So ionic compounds have high melting and boiling points.

  
  

  Ionic compounds are brittle due to the strong bond between the positive and negative ions that make up the molecules. These positive and negative bonds create crystals in rigid, lattice structures. Applying pressure shifts the alignment of the ions and results in brittleness. It takes a lot of energy to break them apart from each other.

  
  

  Ionic solids are brittle and hard because the electrostatic attractions in the solid again hold the ions in definite positions.
• Explain why metals tend to form positive ions and non-metals tend to form negative ions. Predict the charge on ion for common ions.

  Octec ( the tendency of atoms to prefer to have eight electrons in the valence shell) is not as full so they would rather give up than give.
• Explain the difference between covalent and ionic bonding using the idea that there is a continuum of bonding ranging from pure covalent to more or less ionic.

  Covalent bonds form between two nonmetals. Both have relatively high electronegativity, and there are valence electrons which are fought over between the two atoms in the bond. With ionic bonding, a metal and a nonmetal are involved. The metal gives 1 to 3 valence electrons to the nonmetal, forming two ions.
• Use intermolecular forces to explain properties of molecular compounds. Predict/rank relative melting and boiling points of given compounds.

  The stronger the IMF, the higher the relative melting or boiling point. IMF strength depends on they type of IMF (LDF is weakest, H bond is strongest), as well as the relative strength of that type of IMF. For example, the more polar the molecules are in a dipole-dipole interaction, the stronger it is. The more H bonds a molecule can form, the stronger it is.
• Use molecular structure and polarity to predict the types of intermolecular forces present in molecules, including London dispersion forces, dipole-dipole interactions and hydrogen bonding.

  LDFs: present in all atoms and molecules.

  
  

  Dipole-dipole interactions: present between polar molecules

  
  

  Hydrogen bonding interactions: present when one molecule has a hydrogen bonded to an O, N, or F and another molecule has an O, N, or F with a lone pair (O, N, and F are highly electronegative).
• Explain why rotation is (generally) easy around single (sigma) bond, but more difficult around pi bonds. Predict relative potential energies of different rotations around C-C bonds in hydrocarbons.

  P orbitals overlap when they are parallel to each other (double bonds) which are non rotatable because if you try to it will break the bond between the atoms. Where as sigma bonds are rotatable because it doesn't involve the breaking of a bond.
• Explain why we usually use ΔG instead of the total entropy change to predict whether a process is thermodynamically favorable.

  Because you can't measure the entropy of the universe very well. ΔG (Gibbs free energy) focuses on system and easy to measure that. more accurate
• Explain the second law of thermodynamics in terms of the system and surroundings

  The second law of thermodynamics states for any change, the total entropy of the Universe must INCREASE. You cannot get back what you put in.

  
  

  ΔS_total > 0

  Δ_Stotal = ΔS_system + ΔS_surroundings
• Predict the sign of the entropy change for simple systems.

  Solid to Liquid or Gas: More entropy (ΔS positive) because there is more randomness.

  
  

  Gas to Liquid or Solid: Less entropy (ΔS negative) because there is less randomness.
• Explain the role of probability in entropy changes.

  Higher probability = higher entropy.

  Lower probability = lower entropy.

  
  

  The higher probability the higher the number of possible arrangements, so higher temperature. More molecules will be randomly distributed and will move towards a random distribution from an orderly distribution.
• For phase changes, identify the direction of thermal energy change and the sign of q or ΔH.

  Melting and Vaporization: endothermic because it requires heat from the surroundings to overcome the intermolecular forces between its particles in the solid.

  ΔH is positive because the system gains heat.

  
  

  Freezing and Condensation: exothermic because the water loses heat to its surroundings as it freezes. Condensation is exothermic because it releases heat into the atmosphere.

  ΔH is negative because the system loses heat.
• For exothermic and endothermic processes, identify the direction of thermal energy change and the sign of q or ΔH.

  Exothermic: thermal energy leaves system, and enters surroundings.

  △H is negative because the system loses heat.

  
  

  Endothermic: thermal energy leaves surroundings, and enters system.

  △H is positive because the system gains heat from the surroundings.
• Explain the difference between state and path functions and give examples.

  State function - depends on initial and final states (ex. money you have in the bank).

  
  

  Path function - depends on path moving from initial to final state. Depends on how changes take place. (ex. the transactions - HOW the money got into the bank)
• Define and give examples of open, closed and isolated systems.

  Open system - both energy and matter can flow INTO and OUT OF the system. (ex. biological systems)

  
  

  Closed system - Matter (ex. O2) cannot flow into the system, but heat energy (heat) can flow INTO and OUT OF the system. (closed water bottle, pressure cooker).

  
  

  Isolated system - Matter (ex. O2) cannot enter or leave the system. Heat energy (heat) cannot escape or enter the system. (thermos/hydroflask; keeps things heated or cold - which is hard to do).
• Gibbs free energy (ΔG)

  the energy of a system that is available to do work at a constant temperature and pressure

  
  

  ΔG is dependent on temperature.

  
  

  ΔG Negative: spontaneous (favorable)

  ΔG Positive: Non-spontaneous (non-favorable)

  
  

  ΔG = ΔH - TΔS

  (ΔG = change in Enthalpy - temperature * entropy)
• Enthalpy (ΔH)

  The heat absorbed or emitted during a reaction under constant pressure.

  
  

  Equation: q = m x c x ΔT

  (Mass x specific heat x temperature change)
• Internal Energy (ΔU)

  Sum of all KE and PE of all particles in system. Can't measure its absolute value but can calculate change ΔE (Ef - Ei)

  ΔU = q + wterm-10

  (ΔU = heat + work)
• Entropy (ΔS)


  Entropy (ΔS) is a measure of the randomness or disorder of a system.

  
  

  ΔS+ = spontaneous reaction (no energy required)

  ΔS- = non-spontaneous (energy is required)