Chemistry 139 General Chemistry Prep Chemical Bonds 7
Chemistry& 139 General Chemistry Prep
Chemical Bonds 7. 16 Systematic Procedures For Drawing Lewis Structures 7. 15 Resonance Structures 7. 17 Molecular Geometry 7. 18 Electronegativity 7. 19 Bond Polarity 7. 20 Molecular Polarity
Chemical Bonds 7. 16 Systematic Procedures For Drawing Lewis Structures
In a Lewis structure we assume that the total number of valence electrons in a molecule are shared so each atom has a share of 8 electrons (2 for hydrogen). The first step in this process is to count the total number of valence electrons: e. g. CCl 4 has a total of 4 + 4 x 7 = 32 valence e-
The second step in the process is to determine the arrangement of atoms in the molecule or ion. There a few guidelines that help: • Hydrogen is always an outer atom • The element with the lower group number will be the central atom as it needs to form the most bonds to obtain an octet. • For elements in the same group the element with the higher period number is usually the central atom.
The second step is to determine the arrangement of atoms in the molecule or ion. e. g. in CCl 4 Carbon would be the central atom (group IV A) and the chlorines (group VIIA) would be the outer atoms Cl Cl Cl
The third step is to connect each of the outer atoms to the central atom via a single bond. • Add up the number of electrons shared • Subtract this from the total number of valence electrons. 32 – 4 x 2 = 24 e-
The fourth step is to add “lone pairs” to complete the octets of the outer atoms, subtract the number of electrons added from the total: 24 – 6 x 4 = 0 e-
The fifth step is to place any remaining valence electrons as lone pairs on the central atom and subtract this from the total : : Cl : C : : : Cl Cl : : : : Cl : : No remaining e- step 5 not needed
The sixth step is to check if the central atom has a complete octet. If not bring in electron pairs from the outer atoms to form higher order bonds until all atoms have an octet. : : Cl : C : : : Cl Cl : : : : Cl : : Carbon has complete octet step 6 not needed
Write the Lewis structure for CO 2: Step 1: 4 + 2 x 6 = 16 valence electrons Step 2: C in group 4, O in group 6 C is central atom O__C__O Step 3: 16 – 4 = 12 valence electrons left to place O__C__O : : Step 4: 12 – (2 x 6) = 0 valence electrons to place : : Step 5: not needed
: : O=C=O : : Step 6: Form two double bonds to complete octet on carbon Having practiced writing Lewis structures several trends become apparent, typically we observe: • When a halogen is an outer atom it forms one bond • Oxygen atoms most often form two bonds • Nitrogen atoms most often form three bonds • Carbon atoms form four bonds
Chemical Bonds 7. 16 Systematic Procedures for Drawing Lewis Structures In drawing the Lewis structure for covalent compounds we consider all the valence electrons to be pooled. These electrons are then distributed in a systematic fashion to ensure each atom receives a share of 8 (or for H 2) electrons
Chemical Bonds 7. 15 Resonance Structures
Lets draw the Lewis Structure for O 3: Step 1: 3 x 6 = 18 valence electrons Step 2: Oxygen is the central atom Step 3: O O O 18 – 4 = 14 valence electrons left to place Step 4: . . : O. . 14 -12 = 2 valence electrons left to place O . . O: . .
Step 5: . . : O. . O: . . 2 -2 = 0 valence electrons left to place Step 6: “complete the octet on the central atom by bringing in lone pairs from the outer atoms”. . : O. . O:
In the final step of drawing the Lewis structure of O 3 there are two choices. . . O: . . : O. . O: . O. . . O . . : O. . Or we could draw: . . O: . . : O. . Is one structure more valid than the other?
When we experimentally examine the structure of ozone we find all the bond lengths are the same! This is inconsistent with either of a our two Lewis structures! O: . O. . . O . . : O. . The correct structure is actually an “average” or resonance hybrid of the two structures. We indicate this using a double headed arrow.
Example: Draw the Lewis structure of the nitrate anion. Step 1: NO 3 - # valence e- = 5 + (3 x 6) +1 = 24 e. Step 2: N is in group 5, oxygen is in group 6, N is central atom Step 3: 24 – 6 = 18 e- O O N O
Chemical Bonds 7. 15 Resonance Structures Resonance occurs when more than one correct Lewis structure can be drawn for a molecule. If the resonance structures are equivalent the molecule can be described as a resonance hybrid of these structures.
Chemical Bonds 7. 17 Molecular Geometry
Lewis diagrams and the octet rule are all useful tools in exploring the arrangement of electrons in atoms, molecules and ions. However; • These models approximate reality and we will observe exceptions to the “rules”. • Lewis diagrams show which nuclei electrons are associated with, they do not indicate how the nuclei are arranged in 3 D space.
Valence-shell electron pair repulsion (VSEPR) theory proposes that: “each group of electron pairs around a central atom is located as far away from each other group as possible to minimize repulsions” Where a group of electrons maybe: 1. A lone electron pair 2. A bond of any type (single, double, triple)
We will consider the electron group arrangements where we have 2, 3, or 4 electron groups around the central atom: These electron group arrangements are illustrated below: 2 electron groups 3 electron groups 4 electron groups
With two electron groups we have a linear arrangement of groups and a bond angle of 1800: An example includes A central atom in this configuration must have double or triple bonds to obtain an octet e. g. CO 2 with 4 electron pairs:
With three electron groups we have a trigonal planar arrangement of groups and an ideal bond angle of 120 o: e. g. Formaldehyde: Central atoms in this arrangement with a complete octet will form a double bond
With three electron groups when one electron group is a lone pair we have a bent or angular structure with an ideal bond angle of 120 o: e. g. ozone:
When we have four electron groups around the central atom the molecule or ion has a tetrahedral arrangement. e. g. CH 4: In this arrangement each bond angle is 109. 5 o.
Replacing 1 atom with a lone electron pair results in the trigonal pyrimidal structure: e. g. ammonia with one lone pair
Replacing 2 atoms with lone electron pairs results in the angular or bent structure: e. g. water with two lone pairs
The molecular shape observed depends upon the electron group arrangement and the number of lone pairs:
Chemical Bonds 7. 17 Molecular Geometry VSEPR theory proposes that electron groups in molecules and polyatomic ions arrange themselves in 3 D space so as to minimize repulsions. As a result the number of electron groups a molecule or polyatomic ion has can be used to predict its 3 D shape and its bond angles.
Chemical Bonds 7. 18 Electronegativity
Atoms form covalent bonds with each other by sharing electrons to form molecules. Consider two hydrogen atoms separated by a large distance, each has 1 electron and 1 proton. Now lets bring the two atoms together so the path the electrons move along overlap.
Consider what happens when the electrons are located exactly in between each nucleus. Each nuclei feels an inwards attraction to the two electrons that will hold the two atoms together forming a H 2 molecule. This type of interaction is called covalent bonding.
When we have two identical atoms the electrons are shared equally and lie exactly half way between each atom in the covalent bond. Most bonds involve atoms of different types and the electrons are not shared evenly between the atoms. e. g. consider HCl Shared electrons are closer to chlorine atom than the hydrogen atom. The bond is a + - polar covalent bond. Each bond has a more positive and more negative end e. g. dipole
We can predict what pairs of atoms will form polar covalent bonds with uneven sharing of electrons using the periodic table. H Increasing tendency to pull electrons
There have been several schemes devised to quantitatively express the electronegativity (EN) of the elements. We will use the relative EN values devised by Linus Pauling only person to have independently won two Nobel prizes in different disciplines. PNW local (Oregon). Never did graduate high school (didn’t complete Humanities requirement).
Chemical Bonds 7. 18 Electronegativity is the tendency of an atom to pull electrons towards itself in a covalent bond. Electronegativity increases from the bottom left corner of the periodic table towards the top right of the periodic table.
Chemical Bonds 7. 19 Bond Polarity
When atoms with different EN share electrons the electrons are not shared evenly. e. g. consider HCl The bond has a more positive and more negative end e. g. dipole. This can be indicated using “delta notation” or a “polar arrow”.
Bonds may be classified on the basis of the difference in EN and the types of atoms present
Chemical Bonds 7. 19 Bond Polarity The sharing of electrons between atoms with different electronegativity results in a polar bond. The negative and positive ends of a polar bond can be indicated using a polar arrow or “delta notation”. The difference in electronegativity between two atoms in a bond can be used to classify a bond as nonpolar covalent (ΔEN ≤ 0. 4), polar covalent ( 0. 4 < ΔEN ≤ 1. 5), ionic or polar covalent (1. 5 < ΔEN ≤ 2. 0) or ionic (ΔEN > 2. 0).
Chemical Bonds 7. 20 Molecular Polarity
Molecules as well as bonds can have polarity. Polar molecules have an unsymmetrical distribution of electronic charge Water molecules are polar because they are bent in shape. If they were linear the molecule would not be polar. H O H The effect of the polar bonds would cancel!
Because they are polar water molecules attract each other. This is called dipole-dipole interaction. The dipole-dipole interaction in water is called a “hydrogen bond”, these types of bonds are extremely important in biochemistry.
In general if we have an even (symmetrical) distribution of electron groups around the central atom the molecule will be non-polar. e. g. Two identical polar bonds pointing in opposite directions. e. g. Four identical polar bonds in a tetrahedral arrangement.
If we have an uneven (asymmetrical) distribution of electron groups around the central atom the molecule will be polar. e. g. Two identical polar bonds in a bent arrangement. e. g. Three identical polar bonds in a trigonal pyrimidal arrangement.
Chemical Bonds 7. 20 Molecular Polarity An asymmetrical distribution of electron groups about the central atom gives rise to a polar molecule. Polar molecules are easily recognized because there is always more than one “type” of thing (atom or lone pair) attached to the central atom.
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