Chapter 11

1. Chapter 11 Connecting chemical reactions and equations

2. Chemical Changes • A chemical change is a reaction in which one or more substances are transformed into one or more new substances. • Physical changes do not produce new substances – change of state.

3. Chemical Changes • When chemical changes are taking place, we may observe one or more of the following: • A substance disappears • A gas is given off (effervescence) • A solid is precipitated • A colour change takes place • The temperature changes • A new odour is released • Light is given out

4. Chemical Equations • A chemical equation is a way of summarising a chemical change. • It shows the formulae of the original reactants on the left and the formulae of the new substances (products) on the right. • It also shows the physical state of each substance involved, and it may indicate conditions necessary for the reaction to occur.

5. Chemical Equations • In a chemical reaction, the way in which the atoms are joined together is changed. • Bonds are broken and new ones formed as the reactants are changed into products. • The same atoms (both number and type) are present before and after the reaction – they are just arranged differently

6. Chemical Equations • When Na is heated and plunged into green chlorine gas, the white ionic solid sodium chloride is formed. • Na + Cl2 ----> NaCl • Name the reactants, name the products • The equation above is called a ‘skeleton’ equation, the physical states need to be added and it needs to be balanced • 2Na(s) + Cl2(g) ----> 2NaCl(s)

7. Chemical Equations • (s) solid • (l) liquid • (g) gas • (aq) solution of water

8. Chemical Equations - catalysts • In some chemical reactions, a substance called a catalyst is used. • A catalyst speeds up a reaction without being used up • Because a catalyst is not used up, it is neither a reactant or a product and is written above the arrow in the equation. • 2H2O2(aq) ----> 2H2O(l) + O2(g) MnO2

9. Chemical Equations – dissolving • When a substance such as copper sulphate, CuSO4, is dissolved, water merely causes the ions in the lattice to come apart. • Since the water is not being changed, it is written above the arrow. • CuSO4(s) ----> Cu2+(aq) + SO42-(aq) H2O

10. Chemical Equations - heating • When a reaction requires heat, the word ‘heat’ is written above the arrow to denote that heat has been applied. • H2O(s) ---> H2O(l) Heat

11. Review • Complete the sample problems 11.1 & 11.2 page 258 • Complete the revision questions 1, 2 page 258

12. Balancing Chemical Equations • To represent chemical equations correctly, equations must be balanced. • The number of atoms on both sides of the equation must be the same • Law of conservation of mass – the total mass of the reactants in a chemical reaction is equal to the total mass of the products. • Atoms are not created or destroyed, but are rearranged to form new substances

13. Balancing Chemical Equations • In order to balance an equation, numbers called coefficients are placed in front of the whole formulas.

14. Rules for Balancing Chemical Equations • 1. Write the reactants and products using formula and state for each substance • 2. Count the number of atoms of each element on the left-hand side of the equation. Do the same for the right-hand side and compare for each element. If any of these numbers do not match, the equation is not balanced and you will need proceed to the following steps • 3. Balance by placing coefficients in front of the formulae. Do not change the actual formula. If the substance is present as an element, leave the balancing of it to last • 4. Check all atoms of ions to ensure that they are balanced • 5. Make sure that the coefficients are in the lowest possible ratio

15. Review • Complete the sample problem 11.3 page 260 • Complete revision questions 3 – 5 page 260

16. Types of Chemical Reaction • Chemists have determined several main groups of chemical reactions that help us predict the products of these reactions. • Keep in mind, though, that when we write equations they should be based on experimental data for complete certainty

17. Types of Chemical Reaction • Main types of chemical reactions • Precipitation • Acid/base • Combustion • Acid molecules • Acid Base reaction animation

18. Precipitation • Precipitation occurs when ions in solution combine to form a new compound of low solubility in water • This low-solubility compound forms as solid particles that eventually settle. • It is called a precipitate • In order to predict whether a precipitate will form, we must know which substances are soluble in water and which substances are insoluble

20. Precipitation • From table 11.2 (page 262), we can predict that when a solution of NaCl is mixed with a solution of AgNO3, a precipitate of AgCl will form • NaCl(aq) + AgNO3 ----> NaNO3(aq) + AgCl(s) • The Ag+ ions combine with the Cl- ions to form solid AgCl while the Na+ and NO3- ions remain in solution. • Precipitation animations

21. Review • Complete the sample problem 11.4 page 262 • Complete the revision questions 6 – 8 page 263

22. Acid/Base and Neutralisation reactions • When an acid reacts with a base to form a salt and water, a neutralisation reaction occurs • Acid + Base ----> salt + water • HCl(aq) +NaOH(aq) ----> NaCl(aq) + H2O(l)

23. Acid/Base and Neutralisation reactions • Other common reactions involving acids are: • Acid + Metal ----> salt + hydrogen • This reaction does not occur with Cu, Hg, or Ag • Acid + Metal carbonate ----> salt + water + carbon dioxide • 2HCl(aq) + Na2CO3(aq) ----> 2NaCl(aq) + HCl(l) + CO2(g) • Acid + metal oxide ----> salt + water • 2HCl(aq) + CuO(s) ----> CuCl2(aq) + H2O(l) • Acid + metal hydroxide ----> salt + water • H2SO4(aq) + 2NaOH(aq) ----> Na2SO4(aq) + 2H2O(l)

24. Acid/Base and Neutralisation reactions • Since a neutralisation reaction is often difficult to detect, a chemical indicator can be used. • Indicators are compounds that change colour when an acid or a base have completely reacted

25. Review • Complete the sample problem 11.5 page 263 • Complete the revision questions 9, 10 page 264

26. Combustion reactions • When hydrocarbons burn in a plentiful supply of oxygen, they give off heat to their surroundings and produce carbon dioxide and water • When limited air is available for a combustion reaction, carbon monoxide may be formed in preference to carbon dioxide. • Burning Hydrocarbons

27. Combustion reactions • Octane burnt in air • 2C8H18(l) + 25O2(g) ----> 16CO2(g) + 18H2O(g) • Octane burnt in an engine (limited O2) • 2C8H18(l) + 17O2(g) ----> 16CO(g) + 18H2O(g) • The products of hydrocarbon combustion are in the gaseous state

28. Review • Complete the sample problem 11.6 page 264 • Complete the revision questions 11, 12 page 265

29. Ionic Equations • In all the preceding equations, the elements and compounds have been written in their molecular or formula unit forms. • In aqueous solutions, the reactions are best represented by simpler equations called ionic equations.

30. Ionic Equations • Ionic equations are equations that show only the species that are formed or changed in a reaction • Any ions that remain unchanged in a reaction are included in an ionic equation. • Ions that are present in a reaction but do not react are called spectator ions.

31. Ionic Equations – The Rules • Write the balanced chemical equation • Decide, from the solubility table, which substances are soluble and which will form precipitates • Expand the equation by dissociating all the soluble compounds into their free ions • Check for any molecular substances such as acids and certain bases that react with water to produce ions (hydrolyse). Replace the formulae of these substances by the ions that they form • Cancel all free ions that are unchanged on both sides of the equation (the spectators) • Write the net ionic equation • Note – the equation must be balanced in charge as well as in the number of atoms • Precipitation and neutralisation reactions are often represented by ionic equations

32. Writing ionic equations for precipitation reactions H2O • When a salt is dissolve in water it breaks up, or dissociates, into its constituent ions: • NaCl(s) ----> Na+(aq) + Cl-(aq) • The symbol (aq) indicates that each ion has become surrounded by water molecules • If a second soluble salt, silver nitrate, is added, it will also dissociate • AgNO3(aq) ----> Ag+(aq) + NO3-(aq) • When the ions of both solutions come into contact a white precipitate or AgCl is formed H2O

33. Writing ionic equations for precipitation reactions Na+(aq) + Cl-(aq) + Ag+(aq) + NO3-(aq) ----> AgCl(s) + Na+(aq) + NO3-(aq) Precipitates are insoluble salts and so are not dissociated in solution The net ionic equation does not include the spectator ions, Na+ and NO3- Ag+(aq) + Cl-(aq) ----> AgCl(s) Ionic Equation Practice

35. Review • Follow the steps and complete the sample problem 11.7 page 266 • Complete the revision questions 13 – 17 pages 266, 267

36. Chemical Reactions – by patterns • Prediction of reactions by patterns • Only a prediction – exceptions can occur

37. Chemical Reactions – by patterns • A list of common patterns • Acid + metal hydroxide (base) ---> salt + water • Acid + basic oxide ---> salt + water • Acidic oxide + base ---> salt + water • Acid + metal ---> salt + hydrogen (no reaction for Ag, Cu, Pt or Au) • Acid + metal carbonate ---> salt + carbon dioxide + water • Metal carbonate ---> metal oxide + carbon dioxide • Hydrocarbon + (plentiful) oxygen ---> carbon dioxide + water Heat