NOTES PACKET COLLIER CHEMISTRY PRE-AP
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1 SECOND NINE WEEKS NOTES PACKET COLLIER CHEMISTRY PRE-AP 1
2 2
3 UNIT 5 CHEMICAL NAMING & BALANCING Chapter 6, 15.1,
4 NOMENCLATURE: Atoms of elements combine to form that are represented by. All compounds can be in such a way that others can understand the make-up of the compound. Chemical formulas are represented by using the symbols with that show the number of each element in a given compound 4
5 Chemical Bonding Molecules the smallest electrically neutral unit of a substance. 1. are the building blocks of substances. 2. Atoms of most elements will different compounds (except ) 3. Molecules are made up of that act as one unit. Ex: 4. Molecular compounds can be called or molecules. 5. Molecular Compounds are composed of molecules that are. 5
6 Ions & Ionic Compounds 1. are atoms or a group of atoms that have 2. A loss of electrons gives the atom a positive charge; therefore a gain of an electron would give the atom a negative charge. 3. Atoms of elements tend to electrons to become and tend to electrons to become. (+) ions also called. Ex: (-) ions =. Ex: 4. Ionic compounds consist of a bonded to a. The charge of the compound is. 6
7 OXIDATION NUMBER This is a or number assigned to an atom according to a set of rules The oxidation number of a ion is equal to its ionic charge Ex. Br- = Mg +2 = The oxidation number of in a compound is except in a metal hydride where it is The oxidation number of is except in where it is The oxidation number of the atom in its form is 7
8 Electrons in Ionic Bonding Valence Electrons Determines the properties of an element. 2. The determines the number of valence electrons 3. Electron Dot Structures that show only valence electrons. Octet rule -- atoms tend to achieve the. **Atoms will gain or lose electrons to achieve this. 8
9 Ionic Bonds -- the of attraction between. ( & ) Covalent Bonding -- a bond in which two atoms a pair of electrons. ( ) 1. Structural formula H H 2. Sharing of electrons occur if atoms involved achieve the rule. 3. and Covalent Bonding means that more than one pair of electrons are shared. Example: 9
10 Chemical Compounds Chemical Formula s -- Shows the and of atoms in a compound. - show the number of atoms. Ex: H 2 O Ionic compounds consist of a ion with a ion in the lowest whole number ratio. Ex: Molecular compounds typically consist of covalently bonded together Ionic Charges ions consist of only 1 ion. ions consist of tightly bound atoms that behave as a single unit, that has either lost or gained electrons to become (+) or (-). Ex: NO 3-, SO
11 Writing Formulas Ionic Compounds Criss Cross Method Balancing Formulas 1. Write and for each ion present. Compare charges: If charges equal then simply write symbols together. If charges don t equal 0 then and write them as a. Criss-Cross Method Al +3 + O -2 Al 2 O 3 Examples: Sodium Chloride: Aluminum Sulfide: 11
12 Binary Molecular Compounds - Consists of two nonmetals Prefixes 1. Write the symbols of the elements, change to subscripts. Mono = Hexa = Di = Hepta = Tri = Octa = Tetra = Nona = Penta = Deca = Acids Step 1: 1. Acids consist of a bonded to an (negative ion ). Step 2: 12
13 Examples: - Hydrochloric Acid - Nitric Acid - Nitrous Acid Step 3: for this compound as you would when writing any molecular formula. Hydro- ; -ic = -ide -ic = -ate -ous = -ite 13
14 14
15 Ionic Compounds Naming Compounds 1. Write the of the present. 2. Determine which forms the compound by looking at the ratio of the atoms and compare to the ratio of the charges. 3. Remember that the overall charge of the compound has to. 4. Example: CuO is it copper I or II? The ratio of atoms is 1:1 and the charge on O is 2 so the Cu has to be 2 to have the same ratio. The name of the compound would be 15
16 Binary Molecular Compounds: 1. Write the name of the in the first position and if it has a subscript change it to a. 2. Write the name of the element. Change the ending to, and always use a that corresponds with the number of atoms present in the molecule. Example: PO 5 = S 2 N = 16
17 Acids **If the compound starts with an H then it considered to be an acid. Step 1: Step 2: - If the ending is, place in front of the negative ion name, then drop the ide and add. - If the ending is, drop the -ate ending and add. - If the ending is, drop the ending and add. Examples: H 2 S = H 2 SO 4 = H 2 SO 3 = Hydro- ; -ic = -ide -ic = -ate -ous = -ite 17
18 18
19 Molecular Weights 19
20 20
21 UNIT 6 Chemical Quantities--Chapter 7 Atomic Mass (Weight) The of an expressed in. Examples: Carbon Sulfur Iron Mass of a Mole of a Compound 1 st you must of the compound The formula tells you of each element that is present. Example: SO 3 21
22 After you know the then you calculate the of the compound. Molar mass Find the molecular mass of each compound. 1. Li 2 S 2. FeCl 3 3. Ca(OH) 2 4. N 2 O 5 22
23 Percent Composition Grams of element % mass of element = Grams of Compound X 100 Or % Comp. = Part X 100 Whole Practice Problem Ethane (C 6 H 6 ) g of C 6 H 6 mole of C = g mole of H = 6.06 g Find the % Composition of Carbon & Hydrogen in Ethane (C 6 H 6 ). 23
24 Significant Figures & Scientific Notation All the that can be known precisely in a measurement, plus a digit. Expression of numbers in the form where n is equal to or greater than one and less than 10 H 2 O How many much Hydrogen is in Water? What is Hydrogen s % Composition in Water? What would happen to the % of hydrogen with 2 molecules of Water? 24
25 SIGNIFICANT DIGITS To determine the number of significant digits in a written number complete one of the following: 1. Qualifying statement: The numerical value has no decimal place written. Action to take: Count from the first non-zero digit to the last non-zero digit. Examples: 1. 18,004-5 significant digits 2. 10,040,000 4 sig digs sig dig 2. Qualifying statement: The numerical value contains a decimal place. Action to take: Count from the first non-zero digit to the end of the number. Examples: sig digs 2. 1, sig digs sig digs 25
26 What is a mole? Avogadro s number = Standard Temperature and Pressure (STP) a. Standard temperature is or b. Standard pressure is. 1 mol of any gas at STP = 26
27 Representative Particle Refers to the in a substance: usually atoms, molecules, or formula units. (ions) Example: Fe is composed of atoms. K is composed of atoms. Example: How many moles of Mg is 1.25 x atoms of Mg? Know: Unknown: 27
28 Practice Problems How many moles are 2.80 x atoms of silicon? How many molecules is moles of water? Amt. in Compounds Carbon dioxide has 3 atoms. 1-carbon 2-oxygen Thus one mole of CO 2 contains three times Avogadro s # of atoms. 28
29 Example: How many atoms are in 2.12 mole of propane (C 3 H 8 )? Know: Unknown: Desired conversion: 29
30 Practice Problems How many atoms are there in 1.14 mole SO 3? How many moles are there in 4.65 x molecules of NO 2? Mole -- gram How many grams are in 3.32 mole K? How many grams are in 4.52 x 10-3 moles of C 20 H 42? 30
31 Gram -- Mole How many moles are in 3.70 x 10-1 grams of Boron? How many moles are in 27.4 g TiO 2? 31
32 32
33 Unit 7 Balancing Chemical Reactions & Predicting Products Ch. 8 &
34 Reactions 34
35 Equations An is a description of a chemical reaction indicating the, the and a of their quantities. REMEMBER: (reactants) = (products) Coefficients -- The reaction is said to go to when no reactants remain when the reaction has stopped. 35
36 The reaction may be. When this happens, a complete cycling of events exists. It is a continuous ongoing process. The Law of Conservation of Mass: in reactions atoms are neither nor only The Law of Conservation of Mass. What is the relationship between the mass of the reactants and the mass of the products? How are chemical reactions balanced? 36
37 are used to separate reactants or products on the same side of the arrow. Catalyst: 2H 2 + O 2 2H 2 O 1. What are the reactants? 2. What are the products? 3. What are the coefficients for the whole reaction? 37
38 Rules for Balancing Equations: 1. Determine the for all reactants and products. Indicate physical states in parentheses. 2. Write the formulas for the and of the arrow. Then separate two or more formulas with + signs. 3. of each element for both sides of the equation. 4. the elements one at a time by using. 5. Check each atom to be sure that the equation is balanced. 6. Make sure the are in the that balances. Diatomic Elements 38
39 Classifying reactions Synthesis (Combination) reaction: The can be elements or compounds. The will always be a compound. A + B AB Decomposition reactions: The is always a. Most decomposition Reactions requires energy such as light, heat or electricity. AB A + B 39
40 Single replacement reactions: Single elements changes places with the other ion that is the same as the single element. & A + BC AC + B Double Replacement: AB + CD CB + AD Combustion Reactions: 40
41 Factors Effecting Reaction Rates: Temperature: Concentration: Catalyst: Inhibitor: Neutralization: 41
42 Predicting Products 1. To predict products of any reaction the first step is to it is. 2. Then apply existing rules for the type of reaction it is: Activity Series of metals: an invaluable aid to predicting the products of replacement reactions. It also can be used as an aid in predicting products of some other reactions. Going from bottom to top, the metals: 42
43 Metal Metal Ion Reactivity Lithium Potassium Calcium Sodium Magnesium Aluminum Manganese Zinc Chromium Iron Lead Copper Mercury Silver Platinum Gold Li + K + Ca 2+ Na + Mg 2+ Al 3+ Mn 2+ Zn 2+ Cr 2+,Cr 3+ Fe 2+,Fe 3+ Pb 2+ Cu 2+ Hg 2+ Ag + Pt 2+ Au +,Au 3+ Most Reactive Least Reactive 43
44 Synthesis (Combination): + Combine the 2 elements into a compound Practice: 1. Aluminum + Oxygen 2. Chlorine + Potassium 3. Chromium + Bromine 4. Oxygen + Hydrogen 5. Calcium + Iodine 44
45 Single replacement: + If the element forms a + ion replace the + ion of the compound, and if it is a - ion the replace the ion of the compound. Practice: 1. Aluminum + Magnesium oxide 2. Copper (II) sulfide + chlorine 3. Fluorine + Nickel bromide 45
46 Double Replacement: + Switch the position of the (+) ions. Practice: 1. Aluminum sulfite + Calcium acetate 2. Tin (IV) phosphate + Ammonium dichromate 3. Magnesium chloride + Chromium peroxide 46
47 Metallic Chlorates: Decomposition Reactions + Substitute the specific + ion for the M and follow the word equation. Metallic Carbonates: + Substitute the specific + ion for the M and follow the word equation. Metallic Hydroxides: + Substitute the specific + ion for the M and follow the word equation. 47
48 Oxy Acids: Example: Simple Decomposition: + Split the compound into is simpler parts. 1. Potassium sulfide Practice 2. Zinc oxide 48
49 3. Zinc carbonate 4. Sodium carbonate 5. Calcium hydroxide 6. Aluminum hydroxide 7. Potassium chlorate 8. Calcium chlorate 49
50 50
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