Unit 2: All Biology is Chemistry Lesson 2: Compounds and Chemical Bonding.
Chemical Bonding Chapter 8 & 9. I. Overview/Types of Compounds.
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Transcript of Chemical Bonding Chapter 8 & 9. I. Overview/Types of Compounds.
Chemical Bonding
Chapter 8 & 9
I. Overview/Types of Compounds
Overview of Bonds• Chemical Bonds are attractions
between atoms or ions in a compound• A bond forms when valence electrons
are either shared or transferred from one atom to another
• Most elements are more stable when bonded to other atoms, which is why bonds are formed.
Compounds
Before we start learning details about bonding, let’s review
compounds!
A. Definition of a Compound
• Remember, a compound consists of 2 or more elements chemically bonded together. It is always neutral.
EX: H20, NaCl, Sb3(PO4)5
B. Types of Compounds
2 types of chemical compounds
1. Molecular or Covalent Compound:
• consists of only non-metal atoms that share electrons which form a covalent bond.
– Examples: H20, C8H18
– smallest unit is called a molecule – molecules also include diatomics
(ex. H2)
2. Ionic Compound:• a compound that is formed from
the attraction between ions of opposite charges. It is held together by an ionic bond and usually contains a metal and nonmetal.
Ionic Compounds• Examples: NaCl, KBr• smallest unit is called a formula
unit• Cations are + ions. This happens
when an atom loses electrons. (ca+ion)
• Anions are – ions. This happens when an atom gains electrons. (a negative ion)
Practice1. Determine whether the following are
ionic or molecular covalent compounds
a. N2O5 b. PbNO5
c. KF d. AgCl
e. PCl3
ionic
molecular
molecular
ionic
ionic
always1+
2+ 3+ 3- 2- 1-
Practice2. Write the ion formed by these
elements
a. iodine
b. potassium
c. sulfur
d. gallium
e. strontium
S-2
I-
Sr2+
Ga+3
K+
C. Properties of Compounds
• ionic and molecular covalent compounds typically have very different properties
• this stems from the fact that the ionic attraction between formula units is far stronger than the attraction between covalently-bonded molecules
Comparison of PropertiesCharacteristic Molecular
CompoundsIonic Compounds
representative unit
Molecule Formula Unit
type of elements Nonmetals only Metals and Nonmetals
type of bonding Covalent (sharing)
Ionic (transfer)
physical state at RT
S, L, or G Solid only
melting point <300 °C ** > 300 °C
hardness soft hard hard
conduct electricity?
no yes (when dissolved in water)
Examples Candle wax Rock salt (NaCl)
**would melt in a bunsen burner
Ionic Compounds = transfer of electrons
Na Cl+ -
NaClIs formed
Molecular Compounds = share electrons
H O
H2OH
II. Types of BondsWe will learn about 4 different types of bonds: ionic, non-polar covalent,
polar covalent, and metallicBonds between two elements can be categorized based on the difference between the electronegativities of
those two elements.
A. Ionic Bonds
Ionic Bonds• Ionic bonds are the result of the
transfer of electron. It holds an ionic compound together.
• Remember, an ionic compound is a compound formed from the attraction between a cation and anion, which usually includes a metal and non-metal.
Ionic Bonds
• After a cation and anion are bonded to form an ionic compound, the net charge is zero
• The Electronegativity difference between the two atoms bonded is >1.7
[ ]-
Ionic Bonds
Na Na+
Cl Cl
Ionic Compounds = transfer of electrons
Na Cl+ -
NaClIs formed
Drawing ionic compounds
[ ]Cl + -
NaRemember when drawing ionic
compounds, only the anion is in brackets & has its valence electrons showing
Another Example
2+[
]O
2-
Ca
Ionic Bonds Strongest type of bond between
atoms. This results in high melting points and high boiling points.
Formula units form a crystalline structure
a120o
c
a
B. Covalent Bonds
• A covalent bond is formed by the sharing of electron-pairs within molecules
• Covalent bonds are formed between non-metals
Ex. H OH
Covalent Bonds
• Two atoms can share more than one covalent bond between them– Single bonds form between two atoms that
SHARE 1 PAIR of electrons. – Double bonds form between two atoms that
SHARE 2 PAIRS of electrons.
– Triple bonds form between two atoms that
SHARE 3 PAIRS of electrons.
Triple Bonds have the shortest bond length
Bond Length Bond Energy
H H Longest Weakest(163 kJ/mol)
O O Medium (418 kJ/mol)
N ≡ N Shortest Strongest (945 kJ/mol)
• There are two types of covalent bonds: non-polar covalent and polar covalent
Non-Polar Covalent Bonds
• electronegativity difference between atoms: 0-0.3
• even distribution of electrons • symmetrical
observed in all diatomic atoms (BrINClHOF elements)
Br I ClN HOF
Polar vs Non-PolarCovalent Bonds
Polar Covalent Bonds
• electronegativity difference between atoms: 0.3-1.7
• polar means “having opposite ends”
• unequal sharing of charges
Remember Electronegativity?
© 2003 Prentice Hall
Polar Covalent Bonds
• A dipole is formed
• the partial charges can be shown with the following symbols: δ - and δ+
Unequal sharing of electrons
© 2003 Prentice Hall
Polar Covalent Bonds
• electrons are more attracted to the more electronegative element
• Opposites attract concept is observed
Polar or Non-Polar?
Has an even distribution of charge
Non-Polar!!
Polar or Non-Polar?
Has a ΔEN = 0
Non-Polar!!
Polar or Non-Polar?
Has a δ+ and δ- end
Polar!!
Polar or Non-Polar?
I2
Non-Polar!!
Polar Means:
“Having Opposite Ends”
(like N and S Poles)
C. Metallic Bonds
C. Metallic Bonds
• A metallic bond is the bond that holds two metals together. The chemical bonding that results from the attraction between metal atoms and the surrounding sea of electrons.
Metallic Bonds• The bonding
electrons are shared over the whole metal sample. This is known as electron delocalization.
• Because of the nature of metallic bonding, metals are malleable and ductile.
• -malleable = ability of a substance to be hammered or beaten into thin sheets
• -ductile = ability of a substance to be drawn or pulled through a small opening to produce a wire
Summary ChartDon’t forget your general rules of Ionic
compounds = Metal + Non-Metal
Type of Bond (between 2 atoms)
ΔEN(Paulings)
Metallic bondnot quantified
using EN
Ionic(transfer of electrons)
>1.7
Polar Covalent(unequal sharing of
electrons)0.4 – 1.7
Non-Polar Covalent(equal sharing of electrons)
0 – 0.4diatomicsUse a table/chart with electronegativity values to determine the difference
Electronegativity Values
0.3 1.7
© 2003 Prentice Hall
H – F ΔEN1.9
H – Cl ΔEN0.9
H – BrΔEN0.7
H – IΔEN0.4
Most Polar Least Polar
• Although a compound is an ionic compound, it can still possess covalent bond characteristics and vice versa
Ex. H – F
Ionic Character Increases
Covalent Character Increases
2.1 4.0
ΔEN = 1.9
While this is a molecular compound, it possess ionic bond characteristics
III. Lewis Structures
III. Lewis Structures
• Only for molecular compounds!
• Recall drawing electron dot diagrams using the number of valence electrons (Roman Numerals on PT)
Lewis Structures are drawings!!
• Symbol = Nuclei
• Dots = Valence Electrons
• Dashes = Shared Electron Pair
• Electron Pair = unshared (aka “lone pair”)
= electron pair
1 bond = 2 shared electrons
• Electron pairs that are shared can be replaced with a line to represent a covalent bond. Up to three bonds can be shared between two elements
• Single bonds are formed between two elements that have 2 shared electrons (Ex. )
• Double bonds are formed between two elements that share 4 valance electrons (Ex. )
• Triple bonds are formed between two elements that share 6 valence electrons (Ex: )
Triple Bonds have the shortest bond length
Bond Length Bond Energy
N N Longest Weakest(163 kJ/mol)
N N Medium (418 kJ/mol)
N ≡ N Shortest Strongest (945 kJ/mol)
Steps to Drawing Lewis Structures
• Choose the center atom (Tips: choose the one you have less of, and carbon loves to be the center of attention)
• Draw an electron dot structure for the central atom• Draw the remaining electron dot structures near the
central atom’s unpaired electrons• Create bonds between the unpaired electrons• Rearrange your structure so it looks “neat” (spread the
atoms out)• Add lone pairs so elements 1-5 have “duets” and all
other elements have “octets”
Disclaimer: These rules actually only work for the examples we’re doing in our class. There are far more extensive rules that we will ignore for the sake of simplification.
Octet Rule• Octet Rule: chemical compounds tend to
form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level
-Elements 1-5 need to be surrounded by 2 electrons-All other elements need to be surrounded by 8 electrons
Same as
© 2003 Prentice Hall
Examples
CCl4
C
Cl
Cl
Cl
Cl
Examples
C2H2
C HH C
C HH C
Are the duets & octets complete?
ExamplesCO2
OO C
OO C
Are the duets & octets complete?
C OO
ExamplesCO2
OO C
Are the duets & octets complete?
Exceptions to the octet rule (incomplete octets)
Be HH
B F
F
FBF3
BeH2
Exceptions to the Octet Rule
FF
F
BBe HH
BF3
3 valence &3 bonded atoms
BeH2
2 valence &
2 bonded atoms
What does polar and non-polar mean again?
Review on how to draw structures (Ex. NH3)
N H
H
H
Polar
N H
H
H
S SNon-Polar
EN = 2.1
EN = 3.0EN = 2.1EN = 2.1
Polar & Non-Polar Molecules
• Like bonds, molecules can be polar (uneven distribution of electrons) or non-polar (even distribution)
• Non-polar molecules are symmetrical and have no lone pairs around the center atom
• Polar molecules have lone pairs around the central atom or an uneven distribution of electrons
NP OO C
PracticeIdentify each molecule as polar or
non-polar
FF
F
BNP
B F
F
FB
F
F
F
P
PracticeIdentify each molecule as polar or
non-polar
NP
C2HCl
PracticeIdentify each molecule as polar or
non-polar
P
C ClH C
Polar & Non-Polar Molecules
© 2003 Prentice Hall
Dipoles
• Some molecules have dipoles that cancel out and have a zero net dipole
• The poles can be combined to form a net dipole
H HO
Positive end
Negative end
Polar or Non-polar Molecules?
CCl4
C
Cl
Cl
Cl
Cl EN = 2.5
EN = 3.0
EN = 3.0
EN = 3.0
EN = 3.0
C2H2
C HH C
Polar or Non-polar Molecules?
CO2
OO C
Polar or Non-polar Molecules?
N H
H
H
A better way to draw it
NH
HH
Br BrBr Br
B F
F
F
Draw BF3
B
F
F
F
IV. VSEPR Theory
VSEPR Theory
• VSEPR stands for Valence Shell Electron Pair Repulsion
• VSEPR is used to describe the 3D orientation of the electron regions.
• Shared electrons that make up covalent bonds and lone pairs will repel each other as much as possible
B
F
F
F NH
HH
BF3 vs NH3
Shape: Trigonal Planar
Shape: Trigonal pyramidal
Lone pairs repel more than bonds do
BF3 vs NH3
Shape: Trigonal Planar Shape: trigonal pyramidal
Look at the chart on page 7 of your notes
Ahhh! Memorization!?
• You will be responsible for predicting shapes with a * next to the name of shape without any notes
• You should be able to determine the shape of non - * compounds given this chart
1 lone pair
Trigonal planar*
/bent*
Linear*
1 lone pair 2 lone pairs
tetrahedral*
/bent*