Ch 15 & 16 Notes.ppt

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Warm-up

Thursday, January 31

st

_{On your way in, you should have received a}

card with an ion on it. Using your card, do the following…

Find a person or people who balance out the charge on your card and stand by them.

Write the compound you form with your group.

Name the compound.

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Chapter 15

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Electron Configuration in Ionic

Bonding

 Valence Electrons – the electrons in the

highest occupied energy level of an element’s atoms

 To find the number of valence electrons look at the element’s group number

1

(4)



Electron dot structure

-

System of arranging dots

representing

valence

electrons.



G.N. Lewis developed this

system and so they are also

referred to as Lewis Dot

Structures.

(5)

Example:

 Write the electron configuration for phosphorous.

1s22s22p63s23p3

 Write the orbital diagram for phosphorous.

___ ___ ___ ___ ___ ___ ___ ___ ___ 1s 2s 2p 3s 3p

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Practice

Write Lewis dot structures for the

following.

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Octet Rule

Octet rule – in forming compounds, elements tend to achieve the electron configuration of a noble gas.

Metals will typically lose electrons and therefore become positive

Na ___ ___ ___ ___ ___ ___

1s 2s 2p 3s

Ne ___ ___ ___ ___ ___ ___

1s 2s 2p 3s

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Octet Rule

_{Nonmetals will typically gain electrons and}

therefore become negative

O ___ ___ ___ ___ ___ 1s 2s 2p

Ne ___ ___ ___ ___ ___ 1s 2s 2p

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2-_{Electron dot symbols, Lewis dot structures,}

can be used to represent ions.

EXAMPLES:

_{Na Cl F Mg O Al}

Electron Configuration in Ionic

Bonding

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15.2 Ionic Bonds

_{Ionic Bond}_{- A chemical bond formed by the}

electrostatic attraction between a cation and an anion when there is a transfer of electrons.

_{Ionic compounds are generally a metal and a}

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Example:

When potassium reacts with chlorine what kind of compound is formed?

Ionic

_{Is there a transfer of electrons?}

Yes

_{Using the Lewis structures show what}

happens to K and Cl when they combine to form salt.

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Properties of Ionic Compounds

_{High melting points}

_{Solids at room temperature}

_{Crystalline in structure, referred to as a crystal}

lattice

MP:

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Bellwork

Friday, February 1

st

_{Draw Lewis structures for atoms of}

magnesium and sulfur.

_{Show how these atoms could combine to}

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15.3 Bonding in Metals

_{Metallic bonds}_{– consist of the attraction of}

free-floating valence electrons for the positively charged metal ions.

These forces hold metals together

Free Floating valence

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Properties of Metallic Compounds:

_{Good conductors}

_{Ductile – can be drawn into a wire}

_{Malleable – can be hammered or forced into}

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Crystalline Structure of Metals

_{Metals and ionic compounds are in compact orderly}

patterns

_{Identically sized spheres have several arrangements}

Simple cubic

Body-centered cubic

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_Alloys_{– mixtures of two of more elements, at}

least two of which are metals.

_{Steels are the most important alloys today}

Consist of Fe, C, B, Cr, Mn, Mo, Ni, W, V

Ex: Brass, steel, 14 K gold, sterling silver, & cast iron

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Chapter 16

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16.1 Covalent Bonds

_{Covalent bond}_{– Occurs when a pair of}

electrons is shared between two atoms.

Often between two nonmetals.

Single covalent bond – two atoms share a single pair of electrons

Double covalent bond – two atoms share two pairs of electrons

Triple covalent bond – two atoms share three pairs of electrons

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 Lewis dot structures can be useful for

representing covalent bonds between elements in a covalent compound.

 H + H 

 Cl + Cl 

 O + O 

 N + N  _{N N}

O O Cl Cl H H

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Rules for writing electron dot

structures: (use pencil!!!)

1. Add up the valence electrons from all the atoms in the compound.

Don’t try to keep track of which electrons come from which atoms.

If you are working with an ion, you must add or subtract electrons to account for the charge.

H₂O

(2)H + (1) O

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-Rules for writing electron dot

structures:

 _{2. Put the element that you have the fewest of}

as the central element. (Make it symmetrical)

 _{Put the elements in spatial order.} H₂O

(2) H + (1) O

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Rules for writing electron dot

structures:

 _{3. Use a pair of electrons to form a bond}

between each pair of atoms.

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Rules for writing electron dot

structures:

 _{4. Arrange the remaining electrons to satisfy}

the duet rule for hydrogen and the octet rule for all remaining atoms.

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Rules for writing electron dot

structures:

 _{5. Count the number of electrons represented}

in the drawn molecule.

 _{If two too many electrons are represented:}

draw a double bond between two elements

remove a pair of electrons from each element taking place in the bond.

H O H

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EXAMPLES:

 _CH₄

 1. (1) C + (4) H

(1)(4e-) + (4)(1e-) = 8e

- _{2. Spatial order}

 3. Draw bonds

 4. Octet rule satisfied?

 5. # of e- match?

C H

H H

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EXAMPLES:

 _CO₂

 1. (1) C + (2) O

(1)(4e-) + (2)(6e-) = 16e

 3. Draw Bonds

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EXAMPLES:

 _NH₃

 1. (1) N + (3)H

(1)(5e-) + (3)(1e-) = 8e

- 2. Spatial order

 3. Draw bonds

N H

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EXAMPLES:

 CCl₄

 _1. _{(1) C + (4) Cl}

(1)(4e-) + (4)(7e-) = 32e

- _{2. Spatial Order}

 _{3. Draw bonds}

 _{4. Octet rule satisfied?}  _{5. # of e}- match?

C Cl

Cl Cl

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EXAMPLES:

 _NH₄+

 1. (1) N + (4) H - (1)(+)

(1)(5e-)+ (4)(1e-) - (1)(1e-) = 8e

- 2. Spatial order

N H

H H

H

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EXAMPLES:

 SO₄

2- _{1. (1) S +} _{(4) O + (2)(-)}

(1)(6e-)+ (4)(6e-) + (2)(1e-) = 32e

S O

O O

O

(32)

2-EXAMPLES:

 CN

- 1. (1) C + (1) N + (1)(-)

(1)(4e-) + (1)(5e-)+ (1)(1e-) = 10e

 _{3. Draw Bonds}

C N

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-EXAMPLES:

 CO₃

2- _{1. (1) C +} _{(3) O + (2)(-)}

(1)(4e-)+ (3)(6e-) + (2)(1e-) = 24e

C O

O O

(34)

2-Bellwork

Monday, February 4

th

_{Draw the following Lewis dot structures.}

CCl₄ NH₄+

SO₄2- CN

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2-Exceptions to the Octet Rule:

Elements that can have extra electrons!

B F

F

P Cl

Cl

S

F

Boron Trifluoride 3 bonds on B(6e-)

Phosphorous Pentachloride 5 bonds on P (10 e-)

Sulfur

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REMEMBER:

“P B S” bonded to ANY

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Resonance Structures

Structures that can occur when it is possible to write two or more valid electron dot structures that satisfy the octet rule.

EXAMPLES:

CO₃

2-C O

O O

[ ]

2-C O

O O

[ ]

2-C O

O O

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2-Resonance Structures

EXAMPLE:

 NO₃

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-VSEPR

_{Valence Shell Electron Pair Repulsion}

theory

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VSEPR:

Regions of electron density (where pairs of electrons are found) can be used to determine the shape of the molecule.

CO2

Here there are two regions of electron density.

The regions want to be as far apart as possible, so it is linear.

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EXAMPLES:

 CH₄

 There are four electron pairs.

 You would expect that the bond angles would be 90° but…

 Because the molecule is three-dimensional, the angles are 109.5°.

 The molecule is of tetrahedral arrangement.

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EXAMPLES:

 NH₃

_{Four regions of electron density}

_{But one of the electron pairs is a lone pair} _{The shape is called}_{trigonal pyramidal}

N H

H H

1

2

3

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EXAMPLES:

H₂O

_{Four regions of electron density} _{But two are lone pairs}

_{This structure is referred to as}_bent

O H H

1

2

3

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EXAMPLES:

CO₃

2-_{Three regions of electron density}

_{This structure is referred to as}_trigonal

planar

[ ]

C O

O O

1

2

3

(45)

2-Practice determining molecular

shape:

H₂S

4 regions of e- density

_{2 lone pairs}

_bent

S

H

(46)

Practice determining molecular

shape:

SO₂

_{3 areas of e}- density

_{1 lone pair} _bent

S

O O

(47)

shape:

CCl₄

4 areas of e- density

tetrahedral

C Cl

Cl Cl

(48)

shape:

BF₃

trigonal planar

B

F F

F

B F

F

(49)

Practice determining molecular

shape:

NF₃

1 lone pair

pyramidal

N H

H

(50)

Bellwork

Tuesday, February 4

th

_{Using the play-doh at your table, build a}

correctly shaped molecule of one of the following.

BF₃

CH₄

HBr

(51)

16.3

Polar Bonds and Molecules

_{In covalent bonds, the sharing of electrons}

can be equal

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Nonpolar Covalent Bonds

 Nonpolar covalent bond - This is a covalent bond in which the electrons are shared equally.

 EXAMPLES:

 _H₂

 _Br₂

 _O₂

 _N₂

 _Cl₂

 _I₂

(53)

Polar Bonds and Molecules

_{If the sharing is unequal, the bond is}

referred to as a dipole.

_{A dipole has 2 separated, equal but}

opposite charges.

_{“∂” means partial}

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Polar Bonds and Molecules

_{Polar covalent bond}_{- a covalent bond}

that has a dipole

_{It usually occurs when 2 different elements}

form a covalent bond.

EXAMPLE:

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_{Electronegativity}_{- This is the measure of the}

attraction an atom has for a shared pair of electrons in a bond.

_{Electronegativity values increase across a}

period and up a group.

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Electronegativity

 Electronegativity values can be used to

determine the degree of electron sharing which demonstrates the type of bond that will occur.

 Difference > 2  Ionic

 Difference 0.1 - 1.9  Polar Covalent

 Difference = 0 Nonpolar Covalent

Ionic >

Polar C

ovalent

> Non

polar Co

valent

2

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Examples:

_{Identify the type of bond for each of the}

following compounds:

_HBr

Br = 2.8

H = 2.1

.1 < < 1.90.7 _{Polar Covalent}

(58)

Examples:

NaF

F = 4.0 Na = 0.9

3.1

_N₂

N = 3.0

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Molecular Polarity

 _{If there is only one bond in the molecule, the bond}

type and polarity will be the same.

 If the molecule consists of more than 2 atoms, you must consider the shape. To determine its

polarity, consider the following:

Lone pairs on central atom

If so… it is polar

Spatial arrangement of atoms

Do bonds cancel each other out (symmetrical)?

• If so… nonpolar

Do all bonds around the central element have the same difference of electronegativity?

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Attractions Between Molecules

_{Van der Waals forces}_{– the weakest of the}

intermolecular forces. These include London dispersion and dipole-dipole forces.

London dispersion forces – between nonpolar

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Attractions Between Molecules

_{van der Waals}

forces(cont.)-Dipole interactions – between polar molecules

(62)

Attractions Between Molecules

 Hydrogen bonds – attractive forces in which

hydrogen, covalently bonded to a very electronegative atom (N, O, or F) is also weakly bonded to an

(63)

_{Ionic Bonding-occurs between metals and}

nonmetals when electron are transferred from one atom to another.

_{These bonds are very strong.}

(64)

Summary of the Strengths of

Attractive Forces

Ionic bonds

hydrogen bonds