CH.6 - IONIC COMPOUNDS - PERIODIC TRENDS...

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CHEMISTRY - MCMURRY 7E

CH.6 - IONIC COMPOUNDS - PERIODIC TRENDS AND BONDING THEORY

CONCEPT: ATOMIC PROPERTIES AND CHEMICAL BONDS

Before we examine the types of chemical bonding, we should ask why atoms bond at all.

• Generally, the reason is that ionic bonding ____________ the potential energy between positive and negative ions.

• Generally, the reason covalent bonds form is to follow the ____________ rule, in which the element is then

surrounded by 8 valence electrons.

There are three models of chemical bonding:

In ____________________ bonding, metals connect to non-metals.

• __________________ transfers an electron to the ________________ , creating ions with opposite charges that

are attracted to each other.

Li F Li F Li F

In _______________ bonding, non-metals connect to non-metals.

• In it the nonmetals __________________ electron pairs between their nuclei.

ClCl

In _______________ bonding, metal atoms “pool” their valence electrons to form an electron “sea” that holds the metal-ion

together



CONCEPT: CHEMICAL BONDS (PRACTICE)

EXAMPLE: Describe each of the following as either a(n): atomic element, molecular element, molecular compound or ionic compound.

atomic element ––

molecular element ––

molecular compound ––

ionic compound ––

a. Iodine

b. NH3

c. Graphite

d. Na3P

e. Ag2(SO4)2



CONCEPT: THE IONIC-BONDING MODEL

The central idea of ionic bonding is that the metal transfers an electron(s) to a nonmetal.

• The metal then becomes a(n) ____________ (positive ion). and the nonmetal becomes a(n) _____________ (negative ion).

• Their opposite charges cause them to combine into a crystalline solid.

PRACTICE: Determine the molecular formula of the compound formed from each of the following ions.

a. K+ & P3-

b. Sn4+ & O2-

c. Al3+ & CO32-



CONCEPT: ENERGY CONSIDERATIONS IN IONIC BONDING

________________________ is the enthalpy change that occurs when 1 mol of ionic solid separates into gaseous ions. It

tells us the strength of ionic interactions and has an influence in melting point, hardness, solubility and other properties.

Li+ (g) + F – (g) LiF (s) ∆H = 1050 kJ/mol

In order to calculate the energy of an ionic bond we use the following equation;

Ionic Bond Energy =

Radius = __________________________________

EXAMPLE: For each pair, choose the compound with the lower lattice energy.

a. BaO or MgO

b. LiCl or CaS

PRACTICE 1: Choose the compound with the lower lattice energy.

a. AlN or KBr

PRACTICE 2: Choose the compound with the higher lattice energy.

a. CsF or LiCl



CONCEPT: BORN-HABER CYCLE

The Born-Haber cycle is used a method to calculate the ________________________ or _______________________ of a compound.

• It looks mainly at the formation of an ionic compound from gaseous ions.

• The metal being from Groups _______ or _______ and the nonmetallic element being a ________________ or ________________ .

M (s)+ 12X2

ΔHfo

⎯ →⎯⎯ MX (s)

M (g) X (g)

M+ (g) X– (g) HX (s)

1

2

3

4

5

ΔHfo = 1 2 3 4 5+ +++

1

2

3

4

5

=

=

=

=

=

+



PRACTICE: BORN-HABER CYCLE

EXAMPLE: Using the Born-Haber Cycle, demonstrate the formation of cesium chloride, CsCl, and calculate its heat of formation.

ΔHSublimation = 79kJmol

IE1 = 376kJmol

ΔHDissociation =122kJmol

EA = −349 kJmol

U = −661 kJmol



CONCEPT: DIPOLE ARROWS

Before drawing covalent compounds we first need to understand the idea of polarity and its connection to electronegativity.

• Polarity arises whenever two elements are connected to each other and there is a significant difference in their

electronegativities.

• Generally, electronegativity ________________ going from left to right of a period and ________________ going

down a group.

To show this difference in electronegativity we use a dipole arrow.

The dipole arrow points towards the ________________ electronegative element.

The Effect of Electronegativity Difference on Bond Classification

Electronegativity Difference (ΔEN)

Bond Classification

Example

Zero (0.0)

Pure Covalent

Small (0.1 – 0.4)

Nonpolar Covalent

Intermediate (0.4 – 1.7)

Polar Covalent

Large (Greater than 1.7)

Ionic



PRACTICE: DIPOLE ARROWS

EXAMPLE: Based on each of the given bonds determine the direction of the dipole arrow and the polarity that may arise.

a. H Cl

b. S O

c. Br B Br

PRACTICE 1: Based on the given bond determine the direction of the dipole arrow and the polarity that may arise.

a. H C


a. N F


a. H N H



CONCEPT: CHEMICAL BOND IDENTIFICATION

PRACTICE: Answer each of the following questions dealing with the following compounds.

KBr NH3 F2 CaO NaClO

a. Which of the following compound(s) contains a polar covalent bond?

b. Which of the following compound(s) contains a pure covalent bond?

c. Which of the following compound(s) contains a polar ionic bond?

d. Which of the following compound(s) contains both a polar ionic bond and a polar covalent bond?



CONCEPT: ELECTRON-DOT SYMBOLS

Before we look at the first two bonding models, we have to figure out how to depict the valence electrons of bonding atoms.

• In the _________ electron-dot symbol, the element symbol represents the nucleus and inner electrons, and the

surrounding dots represent the ________________ electrons.

EXAMPLE: Draw the electron-dot symbol for each of the following elements.

1A 2A 3A 4A 5A 6A 7A 8A

Li

Be

B

C

N

O

F

Ne

It’s easy to write the Lewis symbol for any Main-Group element:

1) Remember that Group Number equals Valence Electron Number.

2) Place one dot at a time on the four sides (top, right, bottom, left) of the element symbol.

3) Keep adding dots, pairing them up until you have reach the number of total valence electrons for that element.

PRACTICE 1: Draw the electron-dot symbol for the following ion.

Mg2+


N3-


Cr1+



CONCEPT: CHEMICAL BONDING I

Rules for Drawing

1. Least electronegative element goes into the center. Important Facts to Know:

(a) Electronegativity increases across any Period going from left to right and up any Group going from bottom to top.

(b) Hydrogen and Fluorine ________________ go in the center and they only make _________ BOND.

2. Number of valence electrons equals group number.

3. Carbon must make _____ bonds, except in rare occasions when it makes _____ bonds.

• If the carbon atom were positive or negative then it would make _____ bonds

4. Nitrogen likes to make _____ bonds.

5. Oxygen likes to make _____ bonds.

6. Halogens (Group 7A), when not in the center, make _____ bond.

7. Expanded Valence Shell Theory: Nonmetals starting from Period _____ to _____ can have more than 8 valence

electrons around them when in the center.



CONCEPT: INCOMPLETE OCTETS

Nonmetals form covalent bonds to generally follow the ___________ rule, in which the element is surrounded by 8 valence

electrons.

• Sometimes elements form compounds in which they have ____________________ 8 valence electrons.

• These elements are said to have an incomplete octet or to be ________________________________________ .

EXAMPLE: Draw the following molecular compound.

BH3

PRACTICE: Draw the following molecular compound.

BeCl2



CONCEPT: EXPANDED OCTETS

Expanded Valence Shell Theory: Nonmetals starting from Period _____ to _____ can have more than 8 valence

electrons around them when in the center.

EXAMPLE: Draw each of the following molecular compounds.

IF3 KrF5+

PRACTICE 1: Draw the following molecular compound.

SBr4


I3–



CONCEPT: POLYATOMIC IONS

Shortcut: If you have _____, _____, _____, _____, __________________ or __________________ connected to oxygen

then the negative charge tells us how many oxygens are single bonded.

• The remaining oxygens are _______________________ bonded to the central element.

EXAMPLE: Draw each of the following molecular compounds.

SO42-

PO43- H2SO4


SeO42-


XeO64-



CONCEPT: FORMAL CHARGE

Structures and polyatomic ions that break the octet rule often have ________________ Lewis Structures.

• The purpose of using the formal charge formula is to determine which Lewis structure is the best answer.

Formal Charge =

a) Use formal charge formula to check to see if you drew your compound correctly.

b) Formal charges must be either _____, ______, ______.

c) If you add up all the formula charges in your compound that will equal the overall charge of the compound.

EXAMPLE: Calculate the formal charge for each of the following element designated for each of the following.

a. The carbon atom in

b. The sulfur atom in

PRACTICE: Calculate the formal charge for each of the following element designated in the following compound.

a. Both oxygen atoms in:

!A B



CONCEPT: RESONANCE STRUCTURES

Resonance structures are used to represent bonding in a molecule or ion when a single Lewis structure cannot correctly

describe the Lewis structure.

EXAMPLE: Determine all the possible Lewis structures possible for NO2–. Determine its resonance hybrid.

EXAMPLE: Determine the remaining resonance structures possible for the following compound, CO32-.

O

C OO



CH.6 - IONIC COMPOUNDS - PERIODIC TRENDS...

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