Physical Science Chapter 20 Chemical Bonds 1

Let’s Review: Transition Elements 1 18 1 2 13 14 15 16 17 2 3 3 6 4 5 10 11 12 9 4 5 6 7

20:1 Stability in Bonding— Formulas The chemical symbols Na and Cl represent sodium and chlorine. When written as NaCl, this is a chemical formula.

20:1 Stability in Bonding— Formulas A. Some elements combine chemically and no longer have the same properties they did before forming a compound.

20:1 Stability in Bonding— Formulas A chemical formula is composed of symbols and subscripts indicating the number of atoms of an element in a compound. “Subscript” means written below.

20:1 Stability in Bonding— Formulas Familiar Compounds: Let’s look at Table 1 on page 603. Add magnesium oxide to this list. Magnesium oxide MgO (used in laxatives, antacids, fireproofing, etc.)

20:1 Stability in Bonding— Formulas Remember: A compound is a substance formed from two or more elements in which the exact combination and proportion of elements is always the same. Video 2—Atoms and Bonding 1.40 Video 3—Electrons and Energy Levels 1.01

20:1 Stability in Bonding— Formulas C. Atoms form compounds when the compound is more stable than the separate atoms. Video 4—Stability and Chemical Bonds 1.37

1. Noble gases. are more. chemically. stable than. other. elements 1. Noble gases are more chemically stable than other elements because they have a complete outer energy level.

20:1 Stability in Bonding— Atomic Stability Atoms combine when the compound formed is more stable than the separate atoms. Most elements can form compounds, however the noble gases seldom form compounds.

19:1 Stability in Bonding— Unique Noble Gases To understand the stability of the noble gases, you should look at electron dot diagrams. Pg. 604-605. Group 1 has 1 electron in its outer shell. Group 2 has 2 in the outer shell. Group 13 has 3 in outer shell. Group 14 has 4 in outer shell, etc. Group 18 (the noble gases) has 8 in outer shell. It is stable. 11

Title a Page in Your Notebook: “Electron Dot Diagrams”. Hydrogen Carbon Oxygen Nitrogen Helium Fluorine 7. Chlorine 8. Sodium Sulfur Boron Draw electron dot diagrams for these atoms.

20:1 Stability in Bonding— Unique Noble Gases The noble gases are stable because they have 8 electrons in their outer energy level. That makes it full. Kr Ne He Xe Ar Rn

Lewis Theory Between 1916 and 1919, Gilbert Newton Lewis, Walther Kossel, and Irving Langmuir came up with a theory to explain chemical bonding. This theory would be later called Lewis Theory and it is based on the following principles:   Valence electrons, or the electrons in the outermost electron shell, have an essential role in chemical bonding. Ionic bonds are formed between atoms when electrons are transferred from one atom to another. Ionic bond is a bond between nonmetals and metals. Covalent bonds are formed between atoms when pairs of electrons are shared between atoms. A covalent bond is between two nonmetals. Electrons are transferred/shared so that each atom may reach a more stable electron configuration. The noble gas configuration, which contains 8 valence electrons, illustrates this principle. This is called octet rule.

20:1 Stability in Bonding— Unique Noble Gases 2. An atom is chemically stable when its outer energy level is complete. Most atoms require 8 electrons in its outer energy level to be stable. Remember that helium is stable with two electrons. Why?

20:1 Stability in Bonding— Unique Noble Gases 3. Elements that do not have full outer energy levels are more stable when they form compounds. 4. Atoms can lose, gain or share electrons to get a stable outer energy level.

20:1 Stability in Bonding— Energy Levels 5. A chemical bond is the force that holds atoms together in a compound.

20:1 Stability in Bonding— Energy Levels Atoms form compounds when the compound is more stable than the separate atoms. Na Cl Sodium and Chlorine can share an electron and each will become more stable. This new compound is called sodium chloride, commonly known as table salt. Multiples

20:2 Types of Bonds— A Bond Forms A. An ion is a charged particle because it has more or fewer electrons than protons. K I

20:2 Types of Bonds— Gain or Lose Electrons 1. When an atom loses an electron, it becomes a positively charged ion; a superscript indicates the charge. 2. When an atom gains an electron, it becomes a negatively charged ion.

20:2 Types of Bonds— The Ionic Bond B. An ionic compound is held together by the ionic bond —the force of attraction between opposite charges of the ions. An ionic bond is formed when an equal exchange of electrons occurs. Video 9—Ions and Ionic Bonding 2.14

20:2 Types of Bonds— The Ionic Bond 1. The result of this ionic bond is a neutral compound. 2. The sum of the charges on the ions is zero.

20:2 Types of Bonds— Sharing Electrons C. Molecules are neutral particles formed as a result of sharing electrons. Particles formed from covalent bonding of atoms are called molecules. Video 10—Covalent Bonds 3.33

20:2 Types of Bonds— Sharing Electrons 1. A covalent bond is the force of attraction between atoms sharing electrons. 2. Atoms can form double or triple bonds depending upon whether they share two or three pairs of electrons.

20:2 Types of Bonds— Unequal Sharing 3. Electrons shared in a molecule are held more closely to the atoms with the larger nucleus. 4. A polar molecule has one end that is slightly negative and one end that is slightly positive although the overall molecule is neutral.

20:2 Types of Bonds— Unequal Sharing 5. In a nonpolar molecule electrons are shared equally. Because water has a slight positive charge at one end and a slight negative end at the other end, it is a polar molecule. Demonstration of water’s polarity.

20:3 Writing Formulas Chemists use symbols from the periodic table to write formulas for compounds. A binary compound is composed of two elements. In order to write formulas you need to know which elements are involved and what number of electrons they lose, gain, or share in order to become stable.

Oxidation number is sometimes called the valence number. 20:3 Writing Formulas The oxidation number tells you the number of electrons that are lost, gained or shared to become stable. The sum of the oxidation numbers in a neutral compound is always zero. Oxidation number is sometimes called the valence number. Video 11—Oxidation Numbers 2.54

20:3 Writing Formulas 3. Use oxidation numbers and their least common multiples to write formulas. a. When writing formulas, remember that the compound is neutral. The sum of the oxidation numbers must equal 0.

The formula SO42- stands for sulfate. 20:3 Writing Formulas For ionic compounds the oxidation number is the same as the charge on the ion. For example, a sodium ion has a charge of 1+ and an oxidation number of 1+. A chlorine ion has a charge of 1- and an oxidation number of 1-. The formula SO42- stands for sulfate.

20:3 Writing Formulas b. A formula must have the correct number of positive and negative ions so the charges balance.

20:3 Writing Formulas 4. Use the name of the first element, the root name of the second element, and the suffix -ide to write the name of a binary ionic compound.

Oxidation (Valence) Number handout 20:3 Writing Formulas Oxidation (Valence) Number handout The numbers with positive or negative signs in the handout and in Fig. 16 are the oxidation numbers for these elements. Notice how they fit with the periodic table groupings.

20:3 Writing Formulas 5. The elements in the PT can have more than one oxidation number. When naming these compounds, the oxidation number is expressed in the name with a Roman numeral. For example, the oxidation number of iron in iron(III) oxide is 3+.

20:3 Writing Formulas Polyatomic ion —positively or negatively charged, covalently bonded group of atoms. 1. The compound contains three or more elements.

20:3 Writing Formulas To write names, write the name of the positive ion first; then write the name of the negative ion.

20:3 Writing Formulas 3. To write formulas, use the oxidation numbers, their least common multiple, and put parentheses around the polyatomic ion before adding a subscript.

20:3 Writing Formulas— Writing Names D. Writing Names: Write the name of the positive ion. Use the tables and check to see if the positive ion is capable of forming more than one oxidation number. Write the charge of the positive ion using Roman numerals. If only one O#, go to step 3.

20:3 Writing Formulas— Writing Names Write the root name of the negative ion. The root is the first part of the element’s name. For chlorine the root is chlor--- for oxygen it is ox--. Add the ending –ide to the root.

Count the Number of Elements: To Name Compounds Count the Number of Elements: If there are 3 or more elements: Write the name of the first element or polyatomic ion; Locate the polyatomic ion’s name on the “Valence Sheet and write its name. If there are only 2 elements: Write the name of the first element; Write the root of the second element and… Add “ide.”

20:3 Writing Formulas— Writing Names Problem Solving Activity

Write these compounds into your notebook and name them. H2O NaCl CaI2 SnBr4 ZnCl2 FePO4 FeS K2O Hydrogen oxide HF 10. AgNO3 11. Co(NO2) 2 12. HI NH4OH 14. MnSO4 Hydrogen fluoride Silver nitrate Sodium chloride Cobalt nitrite Calcium iodide Tin IV bromide Hydrogen iodide Zinc chloride Ammonium hydroxide Ferric phosphate or iron III phosphate Iron II sulfide or Ferrous sulfide Manganese sulfate Potassium oxide

20:3 Writing Formulas— Writing Names E. A hydrate is a compound with water chemically attached to its ions. Anhydrous is a term that means “without water.”

20:3 Writing Formulas— Writing Names 1. When writing the formula for a hydrate it must show that the water is chemically attached to its ions. An example of this kind of formula is CoCl2 6H2O.

20:3 Writing Formulas F. Name binary covalent compounds by using prefixes to indicate how many atoms of each element are in the compound.

Writing Formulas Activity 2 Copy these compounds into your notebook and write their balanced formula. Aluminum oxide Hydrogen iodide Hydrogen fluoride Potassium fluoride Iron II oxide Sodium bromide Tin IV bromide Copper II oxide Hydrogen oxide Sodium chloride calcium iodide zinc chloride lithium sulfide Lithium phosphate

Homework due and test soon.