Chemistry Exam PDF

Thermodynamics → the study of energy and energy transfers Thermochemistry → a branch of thermodynamics that deals with the energy involved in chemical reactions Laws of Thermodynamics 1.​ law of conservation of energy -​ The total energy of the universe is constant -​ Energy...

Thermodynamics → the study of energy and energy transfers Thermochemistry → a branch of thermodynamics that deals with the energy involved in chemical reactions Laws of Thermodynamics 1.​ law of conservation of energy -​ The total energy of the universe is constant -​ Energy cannot be destroyed nor created -​ ∆ ∑ = 0 𝑢𝑛𝑖𝑣𝑒𝑟𝑠𝑒 2.​ Energy can be transferred -​ Energy can be transferred from one substance to another in various forms -​ To interpret energy changes one must define what part of the universe is being referred to -​ The system System → the part of the universe that is being studied and observed Surroundings → everything else in universe -​ 𝑢𝑛𝑖𝑣𝑒𝑟𝑠𝑒 = 𝑠𝑦𝑠𝑡𝑒𝑚 + 𝑠𝑢𝑟𝑟𝑜𝑢𝑛𝑑𝑖𝑛𝑔 -​ ∆ ∑ = ∆ ∑ + ∆ ∑ =0 𝑢𝑛𝑖𝑣𝑒𝑟𝑠𝑒 𝑠𝑦𝑠𝑡𝑒𝑚 𝑠𝑢𝑟𝑟𝑜𝑢𝑛𝑑𝑖𝑛𝑔𝑠 -​ ∆ ∑ =− ∆ ∑ 𝑠𝑦𝑠𝑡𝑒𝑚 𝑠𝑢𝑟𝑟𝑜𝑢𝑛𝑑𝑖𝑛𝑔𝑠 Thermal Energy -​ Energy available from a substance as a result of the movement of its molecules -​ Heat -​ Is expressed in Joules (J) -​ Heat is transferred spontaneously from a warmer object → a cooler object Temperature (T) -​ is the measure of the average kinetic energy of the particles that make up a system -​ Measured in℃ (a relative scale designed so that water's boiling point = 100℃ and melting point = 0℃) -​ Measured in kelvins (K) (an absolute scale designed so that 0K represents a substance with no kinetic energy (only theoretical)) -​ 𝑇𝐾𝑒𝑙𝑣𝑖𝑛 = 𝑇𝐶𝑒𝑙𝑠𝑖𝑢𝑠 + 273. 15° Types of Thermal Changes -​ Exothermic → thermal energy flows out of the system (energy written at the product side) -​ − ∆𝐻 -​ Endothermic → thermal energy is absorbed by the system (energy written at the reaction side) -​ + ∆𝐻 Chemical System → system + heat 3 types of systems 1.​ Open system → matter and energy can move in/out of the system 2.​ Closed system → only energy can move in/out of the system 3.​ Isolated system → neither matter nor energy can move in/out of the system Calorimetry → the process of measuring energy in a chemical system using a calorimeter (2 types: bomb and coffee cup) Calculating Heat → 𝑄 = 𝑚𝑐∆𝑇 -​ Q → heat (J) -​ c → specific heat capacity (𝐽/𝑔 ℃) -​ The amount of heat (J) required to raise the temperature of 1g of a substance by 1℃/1K) -​ ∆𝑇 → temperature (K/℃) -​ ∆𝑇 = 𝑇𝑓𝑖𝑛𝑎𝑙 − 𝑇𝑖𝑛𝑖𝑡𝑖𝑎𝑙 Enthalpy [H] -​ The amount of heat contained in a substance due to its thermal energy -​ It is the heat content off a substance -​ Chemists do not work with the enthalpy of a substance, but the change in enthalpy before and after a reaction -​ Measured in Joules Enthalpy Changes -​ ∆𝐻= enthalpy change is the difference in enthalpies between the reactants and products in a chemical reaction -​ Energy is absorbed/released to the surroundings when a system changes from reactants → products  -​ ∆𝐻 results from chemical bonds (sources of energy) being broken in reactants and re-formed in different ways in the products Enthalpy → The measure of the total energy of a thermodynamics system (represented by H) -​ In chemical reactions only the change in enthalpy is required -​ ∆𝐻 = 𝑐ℎ𝑎𝑛𝑔𝑒 𝑖𝑛 𝑒𝑛𝑡ℎ𝑎𝑙𝑝𝑦 Thermochemical equation → an equation indicating the amount of heat released/absorbed during a chemical reaction -​ Can be indicated in 4 different ways 1.​ Placing an energy value in an equation 2.​ Add an ∆𝐻 beside the equation 3.​ Stating molar enthalpies (∆𝐻 𝑥) of a reaction 4.​ Potential energy diagrams Exothermic Reactions → products have less potential energy than reactants, therefore excess energy is released into surroundings Endothermic Reactions → products have more potential energy than reactants, therefore energy must be absorbed by the surroundings to form them Enthalpy → the measure of the total energy of a thermodynamics system Molar Enthalpy (ΔHx) → the enthalpy change associated with a physical, chemical, or nuclear change involving one mole of a substance -​ A thermochemical equation represents the energy change that accompanies a chemical reaction -​ The enthalpy change/mol of a substance undergoing change is called the molar enthalpy (ΔHx), where x represents the type of the change occurring -​ The amount of energy involved in a change depends on the quantity of matter undergoing the change -​ To calculate an enthalpy change, ΔH (kJ), for some amount other than a mole, use the formula ΔH = n ΔHx -​ where n = # of moles and ΔHx = molar enthalpy (kJ/mol) Calorimetry → the analysis is based on the law of conservation of energy: the total energy change of a chemical system is equal to the total energy change of the surroundings -​ ΔHsystem = - qsurroundings 2 rules of solving problems using Hess’ Law 1.​ If a chemical equation is reversed the sign of ∆𝐻 changes 2.​ Iff coefficients in a chemical equation are altered by multiplying or dividing by a constant factor, then the ∆𝐻 is altered in the same way 3.​ All reactions take place at different speeds 4.​ Speed at which reaction occurs is called the rate of reaction 5.​ Can be measured by how quickly or slowly a reactant is consumed or a product is produced Kinetic Molecular Theory -​ Suggests particles that make up any substance are constantly moving -​ The faster particles of one reactant move, the more likely they are to collide with particles of the other reactant -​ This increases likelihood of molecules coming apart and combining to form new molecules Collision Model of how a chemical reaction occurs -​ Rate of reaction is affected by the number of collisions of reactant molecules Two Ways to make a reaction go faster 1.​ increase the number of collisions 2.​ increase the fraction of collisions that are effective -​ an effective collision requires the correct angle and speed 5 factors that affect the rate of reaction 1.​ Temperature  -​ As temperature increases, the average speed of the molecules increases -​ This allows for more collisions, and may allow molecules to hit hard enough for chemical bonds to break and new molecules to form 2.​ Concentration -​ When more molecules are packed into smaller space, they are more likely to collide with one another -​ Thus, increasing the concentration of reactants in a container increases the number of collisions b/w molecules 3.​ Surface area -​ Definition: amount of area of a sample of matter that is visible and able to react -​ Decreasing the reactant size by breaking it into pieces will increase its surface area -​ The more area of contact (higher surface area), the faster a reaction will proceed because more particles are available to collide with one another 4.​ Catalysts -​ Definition: a substance that speeds up the rate of a chemical reaction without being consumed. It is NOT a reactant! -​ Provide an easier way for a chemical reaction to occur by allowing for more effective collisions to be made -​ Decrease the amount of collision energy molecules need to break bonds and form new molecules -​ Enzymes - catalysts found in biological systems -​ they speed up the formation of molecules needed by the body -​ are specific to the substrate they bind to -​ enzymes can also speed up the breaking down of substances in the body -​ common in digestion 5.​ The nature of the reactants -​ Elements in same periodic table group may react similarly but their reaction rate may vary -​ e.g. Zn, Fe, Pb react with O2 but at very different rates -​ Activity series of metals demonstrates this characteristic of materials -​ Monatomic ion reaction is extremely fast -​ 5 Fe2+ + MnO- + 5 H+ → 5 Fe3+ + Mn2+ + 4 H2C -​ Fe2+ + Ce4+ → Ce3+ + Fe3+ -​ involve only transfer of e- -​ Molecular substance reaction is much slower -​ 5 C2O42- + 2 MnO4- + 16 H+ → 2 Mn2+ + 8 H20 + 10CO2 -​ involve the breaking & reforming of several chemical bond Collision theory  -​ Recall that as temperature increases, the average kinetic energy increases. Therefore: 1.​ more collisions will take place 2.​ collisions will be more effective Maxwell-Boltzmann (Kinetic Energy) Distribution -​ Recall that temperature is the average kinetic energy of the particles in matter. -​ A Maxwell-Boltzmann curve is a probability distribution representing the kinetic energy (speed) of particles in a gas. -​ These distributions appear as smooth curves due to the number of particles involved. -​ As temperature increases, the kinetic energy distributions flatten & shift to the right. Activation Energy -​ The activation energy (Ea) is the energy threshold required for a successful collision. -​ Depends on the KE of particles -​ Is unique to the reaction -​ Slow reaction has high activation energy -​ Fast reaction has a low activation energy Temperature and Rate -​ At high temperature, a greater proportion of the particles will possess the required activation energy: Potential Energy Diagrams -​ These graphs show the relative potential energy (EP or P.E.) vs. reaction progress. -​ The activation energy represents a high-energy barrier that is required to form an activated complex. -​ Activated complex/transition state → the temporary arrangement of atoms w/ partial bonds -​ If particles have enough energy the activated complex becomes products, if not, it breaks up into reactants -​ The greater the energy barrier (Ea) the lower the rate of reaction. -​ Only the sign (+ or -) of the enthalpy change (ΔH) changes when a reaction is reversed, but the activation energy (Ea) is NOT the same for the forward and reverse reactions! -​ Exothermic reactions will possess a lower Ea than the reverse (endothermic) reaction. Endothermic Reactions -​ convert kinetic energy into potential energy (Ep) (decrease temperature) Exothermic Reactions -​ The reverse of an endothermic reaction is an exothermic reaction. ∆𝐻 = 𝐸𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠 − 𝐸𝑟𝑒𝑎𝑐𝑡𝑖𝑜𝑛𝑠 𝐸𝑎 = 𝑃𝐸 𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠 − 𝑃𝐸𝑎𝑐𝑡𝑖𝑣𝑎𝑡𝑒𝑑 𝑐𝑜𝑚𝑝𝑙𝑒𝑥 A + B​ → AB​​ -​ This symbol means that > 99% of the product AB is formed, but in the reverse direction: AB → no reaction -​ Therefore in the opposite direction < 1% of AB reacts. Equilibrium Systems -​ Many reactions actually take place in both the forward ans reverse directions at the same time -​ These reactions can be shown together using a double arrow -​ A + B ⇌ AB -​ If left long enough, chemical systems may reach a stable state called dynamic equilibrium Dynamic Equilibrium -​ Rate of the forward reaction is the same as the rate of the reverse reaction -​ The reaction is still continuing, however, products are forming and “reforming” at the same rate -​ Therefore, in this state, the concentrations of the reactants and products remains constants -​ 2 types 1.​ Phase equilibrium 2.​ Chemical equilibrium Properties of a System in Equilibrium 1.​ The system is closed 2.​ The forward reaction rate equals the reverse reaction rate 3.​ The concentration of the reactants and products are constant 4.​ The temperature and pressure remain constant 5.​ The same equilibrium state can be reached by starting with either the reactants or the products Phase Equilibrium  -​ A sealed flask containing water will contain a mixture of liquid and vapour water that will eventually establish a stable or balanced condition known as equilibrium -​ H₂O₍ ₎⇌H₂₍g₎ Chemical equilibrium 1.​ Reaction Begins -​ No products yet formed -​ High rate of collisions between A+B -​ Rate of forward reaction high 2.​ Products formed 3.​ Rate of Forward Reaction Equal to rate of reverse reaction Chemical Reactions -​ Scientists use the following to study reactions: -​ Thermodynamics: Determines whether a reaction will occur at a certain temperature & when equilibrium will be reached -​ Rate of reaction: Determines the time it takes for a certain concentration of product to form -​ Equilibrium Constant: The extent of a reaction, the relative concentrations of products to reactants at equilibrium Law of Chemical Equilibrium -​ At equilibrium, there is a constant ratio between the concentrations of the products & reactants Equilibrium Constant -​ For general chemical reaction: ​ ​ aP + bQ ⇄ cR + dS -​ Where: -​ P, Q, R, S represent chemical formulae -​ a, b, c, d represent respective coefficients in chemical equation 𝑐 𝑑 [𝑅] [𝑆] -​ 𝐾ₑ𝑞 = 𝑎 𝑏 [𝑃] [𝑄] -​ Keq represents equilibrium constant -​ Kc represents equilibrium in moles per liter How do you know when a reaction has reached equilibrium? -​ Substitute concentrations of reactants & products into an expression that is identical to equilibrium expression -​ However, expression is given a different name→ reaction quotient Q  -​ Compare the Q value to the K value aP + bQ ⇄ cR + dS -​ Reaction quotient expression:​ 𝑐 𝑑 [𝑅] [𝑆] ​ 𝑄𝑒𝑞 = 𝑎 𝑏 [𝑃] [𝑄] -​ if Qeq = Keq, the system is in equilibrium -​ If Qeq > Keq, the system’s products [right side] are greater than reactants [left side], & must move towards the left to attain equilibrium -​ If Qeq < Keq, the system’s products [right side] are less than reactants [left side], & must move towards the right to attain equilibrium -​ **Qc is most commonly used (as Kc) because concentrations are in moles/litre Le Châtelier’s Principle -​ When a reaction in equilibrium is disturbed, the reaction will do the opposite of the disturbance in order to re-establish equilibrium -​ Ex: If products are removed from the system, more products must be formed to relieve the change. -​ Stressors include: -​ Change in concentration of any substance -​ Change in temperature -​ Change in pressure/volume -​ For Le Châtelier's Principle to work, the chemical system MUST -​ Already be at equilibrium (i.e. concentrations of substances are constant) -​ must be exposed to a stressor (ie. change in concentration, pressure/volume, or temperature) -​ if something is increased, the reaction will shift to the left or to the right in order to decrease it (or vice-versa) -​ If a stress is applied to a system at equilibrium, then the system readjusts, if possible, to reduce the stress. Effect of Ions on Aqueous Equilibrium Systems -​ Common Ion Effect: adding an ion to a solution in which the ion is already present -​ same effect as increasing the concentration -​ Equilibrium shifts away from the added ion  -​ Precipitation of insoluble salts is used to identify the presence of unknown ions Effect of Temperature Change on Equilibrium Systems -​ Heat is considered like a reactant or product -​ In endothermic reactions (ΔH is positive, or heat is a reactant) -​ An increase in temp shifts equilibrium to R, forming more products; -​ Kc increases -​ A decrease in temp shifts equilibrium to L, forming more reactants; -​ Kc decreases -​ In exothermic reactions (ΔH is negative or heat is a product) -​ An increase in temp shifts equilibrium to L, forming more reactants; -​ Kc decreases -​ A decrease in temp shifts equilibrium to R, forming more products; -​ Kc increases Effect of Changes in Volume and Pressure on Equilibrium Systems -​ Volume / Pressure Change with system gases: -​ Reducing the volume of an equilibrium mixture of gases, at constant temperature, causes a shift in equilibrium in the direction of fewer gas molecules. -​ e.g. ​ 2 SO3 ⇄ 2 SO2 + O2 -​ Decreasing volume (thus increasing pressure) pushes equilibrium to the L where there are only 2 molecules instead of 3 as on the R -​ Changing volume does not change Kc -​ Introduction of Inert Gases: -​ Can increase pressure by injecting more gas which if it reacts with other gases affects equilibrium -​ If injected gas is inert, there is no effect on equilibrium as added gas is not part of system (e.g. N2 [low reactivity], noble gases) -​ If container expands, then same scenario as above Effects of Catalysts on Equilibrium Systems -​ Catalyst speeds up the rate of reaction either by allowing a different reaction mechanism or by providing additional mechanisms -​ Overall lowers activation energy -​ Does not affect equilibrium only time taken to achieve it Calculate Equilibrium Concentrations from Initial Conditions→ ICE box calculations  -​ Calculating Kc from a known set of equilibrium concentrations seems pretty clear. You just plug into the equilibrium expression and solve for Kc. -​ Calculating equilibrium concentrations from a set of initial conditions takes more calculation steps. In this type of problem, the Kc value will be given ICE -​ I= initial concentrations: -​ The first two values were specified in the problem and the last value ([HI] = 0) from the fact that the reaction has not yet started, so no HI could have been produced yet -​ C= change in concentrations: -​ minus sign comes from the fact that the H2 and I2 amounts are going to go down as the reaction proceeds. -​ “x” signifies that some H2 and I2 get used up, but how much is not known. However, an EQUAL amount of each will be used up (from the coefficients of the equation). For every 1 H2 used up, 1 I2 is used up also. -​ The positive signifies that more HI is being made as the reaction proceeds on its way to equilibrium. -​ The 2 is important. HI is being made twice as fast as either H2 or I2 are being used up. -​ In fact, always use the coefficients of the balanced equation as coefficients beside the "x" terms. -​ In problems such as this one, never use more than one unknown. Since we have only one equation (the equilibrium expression) we cannot have two unknowns. -​ E= equilibrium concentrations: -​ Are the initial conditions with the change applied to it: Calculating -​ Units are not included when using or calculating the value of K -​ Let “x” represent the substance with the smallest coefficient in the chemical equation. This helps to avoid fractional values of “x” in the equilibrium expression which make solving expressions more difficult. -​ Noticing perfect squares in equilibrium expression makes solving easier by taking the square root- avoiding the quadratic equation. -​ Be careful to write the terms of the equilibrium expression correctly; e.g. [HI]2 should be written as [2x]2 = 4x2 & not 2x2. Avoid this mistake by using brackets. Acids -​ give off hydrogen ion (H+) when dissolved in water -​ are corrosive to organic tissues and metals -​ conduct electricity -​ - taste sour (ex: vinegar, lemon juice) Bases  -​ give off hydroxide ions (OH-) when dissolved in water -​ are corrosive to organic tissues -​ conduct electricity -​ taste bitter (ex: dark chocolate,) -​ when acids/ bases dissolve in water they dissociate: ionize, break apart into their ions -​ NaOH(aq) → Na+(aq) + OH-(aq) -​ this is a reversible reaction, therefore, it will reach an equilibrium -​ an acid/base is STRONG when there is mostly product at equilibrium -​ an acid/base is WEAK when there is mostly reactant at equilibrium Strength of Acids/Bases -​ Can be determined 2 ways: a.​ by the value of the equilibrium constant -​ high Keq value means more products formed, therefore a stronger acid or base -​ Low value for Keq means less products formed, therefore a weaker acid or base. b.​ By the pH value -​ pH = -log10 [H3O+] (measured in mol/L) Acid-Base Theories -​ Svante Arrhenius -​ acids produce H+ ions in aqueous solutions -​ bases produce OH- ions in aqueous solutions -​ water required, so only allows for aqueous solutions -​ only hydroxide bases are allowed -​ Johannes N. Brønsted-Thomas M. Lowry -​ acids are proton donors -​ bases are proton acceptors -​ only for aqueous solutions -​ bases besides hydroxides are permissible -​ Gilbert N. Lewis -​ acids are electron pair acceptors -​ bases are electron pair donors -​ least restrictive of acid-base definitions Conjugate Acid-Base Pairs -​ acids are proton donors -​ however, in the reverse reaction, the substance accepts a proton -​ therefore, in the reverse reaction, the substance is a base  -​ in the reverse reaction, it is a conjugate base -​ an acid donates a proton -​ therefore, the reactant that gains an extra H on the product side, is the acid -​ the acid in the reactants becomes the conjugate base on the product side -​ the base accepts the proton -​ what is a base on the reactant side, becomes the conjugate acid on the product side -​ Conjugate base: -​ is what remains after the acid has donated a proton -​ this species is a base because it can accept a proton to re-form the original acid -​ Conjugate acid: -​ is what results after the base has accepted a proton. This species is an acid because it can give up a proton to re-form the original base: -​ Since a strong acid completely turns into product, it is unlikely that it will turn back into reactants -​ Conjugate bases of a strong acid are weak bases -​ the stronger an acid, the weaker its conjugate base -​ Since a weak acid barely turns into product, it is highly likely that the product will turn back into reactant -​ the weaker an acid, the stronger its conjugate base -​ according to the Bronsted-Lowry definitions of acids and bases: acids are proton donors, bases are proton acceptors -​ the same substance can be an acid in one reaction, but a base in different reaction -​ such substances are AMPHIPROTIC -​ in some reactions they donate protons, in other reactions they accept protons Auto-Ionization of Water -​ Autoionization of water: reaction between 2 water molecules producing a hydronium ion [H3O+] & a hydroxide ion [OH-] -​ At 25oC, only about 2 water molecules in one billion dissociate. -​ water is amphoteric (amphiprotic) -​ acts as acid or base Hydronium Ions -​ a dissociated hydrogen ion (+) will bond to the negative side of a water molecule -​ the H ion does not have any valence electrons -​ the oxygen atom donates 2 electrons to hydrogen to fill its valence shell -​ the water molecule and H ion are bonded together to form hydronium (H3O+) Acids, Bases, and Water at Equilibrium -​ when an acid dissociates in water, an equilibrium equation can be written for the reaction: HA(aq) + H2O(l) ⇄ H3O+(aq) + A-(aq) -​ Ka = [H3O+] [A-] / [H2O] [HA-] ​ -​ however, this can be simplified in 2 ways: 1.​ the H3O+ ion can be written as H+ 2.​ the concentration of liquids and solids in a reaction remains constant, so they are not included in the equilibrium equation -​ Therefore, the equilibrium equation becomes: -​ Ka = [H+] [A-] / [HA] -​ This is called the acid ionization constant, Ka (used when an acid ionizes in water) -​ the same idea applies to bases, called the base ionization constant, Kb (used when a base ionizes in water) -​ water molecules also dissociate -​ H2O ⇄ H3O+ + OH- -​ therefore, an equilibrium equation can also be written for water: -​ Kw = [H3O+] [OH-] / [H2O] -​ however, this can be simplified to: -​ Kw = [H+] [OH-] -​ Kw = 1.0 x 10-14 mol2/L2 -​ called the ion-product constant of water, Kw -​ this can be used to find the concentration of the H+ or OH- ions in a solution What does Ka and Kb tell us? -​ when the Ka or Kb values are LARGE, the acid or base is STRONG -​ when the Ka or Kb values are SMALL, the acid or base is WEAK -​ when an acid is strong (large Ka), its conjugate base is weak (small Kb) -​ when an acid is weak (small Ka), its conjugate base is strong (large Kb) -​ in a weak acid or base [OH-] and [H3O+] concentrations are very small, therefore, concentrations of these ions in solutions can be expressed as pH -​ if the pH of a substance is known, the pH can also be used to find [H3O+] and [OH-] -​ since the pH scale goes up to 14, the pH and pOH values also add up to 14 -​ pH + pOH = 14.00 To calculate pH, [H+], and [OH−], you can use the following formulas: 1.​ pH= −log⁡[H⁺] −𝑝𝐻 2.​ [H⁺] = 10  3.​ [OH⁻]: You can calculate [OH⁻] using the relationship with the ion product of water (Kw): −14 𝑘𝑤 where Kw​at 25°C is 1 × 10 [OH⁻]= [𝐻⁺] 4.​ pOH: You can also find pOH using:​ pOH=14−pH​ and then calculate [OH⁻] from:​ −𝑝𝑂𝐻 [OH⁻]=10 Models of the Atom -​ A model should explain not just what the material is made of (composition) but also how it is going to behave -​ Rutherford’s atomic model could not explain the chemical properties of elements. -​ Basically, it couldn’t explain why things change color when heated. The Bohr model -​ Niels Bohr (1885-1962) was a Danish physicist and a student of Rutherford’s -​ In 1913, Bohr proposed that an electron is found only in specific circular paths, or orbitals, around the nucleus -​ Bohr’s model was the first model to explain how atoms behaved -​ was the beginning of the Quantum Mechanical Model of the atom -​ energy of electrons is quantized: each electron has a specific, fixed amount of energy (like dollar bills are only in specific, fixed amounts) -​ electrons with the same amount of energy orbit at the same distance from the nucleus → called energy level  -​ The higher the energy level, the farther it is from the nucleus -​ when electrons gain energy they can jump from one energy level to a higher one -​ higher energy level = excited state -​ electrons will return to their original orbital, the ground state orbital -​ in doing so, the electron will give off the specific amount of energy it gained (because energy of electrons is quantized) in the form of light -​ depending on the wavelength of light, we see it as different colours -​ when electrons lose energy they release the quantum of energy absorbed as light -​ the wavelength of the energy given off corresponds to a specific colour of light -​ this creates an emission line spectrum specific to each element -​ line spectra are specific to each element and can be used to identify an element Significance of Emission Spectra -​ lines in the spectrum correspond to the transition of electrons from higher energy levels to lower energy levels -​ proves the existence of quantized energy levels and sublevels in an atom Quantum Mechanical Model -​ Rutherford’s and Bohr’s model described an electron like a particle (like a small baseball) -​ Austrian physicist Erwin Schrödinger stated that the electron is a standing wave -​ Quantum Mechanical Model comes from the mathematical solutions to the Schrödinger equation -​ there are no orbitals in this model -​ orbitals are regions of space where there is a high probability of finding an electron -​ Heisenberg Uncertainty Principle: states that it is impossible to know both the exact position and speed of an electron -​ therefore we can only describe the probability of finding an electron in a specific location -​ The electron cloud of an atom can be compared to a spinning airplane propeller. The propeller blade has the same probability of being anywhere in the blurry region, but you cannot tell its location at any instant  -​ The probability of finding an electron within a certain volume of space surrounding the nucleus is represented as a fuzzy cloud. -​ The cloud is more dense where the probability of finding the electron is high -​ orbitals are represented by probability clouds -​ orbitals can have various shapes -​ orbitals are named with letters -​ according to the Quantum Mechanical model of atoms, electrons are arranged in shells or energy levels -​ each shell or energy level is divided into smaller energy sublevels (designated by s,p,d,f) -​ each sublevel is made up of specific orbitals, which can only hold a maximum of 2 electrons -​ s-orbitals are spherically shaped, therefore there is only 1 orbital -​ p-orbitals are “dumbbell” shaped -​ there is one along each axis, therefore, there are 3 orbitals -​ Four of the five d-orbitals have the same shape but different orientations in space. https://docs.google.com/document/d/1v5BchmXW1sy8Ito-A1kFYxMWSR9LqsePAqCYNpBNa M4/edit?usp=sharing https://docs.google.com/document/d/188lqXRY9SvoM7N9Scxr3vPQe0KrLy4lijzm1QzFN59o/e dit?usp=sharing https://docs.google.com/document/d/1XGLZY6mzxZMwZEW5lhifM7UYeZm_g8YEPRHIP3c_ 6dk/edit?usp=sharing -​ location and energy of electrons in an atom is determined by 4 variables from the Schrödinger equation -​ each electron in an atom has a specific and unique set of quantum numbers -​ quantum numbers for an electron are like their home address 4 Quantum numbers 1.​ Principle Quantum number (n) -​ the energy level -​ larger n-value, farther from the nucleus 2.​ Angular Momentum Quantum Number (l) -​ shape of the orbital -​ a.k.a. Angular Momentum  -​ Determined by n-1 3.​ Magnetic Quantum Number (m1) -​ -l to +l -​ determines how many orbitals are in the sub-level -​ refers to a specific orbital in a sub-level -​ Ex A: when l = 0, ml = 0, therefore, there is only one orbital (this makes sense since l=0 refers to sub-level s) -​ Ex B: when l=2, ml = -2, -1, 0, 1, 2(therefore, in sub-level d, there are 5 orbitals) 4.​ Spin Quantum Number (ms ) -​ represents direction of spin of an electron -​ can either be -½ or + ½ https://docs.google.com/document/d/1v5BchmXW1sy8Ito-A1kFYxMWSR9LqsePAqCYNpBNa M4/edit?usp=sharing -​ the more energy an electron has, the farther away from the nucleus it can be found -​ each energy level, n, has a probability distribution for where an electron can be found -​ there probability distributions are called orbitals -​ each orbital can only contain 2 electrons -​ orbitals are separated into sublevels based on their shape Electron Configuration  -​ The electron configuration of an atom is a shorthand method of writing the location of electrons by sublevel. -​ The sublevel is written followed by a superscript with the number of electrons in the sublevel. -​ If the 2p sublevel contains 2 electrons, it is written 2p2 -​ the arrangement of electrons in an atom can be represented 2 ways: 1.​ Energy Level/Orbital Diagram (2s↑↓↑↓…) 2.​ Electron Configuration Diagram (1s^2 2s^2…) -​ how electrons are arranged in the orbitals -​ there are 3 rules to follow when determining electron configurations: -​ the Aufbau Principle -​ the Pauli Exclusion Principle -​ Hund’s rule Electron Configuration Rules -​ Aufbau Principle states: -​ electrons occupy the orbitals of lowest energy first -​ Sub-levels overlap, therefore follow the diagram to place electrons -​ since orbitals overlap, and electrons fill orbitals with lower energy first, electrons will fill orbitals in a specific order -​ electrons fill lower energy orbitals first, THEN higher energy orbitals -​ Pauli Exclusion Principle states: -​ an atomic orbital may contain a maximum of 2 electrons -​ to occupy the same orbital, two electrons must have opposite spins; that is, the electron spins must be paired -​ Hund’s Rule states: -​ electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible -​ therefore, electrons fill orbitals individually, then double up Filling Diagrams for Sublevels -​ Aufbau Principle -​ follow the arrows to fill the orbitals in the correct order -​ some energy levels (shells) overlap each other, therefore, some lower energy level orbitals end up being filled after the higher energy level orbitals Writing Electron Configurations -​ First, determine how many electrons are in the atom. -​ Ex: Iron has 26 electrons  -​ arrange the energy sublevels according to increasing energy according to Aufbau Filling Diagram: -​ 1s 2s 2p 3s 3p 4s 3d … -​ fill each sublevel with electrons until you have accounted for all the electrons in the atom: -​ Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d 6 -​ sum of the superscripts -​ the atomic number of iron (26) -​ An excited atom has an electron or electrons which are not in the lowest energy state -​ excited atoms are energetically unstable and eventually fall back to their ground state -​ * is used to indicate an excited atom Electron Configurations and the Periodic Table -​ The periodic table can be used as a guide for electron configurations, just like the Aufbau diagram -​ The period number is the value of n (energy level) -​ Groups 1A and 2A have the s-orbital filled, s-block -​ Groups 3A - 8A have the p-orbital filled, p-block -​ Groups 3B - 2B have the d-orbital filled, d-block -​ lanthanides and actinides have the f-orbital filled, f-block -​ use the periodic table to predict which is the last sublevel being filled by a particular element Condensed Electron Configurations -​ electron configuration can be abbreviated by substituting the innermost electrons with the symbol of the preceding noble gas -​ the preceding noble gas with an atomic number less than sodium is neon, Ne -​ Condensed electron configuration: ​ Na: [Ne] 3s1 -​ Neon completes the 2p subshell -​ Sodium marks the beginning of a new row -​ So, we write the condensed electron configuration for sodium as: Na: [Ne] 3s1  -​ [Ne] represents the electron configuration of neon, then add the valence electrons -​ Valence electrons: electrons outside of [Noble Gas] Exceptions to the Electron Configuration Rule -​ Some actual electron configurations differ from those assigned using the Aufbau Principle because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations -​ Electrons are most stable with full sublevels, then next stable is half-filled sublevels, then other configurations -​ Exceptions to the Aufbau Principle are due to subtle electron-electron interactions in orbitals with very similar energies https://docs.google.com/document/d/188lqXRY9SvoM7N9Scxr3vPQe0KrLy4lijzm1QzFN59o/e dit?usp=sharing Factors Affecting the Properties of Elements -​ Many of the properties of the elements are related to the force of attraction between the nucleus and the electrons. -​ The force of attraction is dependent on 3 factors: 1.​ distance between nucleus and electrons -​ As the distance between the positive nucleus and the negative valence electrons increases, the attraction between them decreases -​ therefore, electrons in higher energy levels are held less tightly because they are farther away 2.​ size of the positive charge of the nucleus -​ effective nuclear charge (Zeff): net force of attraction between the nucleus and the electrons -​ As more protons are added to the nucleus, the effective nuclear charge (+) increases (higher atomic number) 3.​ Shielding effect -​ The further away valence electrons are, the less they are attracted to the nucleus, because other electrons in lower energy levels are repelling them, which “shields” the valence electrons from the attraction from the nucleus Atomic Radius → is the distance between the nucleus of an atom and its valence shell  -​ determined by measuring the distance between nucleus of one atom and nucleus of next atom ÷ 2 -​ Down a group: -​ Increases due to: 1.​ more orbitals being added 2.​ as more orbitals are added, Zeff decreases due to the Shielding effect: inner electrons repel outer electrons and decrease the nucleus’ pull on the electrons -​ Across a Period -​ the size of the nucleus (+ charge) increases, but the number of orbitals is the same -​ therefore, greater nuclear charge pulls electrons closer to the nucleus, atomic radius DECREASES across a period Ionic Radius → the radius of an ion -​ like atomic radius, ionic radius increases moving down a group because more orbitals are added -​ however, moving across a period is a bit different: -​ Metallic ions (left side of periodic table) -​ when atoms lose electrons, they are losing their valence shell, therefore the radius is smaller -​ also, when electrons are lost, there is a stronger effective nuclear charge (Zeff), therefore, electrons are pulled in closer -​ Non-metallic ions (right side of periodic table) -​ these atoms GAIN electrons to fill their outer shell, therefore, the effective nuclear charge (Zeff), decreases, and the ionic radius increases Electronegativity → measure of the ability of an atom to attract electrons  -​ determined by larger positive nuclear charge, and smaller atomic radius -​ higher electronegativity means a stronger pull on electrons in the atom and other nearby atoms -​ decreases as you move down a group -​ due to electron shielding -​ increases as you move across a period -​ stronger effective nuclear charge (Zeff) because nucleus gets larger and atomic radius gets smaller -​ noble gases are not electronegative! (zero, or close to zero electronegativity) -​ their electron affinity is very low -​ valence shell is full, therefore, if another electron was added/attracted, it would go in the next energy level, this is a higher energy level configuration for the electron, and therefore, is not an ideal situation for the new electron -​ electrons need more energy to orbit farther from the nucleus, therefore, energy must be added to the atom in order to attract the electron, therefore, no or very low electronegativity Ionization Energy → the energy required to remove one electron from an atom in the gas state -​ First Ionization Energy is the amount of energy required to remove the first electron from an atom -​ every time an electron is removed from an atom, there is a stronger positive charge pulling in the remaining electrons -​ therefore, ionization energy increases for each successive electron removed from the same atom -​ first ionization energy < second < third < etc… -​ moving across a period, the effective nuclear charge (Zeff) increases, therefore, it would take MORE energy to overcome this attraction and remove an electron -​ therefore, ionization energy INCREASES moving across a period -​ when orbitals are added, ionization energy decreases due to increased distance of electrons from the nucleus, and increased Shielding Effect -​ therefore, ionization energy DECREASES moving down a periodic table -​ ionization energy increases moving across the periodic table, until the inert gas -​ explain the factors that cause the large drop after inert gas Electron Affinity → the amount of energy released when an electron is added to a neutral atom  -​ energy is released when the atom attracts an electron because the atom becomes more stable -​ when energy is released, the measurement is negative -​ however, second and third affinities are positive, that is, energy must be put in to add an electron, because after the first electron addition, the atom is now more negative and will repel electrons -​ large negative numbers represent high electron affinity -​ small negative numbers or positive numbers represent low electron affinity Points to Consider... 1.​ Chemists have learned that the “most important requirement for the formation of a stable compound is that the atoms achieve noble gas configurations” 2.​ Only valence e-'s are included when writing Lewis structures 3.​ If the valence e- do not exist for an ion (i.e. cation) do not show any dots; For an anion show all, plus those gained Lewis structures for molecules with covalent bonds (Elements of 1° and 2™ periods) Rules 1.​ Sum the valence e-’s from all the atoms. Don’t worry about keeping track of which e- comes from which atom; It’s the total # of electrons that’s important 2.​ Use a pair of e- to from a bond between each pair of bond atoms 3.​ Arrange remaining e- to satisfy the duet rule_for H and the octet rule for 2% row elements Steps to check 1.​ Total # of e-'s is X 2.​ Octet rule has been obeyed for each atom 3.​ Rearrange the e-'s using trial and error Final Thought: Exceptions to the Octet Rule  -​ The Localized electron (LE) model, assumes “a molecule is composed of atoms that are bound together by sharing pairs of e- using the atomic orbitals of the bound atoms. e -​ Electrons localized on an atom are lone pairs e -​ Electrons found in space between atoms are bonding paris 1.​ Boron (B) — tends to from compounds with fewer than 8e- BF, reacts violently with H,O, NH, (i.e. molecules with available lone pairs) WHY? Because B is e- deficient 2.​ 2. Atoms that EXCEED the octet rule (i.e. elements in period 3 and beyond) Guidelines 1.​ 2nd row elements (C, N, O, F) should always be assumed to obey the octet rule. 2.​ 2nd row elements (B & Be) often have fewer than 8e- around them in their compounds. These e- deficient compounds are very reactive 3.​ The 2nd row elements never exceed the Octet Rule, since valence orbitals (2s & 2p can accommodate only 8e-) 4.​ 3rd row and heavier emenents often satisfy the octet rule but can exceed the Octet Rule by using their empty valence d orbitals 5.​ When writing Lewis structures for a molecule a.​ satisfy the octet rule for the atoms first b.​ if e-’s remain, after the octet rule is satisfied, then place the extras on elements having available d orbitals (elements in Period 3 or beyond)! Resonance -​ is invoked when more than 1 valid Lewis structure can be written for a particular molecule! The arrows do not mean the molecule “flips” from 1 structure to another! They simply reflect that the actual structure is an average of the 3 resonance structures! Odd electron molecules -​ Odd electron molecules are not handled well by the (LE) model, since this model is based on pairs of electrons. We require the use of a more sophisticated model, that relies on the calculation of Formal Charge. Formal Charge -​ Some molecules & polyatomic ions have atoms that can exceed the octet rule, and have many nonequivalent Lewis structures, all of which obey the rules -​ applies to the atom in a molecule, and is the difference between the number of valence e-'s on the free atom & the number of valence e-’s assigned to the atom in the molecule. 1 𝑓𝑜𝑟𝑚𝑎𝑙 𝑐ℎ𝑎𝑟𝑔𝑒 = 𝑣𝑎𝑙𝑒𝑛𝑐𝑒 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑛𝑠 − 𝑢𝑛𝑏𝑜𝑛𝑑𝑒𝑑 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑛 𝑝𝑎𝑖𝑟𝑠 − 2 (𝑏𝑜𝑛𝑑𝑒𝑑 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑛𝑠) -​ 2 assumptions must be made when computing formal charges of an atom in a molecule: 1.​ Lone pair e-’s belong entirely to the atom in question 2.​ Shared e-'s are divided equally between the 2 sharing atoms -​ 2 fundamental assumptions about formal charges to evaluate Lewis structures:  1.​ atoms in molecules try to achieve formal charges as close to zero as possible 2.​ any negative formal charges are expected to residue on the most electronegative atoms -​ The structure with double bonds is preferred b/c it has lower formal charges and the -1 formal charges are on electronegative oxygen atoms -​ Rules Governing Formal Charge 1.​ To calculate the formal charge on an atom 2.​ (The sum of the formal charges of all atoms 3.​ If nonequivalent Lewis structures exist for a species, those with formal charges closest to zero and with any negative formal charges on the most electronegative atoms are considered to best describe the bonding in the molecule or ion. -​ Caution: formal charges are only estimates of charge-and should not be taken as actual atomic charges; some Lewis structures can lead to erroneous predictions Valence Bond Theory – atomic (half filled) orbitals overlap to form a new orbital with a pair of opposite-spin electrons Hybridization – the process of forming hybrid orbitals from the combination of at least 2 different orbitals Hybrid Orbital – an orbital that forms from the combination of at least 2 different orbitals Sigma (𝝈) bonds – is formed when the lobes of 2 orbitals directly overlap end to end ​ Pi (π) Bonds – when the sides of the lobes of 2 orbitals overlap Single covalent bond – σ-bond only Double covalent bond – σ-bond and one π-bond Triple covalent bond – σ-bond and two π-bonds Valence Shell Electron Pair Repulsion Theory -​ If a molecule consists of a central atom bonded to 2 or more other atoms the molecular shape can be predicted by: 1.​ Drawing the Lewis Structure for the molecule. 2.​ Counting the number of bonded atoms and lone electron pairs around the central atom. 3.​ Use VSEPR to predict the shape -​ Lone pairs occupy space around the atom and influence the shape -​ The goal is to maximize the bond angle between bonded atoms and lone pairs - reduces tension in the bonds -​ Double and Triple bonds behave the same as a Single bond when trying to figure out the shape Example 1: BrF5 1.​ Draw the Lewis structure. 2.​ Bonded atoms = 5 Lone pairs = 1 3.​ sp3d2 hybridization, octahedral electron arrangement 4.​ Square pyramidal shape Development of VSEPR Theory -​ Late 1800’s – used microscopes, crystallography, polarimeters  -​ Early 1900’s – used spectrometry -​ Early 1930’s – Gilbert Lewis → Lewis dot diagrams -​ Late 1930’s –Linus Pauling → valence bond theory [VBT] -​ Late 1950’s – Ronald Gillespie & Ronald Nyholm → Valence-Shell-Electron-Pair-Repulsion Theory [VSEPR] VSEPR -​ Only the valence shell electrons of the central atom(s) are important for molecular shape -​ Valence shell electrons are paired or will be paired in a molecule or polyatomic ion. -​ Bonded pairs of electrons & lone pairs of electrons are treated approximately equally. -​ Valence shell electron pairs repel each other electrostatically. -​ Molecular shape is determined by the positions of the electron pairs when they are a maximum distance apart (with lowest repulsion possible.) -​ predicts molecular geometry by the repulsion of paired electrons -​ These electrons repel each other to minimize repulsion. Process of Predicting molecular shape 1.​ Draw Lewis dot diagram for molecule & include electron pairs around central atom. 2.​ Count total # of bonding pairs & lone pairs of electrons around the central atom. 3.​ Use # of pairs of electrons to predict the shape of the molecule. -​ Multiple bonds can be treated like single bonds for describing & predicting the shape of a molecule. -​ E.g. Ethene C2 H4 -​ Count multiple bond as single group of electrons -​ Thus, there are 3 sets of bonded electrons around each central C-atoms -​ AX3 → trigonal planar orientation [120 o ] VSEPR Theory – valence shell electron pair repulsion theory -​ According to this theory -​ Only the valence shell electrons of the central atom are important in determining molecular shape -​ Valence shell electrons are paired/will be paired in a molecule -​ Bonded pairs and lone pairs are treated equally -​ Electron pairs repel electrostatically -​ Molecule shape is determined when electron pairs are positioned as far apart as possible (lowest repulsion) 1.​ Step 1 -​ Count the number of bonds around the central atom. Treat double and triple bonds as single bonds. -​ Count the number of lone pairs around the central atom. -​ Add them together to get the total number of charge clouds. 2.​ Step 2 -​ Predict the shape, based on the total number of charge clouds.  -​ Electron pair arrangement – includes all pairs, bonding and lone -​ Molecular shape – includes only bonding pairs -​ VSEPR theory uses a general formula for molecules “AXE” where: ​ -​ A tells you the number of central atoms -​ X tells you the number of bonding electron pairs -​ E tells you the number of lone electron pairs Charge Lone Bonds Notation Shape Clouds Pairs 2 2 0 AX2 Linear 3 3 0 AX3 Trigonal Planar 3 2 1 AX2E Bent Charge Lone Bonds Notation Shape Clouds Pairs 4 4 0 AX4 Tetrahedral Trigonal 4 3 1 AX3E Pyramid 4 2 2 AX2E2 Bent Charge Lone Bonds Notation Shape Clouds Pairs Trigonal 5 5 0 AX5 Bipyramid 5 4 1 AX4E See-Saw 5 3 2 AX3E2 T-Shape 5 2 3 AX2E3 Linear Charge Lone Bonds Notation Shape Clouds Pairs  6 6 0 AX6 Octahedral Square 6 5 1 AX5E Pyramid Square 6 4 2 AX4E2 Planar 6 3 3 AX3E3 T-Shape 6 2 4 AX2E4 Linear Ex. Consider beryllium hydride, BeH2 (covalent bond) NOTE: Be is an exception, it doesn't need 8 (Boron, Phosphorus, Sulfur, Beryllium, Helium, Hydrogen break this rule) The Lewis structure indicates two bonding pairs and no lone pairs: So the general formula for this molecule is AX2. Recall: Polar Covalent Bonds -​ The electron density is greater around the more electronegative atom. The electrons spend more time around this atom than they do the less electronegative atom -​ This end (δ+)of the molecule, where the electron density is less, is the positive pole  -​ This end (δ-) of the molecule has a negatively charged end or “pole”. Polar Molecules (this is different from a polar bond!) -​ POLAR MOLECULE: a molecule in which the uneven distribution of electrons results in a positive charge at one end and a negative charge at the other end of the entire molecule -​ NON-POLAR MOLECULE: a molecule in which the electrons are equally distributed among the atoms, resulting in no poles forming in the molecule. Water is a polar molecule -​ One end of the molecule (the oxygen atom) has a slightly more negative charge than the rest of the molecule, since the electrons in the bonds are pulled closer toward it. -​ The other end of the molecule (Hydrogen atoms) has a slightly positive charge compared to the rest of the molecule, since the electrons of the bond are far away, leaving more protons (+ charge) than electrons. -​ When a charged object is brought close to a stream of water, the water molecules will position themselves in a way that their oppositely charged pole will be closer to the object. -​ This will cause electrostatic attraction between water molecules and the charged object, since opposite charges attract. Determining if a molecule is polar -​ We can use lewis structures to help us determine if a molecule is polar. -​ A polar molecule must have at least 1 polar covalent bond in it. If there are no polar covalent bonds in the molecule it cannot be polar. -​ When drawing dipoles in polar bonds, we can add arrows that point in the direction of the negative dipole. -​ if the arrows are pointing toward the same direction, the molecule is polar. -​ When drawing lewis structures from here on out, the shape is important. -​ Whenever drawing the molecule remember that electrons will repel each other since they are negatively charged. -​ Lone pairs need more space than bonding pairs (lines), so they will repel bonds further away from them. -​ Bonding pairs will repel each other as well, but if we look at H2O on the right… -​ Even though the bonding pairs are closer together (in the bent shape of water), that’s because the lone pairs are pushing them strongly. -​ For water (H2O): There are 2 lone pairs on the Oxygen atom, so they will repel the bonds further away and form a bent shape.  -​ For CO2, there are no lone pairs on the carbon atom, therefore the bonding electrons will be equally spread out. -​ Elements in group 16 have 6 valence electrons and a bonding capacity of 2. -​ If there are 2 bonds on a group 16 atom, it will look bent instead of linear. -​ This gives the lone pairs more space to move, and in the process, these lone pairs will repel the bonding electrons. -​ To be a polar molecule, there has to be at least 1 polar bond in it. -​ But depending on how these polar bonds are arranged, the molecule itself may not be polar! -​ In other words, the shape of the molecule determines if it is polar or not. -​ The shape will determine how the dipole arrows add up. -​ It has to do with the dipole arrows drawn for polar bonds. -​ If the dipole arrows point in opposite directions, they will cancel each other out, and the overall polarity of the molecule will be zero, even though it contains polar bonds. -​ For example, since CO2 has 2 polar bonds that have dipoles pointing in opposite directions, it is a nonpolar molecule overall. Both ends of the CO2 molecule are slightly negative, and since they have the same partial charge, the molecule cannot be polar. Symmetry -​ Generally, if a molecule is symmetrical, it will be non-polar, even though it has polar covalent bonds in it. Steps to Determine the polarity of a molecule 1.​ Draw the Lewis structure for the molecule. 2.​ Determine the electronegativity difference for each covalent bond in the molecule. 3.​ Indicate whether each bond is polar (0.4 < ΔEN < 1.7) or non-polar (ΔEN < 0.4). 4.​ If there are polar covalent bonds in the molecule, indicate the partial charges (dipoles) -​ Write “δ+” by the atom with the lower electronegativity -​ Write “δ-” by the atom with the higher electronegativity. 5.​ Interpret your diagram based on SYMMETRY. -​ If the central atom has lone pairs, the molecule is considered to be asymmetrical  -​ If the central atom does not have lone pairs, or an even number of the same type of bond, the molecule is symmetrical -​ NOTE: A molecule may be nonpolar but still contain polar covalent bonds! -​ If it is symmetrical, it will be a nonpolar molecule. Intermolecular Forces -​ InterMolecular Forces are interactions between charges, partial charges and temporary charges on atoms, molecules or ions -​ list of intermolecular forces 1.​ London Dispersion Forces -​ An attractive force acting resulting from the instantaneous or temporary dipole (concentration of electrons at one end of the molecule) -​ This is present in all molecules -​ As the total number of electrons in an atom or molecule increases, the volume of the electron cloud increases and the greater dispersion forces result in increasing boiling points -​ Both these molecules have the same # of electrons and molar mass but different boiling points -​ This is because of their shape -​ The pentane molecules are long and can interact with 1 another along their entire length 2.​ Dipole-Dipole Forces -​ An attractive force acting between polar molecules where the positive end of one molecule is attracted to the negative end of a second molecule -​ If you have 2 polar molecules with equal or a similar number of total electrons: -​ ↑ the electronegativity of the bond ↑ the polarity of the bond ↑the strength of the dipole-dipole force ↑ the boiling point 3.​ Hydrogen bonds -​ A relatively strong force between a positive ‘hydrogen’ atom of one molecule and a highly electronegative atom (oxygen (O), nitrogen (N) and fluorine (F)) in an adjacent molecule 4.​ Ion-Induced Dipole -​ Polar compounds do not mix with nonpolar molecules. They are said to be immiscible -​ 2 polar compounds or 2 non-polar compounds will mix with each other. They are said to be miscible -​ As an ion approaches a non-polar atom or molecules, it influences the electron cloud of the molecule -​ Polarizability allows us to better understand the interactions between nonpolar molecules and other species -​ Polarizability - the ability for an electron cloud to be distorted -​ The ion approaches the non-polar molecule and causes an induced dipole to occur -​ ↑ charge on the ion ↑ polarizability ↑ force of attraction -​ ↑ size of the electron cloud of the non-polar molecule ↑ polarizability since electrons are more weakly held ↑ force of attraction 5.​ Dipole-Induced Dipole -​ Intermolecular force between polar and nonpolar molecules in dilute mixtures -​ A polar molecule approaches a nonpolar molecule -​ A negative or positive dipole in the polar molecule polarizes the electron cloud of the non-polar molecule -​ The non-polar molecule now has an induced dipole so it can interact with the polar molecule -​ ↑ size of the dipole in the polar molecule ↑ force of attraction -​ ↑ size of the electron cloud of the non-polar molecule↑ polarizability of the electron cloud in the non-polar molecule ↑ force of attraction 6.​ Ion-Dipole Force -​ Intermolecular Force between ions and polar molecules. -​ Occur in aqueous solutions when ion become hydrated by water molecule -​ Strongest of all the IMF in this lesson ISOMER - a molecule that has the same molecular formula but different molecular structure ** since structure determines chemical properties, this is a BIG DEAL! CONSTITUTIONAL ISOMERS -​ a.k.a structural isomers -​ Same molecular formula, but different arrangement/sequence of atoms STEREOISOMERS  -​ Molecules that have same arrangement/sequence of atoms, but a different 3-dimensional arrangement -​ 2 types 1.​ ENANTIOMERS – mirror image -​ molecular structures are not identical but can be turned to face each other and they would seem like mirror images, much like your hands -​ Hint: occurs when C have different “things” attached to each side 2.​ DIASTEREOMERS – non-mirror image -​ involve a double-bonded carbon -​ double bonds cannot rotate, therefore, these arrangements are fixed in place -​ CIS - both molecular groups are on the same side of the carbon bond -​ TRANS - molecular groups are on the opposite sides of the carbon bond Type of Formula Diagram Description Shows # and types of atoms present Empirical C5H12 Expanded Shows grouping of atoms. Brackets Molecular CH3CH(CH3)CH2CH3 indicate the location of chained branches Gives all atoms and the location of Structural bonds. Straight lines represent chemical bonds Does not show C-H bonds (because Condensed assumed present) but shows all other bonds Lines to represent bonds. Each end of Line Structural line represents C atom bonded to max H atoms needed to give it 4 bonds (unless stated otherwise) Every end or point on a line diagram represents a carbon atom. Hydrogen atoms are not shown. Organic Compounds – molecular compounds containing carbon with the exception of CO(g), CO2(g), and HCN(g) Hydrocarbons – organic compounds containing only carbon and hydrogen divided into categories -​ Types of hydrocarbons 1.​ Alkanes -​ contain only carbon-carbon single bonds​ ​ ​ ​ ​ 2.​ Alkenes -​ contain one or more carbon-carbon double bonds 3.​ Alkynes -​ contain one or more carbon-carbon triple bonds​ ​ ​ 4.​ Aromatic hydrocarbons -​ contain one or more benzene rings -​ Drawing Hydrocarbons 1.​ Space Fill Models 2.​ Structural Formula 3.​ Carbon Skeletons 4.​ Line Structures Organic Nomenclature -​ The systematic method of naming organic compounds has been established by the INTERNATIONAL UNION OF PURE AND APPLIED CHEMISTRY (IUPAC) -​ it is based on a series of rules that helps determine the structure of a compound from its name -​ the name indicates: -​ the number and arrangement of carbon atoms -​ the type(s) of functional group(s) in the compound Structure of Hydrocarbons -​ Within these hydrocarbons, the backbone of the molecule made up of carbon-carbon bonds may form: -​ a straight chain -​ a cyclic (ring) structure​ -​ 1 or more branched chains* -​ * Alkyl group - A hydrocarbon branch that is attached to the main structure of the molecule Straight Chain Alkanes (Named with the suffix (ending) –ane) -​ The prefix (beginning) tells us the # of carbons that are in the longest chain in the molecule -​ The alkyl branches in the chain are named with the prefix for the # of carbons in that branch, followed by the suffix –yl Naming Branched Chain Alkanes 1.​ identify the longest continuous chain of carbon atoms and name as in the previous table 2.​ identify all the branched groups (substituents) 3.​ Number the parent chain from the end so that the substituents are attached to the carbon atom with the lowest possible number.​ ​ ​ -​ If there are 2 or more groups and the numbering is a tie, the group that comes first alphabetically gets the lowest number 4.​ If the same substituent is present more than once, use the prefix (di, tri-, tetra-) to indicate this and include a number to indicate each substituents location 5.​ When writing the final name, list the substituents in the alphabetical order, ignoring any prefixes. Separate words by hyphens; separate numbers by commas. Other Branched Chains -​ Benzene rings can also be attached to a parent chain. -​ Then it is referred to as phenyl Naming Alkenes and Alkynes -​ IUPAC nomenclature rules for alkenes and alkynes are similar to alkanes. 1.​ Name the parent compound. Find the longest chain containing the double or triple bond, and name the parent compound by adding the suffix –ene or –yne to the name of the main chain. 2.​ Number the carbon atoms in the parent chain, beginning at the end nearest to the double or triple bond. If the multiple bond is an equal distance from both ends, begin numbering at the end nearer the first branch point. The number indicates which carbon the multiple bond is AFTER. (i.e. between 2 and 3 is 2-) 3.​ Assign numbers and names to the branching substituents, and list the substituents alphabetically. Use commas to separate numbers, and hyphens to separate words from numbers. 4.​ Indicate the position of the multiple-bond carbon. If more than one multiple bond is present, identify the position of each multiple bond and use the appropriate ending diene, triene, tetraene, and so forth. 5.​ Assemble the name. -​ When the carbon chain has 4 or more C atoms, number the chain to give the lowest number to the double or triple bond. Assigning Priority -​ Alkenes and alkynes are considered to have equal priority -​ In a molecule with both a double and a triple bond, whichever is closer to the end of the chain determines the direction of numbering. -​ In the case where each would have the same position number, the double bond takes the lower number. -​ In the name, “ene” comes before “yne” because of alphabetization. Naming Alkenes -​ Named the same as alkanes but using the suffix “-ene” -​ if structural isomers (different arrangement of the bonds) are possible, a number is used to indicate the carbon at which the double bond starts -​ if two double bonds are present, the ending “-diene” is used and two numbers are needed the specify their positions -​ if three double bonds are present, the ending -triene is used and three numbers are needed the specify their positions Naming Alkynes -​ Named the same as alkenes but using the suffix “-yne” Aromatic Hydrocarbons -​ Benzene is a good solvent but extremely “carcinogenic” -​ Examples of Aromatic Hydrocarbons: -​ Toluene (methyl-benzene) -​ the next simplest aromatic compound is also a good solvent and comparatively less harmful -​ 3,4 - benzopyrene -​ a carcinogenic compound found in both cigarette smoke and char-broiled foods -​ Naming Aromatic Hydrocarbons -​ In naming simple aromatic hydrocarbons, the benzene ring to be the parent molecule with alkyl groups named as branches -​ When 2 or more groups are attached to a benzene ring, a numbering system is used to indicate the positioning of the groups. ortho-​ ​ meta- para- ​ Functional Groups – any atom, group of atoms/organization of bonds that determines specific properties of an organic molecule -​ Functional groups are important: -​ For nomenclature of molecules -​ Especially, to help determine the chemical and physical properties of a particular compound Alkyl Halides -​ Group of organic molecules in which halogens have been added to or substituted for hydrogen in a molecule -​ When naming organic halides, consider the halogen as a branch: -​ F ​ fluoro- -​ Cl ​ chloro- -​ Br ​ bromo- -​ I iodo- -​ CFCs are Alkyl Halides Alcohols -​ Alcohols : contain 1 or more –OH groups also known as a hydroxyl groups -​ Naming Alcohols 1.​ identify the longest carbon chain containing the hydroxyl group -> pentane 2.​ replace -e of the name with - OL  3.​ identify the location of the hydroxyl group giving it the lowest number possible -​ OH- group locations Polyalcohols -​ Contain more than 1 –OH (hydroxyl) group -​ When a compound has 2 –OH groups, use the suffix “-diol” -​ When a compound has 3 –OH groups, use the suffix “-triol” -​ Make sure that you do not drop the “e” from the parent chain ex. propane-1,2-diol Cyclic Alcohols -​ Are cyclic compounds with –OH groups attached -​ If there is one -OH gr present, it is not numbered since the lowest possible number will always be on C1 -​ Usually are large molecules known mostly by their common names -​ For ex. menthol and cholesterol Aromatic Alcohols -​ Are aromatic compounds with –OH groups attached -​ Ex. hydroxybenzene or phenol Ethers -​ An organic compounds with 2 alkyl groups attached to an oxygen atom -​ the alkyl groups can be the same or different -​ Cannot form hydrogen bonds -​ The polar C-O bond and their shape make them more polar than hydrocarbons -​ The boiling point of ethers are slightly higher than hydrocarbons -​ The boiling point of ethers is lower than alcohols -​ Mix with both polar and non-polar substances, so these are great solvents -​ Naming Ethers  1.​ IUPAC a.​ identify the 2 alkyl groups b.​ Write the name of the shorter alkyl group, then the suffix - oxy -​ c.​ After this put the name of the longer alkyl group but as if it were an alkane. 2.​ Non-IUPAC a.​ Identify the alkyl groups: b.​ Place the names of the substituents in alphabetical order then a space then the word ether. Aldehydes and Ketones -​ Both these groups contain the carbonyl functional group -​ In aldehydes -​ the carbon atom containing the carbonyl group is bonded to a hydrogen atom and is found at the end of the parent chain -​ Naming 1.​ identify the longest carbon chain containing the carbonyl group and assign the base name. Changing the ending to -al: butane => butanal 2.​ number the chain starting at the end with the carbonyl group 3.​ identify any alkyl groups using the IUPAC rules: 3-methylbutanal -​ In ketones -​ the carbon atom of the carbonyl group is 'sandwiched' between two alkyl groups and is found within the parent chain -​ Naming 1.​ identify the longest carbon chain containing the carbonyl group and assign the base name changing the ending to ­one: pentane => pentanone 2.​ number the chain from the end that will give the carbonyl group the lowest number and indicate its location: pentan-2-one 3.​ Identify any alkyl groups using the IUPAC rules: 4-methylpentan-2-one 4.​ Carboxylic acids -​ Contain the carboxyl group : R-COOH -​ The carboxyl group is always at the end of the chain -​ Are generally weak acids found in : -​ Citrus fruits, crab apples and vinegars -​ Lactic acid, the compound that builds up in your muscles that makes your muscles hurt after a good workout -​ The most common carboxylic acid is ethanoic acid or acetic acid. Vinegar sold in stores is usually 5% acetic acid. -​ Another common acid is aspirin, the brand name for acetylsalicylic acid(ASA)  -​ Naming 1.​ Identify the longest carbon chain and changing the base name ending to –oic acid: pentane -> pentanoic acid (don’t forget to drop the “e”) 2.​ If there are any branches, number these branches from the carboxyl group. pentanoic acid -> 2,3-dimethylpentanoic acid -​ *Remember: you don’t have to number the carbon where carboxyl group is since it will only be at the end of the chain* Formic Acid -​ Produced by venomous/poisonous insects Esters -​ can occur naturally in many plants like fruits and flowers -​ are responsible for and can mimic pleasant odours -​ Ex. bananas, peaches, wintergreen -​ Synthetically made esters are often added as flavourings to processed foods and as scents to cosmetics -​ Structures -​ Esters are produced from the reaction of carboxylic acids and alcohols often called esterification -​ Esters contain two alkyl groups separated by the ester linkage -​ Naming 1.​ Identify the parent carboxylic acid (containing the – COO – ) and change the ending to -oate​​ 2 carbons (ethane)-> ethanoate 2.​ Identify the alkyl group that is attached to the opposite side of the oxygen ester linkage -​ This was the part that originally comes from an alcohol group during the reaction ​ 5 carbons (pentane) 3.​ Change the ending to the alkyl to –yl and indicated where on the alkyl group the oxygen is​ 5 carbons (pentane) -> 1-pentyl -​ Final Name – 1-pentyl ethanoate (don’t forget the space) Amines -​ Amines are a group of organic molecules in which one or more of the hydrogen atoms in the ammonia molecules is substituted by an alkyl or aromatic group -​ Amines have unpleasant odors that can be fishy -​ Naming 1.​ IUPAC a.​ locate the longest chain that is bonded to the nitrogen atom b.​ name the parent chain: drop the "e" and add "amine" : c.​ number the carbons starting nearest the nitrogen atom d.​ indicate the nitrogen atom of the main chain is 3 or more carbons long: e.​ if a substituent is attached the nitrogen atom, indicate this with an N-:  -​ if there are substituents attached to the N (and not a C), they are written with an “N” and not a number 2.​ Non-IUPAC a.​ Identify the alkyl groups: b.​ Place the names of the substituents in alphabetical order before the word amine Amides -​ The amide linkage is similar to the ester linkage with nitrogen in place of the oxygen joining the radical groups -​ Carboxylic acids react with ammonia or primary or secondary amines to produce amides -​ Amides are important in biological systems -​ Basic building block of proteins -​ Amide linkage is sometimes referred to the peptide bonds -​ Naming 1.​ Name the longest C-chain, add “amide” to the end 2.​ The C with the double-bonded O is C-1. 3.​ Name and number the side groups in the parent chain. 4.​ For secondary and tertiary amides, determine the name of the side groups and precede the name with “N”. 5.​ Attach all the side groups to the parent chain in alphabetical order Priority List ​ Types of Reactions -​ during chemical reactions, existing bonds are broken and new ones are formed 1.​ Addition -​ occur around double or triple bonds only -​ involves a double bond becoming a single bond by adding 2 atoms -​ can also involve a triple bond becoming a double bond -​ Types -​ Hydrogenation (H2) -​ similar to halogenation but each carbon receives a H -​ Halogenation (Cl2, Br2, I2) -​ each carbon in the double bond receives a halogen atom -​ Hydrohalogenation (HBr, HCl) -​ HBr or HCl is added across a double bond -​ Markinov’s rule – “the rich get richer” -​ Major Vs. Minor Product (the product formed by following Markovnikov’s rule  is the MAJOR PRODUCT and will be most commonly formed, however, the other product will also form, but is less common: MINOR PRODUCT) -​ Hydration (H2O) -​ adding water to C-atoms in a double bond -​ also follows Markovnikov’s rule 2.​ elimination -​ 2 atoms are removed from an organic molecule and a double bond forms between the two carbon atoms where the atoms were removed -​ The reverse of an addition reaction -​ KOH speeds up the reaction -​ alkanes can undergo elimination reactions to produce alkenes when they are heated in the presence of a strong base or strong acid -​ Zaitsev’s Rule -​ hydrocarbon chains become more stable when carbons are attached to more carbon atoms (instead of other atoms like H) -​ therefore, when removing a H-atom, remove it from the carbon that is already attached to more carbons -​ in other words: take the hydrogen away from the carbon that already has less hydrogens attached to it -​ “The poor get poorer” 3.​ Substitution -​ an atom or side chain is replaced by a different atom or side chain -​ Look for: two compounds react to form two different compounds -​ Look for: same number of bonds between the carbon atoms in the products as in the reactants (hv - energy from a photon) -​ Alcohols and haloalkanes usually undergo substitution reactions -​ When an alcohol reacts with an acid that contains a halogen, halogen atom is substituted for the hydroxyl group of the alcohol -​ A haloalkane can undergo a substitution reaction with a hydroxide ion to produce an alcohol. -​ W/ Benzene Rings -​ Like hydrocarbons, aromatic hydrocarbons are fairly stable  -​ They will undergo substitution reactions with chlorine and bromine in the presence of a catalyst. -​ In an addition reaction, there is only one product -​ In a substitution reaction, there are 2 reactants and 2 products and one of the atoms in the reactant has been switched with a H in the HC chain. 4.​ condensation – to condense -​ Reactions in which two large molecules combine and form one larger molecule and water is a product -​ Think: condensation outside on your window leaves drops of water -​ The formation of an amide is a condensation reaction -​ Carboxylic acid + amine are the large molecules -​ Amide + water are the products 5.​ Esterification -​ A special type of condensation reaction -​ A carboxylic acid reacts with an alcohol to form an ester and a molecule of water -​ Very useful reaction, many artificial scents and flavorings are made using this reaction 6.​ Hydrolysis -​ The reverse of a condensation reaction -​ Hydrolysis means “to break apart using water” -​ Hydroxyl group in a water molecule is added to one side of a bond such as an ester bond -​ The hydrogen atom of the water molecule is added to the other side of the bond -​ The bond is then broken 7.​ Oxidation -​ A reaction in which a carbon atom forms more bonds to oxygen atoms -​ Or fewer bonds to hydrogen atoms -​ Oxidation reactions that involve the formation of a double bond may also be classified as an elimination reaction -​ **To identify an oxidation, count and compare the number of C-H and C-O bonds in both the reactant and the product. 8.​ reduction -​ A reaction in which a carbon atom forms fewer bonds to oxygen, or more bonds to hydrogen -​ Often, a C=O bond or C=C bond is reduced to a single bond by reduction, therefore, can also be classified as an addition reaction  -​ Aldehydes, ketones, and carboxylic acids all become alcohols when reduced -​ Reducing agents: a substance that causes another substance to be reduced ex: -​ LiAlH4 (lithium aluminum hydride) -​ H2/Pt (hydrogen gas over a platinum catalyst) -​ The symbol [H] is used to symbolize a reducing agent -​ **To identify a reduction, count and compare the number of C-H and C-O bonds in the reactant and product 9.​ Combustion – A type of reaction in which a hydrocarbon reacts with oxygen to produce water and carbon dioxide -​ Complete -​ products are water and carbon dioxide -​ incomplete -​ products are water, carbon dioxide, carbon (soot), carbon monoxide -​ Incomplete combustion occurs when there is not enough oxygen present to fully react with a hydrocarbon