Chemicals Equations and Balancing Chemical Reactions.pptx

Balancing Reactions

1. Balancing Chemical Reactions

2. Conservation of Mass During a chemical reaction, atoms are neither created nor destroyed (Conservation of Mass). Hydrogen and oxygen gas react to form water: H 2 (g) + O 2 (g)  H 2 O (l)

3. Conservation of Mass H 2 (g) + O 2 (g)  H 2 O (l) What is wrong with this equation above? Doesn’t it appear that one oxygen atom “went missing”? According to conservation of mass, the proper way to write this reaction is: 2 H 2 (g) + 1 O 2 (g)  2 H 2 O (l) The red coefficients represent the # of molecules (or the # of moles) of each reactant or product.

4. Not All Properties are Conserved During Chemical Reactions! CONSERVED NOT CONSERVED Mass Types of atoms Number of each atom Color Physical state (solid, liquid, gas) Volume Number of moles of reactants/products

5. Counting Atoms SnO 2 + 2H 2 → Sn + 2H 2 O SUBSCRIPT COEFFICIENT

6. Rules for Counting Atoms Coefficients propagate to the right through the entire compound, whether or not parentheses are present. 2) Subscripts affect only the element to the left of the subscript, unless… 3) If a subscript occurs to the right of a parentheses, the subscript propagates to the left through the parentheses. 4) When a coefficient and subscript “meet”, you must multiply the two.

7. Examples of Counting Atoms SnO 2 + 2H 2 → Sn + 2H 2 O 2 C 4 H 10 + 13 O 2 → 8 CO 2 + 10 H 2 O Cu + 2AgNO 3 → Cu(NO 3 ) 2 + 2Ag 3Pb(NO 3 ) 2 + 2AlCl 3 → 3PbCl 2 + 2Al(NO 3 ) 3

8. Classwork Complete “The Count” worksheet on counting atoms in chemical reactions.

9. Warm-Up 2Ca 3 (PO 4 ) 2 + 6 SiO 2 + 10C  6 CaSiO 3 + P 4 +10CO Atom # Atoms on Left Side # Atoms on Right Side Ca P O Si C

10. Rules for Balancing Chemical Reactions __H 2 + __ O 2  __H 2 O Balancing is about finding the right coefficients!

11. Rules for Balancing Chemical Reactions You can change the coefficients, but NEVER the subscripts! __H 2 + __ O 2  __H 2 O Off Limits!

12. Rules for Balancing Chemical Reactions 2) The coefficients must reduced to represent the lowest possible numbers. 4H 2 + 2 O 2  4H 2 O

13. Rules for Balancing Chemical Reactions 3) It is OK to use fraction coefficients, but you must get rid of them in the end (multiply through by denominator). H 2 + ½ O 2  H 2 O

14. Rules for Balancing Chemical Reactions 4) Often, it is helpful to save the following elements until the end (do other elements first): H, C, O

15. Rules for Balancing Chemical Reactions 5) Do a final balance check for each element! 2H 2 + O 2  2H 2 O

16. Practice K + Br  KBr HgO  Hg + O 2 3) Na + H 2 O  NaOH + H 2

17. Practice 4) CaO + H 2 O  Ca(OH) 2 5) Al + HCl  AlCl 3 + H 2

18. Energy Changes Accompanying Chemical Reactions All chemical reactions involve a net release or absorption of energy. Therefore, heat energy moves between the chemical system and the surroundings. This exchange of heat can be monitored by keeping track of changes in temperature of the surroundings (calorimetry). Remember,  q = mc p T where  q = change in heat (in Joules) m = mass of H 2 O (in grams) c p = specific heat capacity of H 2 O (J/g ◦C ) T = change in temperature of H 2 O (in ◦C)

19. Where does the energy come from during a chemical reaction? During chemical reactions, bonds are broken and new bonds are formed. The heat energy that moves between the system and surroundings during chemical reactions is basically the energy that is used to break bonds and the energy that is released when bonds form. (i.e. bond energy) The energy change that accompanies any chemical reaction is called the enthalpy (heat) of reaction or H rxn . H rxn = H final – H initial H simply means that the energy changes during chemical reactions are generally measured at “standard state” conditions of 298 K (25 ◦ C) and 1 atm pressure. It is important to note that absolute amounts of energy within a chemical system cannot be measured. We can only measure changes in energy within a chemical system. Hence we use the “” sign.

20. Exothermic Reactions A chemical reaction is exothermic if energy is given off by the system to the surroundings (the energy exits): Reactants  Products + Energy Released The temperature of the surroundings (including the temperature probe) increases during exothermic reactions because the system releases energy. The H rxn is negative because H final is less than H initial . In other words, the system lost energy. (sign goes with the system) The majority of chemical reactions are exothermic because nature favors a low chemical potential energy. System Surroundings reactants products Chemical Potential Energy (H) Reaction progress H rxn is (-)

21. Example: An Exothermic Reaction The “Smashing” Thermite Reaction: 2Al(s) + Fe 2 O 3 (s)  2Fe (s) + Al 2 O 3 (s) Reaction Progress Chemical Potential Energy (H)

22. Endothermic Reactions A chemical reaction is endothermic if energy is absorbed by the system from the surroundings (the energy enters): Reactants + Energy Absorbed  Products The temperature of the surroundings (including the temperature probe) decreases during endothermic reactions because the system absorbs energy. The H rxn is positive because H final is more than H initial . In other words, the system gained energy. (sign goes with the system) Endothermic chemical reactions are generally unfavorable but may occur only if they are accompanied by an increase in entropy or disorder of the system (due to more particles formed, liquids/gases formed, mixtures formed, volume of gas increases). System Surroundings reactants products Chemical Potential Energy (H) Reaction progress H rxn is (+)

23. Example: An Endothermic Reaction Ba(OH) 2 8H 2 O (s) + 2NH 4 (NO 3 ) (s)  Ba(NO 3 ) 2 (aq) + 2NH 3 (g) + 10 H 2 O (l) Reaction Progress Chemical Potential Energy (H)

24. Do you have to actually perform and observe a chemical reaction to know if it is exothermic or endothermic? No – you can calculate H rxn from data that has already been measured and tabulated by thermo-chemists (see handout). H f = standard heat of formation for a compound (in kJ/mol). It is determined by forming the compound from its elements in their stable forms at conditions of 298K and 1 atm of pressure inside of a calorimeter. For most compounds, H f is negative because bond formation is exothermic! H f of an element is always 0 kJ/mol by def.

25. H rxn = nH f (products) - nH f (reactants) Not as hard as it looks  Basically, you just 1) multiply the coefficient of each product times its standard heat of formation and add together for all products 2) multiply the coefficient of each reactant times its standard heat of formation and add together for all reactants 3) take the difference of 1 and 2 (always products - reactants ) 4) If the difference is (-) the reaction is exothermic; if the difference is (+) the reaction is endothermic.

26. Try this… Calculate the H rxn for the thermite reaction using tabulated data (see handout): 2Al (s) + Fe 2 O 3 (s)  2Fe (s) + Al 2 O 3 (s) H rxn = nH f (products) - nH f (reactants)

27. Try this… Calculate the H rxn for this reaction based on tabulated data: Ba(OH) 2 8H 2 O (s) + 2NH 4 (NO 3 ) (s)  Ba(NO 3 ) 2 (aq) + 2NH 3 (g) + 10 H 2 O (l) H rxn = nH f (products) - nH f (reactants) Compound H f (kcal/mol) NH 4 (NO 3 ) (s) -87.73 Ba(OH) 2 8H 2 O (s) -798.8 Ba(NO 3 ) 2 (aq) -227.62 NH 3 (g) -11.02 H 2 O (l) -68.32 1 kcal = 4.184 kJ

28. Summarizing H rxn If H rxn is ( - ) the reaction is exothermic and the bonds formed are stronger and more stable than the bonds broken . If H rxn is ( + ) the reaction is endothermic and the bonds formed are weaker and less stable than the bonds broken . However, if the entropy of the system has increased to sufficiently to counteract this increase in enthalpy, then the reaction can still occur.

29. Bond Enthalpies Another way to determining an enthalpy change ( H rxn ) for a chemical reaction is to compute the difference in bond enthalpies between reactants and products The energy to required to break a covalent bond in the gaseous phase is called a bond enthalpy (bond dissociation energy). Bond enthalpy tables give the average energy to break a chemical bond. Actually there are slight variations depending on the environment in which the chemical bond is located

30. Bond Enthalpy Table The average bond enthalpies for several types of chemical bonds are shown in the table below:

31. Bond Enthalpies Bond enthalpies can be used to calculate the enthalpy change ( H rxn ) for a chemical reaction. Energy is required to break chemical bonds . Therefore when a chemical bond is broken its enthalpy change carries a positive sign . Energy is released when chemical bonds form . When a chemical bond is formed its enthalpy change is expressed as a negative value. By combining the enthalpy required and the enthalpy released for the breaking and forming chemical bonds, one can calculate the overall enthalpy change for a chemical reaction.

32. Bond Enthalpy Calculations Example : Calculate the enthalpy change ( H rxn ) for the reaction N 2 + 3 H 2  2 NH 3 Bonds broken (energy in) N≡N: = 945 H-H: 3(435) = 1305 Total = 2250 kJ/mol Bonds formed (energy out) 2x3 = 6 N-H: 6 (390) = - 2340 kJ/mol H rxn = [energy used for breaking bonds] + [energy released in forming bonds] Net enthalpy change ( H rxn ) = + 2250 + (-2340) = - 90 kJ/mol (  exothermic reaction) H - H H - H H - H You may have to draw a Lewis Structure to know what type of bonds are present!

33. Another Way to Think About It Chemical Potential Energy (H) of System Start + 2250 kJ/mol (energy in when bonds break) -2340 kJ/mol (energy out when bonds form) H rxn = -90 kJ/mol (net) released by the system to the surroundings