covalent bonding chapter 9. the covalent bond section 9.1 why do atoms bond? to achieve full outer...
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Covalent Bonding
Chapter 9
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The Covalent BondSection 9.1 Why do atoms bond?
To achieve full outer electron shells Octet Rule- atoms gain, lose, or share electrons
to achieve the electron configuration of noble gases
Gain and Lose IONIC BONDING Share COVALENT BONDING
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What is a covalent bond?
A chemical bond that results in the SHARING of valence electrons
Occurs between 2 or more nonmetals
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A molecule is formed when two or more atoms bond covalently Examples: sugars, DNA, proteins, fats,
carbohydrates, cotton, synthetic fibers
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Formation of a covalent bond REMEMBER: Hydrogen (H2), Nitrogen (N2),
Oxygen (O2), Fluorine (F2),Chlorine (Cl2) Bromine (Br2) and Iodine (I2) occur in nature as diatomic molecules
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Covalent Bonding
An attractive force occurs between the protons of one atom and the electrons of the other atom
When a single pair of electrons is shared, such as in the hydrogen molecule, a single covalent bond forms
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Lewis Structures
Use electron-dot diagrams to show how electrons are arranged in molecules
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Group 7A (Halogens) have 7 VE One more VE is necessary A single covalent bond will form
Group 6A have 6 VE Two more VE are necessary Two covalent bonds will form
Group 5A have 5 VE Three more VE are necessary Three covalent bonds will form
Group 4A have 4 VE Four more VE are necessary Four covalent bonds will form
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Sigma Bonds- single covalent bonds- when electron pairs are centered between two atoms
Multiple Bonds In many molecules, atoms
attain noble gas configuration by sharing more than one pair of electrons between two atoms
Carbon, Nitrogen, Oxygen, and Sulfur most often form multiple bonds
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Strength of Covalent Bonds
The strength of covalent bonds depends on how much distance separates both nuclei
The distance between the two bonding nuclei at the position of maximum attraction is called bond length Determined by the size of the atoms and how
many electron pairs are shared Bond length decreases as the number of bonds
increases (triple bond has a shorter bond length than a single bond)
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Energy Changes
An energy change accompanies the forming or breaking of a bond between atoms in a molecule. Energy is released when a bond forms Energy must be added to break the bonds of a
molecule The amount of energy required to break a
specific covalent bond is called bond dissociation energy
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Bond dissociation energy indicates the strength of a chemical bond because a direct relationship exists between bond energy and bond length
In chemical reactions, bonds in reactant molecules are broken and new bonds are formed as product molecules form
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Endothermic reactions occur when a greater amount of energy is required to break the existing bonds in the reactants than is released when the new bonds from in the product molecules
Exothermic reactions occur when more energy is released forming new bonds than is required to break bonds in the initial reactants
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Checkpoint
What is a covalent bond? How does it differ from an ionic bond?
What type of elements form covalent bonds? Draw the Lewis Structures for each of these
molecules: PH3
H2S
CCl4
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Naming Molecules Section 9.2 Naming Binary Molecular Compounds
1. The first element in the formula is always named first, using the entire element name.
2. The second element in the formula is named using the root of the element and adding the suffix – ide.
3. Prefixes are used to indicate the number of atoms of each type that are present in the compound.
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Prefixes in Covalent Compounds
Number of Atoms
Prefix Number of Atoms
Prefix
1 Mono- 6 Hexa-
2 Di- 7 Hepta-
3 Tri- 8 Octa-
4 Tetra- 9 Nona-
5 Penta- 10 Deca-
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Practice Naming
CCl4 As2O3
CO SO2
NF3
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Naming Acids
Binary Acids Use the prefix hydro- to name the hydrogen part
of the compound The rest of the name consists of a form of the root
of the second element plus the suffix –ic, followed by the word acid.
Examples: HCl Hydrochloric Acid HCN Hydrocyanic acid (even though there are more
than 2 elements present, if no oxygen is present- the acid is named as a binary)
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Oxyacids
Acids that contain an oxyanion (polyatomic ion that contains oxygen)
The name of the oxyacids consists of a form of the root of the anion, a suffix, and the word acid If the anion suffix is –ate, it is replaced with the
suffix –ic If the anion suffix is –ite, it is replaced with the
suffix –ous.
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Practice Problems
HI HClO3
HClO2
H2SO4
H2S
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Checkpoint 1. Write the molecular
formula for each of the following compounds
Disulfur trioxide Iodic acid Dinitrogen monoxide Hydrofluoric acid Phosphorus pentachloride
2. What is the difference between a binary acid and oxyacid?
3. Complete the following table
Formulas and Names of Covalent
Compounds Formula Name
PCl5
Hydrobromic acid
H3PO4
Oxygen difluoride
SO2
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Molecular Structures Section 9.3 Structural Formula- uses letter symbols and
bonds to show relative positions of atoms Lewis Structure Procedure
1. Predict the location of certain atoms Hydrogen is always a terminal, or end, atom. Because it
can share only one pair of electrons, hydrogen can be connected to only one other atom
The atom with the least attraction of shared electrons in the molecule is the central atom. This element usually is the one closer to the left on the periodic table. The central atom is located in the center of the molecule, and all other atoms become terminal atoms.
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2. Find the total number of electrons available for bonding. This total is the number of valence electrons in the atoms in the molecule.
3. Determine the number of bonding pairs by dividing the number of electrons available for bonding by two.
4. Place one bonding pair (single bond) between the central atom and each of the terminal atoms.
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5. Subtract the number of pairs you used in step 4 from the number of bonding pairs you determined in step 3. The remaining electron pairs include lone pairs as well as pairs used in double and triple bonds. Place lone pairs around each terminal atom bonded to the central atom to satisfy the octet rule. Any remaining pairs are assigned to the central atom.
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6. If the central atom is not surrounded by 4 electron pairs, it does not have an octet. You must convert one or two of the lone pairs on the terminal atoms to a double bond or a triple bond between the terminal atom and the central atom. These pairs are still associated with the terminal atom as well as with the central atom. Remember that, in general, carbon, nitrogen, oxygen, and sulfur can form double or triple bonds with the same element or with another element.
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Ball-and-stick molecular model
Space-filling Structure Lewis Structure
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Examples
NH3
PO4-3
CO2
BH3
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Resonance Structures
Resonance is a condition that occurs when more than one valid Lewis structure can be written for a molecule or ion
Nitrate Ion Resonance Structures:
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Exceptions to the Octet Rule1. A small group of molecules has an odd
number of valence electrons and cannot form an octet around each atom. (i.e. NO2 , ClO2, and NO)
2. Some compounds form with fewer than eight electrons present around an atom. (Example: Boron)
When one atom donates a pair of electrons to be shared with an atom or ion that needs two electrons to become stable, a coordinate covalent bond forms
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3. Some central atoms contain more than eight valence electrons (expanded octet)
Examples: PCl5, SF6, and XeF4
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Molecular ShapeSection 9.4 The shape of the molecule determines many
of its physical and chemical properties. Molecular shape is determined by the overlap
of orbitals that share electrons Valence Shell Electron Pair Repulsion model
or VSEPR model Based on an arrangement that minimizes the
repulsion of shared and unshared pairs of electrons around the central atom
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VSEPR Model The angle formed by any two terminal atoms
and the central atom is a bond angle. Shared electron pairs repel one another Lone pairs of electrons occupy a slightly larger
orbital than shared electrons Shared bonding orbitals are pushed slightly
together by lone pairs
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Look at the VSEPR cheat sheet
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Hybridization
A hybrid results from combining two of the same type of object, and it has characteristics of both. Hybridization- a process in which atomic orbitals
are mixed to form new, identical hybrid orbitals.
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Practice Problems Determine the molecular geometry and bond
angle for the following: BF3
NH4+
OCl2
CF4
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Electronegativity and Polarity Section 9.5 Electron affinity is a measure of the
tendency of an atom to accept an electron
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The character and type of a chemical bond can be predicted using the electronegativity difference of the elements that are bonded Polar Covalent- Unequal sharing Nonpolar Covalent- Equal sharing
Generally- Ionic bond form when the
electronegativity difference is greater
than 1.70
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Polar Covalent Bonds Polar covalent bonds form because not all
atoms that share electrons attract them equally The shared pair of electrons is pulled toward one
of the atoms Partial charges occur at the ends of the bond
Partially negative
Partially positive
The resulting polar bond is referred to as a dipole (two poles)
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Molecular Polarity
Molecules are either polar or nonpolar, depending on the location and nature of the covalent bonds they contain. A polar molecule has a partial negative charge on
one side, while the other side of the molecule has a partial positive charge
Note:Note: symmetric molecules are usual nonpolar and molecules that are asymmetric are usually polar
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Polar Molecule or not?
Compare H2O and CCl4
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Solubility of polar molecules
The ability of a substance to dissolve in another substance is known as the physical property solubility
The bond type and the shape of the molecules present determine solubility
“Likes dissolve likes” Polar compounds are usually soluble in polar
substances Nonpolar molecules only dissolve in nonpolar
substances
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Properties of Covalent Compounds 1. Lower melting and boiling points (indicating
weak bond strength)
2. Many are liquids or gases at room temperature
3. Do not conduct electricity
4. Many do not dissolve in water (polar)
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Practice Problems
Decide whether each of the following molecules is polar or nonpolar SCl2
H2S
CF4
CS2
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Vocabulary on Test
Structural formula molecule VSEPR model Coordinate covalent
bond hybridization oxyacid electronegativity
polar covalent covalent bond resonance endothermic exothermic terminal atom Sigma bond