Unit 11: Chemical Reactions Chapter 10 in text

Chemical Formulas 1. Empirical formula 2. Molecular formula Identifies elements present in terms of simplest whole number _________ Examples: table salt, NaCl; water, ________ 2. Molecular formula Identifies ________ number of atoms in a molecular compound Example: water, H2O; Glucose ____________ Not table salt, NaCl (ionic compound) ratio H2O exact C6H12O6

Chemical Formulas 3. Structural formula Represents ________________ of atoms within a molecule Related to _______ structure of molecule arrangement 3 – D

Empirical or molecular formula? Ionic - lacking discrete unit, or molecule Composed of __________ and nonmetallic elements Electronegativity difference > _______ Molecular ____________ compounds Usually nonmetals bound to _______________ Molecular and empirical formulas can be different Glucose – molecular C6H12O6 vs. empirical _____________ metallic 1.7 Covalent other nonmetals CH2O

Molecular & Formula Weights ___________ of atomic weights of all atoms in chemical formula Molecular weight Formula weight of a _______________ substance Term often used for non-molecular (________) substances Total sum molecular ionic

Percent composition of compounds Finding the mass percentage of an individual _________ from the formula weight element

Calculate Empirical Formula Divide the percent composition or mass of each element by their ______________. Divide the number of each element you just found by the ___________ number you just found. If you have a ________ such as 0.5, multiply all quotients by 2 0.25 multiply by 4, and 0.33 multiply by 3. The answers will be the ____________ of the empirical formula. The empirical formula should be the simplest __________ number ratio of the subscripts. atomic mass smallest decimal subscripts whole

Empirical Formulas Divide each mass by the atomic mass 2.3g Na / 23 = __________ 0.8 g O/ 16 = ___________ Divide the quotient of each element by the smallest quotient. Na - 0.1/_______ = _____ O - 0.05/______ = _____ The products will be the subscripts of the empirical formula. ___________ 0.1 0.05 2 0.05 0.05 1 Na2O

Empirical Formulas Divide each mass percentage by the mass number Divide the quotient of each element by the smallest quotient. S - 1.56/______ = _______ O - 3.125/ _______ = ______ The products will be the subscripts of the empirical formula. __________ 1.56 3.125 1.56 1 1.56 2 SO2 Sulfur dioxide

Chemical reactions Occur through formation and ___________ of chemical bonds between atoms Involve changes in matter, creation of __________ materials and __________ exchanges Chemical equations - concise _______________ of chemical reactions breaking new energy representations

Chemical equations Reactants - substances existing _________ reaction Products - substances existing __________ reaction Word representation only is not precise Chemical symbols and formulas needed for ____________ purposes before after quantitative

Balancing equations Law of conservation of mass - atoms are neither ________ nor destroyed in chemical reactions Mass of the products = Mass of reactants No mass defect, like with nuclear reactions Any _________ released was stored in the bonds! Change ____________ in front of chemical formulas, not ___________ within formulas, to balance created energy coefficient subscripts

Types of Reactions Decomposition reactions Combination reactions Replacement reactions - Single ion replaces another ion in a compound Ion exchange reactions - Two cations switch places in their compounds

AB + CD  AD + CB Positive ions change/swap which anion they are bonded with (i.e. A was with B and A is now with D)

Types of chemical reactions Oxidation-reduction (redox) reactions Oxidation - ___________ of electrons by an atom Reduction - ___________ of electrons by an atom __________ agents - substances taking electrons from other substances (oxygen, chlorine) _______________ agents - supply electrons to oxidizing agents (hydrogen, carbon) loss gain Oxidizing Reducing

Chemical Reactions Atoms are conserved – Atoms do not get used up in a __________ reaction Burning a piece of paper is a combustion reaction The atoms in the paper and the __________ in the air react to form the ash, smoke, and the CO2 that exist after the reaction occurs Mass is ____________ – Same number and type of atoms, same mass is present Law of combining volumes (gases) Gases at the same temperature and pressure contain equal numbers of molecules chemical oxygen conserved

Units of Measurement Used with Chemical Equations Atomic mass unit = 1 a.m.u. (1 u) Roughly the mass of one _________ or one ___________ 1 u = 1.661 x 10 -27 kg Defined as 1/12 mass of carbon-12 ______ The Periodic Table lists the atomic masses of each element in ______ Since we deal with more than one atom at a time, we need a way to go from the macroscopic scale to the atomic scale and vice-versa… proton neutron atom a.m.u.

The Mole The Mole – one mole of carbon-12 has a mass of 12 grams and contains 6.02 x 1023 atoms of carbon A mole is a number of something, like a dozen means ____ of something or a gross is _____ of something – Could you have a mole of cars? ______ 6.02 x 1023 is known as Avogadro’s number and is the number of things in a mole of that thing One mole of cars would be ____________ cars A mole of carbon-12 atoms is defined as having a mass of _____________ 12 144 Yes! 6.02 x 1023 12 grams

Molar weights ____________ atomic weight - mass in ___________ equal to atomic weight ______________ formula weight - mass in grams equal to formula weight (ionic) Gram molecular weight - mass in grams equal to molecular weight of the _____________ Gram grams Gram compound

Quantitative Use of Equations Possible interpretations: Molecular ratios of reactants and products Mole ratios of reactants and products Mass ratios of reactants and products

Stoichiometry The study of the amount of reactants needed and products generated in a _________ _________. There are _________ steps needed to solve a Stoichiometry problem. If your initial units are __________ and your final units are moles, skip steps ___ and ___. You only need the mole-mole ratio from the coefficients of the chemical reaction. chemical reaction five moles 2 4

Stoichiometry Write a complete and _____________ chemical equation. Convert quantity of given into __________. Multiply by the mole _____ taken from the ________________ in the chemical equation. Convert from moles into the desired unit. _____________ units to ensure that correct conversion factors were used. balanced moles ratio coefficients Cancel out

Stoichiometry What volume of ammonia is produced if 25.0g of nitrogen (diatomic molecule) is reacted with excess hydrogen? Write a complete and _____________ chemical equation. _______________________________________ Convert mass of nitrogen into moles by dividing by the molar mass. balanced N2 + 3H2  2NH3 x 1mol N2 28 g N2 25.0 g N2

Stoichiometry Multiply by the mole ratio. ________________________________________ Convert moles into the desired units Cancel units. Answer: _______________ x 1mol N2 28 g N2 x 2 mol NH3 1 mol N2 25.0 g N2 x 1mol N2 28 g N2 x 2 mol NH3 1 mol N2 x 17.0 g NH3 1 mol NH3 25.0 g N2 x 1mol N2 28 g N2 x 2 mol NH3 1 mol N2 x 17.0 g NH3 1 mol NH3 25.0 g N2 30.4 g NH3

Limiting Reagents A limiting reagent is a reactant that is in less molar quantity or the reactant that determines the quantity of _________________. For example, if we react 2.5 grams of sodium with 300,000 grams of chlorine, we wouldn’t get 300,002.5 grams of sodium chloride We would run out of sodium atoms before we ran out of chlorine atoms Therefore, the sodium “limits” the amount of sodium chloride that can be produced and is our limiting reagent product produced

Limiting Reagents Problem: How many grams of NH3 can be produced from the reaction of 20 g of N2 and 6 g of H2? Calculate the moles of each reactant. _______________________________________ 20 g N2 x 1 mol N2 28 g N2 = 0.713 mol N2 x 1 mol H2 2 g H2 6 g H2 = 3 mol H2

Limiting Reagents According to the equation, N2 + 3H2  2NH3 For every 1 mole of nitrogen, I would need ________ moles of hydrogen. However, I have been given 0.713 moles of N2 and 3 moles of H2. To use all the N2, I need 3 x 0.713 = 2.14 moles of H2. There would be __________________ moles H2 remaining. Therefore, nitrogen is the ___________ reagent and hydrogen is the __________ reagent. three 3 – 2.14 = 0.86 limiting excess

Limiting Reagents This means my Stoichiometry must be based on the quantity of nitrogen since I will need all of it producing the Ammonium. _______________________________________ 0.86 mol H2 x 2 g H2/1 mol H2 = 1.72 g H2 remains How much hydrogen was used? Answer: ______________________ x 1mol N2 28 g N2 x 2 mol NH3 1 mol N2 X 17.0 g NH3 1 mol NH3 20.0 g N2 = 24.28 g NH3 6 g H2 – 1.72 g H2 = 4.28 g H2