Chemical Bonding read chapter 12 - PowerPoint Presentation

Slide1

Chemical Bonding

read chapter 12

Slide2

Types of Chemical Bonds

Ionic bonds

: attraction between positive and negative ions (e- are lost or gained

)

m

etal + nonmetal (K

2

O)

Covalent bonds

:

sharing

of electrons between two

atoms

n

onmetal + nonmetal (N

5

O

2

)

Metallic

bonds

: metal atoms give up valence e- that are then free to move throughout the

material

m

etal + metal (steel)

Goofy song 5 min

https://www.youtube.com/watch?v=ljvX-

RMv_lw

Slide3

Bond Types

Ionic compounds conduct electricity when dissolved in water, because the dissociated ions can carry charge through the solution. Molecular compounds don't dissociate into ions and so don't conduct electricity in solution.

Slide4

Chemical Bonding

Why do bonds form??? To complete their valence shell



“to satisfy their octet.”

Remember the Octet Rule: when elements bond they transfer/share electrons so both have a noble gas electron configuration.

Bonding only involves the

valence electrons

(To discuss bond types we have to visualize electrons)

Slide5

Lewis Dot Diagram Structures

Use dots to show all the

valence electrons

(outer electrons - same as group # except He)

For example, Arsenic – group 5

As = [

Ar

]4s2 3d

10 4p3 - s and p are the only outer e-

Slide6

Lewis Dot Diagram Structures

The rules:

There are 4 sides around the symbol

No side can have more than 2 dots

When filling the sides of the element symbol, each side gets one dot before doubling up (except H and He)

Slide7

YOU TRY! Write the Lewis Structure for

A)

Xe

B) Br

C) Ba

Slide8

Lewis Structures…

https://www.youtube.com/watch?v=a8LF7JEb0IA&index=24&list=

PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr

(12 min)

Slide9

Covalent Bonding

Bonds form when elements share electrons to end up with 8 valence electrons

(except H and He want 2)

Can have more than one bond

Single bond shares 1 pair of electrons

Double bond shares 2 pairs

Triple bond shares 3 pairs (very strong)

Slide10

Lewis Structures & Covalent Bonds

Covalent Bonds – arrange elements so everyone has 8 valence electrons (except H and He want 2)

CH

4

PCl

3

SO

2

Slide11

Lewis “Dot” Diagram Rules

Hydrogen is

always

terminal (Halogens are usually terminal)

LOWEST electronegativity is central molecule (often the oddball

) Carbon is always central

!Elements like symmetry! 1) Find total # valence electrons by adding up group #’s of elements (

add for negative charge, subtract for positive charge)2) Place one pair of electrons between each pair of bonded atoms

3) Place lone pairs around each terminal (except H) to satisfy octet rule. 4)

Left over pairs are assigned to central atomIf central atom is not yet surrounded by 4 electron pairs, convert to double bonds. BUT

ONLY C, N, O, P, S!!

Slide12

YOU TRY!

Draw Lewis Structures for:

A) phosphorus

trifluoride

B) N

2

H

2C) CH2O

Slide13

Strange bonding

Some bonds have a negative charge so an ADDITIONAL ELECTRON is added

(or a positive charge so an electron is eliminated)

INCLUDE BRACKETS AND CHARGE

H3O

+

Slide14

YOU TRY

Write the Lewis structures for the following compounds:

A) SO

4

2-

*B) ClO

3 –C) NOF

Slide15

Exceptions to the Octet Rule

Beryllium (can share) and has a maximum of 4 valence electrons

Boron can bond with only 6 valence electrons

(Fluorine and all halogens will

NOT

double bond)

Slide16

*Exceptions to the Octet –

Expanded Octet

*You will NOT be assessed on this*

Some central atoms can have more than 8 electrons. The terminal electrons (halogens) are “happy” and it results in the central atom having more.

Slide17

Covalent Bonding

But nothing is really equal…

In

c

ovalent bonding one element is more

electronegative

than another and the electrons move closer to one element.

(fluorine MOST electronegative)

Slide18

Covalent Bonding

Types

of covalent

bonds:

A

. Nonpolar covalent: e- are shared evenly (H-H

) (“NOT pushing’”)

B. Polar covalent: bonded atoms have unequal attraction for e- (H-Cl and H-O-H) (“pushing”)

Slide19

Molecular Polarity

For a compound with only 1 bond, electronegativity difference

is compared

Calculate the electronegativity

difference

(p 403)

Below

0.50 = nonpolar

0.50 – 1.7 = polar (and one is slightly negative/positive so arrow goes from + to -)

Slide20

Molecular Polarity

For a compound with more than 1 bond

Draw Lewis Structures

Complete SYMMETRY (radial symmetry) = non polar

ASYMMETRICAL = POLAR

Asymmetrical = terminal

elements are not all the same-

look for lone e-

Slide21

YOU TRY

Write the Lewis structures for the following molecules. Polar or nonpolar??

A) NF

3

B) carbon monoxide

C)

O

3D) C2F2

Slide22

Ionic Bonding (review)

Ions are formed when an atom, trying to move toward a stable electron configuration (octet rule), loses or gains electrons

Na :

Na

1+

Cl :

Cl 1-Groups and their charges

1A- charge 1+ (Li)

5A- charge 3- (N)2A- charge 2+

(Be) 6A- charge 2- (O) 3A- charge 3+ (B)

7A- charge 1- (F)

Nonmetals form anions (negative

ion); metals form cations (positive

ion)

Slide23

Lewis Structures & Ionic Bonds

Ionic Bonds transfer electrons and a new charge is shown

This type of bond is stronger than a

single bond

covalent compound

(the

electronegativity

difference (p 403

) is greater than 1.7)

Slide24

http://

www.bozemanscience.com/chemical-bonds-covalent-vs-ionic

(Bozeman science 9 minutes)

Slide25

YOU TRY

Write the Lewis structures for the following

ionic compounds:

A

) MgBr

2

B) NaIC) K2

S

Slide26

Hydrogen “Bonds”

Hydrogen “bonds” are not real bonds and happen only between hydrogen

and fluorine/oxygen/nitrogen

(FON)

Hydrogen bonds are weak but it does provide some attraction to neighboring molecules.

Let’s talk about water: H

2

OINTRAmolecular force (covalent bond) keeps the 2 hydrogen atoms bonded to the oxygen atom (strong bond)

INTERmolecular force bonds one water molecule to a neighboring water molecule

(hydrogen bond = weak bond)

https://www.youtube.com/watch?v=Jkxad9yMUWo

(2min)