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Chemical Bonds and Formulas

Chemical Bonds and Formulas.

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Chemical Bonds and Formulas

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  1. Chemical Bonds and Formulas • Have you ever noticed the color of the Statue of Liberty? Although the color is green, it was not chosen to be green at all. In fact, the Statue of Liberty is made of the metal copper, an element (Cu), just like a penny. BUT, copper isn’t green; its copper colored! So why is the Statue of Liberty green?

  2. Some matter is uncombined elements such as copper, sulfur, and oxygen. But often, like many other sets of elements, these elements chemically combine to form a compound. The green coating on the Statue of Liberty and some old pennies is a result of this chemical change (color change)—a new substance is formed, copper sulfate. The new substance has different properties from the original elements.

  3. Example: Copper is shiny copper-colored. Sulfur is a pale-yellow solid. Oxygen is a colorless, odorless gas. Copper sulfate, however, is the green color we see on the Statue of Liberty.

  4. Chemical Formula—tells what elements a compound contains and the exact number of the atoms of each element in a unit of that compound. Example: H2O = water The formula contains the symbol H for hydrogen and O for oxygen. The number 2 is a subscript meaning “written below”. A subscript written after a symbol tells how many atoms of that element are in a unit of the compound. If no subscript is written, there is only one atom of that element.

  5. Look at the formulas for each compound listed in the table below. Fill in the remaining boxes.

  6. Counting Atoms Notes How to count the number of atoms in a compound:

  7. The number of each element in a chemical formula is represented by the subscript to the right of the symbol. If there is no subscript, that means there is only one. H2O 2 atoms of hydrogen 1 atom of oxygen Ca3N2 3 atoms of calcium 2 atoms of nitrogen

  8. Chemical formulas will also have coefficients. These are the large numbers in front of the formula. Any coefficients placed in front of the formula multiply with the number of each element in the formula (subscripts). 2 H2O2 X 2 Hydrogen = 4 Hydrogen 2 X 1 Oxygen = 2 Oxygen 3 Ca3N23 X 3 Calcium = 9 Calcium 3 X 2 Nitrogen = 6 Nitrogen

  9. Chemical formulas can also have parentheses with subscripts. Parts of the formula are sectioned off in parenthesis with subscripts to the right of the parenthesis. These subscripts will multiply with the number of each element inside of the parenthesis only. Mg(OH)2 1 atom of Magnesium 2 x 1 Oxygen = 2 Oxygen 2 x 1 Hydrogen = 2 Hydrogen Ca(NO3)21 atom of Calcium 2 x 1 Nitrogen = 2 Nitrogen 2 x 3 Oxygen = 6 Oxygen

  10. Some chemical formulas can have parentheses, subscripts, and coefficients. For these formulas, take it one step at a time. Multiply the subscripts with the parenthesis first, and then multiply the coefficients. 3 Cu(NO3)23 x 1 Copper = 3 copper 2 x 1 Nitrogen = 2 Nitrogen x 3 = 6 Nitrogen 2 x 3 Oxygen = 6 Oxygen x 3 = 18 Oxygen 2 Al2(SO4)32 x 2 Aluminum = 4 Aluminum 3 x 1 Sulfur = 3 Sulfur x 2 = 6 Sulfur 3 x 4 Oxygen = 12 Oxygen x 2 = 24 Oxygen

  11. Atomic Stability Atoms combine to form compounds when the compound formed is more stable than the separate atoms. An atom is chemically stable when its outer energy level is full or complete (Octet Rule). Recall that the outer energy levels of helium and hydrogen are stable with 2 electrons. All other elements are stable when their outer energy level contains 8 electrons. The noble gases are unusually stable and rarely form compounds because they each have a complete outer energy level.

  12. Outer Levels—Getting their Fill Atoms with partially stable outer energy levels can lose, gain, or share electrons to obtain a stable outer energy level. They do this by combining with other atoms that also have partially complete outer energy levels. As a result, each element achieves stability.

  13. Na Cl + Example: Electrons lost/gained

  14. Example: Electrons shared Chemical Bond—the force that holds atoms together in a compound

  15. Types of Bonds—Ionic and Covalent Ionic Bonding—(gaining/losing electrons) Ion—an atom that has lost or gained electrons; charged particle

  16. Example:When potassium forms a compound with iodine, potassium loses one electron from its fourth level, and the third level becomes a complete outer level. However, the atom is no longer neutral. The potassium atom has become an ion. When a potassium atom loses an electron, the atom becomes positively charged because there is one electron less in the atom that there are protons in the nucleus. Iodine gains an electron thus gaining a negative charge. The 1+ and 1- charge is shown as a superscript, meaning “written above”, written after the element’s symbol— K+ or I-.

  17. Ionic Bond—the force of attraction between the opposite charges of the ions in an ionic compound. In an ionic bond, a transfer of electrons takes place. If an element loses electrons, one or more elements must gain an equal number of electrons to maintain the neutral charge of the compound. Ionic bond (with more than 1 electron transfer)

  18. Covalent Bonding—(sharing electrons) Some atoms of nonmetals are unlikely to lose or gain electrons because it would take a great deal of energy to gain or lose more than a few electrons at a time. Covalent Bond—the attractions that forms between atoms when they share electrons Single Covalent Bond • Multiple Covalent Bond

  19. Polarity Unequal sharing of electrons in a compound occasionally occurs causing one end of a molecule to be slightly positive and the other end to be slightly negative. Even though the charge is balanced it is not equally distributed. Polar Molecule—a molecule that has a slightly positive end and a slightly negative end although the overall molecule is neutral. Example: Ammonia and water are examples of polar molecules.

  20. Binary Ionic Compounds Binary Compound—a compound composed of two elements Oxidation number—tells how many electrons an atom has gained, lost, or shared to become stable For ionic compounds, the oxidation number is the same as the charge on the ion. Example: A sodium ion has a charge of 1+ and an oxidation of 1+. A chlorine ion has a charge of 1– and an oxidation of 1–.

  21. There are some special ions that can have more than one oxidation number. When naming these compounds, the oxidation number is expressed in the name with a roman numeral. Example: The oxidation number of iron in iron (III) oxide is 3+.

  22. When writing formulas it is important to remember that although the individual ions in a compound carry charges, the compound itself is neutral. A formula must have the right number of positive and negative ions so the charges balance. Example: Sodium chloride is made up of a sodium ion with 1+ charge and a chlorine ion with a 1– charge. One of each ion put together makes a neutral compound with the formula NaCl.

  23. WRITING FORMULAS FOR IONIC BONDS using the Criss-Cross Method • Assign oxidation numbers for each element involved in the ionic bond. • Cross the numbers so they sit on the bottom right side of the element. • Do NOT include the signs. • Reduce numbers to the lowest common denominator and drop the ones.

  24. COMPLETE THE TABLE: The top two parts have been completed for you as examples. Once the atoms have bonded their charge is ZERO or neutral.

  25. Na+1 bonding with PO4-3 Na3PO4 Polyatomic Ions There are groups of non-metal atoms that act like a single non-metal atom with a single charge. To write formulas with polyatomic ions, put parentheses around the polyatomic ion before adding a subscript.

  26. Law of Conservation of Mass Burning is an example of one the most familiar chemical changes. When you burn a lump of coal, atmospheric oxygen combines with the coal. The combining or recombining of elements differently is called a chemical reaction. The chemical reaction produces carbon dioxide gas, water vapor, and a small amount of ash. This change seems to involve a reduction in the amount of matter. A sizable piece of matter seems to have produced only a small amount of ash. However, careful measurements show that the total mass (grams) of the original coal and oxygen EQUALS the total mass (grams) of the carbon dioxide, water vapor and ash produced.

  27. Measurements of many chemical reactions lead to the LAW OF CONSERVATION OF MASS. This law states that matter cannot be created nor destroyed. This law applies to chemical reactions where elements on the left side of an equation must appear on the right side of the same equation. Not only must the elements be the same on each side, but the number of atoms for each element must be the same on each side, too.

  28. CHEMICAL REACTION represented by a CHEMICAL EQUATION CaCl2 + F2→ CaF2 + Cl2 REACTANTS(grams) = PRODUCTS(grams) Example 2KClO3 2KCl + 3O2 110g  50g + 60g 110g – 50g = 60g

  29. . 2H2 + O2 2H2O 150g + 10g  _____g • . 4Al + 3O2  2Al2O3 _____g + 75g  150g • . 2NaCl  2Na + Cl2 _____g  15g + 15g • . N2 + 3H2 2NH3 75g + _____g  200g

  30. How many grams of sodium iodide (NaI) are produced by the following chemical reaction? 2Na + I2 2NaI 37g + 52g  ________g • Dihydrogen dioxide (H2O2) decomposes into dihydrogen (H2O) monoxide and dioxide (O2). How much dioxide will be produced from the grams of dihydrogen dioxide (H2O2) indicated below? 3000g 2H2O2 +  1100g H2O + ________g O2

  31. Mg + O2 → MgO Mg = 1 Mg = 1 O = 2 O = 1 Since the O atoms are not equal, target those first. Balancing Equations Step-by-Step: Step 1: Determine number of atoms for each element. Mg + O2 → MgO Mg = 1 Mg = 1 O = 2 O = 1 Step 2: Pick an element that is not equal on both sides of the equation.

  32. Adding a 2 in front of MgO will change the number of atoms on the product side of the equation. Mg + O2 → 2MgO Mg = 1 Mg = 12 O = 2 O = 12 2 Mg + O2 → 2MgO Mg = 12 Mg = 12 O = 2 O = 12 Now we need to increase the number of Mg atoms we have on the reactant side. Adding a 2 in front of Mg will give us 2 atoms of Mg and balance the equation. Step 3: Add a coefficient in front of the formula with that element and adjust your counts. Step 4: Continue adding coefficients to get the same number of atoms of each element on each side.

  33. Practice: Na Cl Be F Na Cl Be F Reactants Reactants Products Products Mg Mn O Mg Mn O • 1. __NaCl + __BeF2→ __NaF + __BeCl2 2. __Mg + __Mn2O3→ __MgO + __Mn

  34. 3.__CH4 + __O2→ __CO2 + __H2O C H O C H O Reactants Products Reactants Products Fe Cl Be P O Fe Cl Be P O 4. __FeCl3 + __Be3(PO4)2→ __BeCl2 + __FePO4

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