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Chemical Reactions

Chemical Reactions. Chapter 10. Part I: Counting Atoms. How Many Atoms in a Molecule?. Counting Atoms. Most substances that we encounter are compounds , not elements. • A chemical compound is a pure substance formed from the combination of two or more different elements.

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Chemical Reactions

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  1. Chemical Reactions Chapter 10

  2. Part I: Counting Atoms How Many Atoms in a Molecule?

  3. Counting Atoms • Most substances that we encounter are compounds, not elements. • • A chemical compound is a pure substance formed from the combination of two or more different elements. • The properties of the compound may be completely unlike those of the elements that form it. • • The formula for a compound lists the symbols of the individual elements followed by subscripts which indicate the number of atoms of that element. • (If no subscript is given, it is understood to be “1.”) E.g., NaCl, H2O, C12H22O11.

  4. Counting Atoms • A molecular formula gives the actual number of atoms of each element in a molecule of a compound. • Hydrogen peroxide H2O2 • Water H2O • Glucose C6H12O6 • A structural formula uses lines to represent covalent bonds, and shows how the atoms in a molecule are joined together: • H—O—O—H • H—O—H • O=C=O

  5. Counting Atoms • Example: How Many Atoms? C6H12O6 6 C’s + 12 H’s + 6 O’s = 24 atoms K3PO4 3 K’s+1 P’s + 4 O’s = 8 atoms C2H5OH 2 C’s+ 5 H’s + 1 O’s + 1 H = 9 atoms H2O2 2 H’s and 2 O’s = 4 atoms

  6. Al O O O O O O S S S O O O Al O O O Counting Atoms with Polyatomic Ions - Al2(SO4)3 17 ATOMS =

  7. Counting Atoms with Polyatomic Ions • When counting atoms with polyatomic ions; • Count number of atoms in one polyatomic ion • Ions inside the parentheses • Multiply by number of polyatomic groups in the molecule (number outside the parenthesis) • Examples: • Al2(SO4)3 - 2 Al’s + 3(1 S + 4 O’s) = 2 + 3(5) = 2+15 = 17 atoms • Mg(NO3)2 – 1 Mg + 2(1 N + 3 O’s) = 1 + 2(4) = 1+8 = 9 atoms

  8. Hydrates • Hydrates are ionic compounds which also contain a specific number of water molecules associated with each formula unit. The water molecules are called waters of hydration. • The formula for the ionic compound is followed by a raised dot and #H2O • Example: MgSO4•7H2O. • They are named as ionic compounds, followed by a counting prefix and the word “hydrate” • CuSO4•5H2O copper(II) sulfate pentahydrate • BaCl2•6H2O barium chloride hexahydrate • MgSO4•7H2O magnesium sulfate heptahydrate (Epsom salts)

  9. How Many Atoms in a Hydrate? • When counting atoms in the hydrate, count the water atoms also. • Example: • CuSO4•5H2O • 1 Cu + 1 S + 4 O’s + 5(2 H’s + 1 O’s) • =1+1+4+5(3) • = 6+15 = 21 atoms • BaCl2•6H2O • 1 Ba + 2 Cl + 6(2 H’s + 1 O’s) • = 1+2+6(3) • =3 + 18 = 21 Atoms

  10. Part II: Conservation of Mass

  11. Conservation of Mass • In a normal chemical reaction, the mass of substances in a closed system will remain constant, no matter what processes are acting inside the system. • How ever many atoms a reaction starts with, ends with the same number. • Atoms don’t change their identity in a chemical reaction • Number of atoms for EACH ELEMENT STAYS THE SAME in a chemical reaction • The elements just rearrange their organization • The beginning MASS of the reaction EQUALS the ending MASS of the reaction

  12. H H O O O H H H H O H H Conservation of Mass • Total Mass stays the same in a chemical reaction 2g H2 + 16g O2 yields 18g H2O • Number and Identity of Atoms stays the same in a chemical reaction • 2 H2 + 1 O2 yields 2 H2O

  13. Part III: Writing Reactions How Do You Write a Chemical Reaction?

  14. III. Chemical Reactions • Definition – process by which the atoms of one or more substances are rearranged • KEY: new substances are formed • KEY: No Atoms are Gained or Lost • A chemical reaction is the process by which atoms of one or more substances are rearranged into new substances • Chemical change occurs • How do you know?

  15. III. Evidences of a Chemical Reaction 1) gas production 2) light production 3) temperature change (endo/exothermic) 4) precipitate formed (solid from 2 liquids) 5) permanent color change

  16. III. Energy Changes • Energy is stored in compounds as chemical potential energy • due to specific arrangements of atoms. • A chemical reaction changes the potential energy present.

  17. Energy Changes • When energy is lost as heat, it is called an __________________. exothermic reaction These reactions get hotter. • When energy is gained; heat is added for a reaction to occur. These are called ______________________, endothermic reactions These reactions get colder. • Energy in a reaction is shown with: • ΔH (heat) • kJ • Joules • Heat • energy

  18. III. Chemical Reactions • Representing Chemical Reactions: • Reactants – the ‘stuff’ you start with • An ‘arrow’ which means ‘yields’, or ‘becomes’ • Products – the ‘stuff’ you end up with • Principle of “Conservation of Mass” applies to chemical reactions. • Why?

  19. III. Chemical Reactions • Word Equations: • Reactant-A + Reactant-B yields Product-AB • Example: • Sodium(s) + Chlorine(g) → Sodium Chloride(s) • The small letters in paretheses () indicate the state of the reactant or product (solid, liquid, gas, or aqueous solution) • (s) = solid • (l) = liquid • (g) = gas • (aq) = aqueous = dissolved in water

  20. Part IV: Balancing Equations Applying Conservation of Mass to Equations

  21. VI. Chemical Equations • Step 1: Write a Skeleton Equation • Skeleton Equation uses chemical formulas and symbols instead of words: • Words: Sodium + Chlorine gas yields Sodium Chloride • Symbols: Na(s) + Cl2(g) → NaCl • Skeleton Equations are not complete equations, but are the first step in writing a complete equation

  22. IV. Chemical Equations • Chemical Equation is BALANCED • Balanced means that “conservation of mass” is upheld • All atoms in reactants are also in products • No more, no less • Just rearranged

  23. IV. Chemical Equations • Balancing Equations • Use a number before the compound/element symbol to indicate how many of them are needed • Called a COEFFICIENT • Written in front of the atom/compound • KEY: Coefficient is a MULTIPLIER • Number of atoms per molecule is SUBSCRIPT • Change ONLY the COEFFICIENTS to balance the equation

  24. IV. Chemical Reactions • Steps to Balance Equations • Write the skeleton equation • Count the atoms of EACH element in the reactants • Count the atoms of EACH element in the products • Change the coefficients to make the number of atoms of each element equal on both sides of the equation • Write the coefficients in the lowest possible ratio • Check your work • NEVER CHANGE A SUBSCRIPT

  25. 2. Count Number of atoms for each element on both sides IV. Chemical Equations • Write the skeleton equation: • Al + O2→ Al2O3 This is not balanced because the numbers don’t match 3. Multiply coefficients until they match – multiply the entire units 2 Al + O2 → Al2O3 Go to 6 Oxygens

  26. Al + 2 Al2O3 O2 IV. Balancing Equations Al + O2 2 Al2O3 Multiply each atom by 2 4 3 Balanced

  27. IV. Balancing Equations 2 The work of balancing a chemical equation is in many ways a series of trials and errors. Consider the equation given below. Does this represent a balanced chemical equation? N2 + H2 NH3

  28. N2 + H2 2NH3 IV. Balancing Equations 3 To balance this reaction, it is best to choose one kind of atom to balance initially. Let's choose nitrogen in this case. 2 Nitrogen Atoms in Reactants requires 2 Ammonia molecules in Product to balance the nitrogen

  29. + H2 2NH3 IV. Balancing Equations 2 • Once we know what the molecules are (N2, H2, and NH3 in this case) we cannot change them (only how many of them there are). • The nitrogen atoms are now balanced, but there are 6 atoms of hydrogen on the product side • only 2 of them on the reactant side. • The next step requires multiplying the number of reactant hydrogen molecules by three to give: Balanced N2 3H2

  30. IV. Don’t Forget: Diatomic Elements • Definition – 7 elements that NEVER occur as singular atoms (always paired with an the same or different element) H2 O2 F2 Br2 I2 N2 Cl2 Ex: 2 HCl + 2K  2 KCl + H2

  31. IV. Balancing Equations 3 1. Start with an unbalanced equation 2. Draw boxes around the compounds so you don’t mess with the groups Don’t be threatened by how complex it looks!

  32. IV. Balancing Equations 2 3. Make an element inventory – count number of atoms for each element on each side of the equation

  33. 1 2 Element Reactant Product Balanced? Na 1 2 O 1 1 2 1 H 3 3 4 2 SO4 1 1 IV. Balancing Equations 3 • 4. Write coefficients in front of each of the boxes until the inventory for each element is the same both before and after the reaction • Save Oxygen and Hydrogen for last, Treat Polyatomic like an atom. • Let’s start with Sodium • We have 2 in products, so I need 2 in reactants Multiply reactant with sodium by 2 and recount atoms N Y Y N N Y

  34. 1 2 Element Reactant Product Balanced? Na 2 2 O 2 1 H 4 2 4 2 SO4 1 1 IV. Balancing Equations 3 • Inventory Shows: • Reactant side has FOUR hydrogen atoms • Product side has TWO hydrogen atoms • Using your amazing powers of mathematics • two hydrogen multiplied two becomes four hydrogen Balanced Y Y N Y N Y

  35. Helpful Hints • Balance hydrogen and oxygen last • Balance polyatomic ions as a group if present on both reactants and products • You can consider a polyatomic ion as a single element • If the balancing starts to get very complex: • Stop • Start over • Select a different atom to balance first.

  36. Before Example Using PolyAtomics MgCl2 + NaOH  Mg(OH)2 + NaCl 1 Mg 1 Na 1 Mg 1 Na 2 Cl 1 OH 1 Cl 2 OH • After MgCl2 + 2 NaOH  Mg(OH)2 + 2 NaCl 1 Mg 2 Na 1 Mg 2 Na 2 Cl 2 (OH) 2 Cl 2 (OH)

  37. Types of Chemical Reactions Part V

  38. Classifying Chemical Reactions • Synthesis • Decomposition • Single replacement • Double Replacement • Combustion

  39. Synthesis • Definition – two or more substances react to form ONE product • Product is usually bigger or more complex than either reactant A + B  AB

  40. Hey baby let’s get jiggy.

  41. Synthesis • reaction of two elements 2 3 2 ___Al + ___Cl2 ___AlCl3 Al3+ Cl1-

  42. Decomposition • definition – one substance breaks down into two or more simpler products AB  A + B

  43. Break yoself fool!

  44. Decomposition • Example reaction: 2 2 3 __ NaN3 (s)  ___ Na (s) + ___ N2 (g) 2 2 1 __ CaO (s)  ___ Ca (s) + ___ O2 (g)

  45. Single Replacement Reactions • Definition – one element replaces another element in a compound to form new compound A + BX  AX + B

  46. I’m gon’ dance with yo’ lady

  47. Double Replacement • Defn – exchange of cations between two ionic compounds A B + C D  AD + CB switch

  48. 3 possible products of double replacement reactions • Precipitate • Gas • Water

  49. Reactivity Series (or Activity Series) • More active will replace less active • Less active will NOT replace more active • metals Li K Ca Na Mg Al Mn Zn Fe Ni Sn Pb Cu Ag Au most active least active • halogens F Cl Br I most active least active

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