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Describing Chemical Change: Equations and Types of Reactions

Learn how to write balanced chemical equations and identify different types of chemical reactions. Explore synthesis, decomposition, single-replacement, double-replacement, and combustion reactions.

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Describing Chemical Change: Equations and Types of Reactions

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  1. Chemical Reactions Chemistry

  2. Section 8.1 Describing Chemical Change • Objectives: • Write equations describing chemical reactions using appropriate symbols • Write balanced chemical equations when given the names or formulas of the reactants and products in a chemical reaction

  3. Chemical Reaction • The process by which one or more substances are changed into one or more different substances. • Reactants → Products

  4. Chemical Equation • Describes chemical reactions • Uses symbols and formulas to identify relative amounts of reactants and products.

  5. Indications of a Chemical Reaction • Evolution of heat and light. • Production of a gas. • Formation of a precipitate. • A solid that is formed from solutions in a chemical reaction. • Color change.

  6. Characteristics of Chemical Equations • The equation must represent known facts. • All reactants and products must be identified. • The equation must contain the correct formulas for the reactants and products. • Sodium chloride = NaCl • Water = H2O • Carbon = C

  7. Characteristics of Chemical Equations • The law of conservation of mass must be satisfied. • Atoms are neither created nor destroyed. • Coefficients are used to balance the number of atoms.

  8. Word and Formula Equations • Word Equation – an equation in which the reactants and the products in a chemical reaction are represented by words. • Methane + oxygen → carbon dioxide + water • Has only descriptive (qualitative) meaning. • Formula Equation – represents the reactants and products of a chemical reaction by their symbols or formulas. • CH4 + O2 → CO2 + H2O • Also a qualitative statement.

  9. Additional Symbols Used in Chemical Equations • ↔ (reversible reaction) • ↑ (gaseous product) • ↓ (precipitate) • → (yields) • (s) = solid • (l) = liquid • (g) = gas • (aq) = aqueous

  10. Catalyst – A substance that changes the rate of a chemical reaction, but can be recovered unchanged. • Reversible Reaction – a chemical reaction in which the products re-form from the original reactants.

  11. Significance of a Chemical Equation • 1. The coefficients of a chemical reaction indicate relative, not absolute, amounts of reactants and products. • 2. The relative masses of the reactants and products of a chemical reaction can be determined from the reaction’s coefficients. • 3. The reverse reaction for a chemical equation has the same relative amounts of substances as the forward reaction.

  12. Balancing Equations • The same number of atoms on each side of the yield symbol. • Only the coefficients can change to balance. CH4 + O2 → CO2 + H2O CH4 + 2 O2 → CO2 + 2 H2O

  13. Balancing Chemical Equations • Balance the different types of atoms one at a time. • First balance the atoms of elements that are combined and that appear only once on each side of the equation. • Balance polyatomic ions that appear on both sides of the equation as single units. • Balance H atoms and O atoms after atoms of all other elements have been balanced.

  14. Ex #1: Write a word and a balanced chemical equation for the following: hydrogen gas reacts with fluorine gas to produce hydrogen fluoride gas. • Word • hydrogen + fluorine → hydrogen fluoride • Balanced Chemical Equation • H2(g) + F2 (g) → HF (g) • H2(g) + F2 (g) → 2 HF (g)

  15. Ex #2: Write a balanced chemical equation for the following: When solid copper reacts with aqueous silver nitrate, the products are aqueous copper (II) nitrate and solid silver. Cu(s) + AgNO3 (aq)→ Cu(NO3)2 (aq) + Ag(s) Cu(s) + 2AgNO3 (aq)→ Cu(NO3)2 (aq) + 2Ag(s)

  16. Ex #3: Balance the following equations: • ___AgNO3 + ___H2S → ___Ag2S + ___HNO3 • ___MnO2 + ___HCl → ___MnCl2 + ___H2O + ___ Cl2 • ___Zn(OH)2 + ___H3PO4→ ___Zn3(PO4)2 + ___H2O • ___CO + ___Fe2O3→ ___Fe + ___CO2

  17. Section 8.1 Describing Chemical Change • Did We Meet Our Objectives? • Write equations describing chemical reactions using appropriate symbols • Write balanced chemical equations when given the names or formulas of the reactants and products in a chemical reaction.

  18. Section 8-2 Types of Chemical Reactions

  19. Section 8.2 Types of Chemical Reactions • Objectives: • Identify a reaction as combination, decomposition, single-replacement, double-replacement, or combustion • Predict the products of combination, decomposition, single-replacement, double-replacement, and combustion reactions

  20. Synthesis Reactions • Two or more substances combine to form a new compound. • Also known as composition reaction or combination reaction. • A + X → AX • A and X are elements, while AX is a compound.

  21. Ex #4: Complete and balance the following combination reactions. • Be + O2 → • Cu + S →

  22. Combination Reactions with Water • Some nonmetals oxides react with water to produce an acid. • SO2 + H2O → H2SO3

  23. Ex #5: Complete and balance the following combination reaction. • SO3 + H2O →

  24. Decomposition Reactions • A single compound undergoes a reaction that produces two or more simpler substances. • AX → A + X • AX is a compound and A and X are either elements or simpler compounds. • Most decomposition takes place only when energy in the form of electricity or heat is added.

  25. Decomposition Continued…… • Decomposition of Binary Compounds • Will produce elements. • 2 H2O → 2 H2 + O2 • The decomposition of a substance by an electric current is called electrolysis. • Oxides of less-active metals decompose into their elements when heated. • 2 HgO → 2 Hg + O2

  26. Ex #6: Complete and balance the following decomposition reactions. • HI → • PbO2 →

  27. Single-Replacement Reactions • One element replaces a similar element in a compound. • Also known as displacement reaction. • Many take place in aqueous solutions. • A + BX → AX + B or • Y + BX → BY + X

  28. Single-Replacement Reactions Continued…… • Replacement of a Metal in a Compound by Another Metal • If the single metal is more reactive than the compound metal, the single metal will replace it. • 2 Al + 3 Pb(NO3)2→ 3 Pb + 2 Al(NO3)3

  29. Ex #7: Complete and balance the following single-replacement reactions. • Mg + Zn(NO3)2 → • Cl2 + NaBr →

  30. Activity Series • A list of elements organized according to the ease with which the elements undergo certain chemical reactions. • For metals – greater activity means greater ease of loss of electrons. • For nonmetals – greater activity means greater ease of gain of electrons.

  31. Ex #8: Use the activity series to predict whether each will react. • Zn(s) + H2O(l)→ (at 50.0 °C) • No reaction • Sn(s) + O2(g) → • Yes, a reaction will occur • Cd(s) + Pb(NO3)2(aq) → • Yes, a reaction will occur • Cu(s) + HCl(aq) → • No reaction

  32. Double-Replacement Reactions • The ions of two compounds exchange places in an aqueous solution to form two new compounds. One of the new compounds is usually a precipitate, gaseous product, or water. • AX + BY → AY + BX

  33. Double-Replacement Reactions Continued…….. • Formation of a Precipitate • Cations and Anions form an insoluble compound. • Formation of a Gas • One of the products is an insoluble gas that bubbles out of the mixture. • Formation of Water • Usually an acid and a base will form a salt and water.

  34. Ex #9: Complete and balance the following double displacement reactions. • BaCl2 + K2CO3 → • FeS + HCl →

  35. Combustion Reactions • A substance combines with oxygen, releasing a large amount of energy in the form of light and heat. • Hydrocarbon + O2→ CO2 + H2O • CH4 + 2 O2 → CO2 + 2 H2O

  36. Ex #10: Complete and balance the following combustion reactions. • C6H6 + → • CH4 + → • C11H22 + →

  37. Section 8.2 Types of Chemical Reactions • Did We Meet Our Objectives? • Identify a reaction as combination, decomposition, single-replacement, double-replacement, or combustion • Predict the products of combination, decomposition, single-replacement, double-replacement, and combustion reactions

  38. Section 8.3 Reactions in Aqueous Solutions • Objectives: • Write and balance net ionic equations • Use solubility rules to predict the precipitate formed in double-replacement reactions

  39. Net Ionic Equations • Most ionic compounds dissociate (separate) into cations and anions when they dissolve in water. • A complete ionic equation is an equation that shows the dissolved ionic compounds as their free ions. Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq)→AgCl(s) + Na+(aq) + NO3-(aq)

  40. Net Ionic Equations • If you notice, there are some ions that are the same on both sides of the equation. These are called spectator ions. These ions are not directly involved in the reaction. Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq)→AgCl(s) + Na+(aq) + NO3-(aq) • Na+(aq) and NO3-(aq) are spectator ions for the above equation.

  41. Net Ionic Equations • We can cancel the spectator ions in the following equation. The items left provide us with the net ionic equation, the equation that indicates only those particles that actually take part in the equation. Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq)→AgCl(s) + Na+(aq) + NO3-(aq) Ag+(aq) + Cl-(aq)→AgCl(s)

  42. Ex #11: Write the balanced net ionic equation for the following. Pb(ClO4)2(aq) + 2NaI(aq)→ PbI2(s) + 2NaClO4(aq) Complete Ionic Equation: Pb(aq) + 2ClO4(aq) + 2Na(aq) + 2I(aq)→ PbI2(s) + 2Na(aq) + 2ClO4(aq) Spectator Ions: 2Na(aq)and2ClO4(aq) Net Ionic Equation: Pb(aq) + 2I(aq)→ PbI2(s)

  43. Solubility Rules

  44. Ex #12: Write a complete ionic equation and a net ionic equation for the reaction of aqueous solutions of iron (III) nitrate and sodium hydroxide. Fe(NO3)3 + NaOH → NaNO3 + Fe(OH)3 Fe(NO3)3 + 3NaOH → 3NaNO3 + Fe(OH)3 Fe(NO3)3(aq) + 3NaOH(aq)→ 3NaNO3(aq) + Fe(OH)3(s) Complete Ionic Equation: Fe(aq)3+ + 3NO3-(aq) + 3Na(aq) + 3OH-(aq)→ 3Na(aq) + 3NO3-(aq) + Fe(OH)3(s) Net Ionic Equation: Fe(aq)3+ + 3OH-(aq)→ Fe(OH)3(s)

  45. Section 8.3 Reactions in Aqueous Solutions • Did We Meet Our Objectives? • Write and balance net ionic equations • Use solubility rules to predict the precipitate formed in double-replacement reactions

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