chemical bonds ch. 3 chemistry ii milbank high school

Chemical BondsCh. 3

Chemistry IIMilbank High School

Sec. 3.1Stable Electronic Configurations• Objectives

– To determine what electron configuration appears to be the most stable.

– Define the octet rule

Background

• There are about 115 known elements

• There are about 20 million known chemical compounds

• Chemical Bonds: force that holds the atoms of elements together

Sec. 3.1 Stable Electron Configurations

• Which group is the most stable?

• Which group is the most reactive?

• What does this have to do with electrons?

• See Pg. 68

So what group has a stable electron configuration?

• Noble gases!

• They have a stable octet of electrons in the highest main energy level

• Octet Rule: atoms seek an arrangement that will surround them with eight electrons in the outer shell

Atoms and ions

• Atoms are electrically neutral.

• Same number of protons and electrons.

• Ions are atoms, or groups of atoms, with a charge (positive or negative)

• Different numbers of protons and electrons.

• Only electrons can move.

• Gain or lose electrons.

Anion

• A negative ion.• Has gained electrons.• Nonmetals can gain electrons.• Charge is written as a superscript

on the right.

F1-Has gained one electron (-ide is new ending= fluoride)

O2- Gained two electrons (oxide)

CationsPositive ions.Formed by losing electrons.More protons than electrons.Metals can lose electrons

K1+ Has lost one electron (no name change for positive ions)

Ca2+ Has lost two electrons

Ionic Compounds

• Ionic compounds- from joining metal cations and nonmetal anions- they are electrically neutral

• Usually solid crystals

• Melt at high temperatures

Sec. 3.2 Lewis Structures

• Objectives– Determine and then solve Lewis

Structures

Lewis Structures

• A representation of an element in which the chemical symbol stands for the core of the atom and dots are placed around the symbol for its valence electrons (outer shell)

Examples

• What are the Lewis structures (or electron dot structures) of the following elements?– Nitrogen– Argon– Potassium

Sec. 3.3 Intro Ionic Bonds and Ionic Compounds• Ionic Compounds: consists of

oppositely charged ions held together by electrostatic attractions

• Ionic Bonds: attractive forces between positive and negative ions

Sec. 3.3 and Sec 3.4The Sodium-Chlorine Reaction and Other Ionic Bonds

• Objectives– Investigate the sodium-chlorine

reaction– Make some general considerations

about ionic bonds

Background

• Sodium is HIGHLY reactive

• Chlorine not very reactive, other than it is irritating to the respiratory tract

Reaction…

• Cl2 + 2Na 2 Cl- + 2Na+ • See Lewis Structure• Name: NaCl

• More reactive substances become less reactive when in a compound– Release energy in the process

• Sodium atom becomes less reactive by losing an electron

• Chlorine atom becomes less reactive by gaining an electron

Ionic Bonds: Some General Considerations

• Normally, metallic elements in Groups 1A, 2A, and 3A react with nonmetallic elements in Groups 5A, 6A, and 7A to form stable crystalline ionic structures

Sec. 3.5Names of Simple Ions and Ionic Compounds

• Determine how ionic compounds are named

Naming

• Monoatomic positive ions (cations)—add “ion”– Na+ is called a sodium ion

• Monoatomic negative ions (anions)—add “ide” to the usual ending– Cl- is called chloride ion– S2- is called sulfide ion

Naming• See Table 3.2• Charge on a Group 1A element is

usually 1+• Charge on a Group 2A element is

usually 2+• You can calculate the charge by

subtracting 8 from the group number.– Oxygen is 6-8 = 2-

• No simple way to determine charge on B subgroup elements– Roman numerals usually indicate charge – Iron(II) ion means Fe2+

Formulas for Ionic Compounds

• Electrically neutral ions– Potassium (K+) combines with

Bromide ions (Br-)• Name? • KBr• Ratio 1:1

Formulas for Ionic Compounds

• Non-neutral ions– One Calcium ion (Ca2+) combines with

two chloride ions (Cl-)• Ratio 1:2• Name?

• CaCl2

Crossover Method

• The charge number for one ion becomes the subscript for the other

• What is the formula for aluminum oxide?– Charges

• Al3+ and O2-

– Name• Al2O3

Names for Binary Ionic Compounds

• What is the name of the compound Na2S?

• Find the ions (Table 3.2)– Sodium?

• Sodium ion -- Na+

– Sulfur?• Sulfide ion – S2+

– Name: sodium sulfide

Sec. 3.6Covalent Bonds: Shared Electron Pairs

• Determine what a covalent bond is

• Show how to write covalent bonds using Lewis Structures

Covalent Bonds

• Covalent Bonds: bond formed by a shared pair of electrons between atoms

• Molecule: group of atoms that are chemically bonded together– H2 represents a molecule of hydrogen

Examples

• H

• Cl– Bonding Pair (two shared electrons)– Lone Pairs (nonbonding)

Sec. 3.7Multiple Covalent Bonds

• Determine the difference between single, double, and triple bonds

Covalent Bonds

• Single Bond: single pair of shared electrons

• Double Bond: a covalent linkage in which the two atoms share two pairs of electrons

Examples

• CO2

• Sometimes represented by dashes• O=C=O

• In general, many nonmetals often form a number of covalent bonds equal to the eight minus the group number.

• Oxygen.. 8 - 6 = 2 covalent bonds

Sec. 3.8Naming Covalent Compounds

• Determine the correct naming structure for covalent compounds

Names for covalent compounds

• Makes use of prefixes to name compounds

• Example:– N2O4

• dinitrogen tetraoxide

Prefixes

• Mono 1• Di 2• Tri 3• Tetra 4• Penta 5• Hexa 6

• Hepta7

• Octa 8• Nona 9• Deca 10

Sec. 3.9Unequal Sharing: Polar Covalent Bonds

• Show another way in which atoms combine

• Determine the difference between nonpolar and polar covalent bonding

Hydrogen/Chlorine Reaction

• Both need an electron to achieve a noble gas configuration

• They share a pair and form a covalent bond

• Chlorine atom have a greater attraction for a shared pair of electrons than do hydrogen atoms– More electronegative than hydrogen– They hold the shared electrons more tightly

Polar vs Nonpolar

• Nonpolar Covalent Bond: covalent bond in which electrons are shared equally

• Polar Covalent Bond: electrons are drawn more closely to the more electronegative atom, creating a separation of charge

Representation

• Use δ+ and δ- to indicate that one end is partially positive and one end is partially negative

• Example H—Cl

Sec 3.10Electronegativity

• Define electronegativity

• Show electronegatitive tendencies using the Periodic Table

Electronegativity

• The measure of an atom’s tendency to attract electrons in a covalent bond to itself

• The greater the electronegativity of an atom in a molecule, the more strongly the atom attracts the electrons in a covalent bond

Periodic Table Tendencies

• Within a period, elements generally become more electronegative from left to right

• Within a group, electronegativity decreases from top to bottom

Sec. 3.11Rules for Writing Lewis Structures

• Discuss the rules for writing Lewis Structures

• Write Lewis Structures

• Go through the individual steps used to write a Lewis formula

Lewis Structures

• Skeletal structure—model of a molecule that tells us the order in which the atoms are attached to one another

Rules for Writing Lewis Structures

• 1. Hydrogen atoms form only one bond; they are shown at the end of a sequence of atoms

• 2. Polyatomic molecules and ions often consist of a central atom surrounded by more electronegative atoms– Hydrogen exception—always on

outside, even when bonded to a more electronegative element

Steps after you have the skeletal formula

• 1. Calculate the total number of valence electrons

• 2. Write the skeletal structure• 3. Place electrons about outer atoms so

that each has an octet• 4. Subtract the number of electrons

assigned so far from the total calculated in Step 1. Assign the rest to the central atom

• 5. Move lone pairs to form a multiple bond on central atom.

Examples

• Example 3.12

• Example 3.13

Sec. 3.12The VESPR Theory

• Predict and construct shapes of many molecules by using the VESPR theory

VSEPR Theory

• Valence shell electron theory—describes the geometrical shape of a molecule or polyatomic ion based on the mutual repulsions among electon groups surrounding the central atom(s) in the structure

• Minimizes repulsion between the like-charged particles

Steps in determining shapes

• 1. Draw a Lewis Structure– Use dots to indicate lone pairs

• 2. To determine the shape, use the steric number (number of atoms bonded to the central atom + number of lone pairs on the central atom)

• 3. Draw shape based upon steric number

• 4. Place electron pairs as far apart as possible.– See Table 3.4

Examples

• Use table 3.4

• Examples 3.14– BH3

– SCl2

Sec. 3.13Polar and Nonpolar Molecules

• Define dipole

• Show the correct way to represent dipoles

Dipoles

• A molecule is a dipole if it has a positive and negative end

• Represented by an arrow with a plus end

H—Cl

Sec. 3.14Polyatomic Ions

• Define polyatomic ions

• Determine formulas for polyatomic ions

Polyatomic Ions

• Charged particles containing two or more covalently bonded atoms

• Table 3.5 – some common polyatomic ions

• Example 3.15

Sec. 3.15 Exceptions to the Octet Rule

• Determine what a free radical is

• Show expanded valence shells

Three exceptions to the octet rule

• 1. Molecules in which the total number of valence electrons is an odd number

• 2. Molecules in which the total number of valence electrons is too low to allow all atoms to have a filled valence shell

• 3. Molecules involving elements from the third and higher periods can have an expanded valence shell

Molecules with odd numbers of valence electrons

• Free Radicals—a highly reactive atom or molecular fragment characterized by having one or more unpaired electrons

• See examples on pg. 91

Expanded Valence Shells

• Refer to situations in which the central atom in a Lewis Structure is able to accommodate more than the usual octet of electrons in its valence shell

• Usually encountered when the central atom is a nonmetal of the third period or beyond