Chapter 11 Matter and Change 11.1 Describing Chemical Reactions 11.2 Types of Chemical Reactions

1. Chapter 11 Matter and Change 11.1 Describing Chemical Reactions 11.2 Types of Chemical Reactions 11.3 Reactions in Aqueous Solution

2. Introduction to Chemical Equations Parts of a Chemical Formula: 2Al2(SO4)3 Introduction to Chemical Equations Element Symbols Subscripts Coefficient Coefficient: Number of moles of molecules or formula units of that compound; Symbols: Elements in the compound; Subscripts: Proportion of atoms in one mole or one molecule or f.u. of the compound. Remember that these can NEVER change in a formula!!

3. Introduction to Chemical Equations • Interpreting Chemical Equations • Chemical equations describe actual chemical reactions • Starting substances are called reactants, and are on the left side of the arrow • Ending substances are called products, and are on the right side of the arrow • Recall that all chemical equations must demonstrate the Law of Conservation of Matter by being balanced!! (Chapter 2)

4. Introduction to Chemical Equations • Diatomic Elements--exist naturally as two atom molecules: • Fluorine: F2 Hydrogen: H2 • Chlorine: Cl2 Oxygen: O2 • Bromine: Br2 Nitrogen : N2 • Iodine: I2 • Fondly known as “HOBrFINCl” • Be sure to use these diatomic symbols when writing chemical equations with these individual elements, or equations will never balance!

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6. Introduction to Chemical Equations Skeleton (Word) Equations How could you describe the rusting of iron? • You could say, “Iron reacts with oxygen to produce iron(III) oxide (rust).” • It is quicker to identify the reactants and product by means of a word equation. Iron + oxygen → iron (III) oxide

7. Introduction to Chemical Equations Skeleton (Word) Equations The production of a new substance, a gas, is evidence of a chemical change. • Two new substances are produced in this reaction, oxygen gas and liquid water. • You could describe this reaction by saying “hydrogen peroxide decomposes to form water and oxygen gas.” Hydrogen peroxide → water + oxygen gas

8. Introduction to Chemical Equations MnO2 H2O2 (aq) H2O (l) + O2 (g) Chemical Equations In many chemical reactions, a catalyst is added to the reaction mixture. • A catalystis a substance that speeds up the reaction but is not used up in the reaction. • A catalyst is neither a reactant nor a product, so its formula is written above the arrow in a chemical equation.

9. Introduction to Chemical Equations • Translating word equations to skeleton equations: • Word equations are translated into “skeleton equations” by using chemical formulas and common reaction symbols. • Skeleton equations are not balanced. • Draw a box around names of chemicals and circle descriptions of physical states. • The verb phrase will determine the placement of the reaction arrow. • Write formulae/symbols for all chemicals (Ch. 9!!) • Write symbols for physical state as subscripts after formulae.

10. Introduction to Chemical Equations • Examples • 1. Liquid hydrogen peroxide (H2O2) decomposes to form water vapor and oxygen gas in the presence of the catalyst manganese (IV) oxide. • 2. Solid calcium carbide (CaC2) reacts with water to form ethyne (C2H2) gas and aqueous calcium hydroxide.

11. Introduction to Chemical Equations • 3. Ethyne gas reacts with oxygen in the air in the presence of a flame to produce carbon dioxide gas and water vapor. • 4. Aqueous solutions of lead (II) nitrate and sodium iodide react to form solid lead (II) iodide and aqueous sodium nitrate.

12. We can use everyday words to describe chemical reactions. What is the advantage of using an equation? A chemical equation for a reaction is easier to read quickly. • It shows all of the relevant information: • quantities of reactants and products, • the direction of the reaction, and • any catalysts needed

13. Balancing Chemical Equations A chemical reaction is described by a balanced equationin which each side of the equation has the same number of atoms of each element and mass is conserved. • As reactants are converted to products, the bonds holding the atoms together are broken, and new bonds are formed. • The atoms themselves are neither created nor destroyed; they are merely rearranged. • Recall from Chapter 2: In any chemical change, the TOTAL MASS remains the same!!

14. Balancing Chemical Equations • Balancing Chemical Equations • 1. Be sure all formulas/symbols are correct before attempting to balance the equation! • 2. Subscripts are “Off- Limits”!! • 3. If it is helpful to draw a picture of the molecules, do so!! • 3. Adjust coefficients in front of formulas/symbols only! • 4. The number and type of atoms on each side of reaction must balance! • 5. Coefficients used must be in the lowest ratio possible!!

15. Balancing Chemical Equations C(s)Carbon O2(g)Oxygen + CO2(g)Carbon dioxide Reactants1 carbon atom, 2 oxygen atoms Product1 carbon atom, 2 oxygen atoms Example #1: Carbon burns in the presence of oxygen to produce carbon dioxide. • This equation is balanced—the number of each type of atom is the same on both sides of the equation. • You do not need to change the coefficients.

16. Balancing Chemical Equations Example #2: When hydrogen and oxygen are mixed, the product is water. • The formulas for all the reactants and the product are correct, but this equation is not balanced. • As written, the equation does not obey the law of conservation of mass (check the atom counts!! )

17. Balancing Chemical Equations When hydrogen and oxygen are mixed, the product is water. • If you put the coefficient 2 in front of H2O, oxygen will be balanced. • Now twice as many hydrogen atoms are in the product as are in the reactants.

18. Balancing Chemical Equations When hydrogen and oxygen are mixed, the product is water. • To correct this equation, put the coefficient 2 in front of H2. • The equation is now balanced.

19. Balancing Chemical Equations Hints for Balancing Reactions: • Balance by listing chemicals individually & counting as needed; • Balance using multiples of two where possible; • Utilize least common multiples where possible; • Balance polyatomic ions as a whole unit rather than as individual atoms; • Use fractions as appropriate to balance mathematically, and thenmultiply the whole reaction by the factor that clears the fraction. • Make sure final equation uses lowest possible ratio of coefficients.

20. Practice Problems: _____H2O2_______H2O + _______O2 ___CaC2 + ___H2O____C2H2 +___Ca(OH)2 __ C2H2 + ___O2 ____CO2 + ___H2O

21. Practice Problems: ___Pb(NO3)2 + ____NaI ____NaNO3 + ____PbI2 ___Al + ___HCl ___AlCl3 + ____H2 ___C10H22 + ___O2 ___CO2 + ____H2O

22. Practice Problems: _____HBr _____H2 + _____Br2 ___Mg + _____H3PO4  ___Mg3(PO4)2 + ___H2

23. Balance the following equation. C3H8 (g) + O2 (g)→ CO2 (g) + H2O (l) C3H8 (g) + 5 O2 (g)→ 3 CO2 (g) + 4 H2O (l)

24. BIGIDEA Chemical Reactions • The law of conservation of mass states that mass is neither created nor destroyed. • In order to show that mass is conserved during a reaction, a chemical equation must be balanced.

25. END OF 11.1