Transcript CHEMISTRY

CHEMISTRY – Chapter 6
Chemical Bonding
Chapter 6 – Chemical Bonding
Objectives:
1.
2.
3.
4.
5.
Define chemical bond.
Explain why most atoms form chemical bonds.
State the octet rule.
Discuss ionic, covalent, and metallic bonding.
Classify bonding type according to
electronegativity differences.
6. Draw complete Lewis dot diagrams.
7. Properly assign bonds and show valence
electrons for most compounds.
Chemical Bonding
 There are more than 100 elements
 From these thousands of compounds are
formed w/bonds
 Chemical bond – a mutual electrical attraction
btwn nuclei and valence e- of different atoms
 Nucleus are + and e- are –
 Two atoms coming together
 Why do elements want to bond together?


To become stable
To complete their valence shell
Lewis Dot Diagrams
 The chemical symbol represents the kernel
(center of the atoms)
 Each dot represents a valence e Ions are put in brackets [Na]+
Lewis Dot Diagrams (cont.)
ex.
Na
Mg
Al
S
F
Ar
Si
P
Octet Rule
 The tendency of atoms to rearrange
themselves into a stable octet (8 e-)
 All atoms want to complete their “octet”
and be configured like a Noble Gas.
 Example: Neon
 For this to happen the position of 1 or more
e- in the valence shell is altered
 Lost, gained, or shared
Ionic Bonding
A bond formed between a metal and a non-metal
Metals lose eNon-metals gain ee- in valence shell of metals are TRANSFERRED to
valence shell of non-metals
 Electrostatic force of attraction btwn oppositely
charged particles (+ and -)
 Want a valence shell of s2p6
 Except He and H
 Formula unit – simplest collection of atoms from
which an ionic compound can be formed




Formation of a sodium cation
Sodium atom
Na

[Ne]2s1
11 p+
11 e0
– e
Sodium ion

Na
+
[Ne]
11 p+
10 e1+
Formation of a chloride anion
unpaired electron

:Cl 
+ e

[Ne] 2s2 2p5
9 p+
9 e0
completed octet
1

:Cl:

[Ne] 2s2 2p6
9 p+
10 e1ionic charge
Learning Check
A.
Number of valence electrons in aluminum
1) 1 e2) 2 e3) 3 e-
B.
Change in electrons for octet
1) lose 3e2) gain 3 e3) gain 5 e-
C. Ionic charge of aluminum
1) 32) 5- 3) 3+
Ionic bonds
 Bond resulting from
electrostatic attraction
btwn cations and anions
 electrons are
TRANSFERRED
 + and – attract together
 Cations – lose e- (+)
 Anions – gain e- (-)
 As a result of
electrostatic
attraction
ex. NaCl, MgF2, KCl
Examples
K
+
Mg +
Cl
F
→
→
K Cl
F Mg
→ KCl
F
→ MgF2
Covalent Bonds
 Bond resulting
from the SHARING
of e- pairs btwn 2
atoms
 Orbital of e- from 1
element overlaps
orbital of e- from
another element
 e- are SHARED
Ionic bonds vs. Covalent bonds
Ionic bond
 Strong

+ and -
 High mp

Because of strength
 Hard and brittle
 Conduct electricity
in H2O

Break into ions
Covalent bond
 Weaker
 molecules
 Low mp
 B/o weak bonds
 Gases
 Do NOT conduct
 Don’t break up into
ions
Consider Chloride
HONC 1234
 Remember the HONC 1234 Rule




H = 1 bond
O = 2 bonds
N = 3 bonds
C = 4 bonds
Polar-covalent
 bond in which the atoms have an
unequal attraction for the shared e One end is more + and the other end is
more -
Examples
Nitrogen reacts with fluoride to form
nitrogen trifluoride
Carbon reacts with chloride to form
carbon tetrafluoride
Ionic, covalent, or polar?
 ionic or covalent bonds
can be estimated by
calculating the
difference in
electronegativities
 Difference of 1.7 or
less is covalent
 Difference of 0 to
0.3 are nonpolar
covalent
 Difference of 0.3 to
1.7 is polar
covalent
 Difference of higher
than 1.7 is ionic
Examples
Metallic Bonding
 A metal and a metal
 Outer E level orbitals overlap and allow
e- to roam freely throughout the entire
metal
 Sea of e-
Characteristics of metallic bonds
 Freedom of e- to move accounts for high
electrical and thermal conductivity
 Many orbitals, separated by small E
differences allows metals to absorb a
wide range of light frequencies
 e- fall back to lower E levels and give off E as
light
 Shiny (luster)
Characteristics of metallic bonds
 Malleability – ability to be hammered or
beaten into thin sheets
 Ductility – ability to be drawn or pulled into
a wire
 This can happen b/c metallic bonding is the
same in all directions
 One plane of atoms can slide past another
w/out breaking bonds
 Strength varies w/nuclear charge and # of
e- in the sea
Polyatomic Ions
 Charged group of covalently bonded atoms
 An ion made up of more than one atom
 Very strong covalent bonds hold them
together
 Behave as a SINGLE atom w/a charge
ex.
hydroxide OHnitrate NO3phosphate PO4-3
sulfate SO4-2
carbonate CO3-2
ammonium NH4+
Lewis Dot Diagrams for molecules
and polyatomic ions
Steps
 1. Count total valence e- for all atoms
 CO2
C = 4 e-(1) = 4 eO = 6 e-(2) = 12 eTotal = 16 e-
 2. Count the # of octet e- each atom wants
 All want 8 e- except H
 C is 1 octet = 8
 O is 1 octet = 8(2)
 Total = 24 e-
Lewis Dot Diagrams for molecules
and polyatomic ions
 3. Subtract valence e- from octet e- and this
equals bonding e 4. Divide bonding e- by 2
 This gives you the number of bonds
 5. Draw atoms in the arrangement that they
bond
 6. Find the lone pairs by subtracting bonding efrom valence e 7. Make sure that each e- is satisfied with an
octet
 Bonding: Lewis Dot Structures Practice
Examples