Chapter 6 – INTRO AND *HONORS

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Chapter 6 – INTRO AND *HONORS
6.1 Chemical Bond=
Why does it occur?
Types of Bonds (An Overview)
1. Ionic Bond
Ion=
Metal=
Nonmetal=
Want to gain or lose valence electron(s) to become like a Noble Gas
WRITING ELECTRON CONFIGURATIONS OF IONS:
1
Ionic Bond=
->usually metal + nonmetal
2. Covalent Bond
->usually nonmetal + nonmetal
nonpolar covalent =
polar covalent=
How can you tell which it is? BY ELECTRONEGATIVITY!!
Electronegativity =
Here’s an electronegativity chart:
2
If the electronegativity difference is
>1.7 =
<.3 =
.3 – 1.7 =
3. Metallic Bond
-> metal + metal
More Details on the Bond Types
6.2 – Covalent Bonds and Molecular Compounds (#2 from Section 6.1)
Covalent Bond=
Molecule=
Diatomic Molecule=
3
Molecular Compound=
Molecular Formula=
Formation of Covalent Bonds:
*Overlapping of Orbitals (remember: want a noble gas
configuration):
The Octet Rule=
4
Drawing Lewis Structures:
Single Bond=
Double Bond=
Triple Bond=
6.3
Ionic Bonding and Ionic Compounds (#1 from Section 6.1)
Ionic Bond=
Crystal Lattice=
5
Formula Unit=
Dot Structure for Ionic Bonds (remember: they want to be like noble gases):
A COMPARISON OF MOLECULAR AND IONIC COMPOUNDS:
Molecular
6.4
Ionic
Metallic Bonding (#3 from Section 6.1)
Metals have LOW electronegativity
Easily give up electrons (losers)
No one is available for “grabbing”
The steps involved:
1. donates valence electrons to the surrounding area
2. electrons are free to move about – “electron sea”
3. all electrons are shared by all of the metal atoms
Picture:
6
Properties of metals:
1.
2.
3.
4.
5.
TEST!!!!!!!!
6.5
Properties of Molecular Compounds (Have covalent bonds!)
A. VSEPR =
Valence electron pairs want to be as far apart as possible!
1.
2.
3.
4.
Draw Lewis Structure
Look at Central Atom
Count # of electron areas (bond + lone pairs)
Use info from VSEPR chart below: (need to memorize for 2, 3, 4 total areas)
7
8
Examples:
Hybridization=
Explains the shapes we see – many elements do this
Carbon hybridizes to form four EQUAL orbitals
9
C
is actually
C
*1s22s22p2 changes to 1s22s12p3
hybridize to form four sp3 hybrid orbitals
B. Types of Molecules
1. Dipole = Molecule with overall charge
2. NonPolar With Polar Sites = Molecules with area of charge which cancel out
3. Nonpolar = Molecule with no areas of charge
How Do You Tell the Difference?
Ask yourself these questions?
1. Is there charge on the molecule?
Yes
No = Nonpolar
2. Can it be sliced?
(so that + separated from -)
Yes
=Dipole
No
=NPWPS
->Non-Honors classes can use models to help them.
Examples:
10
C. Intermolecular Forces (external bonds)
The attraction between molecules
Types of External Bonds
1. Dipole-Dipole Bonds
Occur between :
A. two dipoles (strongest)
Hydrogen Bond=
2. London Force (the weakest external bond)
Occur between two nonpolar molecules (WEAKEST) or Nonpolar with
Polar Sites
The Steps:
A.
B.
C.
D.
11
Properties Based on Number/ Strength of External Bonding:
1. State of Matter
s>l>g
2. Evaporation (*volatility)
slow>fast
3. Thickness (*viscosity)
thick>thin
4. Wetness (*adhesion)
To feel wet the substance must bond to your skin (to the Na+Cl-)
5. Dissolving
LIKE DISSOLVES LIKE
Demonstrations:
12
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