Unit 3 Molecules and Reactions

Transcription

1 Suggested Reading Chapter 2 Atoms, Molecules, and Ions ( ) Chapter 3 Stoichiometry: Calculations with Chemical Equations ( ) Chapter 4 Aqueous Reactions and Solution Stoichiometry ( ) Suggested Problems (Blackboard) Practice Exams for Unit 3 posted!! Suggested Problems (Text) Chapter 2 ( 2.9, 61, 63, 65, 69) Chapter 3 (3. 11, 13, 17, 19, 21, 23, 27, 33, 37, 43, 49, 57, 59, 61, 63, 71, 73, 75, 77, 79) Chapter 4 (4.17, 19, 21, 23, 33, 39, 41, 51, 52, 55, 57, 63, 65, 69)

2 The Mole One of the most fascinating and sometimes frustrating aspects of chemistry and chemical reactions is the fact you can t usually see what is happening. You can t physically see an atom, or a molecule, or a compound. We see a grouping of an extremely large number of these atoms in, for example, a glass of water. What if you needed to dissolve vinegar in water in exactly 1 molecule vinegar to 1 molecule water? You can t see the molecules so you can t separate them or count them, you can t rely on weighing water and vinegar, because we have seen all of the elements have a different weights and different densities, so H 2 O isn t the same as CH 3 CO 2 H. We need a way to compare the elements, compounds, and molecules. That is the basis for the mole. The mole, is essentially a way to count atoms, no matter what their size is. The word mole was introduced around 1896, from the Latin word moles meaning heap or pile. The official definition: A mole is the amount of substance that contains as many elementary particles (atoms or molecules) as there are atoms in 12.0 grams of carbon 12 isotope. First, the mole always contains the same amount of particles, no matter what the substance. How many? 1 mole = X particles This is Avogadro s number, named after Amadeo Avogadro, an Italian lawyer and physist.

3 Moles translate to mass, using the atomic mass number for each element. Thus, the molar mass is equal to the atomic weight of each element. This is essential!! 1 mole = X particles = atomic weight (g/mol) So, Let s look at Oxygen Oxygen s molar mass = 16.0 grams/mole or g/mol = X atoms of O Let s work a problem Compounds, Molecules, and Molecular Formulas The combination of elements in precise, well defined ratios into a pure substance called a compound. Any compound can by decomposed into its individual elements!! Compounds, a pure substance made up of two or more elements in a fixed ratio. The Law of Constant Composition, For example, water is 89% oxygen and 11% hydrogen, not matter where it comes from. The smallest unit of a compound which retains the characteristics of that compound is a molecule. Molecules are groups of two or more atoms held together by the forces of chemical bonds. The composition of a molecule is represented by a molecular formula.

4 Chemical Formulas We just learned the chemical formula (molecular formula) is a the smallest unit in a substance which still possesses the properties of that substance. All molecules has a strict ratio of elements which is maintained through out the substance. There is a term for this relationship, stoichiometry, which is the quantitative relationship between elements in a molecule is composition stoichiometry. We are now going to look at the quantitative relationship between reactants and products, or reaction stoichiometry. Molecular formulas gives us the number of elements in the molecule, but not how they are bonded together. The structural formula shows us the connectivity of the atom. For example, C 3 H 8, is propane. The structural formula shows which atoms are bonding with which atoms by drawing the molecule like this, with each line showing a chemical bond.

5 What if we are a chemist and we were given a sample of colorless liquid, could we find the molecular formula? Yes. Elemental Analysis is an established qualitative analysis that gives us the % composition of the elements in that powder. Let s say we had the elemental analysis done, and the results told us.. H 1.6% N 22.2% O 76.2% Can we get to the molecular formula? Yes. HNO 3 Let s clarify here. The empirical formula is the simplest formula, meaning the lowest possible ratio equivalents, like BH 3 The molecular formula is what the molecule actually exists as., B 2 H 6. So if we have 100g of our sample: We convert each to moles. Then we find the ratio Notice we used moles here. Let s do another problem here.

6 Chemical Equations Remember that a physical change in a substance does not change the chemical composition, or molecular formula of the substance. For example, boiling water does not change the molecular formula of water. However, when you ignite methane (light your stove burner), you are using up the methane and forming carbon dioxide instead, this is a chemical change. First, let s review the Law of Conservation of Matter There is no observable change in quantity of matter in a chemical reaction or a physical change. Second, let s review the Law of Conservation of Energy Energy can not be created or destroyed in a chemical reaction or physical change. It can only be converted from one form to another. Third, the Law of Conservation of Matter and Energy The combined amount of matter and energy in the universe is fixed.

7 So what? What does this mean for us? Chemical equations are used to describe a chemical reaction(or change). Chemical equations follow the Laws of Conservation of Matter and Energy!! Chemical equations show the molecular formulas of the substances reacting, called reactants, and the molecular formulas of the substances being formed, called products. They are written from left to right with reactants always on the right and products always on the left. Balancing Equations Because chemical equations must not create or destroy matter the number of each of the elements must remain the same!! This means each side of the chemical equation must be balanced to equal the other side. Note the coefficients are whole numbers that indicate how many of each molecule is needed to react. The subscript numbers are representative of the actual molecular formula, these numbers do not change. In balancing reactions, we only balance the number of molecules reacting and never change the molecules reacting. These equations are based on experimental findings!! We can not change our observations to fit our equations it is always the other way around. Let s look at the reaction of burning diethyl ether. Unit 3

8 Now that our equation is balanced what does it tell us? A balanced reaction equation tells us the number of moles for each molecule!! Let s look at our stove burner again, CH 4 + 2O 2 CO 2 + 2H 2 O 1 mol 2mol 1mol 2mol This concept is extremely important to remember!! We will use this time and again in solving problems for the rest of the course. Limiting Reactants Concept Unless your reaction has an infinite amount of reactants the reaction will end as soon as one of the reagents is used completely. This reagent is called the limiting reagent. When you light your stove, methane is being supplied and the reaction is proceding, but when you turn of the burner and the flow of methane, the reaction consumes the rest of the methane and then the burner turns off. In this case, the methane is your limiting reagent. Here is our reaction. CH 4 + 2O 2 CO 2 + 2H 2 O If we are given a finite amount of CH 4 and O 2, which will run out faster? Unit 3

9 Percent Yield You will always need to convert to moles for your percent yield!! Percent yield = moles of product/ moles of reactant The book shows something a little different here, but keep it simple always use moles. Theoretical yield simple means the yield if all of your reactant is converted to your product. Always on paper, never in the lab!! Actual yield is the yield of product you actually make in the lab.

10 Binary Compounds When H is bonded to a Group 6A or 7A the resultant molecule is an acid in aqueous solutions, and is called a binary compound. These are similar to the naming of ionic compounds, except H is always the positive ion. So HCl, is hydrogen chloride But when dissolved in water.. HCl is hydrochloric acid Ternary Acids These are the oxoacids. When a molecule consists of hydrogen, oxygen, and one of the metals, this is a ternary acid. Again, we haven t learned enough about bonding to recognize the structural elements of this acid, but there is a chart in your book (pg. 137) which lists the common acids. The important thing to remember here is that ic is assigned arbitrarily to these acids..so use the chart. One less oxygen in the molecular formula means the name changes to ous, i.e. H 2 SO 3 sulfrous acid. One less oxygen in the molecular formula means hypo is added to the beginning, i.e. H 2 SO 2, would be hyposulfurous acid. Basically, this section can be extremely confusing!! Just remember that names of compounds are designed to enable you to write the formula from the name and vice versa.

11 Ionization Reactions Compounds that are water soluble and separate into its ions in solution are said to dissociate upon dissolving, or undergo a dissociation reaction. Here s an example. Another dissociation reaction is the addition of HX, like HCl to water. When the cation is H + the reactant is an acid, and the solution is acidic!! There are strong acids like HCl and weaker acids like HCN. What makes a strong acid, strong? A weak acid, weak? How about NaOH in water? This is a strong base. When a compound dissociates to give OH, the compound is a base and the solution is basic. There are also strong and weak bases. Remember the arrows on the reaction, they can mean different things.i.e. how strong is the reaction and is it reversible? How do you get a reaction that is not reversible? Let us look at our Chloride Lab. We formed a white ppt. (AgCl) this was insoluble in water and stopped reacting!!

12 Aqueous Solutions Molarity We need a way for practical laboratory practices to measure reacting species in solution. One way is to use molarity. M = mol/l Remember you can use molecular weight to calculate moles and density for conversion of liquids to volume. Let s try a problem Reactions!! There are three main forms to write many of these equations. The first is the formula unit equation. This means we write the complete formula for each compound!! The second is the total ionic equation, This means we can write each compound as its reacting ion. The third is the net ionic equation. This is just the reacting ions.

13 Oxidation Reduction Reactions Redox Reactions Oxidation refers to increasing the oxidation number!! This means the atom is losing electrons!! Reduction refers to reducing the oxidation number!! This means gaining electrons. But, remember the oxidation number in a molecule must add up to zero.the same will be true for a reaction. This means, oxidation and reduction must both occur in the same reaction with equal but opposite oxidation states!! Remember the big picture here, redox reactions are the exchange of electrons, we can not gain or lose the total number of electrons present, only move them around. For example,

14 Combination Reactions Reactions in which two or more substances combine to form a compound are called combination reactions. There are three main classes of these reactions; Element/Element Compound/Element Compound/Compound Decomposition reactions These are reactions in which compounds break down into one or more elements and one or more compounds, or two or more compounds. Essentially these are the reverse of combination reactions. Element/Element Compound/Element Compound/Compound

15 Displacement Reactions Reactions in which one element replaces another in a compound are called displacement reactions. For example; How can we predict if this will happen? Active metals displace less active metals!!

16 Metathesis Reactions A reaction in which the anions switch cation partners. For example, An acid base reaction often precipitation forming reactions or gas forming reactions