1 / 35

Chapter 8

Chapter 8. Describing Chemical Change Types of Chemical Reactions Reactions in Aqueous Solution. Chapter 8.1 Describing Chemical Change. Word Equations Chemical Equations Balancing Chemical Equations. Word Equations.

kenneth-garza
Download Presentation

Chapter 8

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter 8 • Describing Chemical Change • Types of Chemical Reactions • Reactions in Aqueous Solution

  2. Chapter 8.1 Describing Chemical Change • Word Equations • Chemical Equations • Balancing Chemical Equations

  3. Word Equations • Reaction – one or more substances (the reactants) change into one or more new substances(the products) • Reactants  Products •  = ? • Yields, gives or reacts

  4. Word Equations • As reactants are converted to products, the bonds holding the atoms together are broken and new bonds are formed. • REMEMBER: The atoms are neither created nor destroyed, just rearranged. (Law of Conservation of Mass)

  5. Word Equations • Rust • Iron reacts with oxygen to produce iron(III) oxide (rust) • Iron + Oxygen  Iron(III) Oxide (reactants) (yields) (products)

  6. Word Equations • Hydrogen peroxide reacts to form water and oxygen gas. • Hydrogen peroxide  water + oxygen MnO2 • H2O2(aq)  H2O(l) + O2(g)

  7. Word Equations • Burning of Methane • Methane + Oxygen  Carbon Dioxide + Water

  8. Chemical Equations • Iron + Oxygen  Iron(III) Oxide • Fe + O2  Fe2O3 • Now add physical states • Fe(s) + O2 (g)  Fe2O3 (s)

  9. Chemical Equations • Reactions with a catalyst • Catalyst – a substance that speeds up the rate of a chemical reaction, but is not used up in the reaction. • A catalyst is written above the arrow

  10. Chemical Equations • Manganese(IV) oxide catalyzes the decomposition of hydrogen peroxide.

  11. Balancing Chemical Reactions • Each side of the equation has the same number of atoms of each element. • C(s) + O2(g) CO 2(g) • 1 carbon, 2 oxygen  1 carbon, 2 oxygen

  12. Balancing Chemical Reactions • H2(g) + O2(g) H2O(l) • 2 hydrogen, 2 oxygen  2 hydrogen, 1 oxygen • Balance • 2 H2(g) + O2(g) 2 H2O(l)

  13. Balancing Chemical Reactions • Rules • 1) Determine the correct formulas • 2) Write the formulas for the reactants on the left and the products on the right. Place a  in between. If there are two or more reactants or products, use a + in between. • 3) Count the number of atoms of each element.

  14. Balancing Chemical Reactions • Rules • 4) Balance the elements one at a time until you have equal numbers of elements on each side • 5) Make sure all numbers are in their smallest whole number ratio.

  15. Chapter 8.2 Types of Chemical Reactions • Classifying Reactions • Combination Reactions • Decomposition Reactions • Single-Replacement Reactions • Double-Replacement Reactions • Combustion Reactions • Predicting Products of Reactions

  16. Classifying Reactions • Identify the five general types or reactions: • Combination Reactions • Decomposition Reactions • Single-Replacement Reactions • Double-Replacement Reactions • Combustion Reactions

  17. Combination Reactions • Two or more substance combine to form a single substance. • General Reaction: • R + S = RS • Example: • 2Mg(s) + O2(g) 2 MgO(s)

  18. Decomposition Reactions • A single compound is broken down into two or more substances. • General Reaction: • RS = R + S • Example: • 2HgO(s) 2 Hg(l) + O2(g)

  19. Single-Replacement Reactions(Single Displacement Reactions) • One element replaces a second element in a compound. • General Reaction: • T + RS = TS + R • Example: • 2K(s) + 2H2O(l) 2KOH(aq) + H2(g)

  20. http://www.kentchemistry.com/links/Kinetics/PredictingSR.htm

  21. Double-Replacement Reactions • Exchange of positive ions between two reacting compounds. • General Reaction: • RS + TU = RU + TS • R+S- + T+U- = R+U- + T+S- • Example: • K2CO3(aq) + BaCl2(aq) 2KCl(aq) + BaCO3(s)

  22. http://www.kentchemistry.com/links/Kinetics/DRFlash.htm

  23. Combustion Reactions • An element or compound reacts with oxygen often producing energy as heat or light. • General Reaction: • CxHy + (x + y/4) O2 xCO2 + (y/2)H2O • Example: • CH4 (g) + 2O2 (g) CO2(g) + 2H2O (g)

  24. Name each type of reaction • 1) 2Mg(s) + O2(g) 2 MgO(s) • 2) 2HgO(s) 2 Hg(l) + O2(g) • 3) CH4 (g) + 2O2 (g) CO2(g) + 2H2O (g) • 4) 2K(s) + 2H2O(l) 2KOH(aq) + H2(g) • 5) K2CO3(aq) + BaCl2(aq) 2KCl(aq) + BaCO3(s)

  25. Name each type of reaction • 1) 2Mg(s) + O2(g) 2 MgO(s) • Combination • 2) 2HgO(s) 2 Hg(l) + O2(g) • Decomposition • 3) CH4 (g) + 2O2 (g) CO2(g) + 2H2O (g) • Combustion • 4) 2K(s) + 2H2O(l) 2KOH(aq) + H2(g) • Single Replacement • 5) K2CO3(aq) + BaCl2(aq) 2KCl(aq) + BaCO3(s) • Double Replacement

  26. Chapter 8.3 Reactions in Aqueous (aq) Solutions • Net Ionic Equations • Predicting the Formation of a Precipitate

  27. Net Ionic Equations • AgNO3(aq) + NaCl (aq) AgCl(s) + NaNO3(aq) • Double Replacement Reaction • Most ionic compounds dissociate (separate) into ions (cations and anions) when they dissolve in water. • When all ions dissociate, we write the equation with the charges on it. = Complete Ionic Equation

  28. Complete Ionic Equation • Ag+(aq) + NO3-(aq) + Na+(aq) Cl-(aq) AgCl(s) + Na+(aq)+ NO3-(aq) • The equation can be simplified by crossing out any ions that do not participate in the reaction. You do this by canceling out ions that appear on both sides. • Ag+(aq) + NO3-(aq) + Na+(aq) Cl-(aq) AgCl(s) + Na+(aq)+ NO3-(aq) • You are left with the Net Ionic Equation: • Ag+(aq) + Cl-(aq) AgCl(s)

  29. Pb(s) + AgNO3(aq) Ag(s) + Pb(NO3)2(aq) • Pb(s) + Ag+(aq)+NO3 -(aq) Ag(s) + Pb2+(aq) +2NO3-(aq) • Pb(s) + Ag+(aq) Ag(s) + Pb2+(aq) • Need to balance charges • Pb(s) + 2Ag+(aq) 2Ag(s) + Pb2+(aq)

  30. Predicting the Formation of a Precipitate • Table F in reference table

More Related