Stoichiometry & Equations

1. Stoichiometry & Equations

2. What is stoichiometry? • The study of quantitative relationships in a balanced chemical equation. • Equations represent reactions.

3. Physical vs. Chemical Change • Physical Change = • Chemical Change = Change in form or appearance. Dissolving and phase changes are 2 kinds of physical changes. Memory Jogger Change in identity. Change in formula representing substance.

4. Chemical Change 2 H2O  2 H2 + O2 • How do I know this is a chemical change? • The formulas DO NOT match.

5. Physical Change • H2O (s)  H2O(l) • NaCl(s)  NaCl(l) • The formulas match! The identity of the substance does not change. Only the form or appearance is altered.

6. Names of some common reactions • Corrosion or rusting • Burning • Rotting • Fermentation • Growing • Most cooking • Generating electricity • Running car engines • Decomposition Memory Jogger

7. Evidence of a Chemical Reaction • Temperature change • Increase or Decrease • Emission of Light Energy • Change in identifying property: Color, mp, bp, density, Hf, Hv, C … • Formation of a Gas • Bubbling, odor • Formation of a Solid • Precipitation Ammonium Dichromate Volcano

8. Chemical Equations A + B  C + D Left Side = Reactants or starting materials Right Side = Products or ending materials “” is read as produces or yields How do we show the physical state of the reactants & products? Memory Jogger (s), (l), (g), (aq)

9. Law of Conservation of Matter • Matter is neither created nor destroyed in a chemical reaction. • Mass reactants = Mass products • Chemical bonds in reactants may break. • New bonds form to produce products • Number of atoms of each element is “constant.” • # of atoms of each element is the same on both sides of the equation. Memory Jogger

10. Coefficients in Chemical Equations • Coefficients are the numbers in front of the formulas. • They apply to everything in the given formula.

11. Coefficients in Chemical Equations • Coefficients connect the microscopic world with the macroscopic world. • Microscopic: Coefficients represent numbers of individual atoms or molecules • Macroscopic: Coefficients give mole ratios! • Moles  connected to mass through formula mass

12. Writing Chemical Equations • Begin with word equation – describe what happens • Next go to skeleton equation – replace names of substances with chemical formulas • Balancethe skeleton equation. The balanced equation must demonstrate the law of conservation of mass.

13. Equation Balancing • Surveythe skeleton equation • Count up the # of each type of atom on reactant side • Count up the # of each type of atom on product side • Use COEFFICIENTS to adjust the numbers of each type of atom. • Check atom counts after every change. • NEVER CHANGE SUBSCRIPTS IN FORMULAS • That’s like changing the identity of the reactant or product.

14. Balanced Equations • Coefficients should be in their lowest possible ratios • Double check your work. Count up the atoms of each element on both sides – they should be equal • Get in the habit of using a table to keep track of elements

15. Example 1 Fe + O2 Fe2O3 • Left Side: 1 Fe and 2 O • Right Side: 2 Fe and 3 O. • Hint: LCM of 2 and 3 = 6. Get both O to 6 Fe + 3 O2 2 Fe2O3 • O’s are balanced. Fix Fe. 4 Fe + 3 O2 2 Fe2O3 • 4 Fe and 6 O each side.

16. Example 2 Na + H2O  NaOH + H2 • Notice: even # H’s on left, odd # on right. • Have to make # H’s on left even. Put a 2 in front of NaOH Na + H2O  2 NaOH + H2 • Put a 2 in front of Na to balance Na’s 2 Na + H2O  2 NaOH + H2 • 2 in front of H2O to balance O and H 2 Na + 2 H2O  2 NaOH + H2

17. Example 3 AgNO3 + MgCl2 Mg(NO3)2 + AgCl • Hint: Treat the NO3- as a unit, because it appears unchanged on both sides of eq. Balance nitrates with a 2. 2 AgNO3 + MgCl2 Mg(NO3)2 + AgCl • Balance Ag’s and Cl with a 2. 2 AgNO3 + MgCl2 Mg(NO3)2 + 2 AgCl

18. Types of Reactions • Synthesis • Decomposition • Single Replacement • Double Replacement • Combustion

19. Synthesis Format: A + B  C Identifying feature: 1 product only Note: A and B may be elements or compounds. C is a compound Ex: 2 Fe(s) + 3 Cl2(g)  2 FeCl3(s) 2 Na(s) + Cl2(g)  2 NaCl(s) CaO(s) + H2O(l)  Ca(OH)2(s)

20. Synthesis Reactions • Burning of coal – sulfur impurities are oxidized to form SO2 (Also combustion) • RainS1O2.mov Al + Br2 Fe + Cl2

21. Decomposition Format: AB  A + B Identifying feature: 1 reactant only Note: A & B may be elements or compounds Ex: 2 NaN3(s)  2 Na(s) + 3 N2(g) NH4NO3(s)  N2O(g) + 2 H2O(g)

22. Single Replacement Format: A + BX  AX + B Identifying Feature: Element + Compound  New Element + New Compound Ex: 2 Li(s) + 2 H2O(l)  2 LiOH(aq) + H2(g) Cu(s) + 2AgNO3(aq)  2Ag(s) + Cu(NO3)2(aq) Cu + 2AgNO3 Zn + 2AgNO3

23. Double Replacement Format: AX + BY  AY + BX Identifying feature: 2 compounds yield 2 new compounds Ex: Ca(OH)2(aq) + 2HCl(aq)  CaCl2(aq) + 2 H2O(l) 2 NaOH(aq) + CuCl2(aq)  2 NaCl(aq) + Cu(OH)2(s)

24. Double Replacement Rxns • Cu(NO3)2 + NaOH • NaI + KNO3 No rxn

25. Combustion or Reaction with O2 Format: A + O2 B ( + C + …) Identifying feature: One reactant is O2. Note: “A” can be an element or a compound. Usually more than 1 product. Ex: CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g) 2 H2(g) + O2(g)  2 H2O(g) C(s) + O2(g)  CO2(g) For the Regents, synthesis trumps combustion!

26. Identify the Rxn Type • CaO + CO2 CaCO3 • 2 H2O  2 H2 + O2 • NaOH + HCl  NaCl + H2O • Zn + 2 HCl  ZnCl2 + H2 • 2 Mg + O2  2 MgO • Mg + H2SO4  MgSO4 + H2 • 2 KClO3  2 KCl + 3 O2 • AgNO3 + NaCl  NaNO3 + AgCl Synthesis Decomposition Dbl. Rep. Single Rep. Synthesis Single Rep. Decomposition Dbl. Rep.

27. Stoichiometry!!! • Use the balanced chemical equation to predict the amount of a given reactant or product under certain conditions.