Chapter 3A & 4A

1. Chapter 3A & 4A Stoichiometry and Reactions

2. Chemical Reactions Chemical equation – a representation of a chemical reaction. Reactants – the starting chemicals. Products – the ending chemicals. CH4 + O2 → CO2 + H2O Reactant Product Conservation of matter – in a chemical reaction, atoms are neither created nor destroyed.

3. Physical States The state of the compound can be stated in the reaction equation using the following symbols. • Solid (s) • Liquid (l) • Gas (g) • Dissolved in water (aq)

4. Balancing Equations C(s) + O2(g) → CO2(g) Reactants Products All atoms present in the reactants must be accounted for among the products C-1 → C-1 O-2 O-2

5. Balancing Rules • Read the description and write the formulas for the reactants and products. • Write the unbalanced formula. • Balance the equation starting with the most complicated molecule. Remember to not change the formula of the chemicals. • Check to confirm balance.

6. Balancing Equaitons Methane reacts with oxygen to produce carbon dioxide and water. CH4(g) + O2(g) → CO2(g) + H2O(l) C-1 C-1 H-4 H-2 O-2 O-3

7. Balancing equations • K + H2O → H2 + KOH • NH3 + O2 → NO + H20 • SiO2 + HF → SiF4 + H2O • CaC2 + H2O → Ca(OH)2 + C2H2

8. Balancing Equations • Hydrogen gas combines with oxygen gas to produce water. • Sodium metal combines with water to produce sodium hydroxide liquid and hydrogen gas. • Aluminum solid combines with chlorine gas to produce aluminum chloride solid. • Hydrochloric acid (aqueous solution) combines with calcium carbonate solid to form calcium chloride, carbon dioxide gas and water.

9. Reaction Types • Composition (synthesis) • Decomposition • Single Replacement • Double replacement (precipitation) • Acid/Base

10. Reaction Types Composition reaction – a reaction that combines substances to form a single substance. A + B → C 2H2 + O2 → 2H2O C + O2 → CO2 2Na + Cl2 → 2NaCl

11. Reaction Types Decomposition reaction – a reaction in which a compound is broken down into two or more simpler substances. AB → A + B 2H2O → 2H2 + O2 2NaCl → 2Na + 2Cl CaCO3 → CaO + CO2

12. Reaction Types Single replacement reaction – a reaction in which one element takes the place of another element A + BC → B + AC Cu + 2AgNO3 → Cu(NO3)2 + Ag 2K + 2H2O → H2 + 2KOH

13. Reaction Types Double replacement reaction – a reaction in which two different compounds exchange positive ions and form two new compounds. AB + CD → AD + CB Pb(NO3)2 + KI → PbI2 + 2KNO3 CaCO3 + 2HCl → CaCl2 + H2CO3

14. Reaction Types • 2K + Cl2 → 2KCl • Fe2O3 + 2Al → Al2O3 + 2Fe • 2Mg + O2 → 2MgO • HNO3 + NaOH → H2O + NaNO3 • KBr + AgNO3 → AgBr + KNO3 • PbO2 → Pb + O2

15. Single Replacement Reaction If the free metal is higher on the metal activity table than a metal in solution then the free metal will exchange places with the metal in solution. Ex. Zn + HBr → NiCl2 + K → Cu + HCl →

16. Double Replacement Reaction Prediction Metathesis reactions are driven by three "driving forces." 1) A product is a precipitate 2) A product is a gas 3) A product is a weak or non-electrolyte We reduced these to; 1) A product is an ionic precipitate 2) A product is a molecular, non-strong acid

17. Double Replacement The solubility table is used to determine if an ionic precipitate is formed from the reactants. Ex. K2S + Fe(NO3)2 Na2SO4 + BaCl2 HCl + MgS KI + NH4NO3

18. Double Replacement Reaction • Molecular equation – equation in which the complete formulas of all reactants and products are given. • Complete ionic equation – equation in which all substances that are strong electrolytes are represented as ions. • Spectator ions – ions which do not participate in the reaction. • Net ionic equation – equation in which the spectator ions are not represented only the components involved in the reaction.

19. Double Replacement Reaction For each of the following reactions, write the balanced molecular equation, the balanced complete ionic equation, and the balanced net ionic equation. • Silver nitrate and sodium chromate AgNO3 + NaCrO4 • Nickel(II) nitrate and potassium carbonate Ni(NO3)2 + K2CO3

20. Electrolytes • Strong Electrolytes – ionic compounds that total dissociate in water and conduct electricity. • Weak electrolytes – molecular acids that can totally or partially dissociate in water. • Non electrolytes – molecular compounds that do not dissociate in water so do not conduct electricity.

21. Acid/Base Reactions Acids – substances that ionize in aqueous solutions to form hydrogen ions. H2SO4(aq) → H+(aq) + HSO4-(aq) Strong acids – strong electrolytes Bases – substances that ionize in aqueous solutions to form hydroxide ions. NaOH(s) → Na+(aq) + OH-(aq)

22. Strong Acids/Bases Strong Acids • HCl –hydrochloric acid • HBr – hydrobromic acid • HI – hydroiodic acid • HClO3 – chloric acid • HClO4 – prechloric acid • HNO3 – nitric acid • H2SO4 – sulfuric acid Strong Bases • Group 1A metal hydroxides (LiOH, NaOH, KOH, RbOH, CsOH) • Group 2a metal hydroxides (Ca(OH)2, Sr(OH)2, Ba(OH)2)

23. Acid/Base Reactions Acid/Base reactions result in the formation of a salt and water Ex. Write the balanced molecular, balanced complete ionic, and balanced net ionic equations for the reaction of aqueous hydrobromic acid and aqueous sodium hydroxide.

24. Acid/Base Reaction Write the balanced molecular, balanced complete ionic, and balanced net ionic equations for the reaction of aqueous sulfuric acid and aqueous potassium hydroxide.