1 The Mole and Chemical Equations. 2 Mole and Chemical Equations Concepts to Master Be able to...

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The Mole and Chemical Equations

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Mole and Chemical EquationsConcepts to Master

• Be able to covert the moles of a substance to atoms (and vice versa).

• Be able to convert grams of a substance to moles (and vice versa).

• Be able to balance equations.• Be able to convert the grams of one

substance to the grams of another substance using a balanced equation.

• Be able to convert the moles of one substance to the moles of another substance using a balanced equation.

• Be able to determine the empirical formula from gram analysis data.

• Be able to determine the empirical formula from % composition.

• Be able to determine the molecular formula from an empirical formula.

• Be able to predict the products of a chemical reaction using Table J (for single replacement) and Table F (for double replacement).

• How do you know that a reaction has come to completion?

• Be able to use Table F to determine solubility.

• Is a combustion reaction endothermic or exothermic?

• Which reactions release energy?• Which reaction absorb energy?• When energy is absorbed, is it a

product or reactant?• What is the conservation of mass

law and why is it important to this chapter?

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Vocab• Addition reaction• Analysis reaction• Aqueous• Atomic Mass• Coefficient• Combination reaction• Combustion• Composition reaction• Decomposition• Double replacement• Empirical Formula• Endothermic• Exothermic• Formula mass• Formula units• Gas• Liquid• Metathesis• Molar Mass• Mole

• Molecular Formula• Phases of matter• Precipitate• Products• Reactants• Single replacement• Solid• Stoichiometry• Structural formula• Synthesis

Labs• Celebrating the Mole• Burning Mg• Balance Beads• Mole and Mass

Relationships

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THE MOLE!!!

• A mole is the number of protons or neutrons needed to make up one gram of any element.

• 1 mole = 602 000 000 000 000 000 000 000 = 6.02x1023 particles

• It doesn't matter if it's atoms, donuts or basketballs, a mole is the same number of things. (Like a dozen of anything is always 12.)

• Atomic weight of an element is the average of the masses of it’s isotopes.– For hydrogen, this value is 1.00794 atomic mass units, for chlorine its 35.453 amu

– But since a mole of amu’s is a gram, we can use these exact same numbers for atomic weights in “grams per mole.” Hydrogen weighs 1.00794 grams per mole of atoms; chlorine weighs 35.453 grams per mole of atoms, etc.

• No one ever deals with just one or two atoms. We deal with enough to see, which means trillions and trillions of them – measurable numbers or fractions of moles.

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Molar Mass = Atomic Mass found on periodic table

Formula Composition Moles g/mol

H2OTotal of 3 atoms in each molecule.

1 oxygen and 2 hydrogen

1 mole of oxygen for every 2 moles of hydrogen

1 (16.00) = 16.00 g/mol

2 (1.01) = 2.02 g/mol

18.02 g/mol

CO2Total of 3 atoms in each molecule.

2 oxygen and 1 carbon

2 moles of oxygen for every 1 mole of carbon

1 (12.00) = 12.00 g/mol

2 (16.00) = 32.00 g/mol

44.00 g/mol

O2Total of 2 atoms in each molecule.

2 oxygen

2 moles of oxygen

2 (16.00) = 32.00 g/mol

C6H12O6Total of 24 atoms in each molecule.

6 oxygen, 12 hydrogen, and 6 carbons

6 moles of oxygen, 12 moles of hydrogen, and 6 moles of carbon

6 (16.00) = 96.00 g/mol

12 (1.01) = 12.12 g/mol

6 (12.00) = 72.00 g/mol

180.12 g/mol

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Mass to Mole Conversions

• How many moles are present in 152 g of H2SO4?

• How many moles are present in 13.3 kg of KNO3?

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Mole to Mass Conversions

• How many grams are present in 252.3 moles of CO2?

• How many grams are present in 1.3 moles of H2O?

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• Use Avagadro’s number in your dimensional analysis

• How many molecules are present in 8.2 moles of ammonia?

• How many molecules are present in 10 moles of water?

Mole to Particle Conversions

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• Use Avagadro’s number in your dimensional analysis• How many moles are present in 2.359 x 1036 molecules

of carbonic acid?

• How many moles are present in 250 x 1023 atoms of neon?

Particle to Mole Conversions

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Molecular and Empirical Formulas

• An empirical formula for a compound shows the simplest whole number ratio for the elements in the compounds.

• A molecular formula shows the molecule as it actually exists In nature.

CMPD Empirical Molecular

Water H2O H2O

Hydrogen peroxide

HO H2O2

glucose CH2O C6H12O6

Methane CH4 CH4

Ethane CH3 C2H6

Octane C4H9 C8H18

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Determining an Empirical Formula from Gram Analysis

A compound was analyzed and found to contain 13.5 g Ca, 10.8 g O, and 0.675 g H. What is the empirical formula of the compound?

1. Start with the number of grams of each element, given in the problem.

2. Convert the mass of each element to moles using the molar mass from the periodic table.

H

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Determining an Empirical Formula from Gram Analysis

3. Divide each mole value by the smallest number of moles calculated. Round to the nearest whole number.

4. This is the mole ratio of the elements and is represented by subscripts in the empirical formula.

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Steps for Determining an Empirical Formula

1. Start with the number of grams of each element, given in the problem.

1. If percentages are given, assume that the total mass is 100 grams so that the mass of each element = the percent given.


3. Divide each mole value by the smallest number of moles calculated.

4. Round to the nearest whole number. This is the mole ratio of the elements and is represented by subscripts in the empirical formula.

5. If the number is too far to round (x.1 ~ x.9), then multiply each solution by the same factor to get the lowest whole number multiple.

1. If one solution is 1.5, then multiply each solution in the problem by 2 to get 3.

2. If one solution is 1.25, then multiply each solution in the problem by 4 to get 5.

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Determining an Empirical Formula from % Composition

NutraSweet is 57.14% C, 6.16% H, 9.52% N, and 27.18% O. Calculate the empirical formula of NutraSweet and find the molecular formula. (The molar mass of NutraSweet is 294.30 g/mol)

1. Start with the number of grams of each element, given in the problem. If percentages are given, assume that the total mass is 100 grams so that the mass of each element = the percent given.


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3. Divide each mole value by the smallest number of moles calculated. Round to the nearest whole number.

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4. If the number is too far to round (x.1 ~ x.9), then multiply each solution by the same factor to get the lowest whole number multiple.

5. This is the mole ratio of the elements and is represented by subscripts in the empirical formula.

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Determining Empirical Formula

• Vitamin C (ascorbic acid) contains 40.92 % C, 4.58 % H, and 54.50 % O, by mass. The experimentally determined molecular mass is 176 amu. What is the empirical and chemical formula for ascorbic acid?

• In 100 grams of ascorbic acid we would have: – 40.92 grams C

– 4.58 grams H

– 54.50 grams O

• This would give us how many moles of each element?

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• Determine the simplest whole number ratio by dividing by the smallest molar amount (3.406 moles in this case - see Oxygen):

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• Multiply the relative amounts of each atom by '3', to get whole numbers for each atom.

C = (1.0)*3 = 3H = (1.333)*3 = 4O = (1.0)*3 = 3

C3H4O3 = Empirical Formula

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Determining Molecular Formula from Empirical Formula

• We are told that the experimentally determined molecular mass is 176 amu or g/mol. What is the molecular mass of our empirical formula?

• (3*12.011) + (4*1.008) + (3*15.999) = 88.062 g/mol

• The molecular mass from our empirical formula is signficantly lower than the experimentally determined value. What is the ratio between the two values?

• (176 amu/88.062 amu) = 2.0

• Multiplied the subscripts in the empirical formula by '2', to get the molecular formula.

• 2* C3H4O3 = C6H8O6

• C6H8O6 = Molecular Formula

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Determining Molecular Formula from Empirical Formula

– Need to know the molar mass of the molecule

– Calculate the mass of the empirical formula

– Divide the mass of empirical formula by molar mass of molecule.

– Multiply all the atoms in the empirical formula (subscripts) by this ratio to find the molecular formula.

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Formulas (in summary)• Formulas are used to express

compounds/molecules and elements

• Types of Formulas– Empirical

• simplest whole number ratio of atoms of elements within a molecule

– Molecular• actual number of atoms within the

molecule as found in nature– Structural

• A formula in which atoms are shown in relation to each other, and bonds are shown as lines or dashes

– Formula Units• Only ionic compounds• Number of ions within the compound

CAFFEINE

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Types of Formulas

Formula Type? # of Formula Units

CaSO4

C6H12O6

CH2O

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– Chemical reactions occur when bonds between the outermost parts of atoms are formed or broken

– Chemical reactions involve changes in matter, the making of new materials with new properties, and energy changes.

– Symbols represent elements.– formulas describe compounds. – chemical equations describe a chemical

reaction.

Chemical Reactions

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Chemical Equations

• Equations explain chemical reactions.– Reactants are the substances that exist before the

chemical reaction takes place (before the arrow).– Products are the new substances that exist after the

chemical reaction takes place (after the arrow).

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Symbols used in Chemical Equationstextbook pg 206

Symbol Explanation

+ Used to separate two reactants or two products

→ Means “yields” or “produces”, separates reactants from products

↔ Used in place of → for reactions that can go either forwards or backwards

(s) Appears after the formula, the compound or molecule is in the SOLID state.

(l) Appears after the formula, the compound or molecule is in the LIQUID state.

(g) Appears after the formula, the compound or molecule is in the GASEOUS state.

(aq) Appears after the formula, the compound or molecule is dissolved in water.

∆Heat

Heat is required or supplied to the reaction

Pt Platinum is used as a catalyst for the reaction

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Law of Conservation of Mass• Mass cannot be created or destroyed. It can only change

form.

• Mass of reactants must equal the mass of the products.

• The number of atoms of each element must be the same on BOTH sides of the arrow.

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Balancing Chemical Reactions• Balanced Equation – the number of atoms of

each element on the reactant side of the arrow (left) equal the number of atoms of each element on the product side of the arrow (right) .

• To do this you can change the COEFFICIENTS ONLY! These are the numbers in front of the chemical formulas (Bold).

• NEVER CHANGE THE SUBSCRIPTS!

4 Al (s) + 3 O2 (g) ---> 2 Al2O3 (s)

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There are four basic steps to balancing a chemical equation.

1. Write the correct formula for the reactants and the products. DO NOT TRY TO BALANCE IT YET! You must write the correct formulas first. And most importantly, once you write them correctly DO NOT CHANGE THE FORMULAS!

2. Find the number of atoms for each element on the left side. Compare those against the number of the atoms of the same element on the right side.

3. Determine where to place coefficients in front of formulas so that the left side has the same number of atoms as the right side for EACH element in order to balance the equation.

4. Check your answer to see if:– The numbers of atoms on both sides of the equation are now

balanced.– The coefficients are in the lowest possible whole number

ratios. (reduced)

Steps to Balancing EquationsSteps to Balancing Equations

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Helpful Hints for balancing equations

• Take one element at a time, working left to right except for H and O. Save H for next to last, and O until last.

• IF everything balances except for O, and there is no way to balance O with a whole number, double all the coefficients and try again. (Because O is diatomic as an element)

• (Shortcut) Polyatomic ions that appear on both sides of the equation should be balanced as independent units

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Balance

5. ______HgO → ______ Hg + ______ O2

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• Which of the following reactions best illustrates the Law of Conservation of Mass?

a) H2O2 → H2O + O2

b) Na + CuS → Na2S + 2 Cu

c) K + AgCl → KCl + Ag

d) NaOH + 2 HCl → NaCl + H2O

• When sodium chloride reacts with calcium oxide to form sodium oxide plus calcium oxide, which of the following equations best illustrates the Law of Conservation of Mass?

a) NaCl + CaO → Na2O + CaCl2

b) NaCl + CaO → 2 Na2O + CaCl2

c) 2 NaCl + CaO → Na2O + CaCl2

d) 3 NaCl + 2 CaO → Na2O + 3 CaCl2

• In the following reaction: 2NaN3 decomposes to form 2Na + 3N2. If 500 grams of NaN3

decomposes to form 323.20 grams of N2. How much Na is produced?

a) 100 grams

b) 176.80 grams

c) 323.20 grams

d) 500 grams

Law of Conservation of Mass

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Types of Chemical Reactions

Composition, Addition, Synthesis or Direct Combination

– 2 1 single product is formed

– General Equation: A + X AX– element + element compound

– compound + element bigger compound

– compound + compound bigger compound

– Examples:

Fe + S FeS

2 KCl + 3 O2 2 KClO3

SO3 + H2O H2SO4

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Burning Ribbon

• List the reactants

• Observations

• Explain the products in terms of your observations

• Write the balanced chemical equation

• Explain why this is considered a synthesis reaction.

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Decomposition, Analysis

– 1 2 Single reactant is used up– 1 substance splits up into 2 or more substances.

– General Equation: AX A + X– compound element + element– big compound compound + element – big compound compound + compound – Examples

2 H2O 2 H2 + O2

NaClO2 NaCl + O2

Ca(OH)2 CaO + H2O


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Elephant Tooth Paste


• Observations



• Explain why this is considered a decomposition reaction.

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Good Bye Gummie Bear


• Observations



• Explain why this is considered a decomposition reaction.

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Combustion, Burning, Rapid Oxidation

Fuel + O2 CO2 + H2O + E

• Fuel is a hydrocarbon• Hydrocarbons are molecules that contain

only carbon and hydrogen• 2C8H18 + 25O2 16CO2 + 18H2O • Lots of energy is released• oxygen is always a reactant


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Methane Bubbles


• Observations



• Explain why this is considered a combustion reaction.

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Double Replacement, Ionic, Exchange of Ions, Metathesis

– 2 2

– compound + compound different compound + different compound

– General Equation – B+Y- + A+X- B+X- + A+Y-

– Compounds change partners– Examples

Cu+2S-2 + 2Na+1Cl-1 Na2+1S-2 + Cu+2Cl2-1

CuS + 2NaCl → Na2S + CuCl2


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Single Replacement– 2 2– element + compound different compound + different

element– General Equation: Bo + A+X- B+X- + Ao

– positives replace positives– the more active metal replaces the less active metal

Or– General Equation: Yo + A+X- A+Y- + Xo

– negatives replace negatives– the more active nonmetal replaces the less active nonmetal

Examples: Zn + CuS ZnS + Cu (metals)

F2 + 2NaCl 2NaF + Cl2 (nonmetals)


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Who’s Replacing Who?


• Observations



• Is this reaction a double or single replacement reaction?

• Explain how you know.

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• Go back to balanced eqns!

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Activity Series – Table J

• Each element from the list displaces any of the elements below it from a compound.

– The larger the space between the elements in the table, the more vigorous the reaction.

• All metals above hydrogen displace hydrogen from HCl or H2SO4.

• Metals above Mg vigorously displace hydrogen from water.

• Elements near the top of the table are never found free in nature.

• Elements near the bottom of the table are often found free in nature.

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Complete the following equations using Table J to predict the products. Write out the equation and be sure it’s balanced.

• magnesium + copper sulfate yields ?

• calcium + zinc nitrate yields?

• copper + aluminum nitrate produces?

• tin + copper nitrate produces?

• iron + tin sulfate yields?

• aluminum + silver nitrate produces?

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• Write the correct formulas for each reactant and product.• Balance the equation• Identify the type of reaction

– Liquid water decomposes to hydrogen and oxygen gases.

– Solid sodium reacts with fluorine gas to produce solid sodium fluoride.

– Aqueous Iron (III) iodide and chlorine gas yields aqueous iron (III) chloride and iodine gas.

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• Write the correct formulas for each reactant and product.• Balance the equation• Identify the type of reaction

– Aqueous calcium bromide reacts with aqueous sodium sulfide to produce the aqueous forms of calcium sulfide and sodium bromide.

– Aqueous calcium chloride combines with aqueous silver nitrate to yield solid silver chloride and aqueous calcium nitrate.

– Solid magnesium reacts in the presence of a nickel catalyst with chlorine gas to produce solid magnesium chloride.

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Predicting When a Solid is Produced from the Reaction between Two Aqueous Substances

• Solubility Guidelines – Table F• Double Replacement Reactions• Examples:

– Ag2SO4(aq) + Ca(NO3)2 (aq)

– AlCl3(aq) + (NH4)3PO4 (aq)

– Na2CO3(aq) + ZnCl2 (aq)

– NaCl(aq) + Pb3(PO4)2(s) →

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How do you know when a rxn goes to completion?

• The phrase go to completion means that the reaction continues until all reactants are used up and the reverse reaction can NOT happen.

• Conditions under which this occurs:– When a gas is formed. (Zn + 2 HCl → ZnCl2 + H2 ↑) The

hydrogen gas bubbles out of the test tube so the reverse reaction cannot occur.

– When a precipitate is formed. (AgNO3 + HCl → AgCl ↓ + HNO3)

The ions that make up silver chloride are no longer present in the reaction so the reverse reaction cannot occur.

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Stoichiometry and Chemical Equations

• Stoichiometry is the quantitative relationship between reactants and products in a chemical reaction.

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4 Al(s) + 3 O4 Al(s) + 3 O22(g) ---> 2 Al(g) ---> 2 Al22OO33(s)(s)

This equation meansThis equation means4 Al atoms + 3 O4 Al atoms + 3 O22 molecules molecules

---produces--->---produces---> 2 molecules of Al2 molecules of Al22OO33

or

4 moles of Al + 3 moles of O4 moles of Al + 3 moles of O22 ---produces--->---produces--->

2 moles of Al2 moles of Al22OO33

Coefficients = MolesStoichiometry means…

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Stoichiometry does NOT mean…

4 grams of Al + 3 grams of O4 grams of Al + 3 grams of O22

---produces---> ---produces--->

2 grams of Al2 grams of Al22OO33

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Conversions Review

• How many moles are present in 73g of HCl? (mass to mole)

• How many grams are present in 2.6 moles of NaCl? (mole to mass)

• How many molecules are present in 5 moles of glucose? (mole to molecule)

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• Using the coefficients from a balanced chemical equation in your dimensional analysis.

• How many moles of H2 are produced when 7 moles of Na are allowed to react according to the following equation?

2 Na + 2 H2O → 2 NaOH + H2

Mole to Moles Conversions

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• Use dimensional analysis to do the following:– Convert grams to moles (use formula mass)

– Convert moles to moles (use coefficients in equation)

– Convert moles to grams (use formula mass)

• How many grams of carbon dioxide are produced upon combustion of 44 grams of CH4?

CH4 + 2O2 → CO2 + 2 H2O

Mass to Mass Conversions

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Stoichiometry Problems

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