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Chemical Reactions

Chemical Reactions. Dr. Molarity Hazlett Mandeville High School. What is a Chemical Reaction?. It is the process by which one substance changes into another The Conservation of Mass Law is followed All reactions ( rxns ) have some mechanism by which they occur

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Chemical Reactions

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  1. Chemical Reactions Dr. Molarity Hazlett Mandeville High School

  2. What is a Chemical Reaction? • It is the process by which one substance changes into another • The Conservation of Mass Law is followed • All reactions (rxns) have some mechanism by which they occur • There is a limit and process by which the change(s) occur

  3. When there is a chemical change: • There may be a color change meaning the new substances will reflect different wavelengths (colors) of light • There may be a change in energy – heat or cold • A gas or vapor may be produced • A precipitate may be formed which is an insoluble solid forming out of the solution • For the chemical reaction to take place the reactants must be able to form chemical bonds

  4. I. Chemical Equations • The chemical equation is the rxn formula • Reactants  Products • Each component will have a phase indicator: • (g) meaning it is in its gaseous phase (not just gassy) • (l) meaning it is in its liquid phase • (s) in its solid phase • And (aq) meaning the substance is in a solution of water, aq meaning aqueous

  5. Must remember which elements are normally diatomic (N, O, F, Cl, Br, I, and H) • All molecules in an equation must be balanced first!! • Remember the criss-cross rule!! • You may not adjust any subscripts from the original formula • You may add and adjust, as you will see, the coefficients in front of each item in the equation

  6. Example • H2 + O2 H2O • This is the skeleton equation • According to the Law of Conservation of Matter, both sides of the arrow must have the SAME number of atoms for each and every element – NO EXCEPTIONS • The  can be treated like an = sign • In reality, it indicates that some sort of process occurred to cause the reaction • So. . . . .

  7. To balance this simple equation: • We ARE NOT ALLOWED TO CHANGE SUBSCRIPTS • We CAN ADJUST COEFFICIENTS ONLY • The subscripts are the numbers after and below each element’s symbol • The coefficients are number in front of a unit (atoms or molecules) and tell how many units there are • The coefficients are multiplied out to each and every unit’s atom they are in front of • So. . . . .

  8. H2 + O2 H2O • There are 2 H and 2 O on the reactant side of the equation (the left side) • There are 2 H and only 1 O on the product side (the right side) • Each side must balance • You may add, adjust, finagle, cram, etc. any coefficient in front of any and/or all units to get the equation to balance • Therefore: 2H2 + O2 2 H2O

  9. 2H2 + O2 2H2O • Now this is balanced! • It means it takes 2 hydrogen molecules and one oxygen molecule to form 2 water molecules • It also means: • 2 mols H2 + 1 mol O2 yields 2 mols H2O • (2 x 2.02 g H) + 32 g O yields (2 x 18.02 g H2O) or 36.04 g of H and O yields 36.04 g H2O

  10. II. Balancing Equations -Let the Whining Commence! • The process of balancing equations is not that hard, although some equations can be tedious and frustrating! • Follow these steps, and your nervous breakdown can be postponed for a few months until you see the final.

  11. Balancing Equations Steps • First identify all the reactants and products in the equations • Remember – subscripts indicate how many of each element’s atoms are present – with 1 being understood • Remember to multiply out all subscripts that are outside a unit in parentheses! • YOU CAN’T CHANGE SUBSCRIPTS • COEFFICIENTS HAVE TO GO IN FRONT OF A UNIT

  12. Let’s take the unbalanced equation of: KClO3 KCl + O2 • List the elements and how many for both sides of the arrow K 1  K 1 Cl 1 Cl 1 O 3 O 2 • Obviously, everything is fine except for oxygen • This is where we have to adjust

  13. We can only use coefficients • So we try to multiply each Oxygen by a number to get them to equal out • These multipliers become coefficients K 1  K 1 Cl 1 Cl 1 O 3 x 2 = 6 O 2 x 3 = 6 • So the new equation is: 2 KClO3  KCl + 3 O2 • This changes the Number of K and Cl now • You have to readjust again. . . . . .

  14. 2 KClO3 KCl + 3 O2 • Now we have: K 2  K 1 Cl 2 Cl 1 O 6 O 6 • Multiply the product KCl by a coefficient of 2 and it balances • Let’s check: 2 KClO3  2KCl + 3O2 K 2  K 2 Cl 2 Cl 2 O 6 O 6 • It’s Balanced! Finally.

  15. Another Example: C2H6 + O2 CO2 + H2O • List the atoms and numbers: C 2 C 1 H 6 H 2 O 2 O 2 + 1 = 3 • Let’s go with C first by multiplying CO2 by a coefficient of 2 C2H6 + O2  2 CO2 + H2O

  16. This gives us: C 2 C 1 x 2 = 2 H 6 H 2 O 2 O 4 + 1 = 5 • Now, let’s balance H by multiplying H2O by 3 • This gives us: C2H6 + O2 2 CO2 + 3 H2O C 2 C 2 H 6 H 2 x 3 = 6 O 2 O 4 + 3 = 7 • It’s still not balanced! • Let’s try readjusting Oxygen to get it the same amount

  17. So, if we change the reactant oxygen to 7 and the product water to 6, we get: C2H6 + 7 O2 2 CO2 + 6 H2O • This also changes our product hydrogen. • Therefore, change the reactant C2H6 and the product CO2 to balance and you get: 2 C2H6 + 7O2  4 CO2 + 6 H2O C 4 C 4 H 12 H 12 O 14 O 14

  18. III. Why Reactions Occur • Molecular Collisions -This occurs especially in gaseswhere most of the volume is free space -Bimolecular collisions at STP occur on average 109 times/sec; and trimolecular is around 105 times/sec -The species’ density, volume, velocity, temperature and pressure all affect this rate

  19. According to Collision Theory, reacting molecules must collide with enough energy to form products • Particle motion is set by the level of chemical PE and KE • Chemical reactions require activation energy • This can be acquired from the environment • Or, additional energy may have to be applied to instigate the reaction

  20. Collision Rate: • Collision Rate = M v2 = x g cm2/mol sec2 K TK • KE ½ m v 2 • The collision formula:

  21. 2. Molecular Orientation • Polarization of molecule matters • This means molecules may repulse/attract each other

  22. 3. Reactant Concentration (RConc) • More concentration, more collisions/rxns • M = Molarity = mols/liter • Increase in T and/or P • This increases KE of molecules, increasing collisions • Surface Area • More area – more collisions (adsorption)

  23. 6. Activation Energy (EA) • Arrhenius (1888) hypothesized that colliding molecules need certain level of energy to break/form bonds and react K = Ae-Ea/RTwhere k is the rate constant; Ea is activation energy; R is gas constant (8.314 J/K mol); T is temp in K; and Ae is collision frequency • This minimum amount of energy is EA • Catalysts • Used to affect rate of rxn • Provides new rxn mechanism (path) for the molecules to react

  24. Adsorption – reacting molecules binding to surface of catalyst • Types of Catalysts: • Homogeneous – in same phase as reactants • Heterogeneous – in different phase from reactants

  25. IV. Predicting Products: Reaction Types • SYNTHESIS (or Direct Combination or Composition) REACTIONS • 2 + reactants join together to form a single product • Resulting compound is based on common oxidation numbers of the reactant elements • There is typically an electron transfer from the element with the lower EN to the one with the higher EN

  26. If two nonmetals involved – a covalent bond formed • If two metals – a metallic bond • If metal with a nonmetal – ionic bond • Metals react w/ Halogens to form either ionic or covalent bonds • Nonmetal oxides + H2O form ternary acids • SO4 + H2O  H2SO4 + O2 (unbalanced) • Can form hydrates • An oxide or sulfide w/ H2O • CuSO4 + H2O  CuSO4•5H2O (unbalanced)

  27. DECOMPOSITION REACTIONS • Compounds break down into components • If it is a Metallic . . . • Carbonates (CO3-2) – into metal oxide and CO2 (g) • CaCO3 CaO + CO2 • Bicarbonates – into metal oxide + H2O + CO2 (g) • Sulfates (SO4-2) – into metal oxide and SO3 • Hydroxides (OH-) – into metal oxides and H2O • Ca(OH)2 CaO + H2O • Chlorates (ClO3-) – into metal chlorides and O2 (g) • 2 KClO3 2 KCl + 3 O2 • Hydrates (H2O) – into anhydrous compound and H2O • Binary Compounds – into the base elements • 2 NaCl 2 Na + Cl2 • And oxy-acids – into H2O and nonmetal oxides • H2CO3 H2O + CO2

  28. REPLACEMENT REACTIONS • Single Replacement (Displacement Rxn) • Key Rule: Metals Replace Metals • A + BC  AC + B • Use Activity Series to determine this • If Nonmetal – a transfer of e- from more reactive to lesser one • Halogens Replace Halogens also • Metals replace H in H2O  Metal OH- + H2 (g) • Metals replace H in Acids  salt + H2(g) • Al + H2SO4  AlSO4 + H2(g) • 2 Sc(s) + 6 HCl (aq)  2 ScCl2(aq) + 3 H2(g)

  29. Double Replacement (Metathesis Reaction) • A coupling substitution rxn • FeCl3 + 3 NaOH 3 NaCl + Fe(OH)3 • Cations exchange anions • No change in oxidation numbers • Better know your ions and polyatomics • Remember the criss-cross rule • Metal Carbonate + acid  salt + water + CO2 • Metal Bicarbonate + acid  salt + water + CO2

  30. COMBUSTION • An exothermic rxn (gives off energy) • Usually find CO2 and H2O in products • O2 usually found in reactants • CH4(g) + 2 O2 CO2(g) + 2 H2O(g) + heat • 2 C4H10(g) + 13 O2(g)  8 CO2(g) + 10 H2O(g)

  31. ACID/BASE REACTIONS • This will be covered more in a later chapter • However: • An acid + base  salt + H2O • Acids lose a H+ ion and the bases lose OH- ion • These make up one of the products, water • Process is called neutralization • Proceeds at varying rates

  32. Precipitate Rxn • This reaction create a solid product

  33. Precipitation Reactions • If a reaction taking place in aqueous solution produces an insoluble product, that product will precipitate, forming a solid product at the bottom of the reaction vessel. • Precipitation often happens with double replacement reactions, and one of the most dramatic of such reactions is the reaction between lead nitrate and potassium iodide: 2KI(aq) + Pb(NO3)2(aq) -> 2KNO3(aq) + PbI(s) • Both of the reactants form colorless solutions, but if these solutions are poured together the lead iodide product is a dramatic yellow solid which precipitates out of the solution. Both of the reactants are ionic compounds and if it were not for the precipitation, the reaction would not proceed - you would just get the four ions in equilibrium. • From experience, empirical rules have been developed to help anticipate when a precipitation reaction will occur.

  34. V. Energy and Reactions • Reactions can be exothermic (give off energy) or endothermic (absorb energy) • Many chemical reactions release energy in the form of heat, light, or sound. These are exothermic reactions. • Exothermic reactions may occur spontaneously and result in higher randomness or entropy (ΔS > 0) of the system. They are denoted by a negative heat flow (heat is lost to the surroundings) and decrease in enthalpy (ΔH < 0). • In the lab, exothermic reactions produce heat or may even be explosive. • There are other chemical reactions that must absorb energy in order to proceed. These are endothermic reactions. • Endothermic reactions cannot occur spontaneously. Work must be done in order to get these reactions to occur. When endothermic reactions absorb energy, a temperature drop is measured during the reaction. • Endothermic reactions are characterized by positive heat flow (into the reaction) and an increase in enthalpy (+ΔH).

  35. Examples of Endothermic and Exothermic Processes • Photosynthesis is an example of an endothermic chemical reaction. In this process, plants use the energy from the sun to convert carbon dioxide and water into glucose and oxygen. This reaction requires 15MJ of energy (sunlight) for every kilogram of glucose that is produced: Sunlight + 6CO2(g) + H2O(l)  C6H12O6(aq) + 6O2(g)

  36. An example of an exothermic reaction is the mixture of sodium and chlorine to yield table salt. This reaction produces 411 kJ of energy for each mole of salt that is produced: 2 Na(s) + Cl2(s)  2 NaCl(s) + Energy

  37. Entropy, Enthalpy and Free Energy • This material will be explored further in the thermochemistry section – bet ya’ can’t wait, huh? • Entropy (S) – randomness or chaos of the system • If +, max. entropy; if -, min. entropy • Enthalpy (H) – amount of heat content / internal energy • H = E + PV or H = U + PV where E and U are internal energy • Gibbs Free Energy (G) – measure of useful work obtained from the system or reaction • G = H - TS • G = U + PV - TS • Where U is internal energy (joules); P is pressure; V is final volume (m3); T is temp (K); S is entropy (J/k); and H is enthalpy (joules) • -TS is energy you get from the system’s environemt by heating • PV is work to give the system at final V and constant P

  38. In Summary:

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