Atomic and bonding theory: checklist of what you need to know for the unit Collection of different possible question types with full solutions detailed review sheet comparison chart for non-polar and polar covalent bonds ionic bonds comparison chart for bohr rutherford and quantum mechanic model word bank flash cards for types of intermolecular forces orbitals diagram properties of solids chart- matching cards orbitals chart matching game images Atomic and bonding theory: Checklist of what you need to know for the unit: Models Ernest Rutherford’s model Niel Bohr’s model Bohr- Rutherford model Quantum mechanic model comparison table with bohr rutherford and quantum mechanics Quantum numbers Orbitals Energy level diagrams Electron configuration “Shorthand form” Valence electrons Anomalous electron configuration Different types of chemical bonds Non-polar covalent bond Polar covalent bond Ionic bond Metallic bond Intermolecular forces Ion dipole Dipole- dipole London dispersion forces Hydrogen bonding VSEPR theory Properties of solids Lewis structure Electronegativity Predicting molar polarity steps to determine polarity Hybridization Lone pair Coordinate covalent bonds Bond naming Theories and principles Dalton’s Theory Rutherford’s Theory (and the problems) Bohr’s Theory (and the problems) Heisenburg’s uncertainty Principle Hund’s Rule Pauli Exclusion Principle Aufbau Principle Atomic and bonding theory 1. Models Ernest Rutherford’s model - In this beehive model, a positively charged nucleus is in the centre of the atom which contains majority of the atom’s mass - It’s surrounded by electrons - Suggests existence of neutrons Bohr- Rutherford model - - This model shows the amount of electrons present in each shell of the atom Presented by Niels Bohr and Ernest Rutherford in 1913 a dense, small nucleus that is circulated by electrons Niel Bohr’s model - Referred to as the “planetary model” - The nucleus remains positive and dense, similar structure to the sun - The electrons have exact energies, they follow similar paths like a planet, by transporting in circular orbits Quantum mechanic model - - Developed in 1929 by Shroedinger and Heisenberg Derived by wave equations that explain energy and motion of an electron surrounding the nucleus Electrons are found in orbitals Electrons can only have specific energies which are related to the orbitals they are placed in Comparison chart between Bohr-Rutherford model and Quantum mechanic model Bohr- Rutherford Model Quantum mechanic model - Electron is a particle that obtains a very small mass - Electron portrays characteristics of both waves and particles - Electron maintains an orbit shape around the nucleus - Motion and path of the electron cannot be identified Electrons are found in the orbitals - Electrons can only have specific energies which are dependent on the orbit - Electrons circulate around the nucleus, similar to how the planets orbit around the sun in the solar system - Electrons can only have specific energies which are dependent on the orbitals Heisenberg uncertainty principle: - Cannot identify where the electron is or how it is transporting - The electrons spin in one two directions Pauli exclusion principle: - maximum of 2 electrons per orbital 2. Theories and principles Dalton’s Theory - matter is made up of atoms which are indestructible and invisible atoms from the same element are identical atoms from different elements are different Problems with Rutherford’s Theory - - due to classical physics, it states that when a charged particles changes direction in space, it will lose energy - therefore electrons should spiral into the nucleus which implodes the atom excited atoms release a line spectrum, not a continuous spectrum when their electrons are excited 3. Quantum numbers Electron address: Each electron in an atom has a specified set of four quantum numbers, this defines where they are most likely to be found. n= The principle quantum number, this indicates the energy level and how far the electron may be from the nucleus. The lesser the quantity of the number (1), the closer it is to the nucleus, the greater the quantity of the number (3), the farther it is from the nucleus. example: n= 1,2,3 l= The secondary quantum number, this indicates the shape. l= 0 “s” l= 1 “p” l= 2 “d” l= 3 “f” ml = magnetic quantum number which indicates the orientation ml = (-l…+l) ms = spin quantum number example: + 1 2 = clockwise ↑↾ - 1 2 = counter- clockwise ↓⇂ 4. Orbitals - regions of space where the electrons are most likely to be identified. Orbitals Orbital name s l amount This is the first energy level. n=1 l = 0…n-1 l = 0…1-1 l = 0 (s) Shape ml = (-l…+l) ml = -0…+0 ml = 0 Only one orientation p This is the second energy level. n=2 l = 0…n-1 l = 0…2-1 l = 0 (s), 1(p) ml = (-l…+l) ml = -1…+1 ml = -1,0 ,+1 Therefore there are 3 orientations. d This is the third energy ml = (-l…+l) level. n=3 l = 0…n-1 l = 0…3-1 l = 0 (s) ,1 (p) ,2 (d) f This is the fourth energy level. n=4 l = 0…n-1 l = 0…4-1 l = 0 (s) ,1 (p) ,2 (d), 3 (f) ml = -2…+2 ml = -2, -1,0 ,+1, +2 There are 5 orientations. ml = (-l…+l) ml = -3…+3 ml = -3, -2, -1,0 ,+1, +2, +3 There are 7 orientations. They are very complex shapes. 5. Energy level diagrams l n ml subshell notation #of orbitals in subshell #of electrons needed to fill shell 1 0 0 1s 1 2 2 0 0 2s 1 2 2 1 -1, 0, +1 2p 3 6 3 0 0 3s 1 2 3 1 -1, 0, +1 3p 3 6 3 2 -2.-1, 0, +1, -1 3d 5 10 4 0 0 4s 1 2 4 1 -1, 0, +1 4p 3 6 4 2 -2.-1, 0, +1, -1 4d 5 10 4 3 -3,-2.-1, 0, +1, -1, +3 7 14 total # of electrons in subshell 2 8 4f 18 32 6. Electron configuration Energy level diagrams and electron configurations are used to represent the electrons by the level of energy that they occupy. It is a linear distribution of electrons from each orbital of an atom. The number and location of the electrons are placed in order of increasing energy. Valence shell electron configuration Representative elements valence electrons are represented in the outer s and p orbitals. - when forming cations they lose their s electrons and then their p orbitals when forming anions they gain p electrons Transition elements valence electrons obtain the greatest energy level s orbital, including the d orbital - when forming cations they lose the highest energy level which is the s orbital then the d orbital Naming- General form Fe has 26 electrons, the electron configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶ When all the superscripts are added together, they would be equivalent to the amount of electrons, which in this case is 26. Shorthand form - use the symbol closest to the previous noble gas element square brackets should be used around this element continue the electron distribution For Fe it would be: [Ar] 4s² 3d⁶ Anomalous electron configurations - it was observed after experiments that these electrons appeared half-filled and filled orbitals are more stable compared to unfilled orbitals For Cu it would be: [ar] 4s¹ 3d¹⁰ 7. Chemical bonds Non-polar covalent bonds - an equal sharing of electrons a non-metal and another non-metal atom are involved in the bonding the atoms would have similarly high and equal electronegativity 𐤃EN= 0…0.5 Polar covalent bonds - unequal sharing of electrons a non-metal and another non-metal atom are involved in the bonding a transitional metal and non- metal are involved in the bonding relatively high but unequal electronegativity 𐤃EN= is 0.5 or greater but less than 1.7 Ionic bonds - electron transfer results in cation and anion being attracted to each other a metal and a non- metal are involved in the bonding one of the atom has a low EN, however the other has a high EN 𐤃EN= 1.7 or greater Metallic bonds - attraction between delocalized electrons and positively charged nuclei a metal and another metal atom lower electronegativity that have a small difference 8. Intermolecular forces Quizlet- flash cards made by me https://quizlet.com/_axrawl?x=1qqt&i=34jmnm 9. VSEPR theory shapes summary This is a visual non-mathematical theory formed by R. Gillespie V= Valence S= Shell E= Electron P= Pair R= Repulsion 1. 2. 3. 4. Bonded and unbonded electron pairs are treated as a negatively charged cloud. Each negatively charged electron cloud repels all the other charged clouds nearby. For minimum potential energy, the electron clouds locate themselves as far as possible. The spatial orientation of the charged clouds depends on various factors: - the amount of clouds that are around the central atom - the location of the cloud - localized (electrons are not free to move), therefore there will be a bonded electron pair between 2 nuclei - delocalized (electrons are free to move), therefore a lone pair of electrons 10. Properties of solids Did a matching puzzle game to memorize. https://puzzel.org/en/matching-pairs/play?p=-MtutTvJMAJTaDOp0u9P 11. Lewis structure ionic bond Covalent bond Multiple covalent bonds Coordinate covalent bond Resonance structures Non-conformist molecules Central atom electrons less than octet Central atoms electrons greater than octet Central atom with an odd number of electrons 12. Predicting molecular polarity - notes done by myself 13. Steps to determine polarity - notes done by myself 14. Lewis structure - notes done by myself 15. Hybridization A video about hybridization can be accessed by this link 16. Intermolecular forces, liquids and solids Kinetic molecular theory that explains the states of matter State of matter Gas Description - - Liquid - Solid - Explanation highly compressible assumes the volume and shape of the container follows the Gas laws and not the identity of their particles - almost compressible assumes the shape of the container does not assume the volume of the container - molecules held together from gas molecules very rigidly - results in molecules still being able to pass each other incompressible definite volume and shape - molecules are packed together the molecules cannot pass by each other easily - - - molecules are drifted apart the molecules have barely any interaction or attraction with each other no bonds Kinetics molecular theory explains the change of state - molecules must get closer to each other to change a gas into a liquid or solid cooling or compressing a substance must be done in order for intermolecular bonds to be made molecules must get farther away from each other for a solid to turn into a liquid or gas heating or reducing pressure must be done so that the intermolecular bonds can be broken Types of forces - Intermolecular forces are the attraction between molecules, they are weaker than intramolecular forces - responsible for physical properties such as changes of state and solubility - Intramolecular forces are the forces between atoms in a molecule gases follow a gas law but liquids and solids follow different laws because gases don’t have intermolecular forces Relationship between boiling points and intermolecular forces - - substance boils when the vapour pressure in the liquid is equivalent to the atmospheric pressure the greater the intermolecular force in a substance the smaller the number of molecules that are able to escape - more energy would be needed to reach a substance’s boiling point overall, relative boiling points are a measure of the strength of intermolecular forces 17. Bond naming Sigma bond - first bond created between two atoms union formed between 2 atoms because of the “end to end” overlap of orbitals electrons are localized between the two nuclei of the atoms involved on the bond axis eg. “s” and “s” or “two hybrids” or “p” and “p” Pi bond - second and third bonds created between two atoms union between 2 atoms due to the side to side overlap of P orbitals electrons are not localized between the nuclei of the atoms - they are concentrated in two separate regions - electrons are available for a reaction because they are not held as strongly Thermochemistry and Kinetics energy chart for types of energy changes practice questions wit solutions notes diagrams examples checklist of what needs to be known Kinetic energy review sheet kahoot checklist of what's in the unit diagrams Thermochemistry First law of thermodynamics Types of energy Temperature Heat Thermal energy Types of energy changes Endothermic exothermic types of energy changes to a system types of systems Changes of state measuring and calculating heat calculating heat the mass of the system the change in temperature specific heat capacity enthalpy calorimetry molar enthalpy communication enthalpy 4 ways of how to the communicate enthalpy Thermochemical equation Hess’s law Standard Enthalpies of formation summation of heat bond energy Thermochemistry 1. First law of thermodynamics The first law of thermodynamics states that energy cannot be created or destroyed. To better explain, the total energy in the universe is constant but the energy can be transferred into different types. 2. Types of energy Types of energy Kinetic Energy associated with a moving object. Thermal Energy associated with the average kinetic energy of the particles of a system. Sound Energy associated with the movement produced by an object's vibrations. Electromagnetic Energy associated with electromagnetic radiation. Chemical potential Energy stored in chemical bonds between atoms. Gravitational potential Energy stored in an object relative to its height above the earth’s surface. Nuclear potential Energy stored in the bonds between the subatomic particles in the nucleus. Elastic potential Energy stores in substances whose shape has been distorted or stretched. 3. - Thermal energy internal energy of a system related to the kinetic energy of all particles that the system is composed of the amount of thermal energy is determined by the avg kinetic energy of the particles and the amount of particles present - the greater the amount of particles, the greater the amount of thermal energy 4. Heat - measure of the transfer of thermal energy between objects - the amount of thermal energy transferred between two objects with different temperatures is referred to as “heat energy” - hotter substance to cooler substance 5. Temperature - measure of the avg kinetic energy of the particles in a given substance - it’s related to thermal energy 6. Types of energy changes When the amount of energy in a system changes, it can either gain energy or lose energy. Exothermic The system loses energy to its surroundings. Endothermic The system gains energy to its surroundings. 7. Types of energy changes to a system Physical change Types of bonds - - Examples - - Chemical change Nuclear change typically intermolecular forces possibility of intramolecular if a covalent network, ionic or metallic solid Intramolecular forces Nuclear forces (strong nuclear force or weak nuclear force) Phase changes (melting, freezing, evaporating and condensation) solubility Chemical reactions (combustion, redox, acid-base) - radioactive decay nuclear fusion nuclear fission Typical amounts of energy ∆E= 10° - 10^2 kJ/mol ∆E= 10^2 - 10^4 kJ/mol ∆E= 10^10 - 10^12 kJ/mol 8. Types of systems Open - open system is avoided since both matter and energy is lost Closed - due to reality, closed system is used since it doesn’t allow matter to escape, only energy Isolated - neither energy or matter is lost to surroundings impossible because to create substance completely isolated from collisions with surrounding particles 9. Changes of state 10. Measuring and calculating heat Exothermic: -q system, +q surroundings Endothermic: +q system, -q surroundings - measures the amount of energy that is gained or lost by a system which goes through a physical, chemical or nuclear change calorimetry is used to measure the changes in energy 11. Calculating Heat Mass of the system (m) - mass is required because it’s a measurement of the # of particles in the system, due to thermal energy being proportional to the # of particles in the system measured in grams typically Change in temperature - change in temperature is calculated by final temperature-initial temperature or ∆t The specific heat capacity (c ) - amount of energy that iq required to raise the temperature of 1 gram of a given substance is referred to as specific heat capacity For water: 1mL = 1g Assumption that if dealing with a dilute solution, the heat capacity is 4.18 J/g°c q= m ᐧ c ᐧ ∆t 12. Enthalpy (H) - certain amount of energy due to the combination of kinetic energy and potential energy that’s present Enthalpy of a system decreases -∆H - heat content decreases energy leaves the system heat flows into the surroundings increase in temperature exothermic changes Enthalpy of a system increases +∆H - heat content increases energy enters the system heat flows from the surroundings into the system endothermic changes 13. Calorimetry - measuring the change in enthalpy of a system - measuring the temperature change of the system when a physical, chemical or nuclear change occurs - bomb calorimeter creates a near isolated system that allows very little energy to escape the system into the surroundings - enthalpy is ∆H - the heat energy released (or absorbed) is equivalent to the enthalpy change ∆H= +- l q surroundings l - positive or negative is decided on weather the enthalpy change is endothermic or exothermic enthalpy expressed per mole (joules/mole) enthalpy can also be expressed per mass (joules/gram) 14. Molar enthalpy - process (physical, chemical, nuclear change) on 1 mole of a substance ∆H= ∆𝐻 𝑛 The n is the # of moles of the substance the process occurred, this can be identified by n= 𝑚 𝑀𝑀 The enthalpy change of a system can be measured from the heat gained/lost by the surroundings. ∆Hmolar= −𝑞 𝑛 ∆H= -q surroundings The heat energy gained/lost by the surroundings can be identified by: q= m ᐧ c ᐧ ∆T Practice questions with examples and solutions are at the end of this workbook. 15. Communicating Enthalpy the change in enthalpy of a system= the heat energy lost or gained by the measured surroundings. ∆Hsystem= +- lql surroundings exothermic reaction: q is in the products! aA+bB-> cC+dD +q - system loses energy to the surroundings the enthalpy (H) of the system decreases, it is negative the heat of the reaction, q, is positive endothermic reaction: q is in the reactants! - system gains energy to surroundings the enthalpy (H) of the system increases, it is positive heat of reaction, q, is negative aA+bB +q -> cC+ dD 16. Four ways of communicating enthalpy - state the enthalpy change - Example: ∆H =-231 kJ/mol - state the enthalpy change following the reaction - Examples: H2 (g)+ ½ O2 (g) -> H2O (g) ∆H=-231 kJ/mol - state the enthalpy change as a part of the reaction - example: H2 (g) + ½ O2 (g) -> H2O (g) + 231 kJ - Draw energy diagrams 17. Thermochemical equations ● the physical states of the materials (s) (l) (aq) ● ∆H° …the “°” designates that energies were measured at standard SATP conditions ● coefficients representing “moles” not molecules, meaning they may be fractions 1. N2 (g) + 3H2 (g) -> 2NH3 (g) 2. 𝑁2 (𝑔) + 3𝐻2 (𝑔) −> 2𝑁𝐻3 (𝑔) 2 ∆H°= -92.4 kJ/ mol of N2 −92.4 2 ∆H°= -46.2 kJ/ mol of NH3 formed 3. ½ N2 (g)+ 3/2 H2 (g) -> NH3 (g) 4. The reverse would be flipping the equation sides and the ∆H°’s sign: NH3 (g) -> ½ N2 (g) + 3/2 H2(g) This is now endothermic. ∆H°= 46.2 kJ/mol 17. Hess’s Law Energy change for any chemical or physical process is independent of the pathway or number of steps required to complete the process, provided that the final and initial reaction conditions are the same. When a reaction can be expressed as the algebraic sum of two or more other reactions, the enthalpy is the algebraic sum of the heats of these reactions. 18. Thermochemistry- Standard enthalpies of formation - the enthalpy change for 1 mole of a substance from elements in their standard state at SATP is called the standard enthalpy of formation Method 1: Additivity of Heats Method 2: Summation of Heats 18. Bond energy - amount of energy required to break a bond endothermic Example of a bond energy problem I did: Practice questions 1. There is 1.50 L of water in a kettle. Calculate the quantity of heat that flows into the water when it is headed from 18.0°C to 98.7°C. 2. The specific heat capacity of porcelain is 1.1J/G°C a) If the temperature of a toilet bowl is 15.0 °C and has a mass of 20.0 kg, how much heat energy would be required to raise the temperature to 20.0 °C 3. What mass of ethylene glycol would evaporate if 474 kJ of heat energy were added to it? 4. Communicate the enthalpy change using the 4 methods for each of the following chemical reactions, assume SATP for all substances. a) the decomposition of aluminum oxide (∆H°= +1676kJ/mol of CO2) b) When 46.2 kJ of heat energy is applied to 1 mole of ammonia, it will decompose to hydrogen and nitrogen. 5. By applying Hess’s Law, calculate the enthalpy of this reaction: 2C(s) + H2(g) ---> C2H2(g) ΔH° = ??? kJ Using the following information: C2H2(g) + 5⁄2O2(g) ---> 2CO2(g) + H2O(ℓ) ΔH° = −1299.5 kJ C(s) + O2(g) ---> CO2(g) ΔH° = −393.5 kJ H2(g) + 1⁄2O2(g) ---> H2O(ℓ) ΔH° = −285.8 kJ Solutions Kinetics energy measuring the rate of reaction average rate of reaction molar ratios chemical kinetics: reaction rates analyzing the progress of a reaction dimensional analysis rate laws and order of reaction rate law and reaction mechanisms collision theory factors that increase reaction rate temperature and reaction rate kinetic energy distribution curves- maxwell boltzmann curves potential energy diagrams endothermic exothermic effect of catalyst effect of a temperature change PE diagram for a reaction mechanism 1. Measuring the rate of reaction - speed or rate of a chemical reaction can be measured by: - disappearance of reactants - or appearance of products - will yield a quantitative measure of the rate of a reaction Homogenous system: all reactants are present in the same state Heterogeneous system: reactants are present in different phases - reactions where gas is produced can be measured using gas collection techniques reactions with ions can be measured with conductivity meters or pH meters for acids or bases reactions with colour changes can be measured by spectrophotometer −[𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠] Δ𝑡 rate= +[𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠] Δ𝑡 or Δ𝑥 Δ𝑡 Chemicals can be measured with a variety of quantitative measures (Δx) : - mass (g) volume (L) moles concentration (mol/L) average rate of reaction= or r= Δ𝑐 Δ𝑡 = 𝑚𝑜𝑙 𝐿𝑥𝑠 𝑐ℎ𝑎𝑛𝑔𝑒 𝑖𝑛 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 𝑐ℎ𝑎𝑛𝑔𝑒 𝑖𝑛 𝑡𝑖𝑚𝑒 2. Molar ratios (note done by me) 3. Chemical Kinetics: Reaction Rates - quickly or slowly reactants are consumed or products are formed in a reaction Average reaction rate= 𝑎𝑚𝑜𝑢𝑛𝑡 𝑜𝑓 𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡 𝑐𝑜𝑛𝑠𝑢𝑚𝑒𝑑 𝑡𝑖𝑚𝑒 −Δ[𝑟] Δ𝑡 Average reaction rate= 𝑎𝑚𝑜𝑢𝑛𝑡 𝑜𝑓 𝑝𝑟𝑜𝑑𝑢𝑐𝑡 𝑝𝑟𝑜𝑑𝑢𝑐𝑒𝑑 𝑡𝑖𝑚𝑒𝑠 +Δ [𝑝] Δ𝑡 =l l = l l 4. Analyzing the progress of reaction (note done by me) m= 𝑦2−𝑦1 𝑥2−𝑥1 5. Dimensional analysis (note done by me) 6. Rate law and order of reactions note done by me (not my note) Video that I found helpful when learning about rate law 7. Rate law and reaction mechanisms - rate law equation provides a quantitative description about how the concentration of reactants (in the aqueous or gas state) affects the initial rate of the reaction One- step (elementary step) reactions - max of 2 particles - max of 3 particles cannot be explained in terms of simpler steps since they involve the direct collision of reactants order of each reactant is determined by the coefficient of each reactant Multi-step reactions - exponents is not equal to coefficients elementary steps by which the overall reaction - slowest elementary step (rate determining step) is the only one that affects the overall reaction rate these are the only steps that appear in the rate law equation for the overall reaction - Rules of a reaction mechanism a) steps have to add up to the overall reaction b) Max of 3 particles in each step c) RDS must be reflected in rate law i) coefficients = orders 8. Collision Theory - reaction might not take place even if molecules with sufficient energy take place orientation is very important for reactions to occur 9. Temperature and reaction rate - - if temperature of a reaction system has increased: - molecules move more quickly - molecule will possesses more energy this leads to: - more collisions - the collisions will be more effective 10. Kinetic energy distribution curves- Maxwell boltzmann curves - not all molecules in a reaction system will have the same kinetic energy at a specific temperature - reactions have an energy threshold or activation energy which is required to complete the reaction at high temperatures, molecules will have more energy for the reaction to proceed rate is greater at higher temperatures a 10 degrees celsius rise in temperature will double the rate of reaction - 11. Potential energy diagrams - for reverse reactions, the ΔH stays the same, however the activation energy is not the same as it is for forward reactions Endothermic reactions - kinetic energy converted into potential energy - decrease in temperature - potential energy increases as molecules approach - the activated complex can be formed if the molecules have enough kinetic energy - activated complex breaks down to form either the stable reactants or stable products Exothermic reactions - converts potential energy into kinetic energy - increase in temperature 12. The effect of a catalyst Distribution curve - catalyst lowers the activation energy no effect on the ΔH Potential energy diagram - catalyst provides an alternative pathway with an activated complex of lower potential energy lowers the activation energy for both the forward and reverse reaction 13. Effect of a temperature change Distribution curve - activation energy is the same, despite change of temp at higher temperatures, more molecules have the required kinetics energy for the reaction - faster reaction - Potential energy diagram - NO EFFECT I made a kahoot, however it’s not letting me share it because I made it on my school account, so I took screenshots of the kahoot and uploaded on google drive, here’s the link: https://drive.google.com/drive/folders/1yL0fLlGqeOBGsjlOC8YKFY3SITgvD1NC?usp=sharing Equilibrium review sheet diagrams + graphs kahoot quiz examples images youtube videos Equilibrium thermodynamics vs kinetics enthalpy and entropy laws of thermodynamics drives in natures predicting entropy spontaneous process and entropy free energy concept factors affecting equilibrium tendency towards minimum potential energy tendency towards maximum randomness Gibbs free energy Introduction to equilibrium dynamic closed system steady state system types of equilibrium chemical equilibria altering equilibrium le chatelier principle concentration temp pressure and volume catalyst equilibrium constant ICE 1. Thermodynamics vs. kinetics - study of energy changes that accompany physical and chemical changes - thermodynamics helps understand if the process will occur - does not give information about the amount of time required for the process 2. Enthalpy and entropy - enthalpy is the relationship of ΔH to q released or absorbed - entropy, S, is a measurement of the randomness or disorder or available arrangements of the system or surroundings 3. - Laws of thermodynamics energy of the universe is constant in any spontaneous process there is always an increase in entropy a substance that is perfectly crystalline at 0 K has an entropy of 0 4. Two drives of nature A chemical reaction will favour the side with: Minimum enthalpy if there are no other factors considered Maximum enthalpy if there are no other factors considered 5. - Predicting entropy ΔS reaction breaks up a larger molecule smaller molecule fragments reaction occurs where there is an increase in the moles of gas in the product process where solid changes to a liquid or gas to a liquid changes to a gas 6. Spontaneous Processes and Entropy - physical or chemical change occurs by itself after it starts, it continues to completion non- spontaneous: - processes occur without requiring an outside force constant supply of energy is required to continue 7. Free energy concept - energy associated with a chemical reaction that can be used to do work ΔG = ΔH- TΔS 8. The factors affecting equilibrium Tendency towards minimum potential energy (enthalpy) ΔH= Hproducts- Hreactants exothermic direction is favoured. -ΔH<0: Forward reaction is exothermic and favoured +ΔH>0: Forward reaction is endothermic and reverse reaction is favoured. Tendency towards maximum randomness (entropy) ΔS= Sproducts- Sreactants +ΔS>0: products are more random and the forward reaction is favoured. -ΔS<0: products are less random and the reverse reaction is favoured. - low temperatures result in enthalpy change having the greatest influence and exothermic reactions are typically spontaneous high temperatures results in the random motion of molecules to increase and the entropy to have more influence on the equilibrium state 9. Gibb’s free energy - enthalpy and entropy are incorporated predict whether the forward or reverse reaction is favoured ΔG= ΔH- TΔS ΔG < 0 (-) -> forward reaction is spontaneous ΔG > 0 (+) -> forward reaction is not spontaneous ΔG = 0 -> equilibrium is reached 10. Introduction to equilibrium Dynamic reactions are still occurring on the microscopic level but there are no macroscopic changes Closed system no gain or loss of materials with the surroundings energy may be exchanged Steady state system not an equilibrium - open system that appear to be in equilibrium but are not since there is no reverse process occurring - constant input of materials and energy that are equivalent to the output of materials and energy Types of equilibrium: Solubility equilibria Equilibrium between a solvent and a solute of a saturated solution. Phase equilibria Equilibrium between different states of matter in a closed system. Chemical reaction equilibria Equilibrium between reactants and products of a chemical reaction in a closed system. 11. Chemical equilibria - system is closed forward rate = reverse rate the concentrations of reactants and products are constant temperature and pressure remain constant same equilibrium state can be reached by starting with reactants or products 12. Equilibrium constant Solids are omitted. Here is a helpful video about calculating concentration at equilibrium https://youtu.be/J4WJCYpTYj8 Why can someone get an A+ using this study guide (how are my conventions helpful)? This study guide has been designed to help any student get an A+ because it has different types of conventions involved throughout the guide to incorporate different learning styles. I have different types of learning styles which involve doing practice questions, reading notes, watching videos, doing activities, re-writing notes, etc. There are also different kinds of learners like hands-on learner, auditory learner or visual learner, which this guide all accommodates for. This study guide helps students learn and revise Atomic and bonding theory, Thermochemistry and Kinetics energy, along with Equilibrium. For hands-on learners, there are practice quizzes, matching games and practice questions. These are especially helpful for thermochemistry and Kinetics energy because a lot of math is involved in it. I found it particularly helpful when I did my homework questions continuously and checked back to the solutions to see if I was on the right track. Hence, I added different levels of questions with descriptive solutions. For auditory learners. There are videos that thoroughly explain the unit. When I tried to understand hybridization, I found it useful to watch different videos which helped me thoroughly comprehend. For visual learners, there are many notes that have been provided on the review shit, videos, notes done on my virtual white board, images, diagrams and charts. This was helpful when explaining how to calculate the concentration at equilibrium because the student can easily follow my steps from the white board. In addition to this, I find it helpful to have important factors or ideas to be highlighted so I can easily distinguish them. I used an appropriate amount of colours so that my study guide is also visually pleasing to look at. My study guide is organized so that if a student wants to look back to a certain topic, they can easily identify it.
