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Chemical Reactions Honor’s

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Presentation on theme: "Chemical Reactions Honor’s"— Presentation transcript:

1 Chemical Reactions Honor’s

2 Reactions and Equations Objectives
1. Recognize evidence of chemical change. 2. Represent chemical reactions with equations. 3. Balance chemical equations. 4. Classify chemical reactions. 5. Identify the characteristics of different classes of chemical reactions. 6. Describe aqueous solutions. 7. Write complete ionic and net ionic equations for chemical reactions in aqueous solutions. 8. Predict whether reactions in aqueous solutions will produce a precipitate, water, or a gas.

3 Reactions and Equations
chemical reaction: process by which one or more substances are rearranged to form different substances; also called a chemical change -reactants: the starting substances -products: resulting, new substances The law of conservation of mass states that the total mass of reactants must equal the total mass of the products.

4 Evidence of Chemical Reactions
While some reactions are hard to detect, most provide evidence they have occurred. 1. temperature change -exothermic-heat released -endothermic-heat absorbed 2. color change -by itself it doesn’t necessarily indicate a chemical change since it is also an indication of a physical change

5 -light, sound, electricity
production of a gas -may see bubbling -may also produce new odor 4. formation of a precipitate, a solid formed as a result of a chemical reaction in solution and that separates from the solution energy change -light, sound, electricity

6 -must always state the physical state of each reactant and product
No matter the method of representation, there is a universal set of symbols all scientists use. -must always state the physical state of each reactant and product -arrow always points to the products; most often reactants are written on the left, products on the right Symbol Meaning + plus; separates 2 or more reactants or products produces, yields, or forms (s) solid (l) liquid (g) gas (aq) aqueous; dissolved in water

7 Chemical Equations chemical equation: uses chemical formulas to show the identities and relative amounts of reactants and products in a chemical reaction. -must show mass being conserved # of reactant atoms = # of product atoms -accomplished by balancing equations Chemical equations are the representation chemists use to describe chemical reactions.

8 Balancing Chemical Equations
coefficient: number written in front of a reactant or product that states the ratio of amounts for each substance -usually a whole number -number ‘1’ is assumed and not written Steps for balancing equations: 1. Write the skeleton equation for the reaction 2. Count the atoms of elements in the reactants/products 3. Place coefficients in front of each substance; change until the equation is balanced 4. Reduce coefficients to smallest possible ratio 5. Check your work.

9 Note: Never change the subscript in the formula because that changes the identity of the substance.
Example: Hydrogen chloride is formed during the reaction between hydrogen and chlorine. 1. Write the skeleton equation for the reaction. hydrogen + chlorine  hydrogen chloride H2 (g) Cl2 (g)  HCl (g)

10 2. Count the atoms of elements in the reactants/products
___H2 (g) + ___Cl2 (g)  ___HCl (g) H Cl 3. Place coefficients in front of each substance; change until the equation is balanced 4. Reduce coefficients to smallest possible ratio H2 (g) + Cl2 (g)  2HCl (g) 5. Check your work.

11 Balancing Equations Practice
__SnS2 + __O2  __SnO2 + __SO2 __C2H6 + __O2  __CO2 + __H2O __Al + __HCl  __AlCl3 + __H2 __N2 + __H2  __NH3 __NaCl + __F2  __NaF + __Cl2 __AgNO3 + __MgCl2  __AgCl + __ Mg(NO3)2 __HCl + __CaCO3  __CaCl2 + __H2O + __CO2 __CO2 + __H2O  __C6H12O6 + __O2

12 Word Equations Word equations use words to represent chemical reactions. Skeleton equation uses chemical formulas to identify the reactants and products. Chemical equations show conservation of mass by placing coefficients in front of each substance to balance it. “Iron and chlorine react to produce iron (III) chloride” word: iron (s) + chlorine (g)  iron (III) chloride (s) skeleton: Fe (s) Cl2 (g)  FeCl3 (s) chemical: 2Fe (s) Cl2 (g)  2FeCl3 (s)

13 Word Equations Know ~8 diatomics: H2, N2, O2, F2, Cl2, Br2, I2, At2
(red are gases, blue is a liquid and black is a solid) ~nonmetal oxides tend to be gases (CO, SO2, …) ~ionic compounds tend to be solids (watch for aq solution) ~be careful with solubility rules in aq solution ~pure metals are solids (except Hg)

14 Chemical Equations Practice 1
Write chemical equations and balance each: 1. In water, iron (III) chloride reacts with sodium hydroxide, producing solid iron (III) hydroxide and sodium chloride.

15 Chemical Equations Practice 1
Write chemical equations and balance each: 2. Liquid carbon disulfide reacts with oxygen, producing carbon dioxide and sulfur dioxide.

16 Chemical Equations Practice 1
Write chemical equations and balance each: 3. Solid zinc and aqueous hydrogen sulfate react to produce hydrogen gas and aqueous zinc sulfate.

17 Classifying Chemical Reactions
There are 4 basic types of reactions: 1. synthesis reaction- 2 or more simple substances combine to form a new, more complex substance -examples: 2Na + Cl2  2NaCl a. combustion reaction- a substance combines with oxygen, releasing a large amount of energy in the form of heat and light -ex: 2CH O2  2CO2 + 4H2O + energy

18 b. polymerization reaction: reaction in which
monomer units (a small, simple organic molecule) are bonded together to form a polymer, a large molecule consisting of many repeating structural units (monomers) -addition polymerization: all atoms present in monomer are present in the polymer product ~ex: ethene (ethylene)  polyethylene H2C=CH2  -CH2-CH2-CH2-CH2- -condensation polymerization: when monomers containing at least 2 functional groups combine with the loss of a small by-product, usually water

19 2. decomposition reaction - a complex substance breaks down into at least 2 or more simpler substances -opposite of synthesis reactions -many require energy to occur (endothermic) -examples: a) H2CO3  H2O + CO2 b) electrolysis: decomposition by electric current

20 3. single-displacement reaction - an uncombined element replaces a similar element that is part of a compound -more reactive element displaces the less reactive one (use activity series; most metals replace H) -example: 2Na + 2H20  2NaOH + H2

21 4. double-displacement reaction- different atoms (usually ions) in 2 different compounds exchange places to form new compounds -one of the compounds formed is usually a precipitate, gas, or liquid ~use solubility rules -example: MgCO3 + 2HCl  MgCl2 + H2CO3 Classify each of the reactions (1-5) from Equations Practice 1. Using this information, you should be able to predict products when given reactants.

22 Reactions in Aqueous Solutions
Some aqueous solutions contain molecules (such as sugar, ethanol). Others contain ionic compounds (or acids) that break apart, or dissociate, in water to form ions. -when two aqueous solutions that contain ions react, a double displacement reaction occurs -strong electrolytes ionize in water: strong acids, strong bases, soluble salts

23 Aqueous Solutions You need to memorize:
-strong acids: HCl, HBr, HI, HNO3, H2SO4, HClO4, HClO3 -strong bases: all group IA hydroxides; group IIA hydroxides that include Ba, Sr, Ca Know: -gases, pure liquids, solids are non-electrolytes -H2CO3 decomposes into H2O & CO2 -NH4OH deomposes into H2O & NH3 Remember: -the products of a double displacement reaction is either a precipitate, liquid or gas. -Use your solubility rules to determine if a precipitate occurs.

24 Aqueous Solutions Chemists use ionic equations to show the driving force of a chemical reaction. -complete ionic equation: shows all particles in a solution -net ionic equation: shows only the particles that participate in the reaction (forms a solid, liquid, gas) -spectator ions: ions that do not participate; they just ‘watch’

25 Net Ionic Equations Example:
2NaOH (aq) + CuCl2(aq)  2NaCl (aq) + Cu(OH)2 (s) 2Na+(aq) + 2OH-(aq) + Cu+2(aq) + 2Cl-(aq)  2Na+(aq) + 2Cl-(aq) + Cu(OH)2 (s) (notice the solid precipitate is not split up into its ions-this is true for gases and liquids produced as well) 2OH-(aq) + Cu+2(aq)  Cu(OH)2 (s) Na+ and Cl- are the spectator ions since they are not part of the formation of Cu(OH)2 (s)

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