Chapter 5 (2nd Half) Covalent Bonding - Lecture Slides Added Notes from the prof PDF

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chapter 5( 2nd half) summery by the teacher + its has notes from the teacher himself...

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Molecules all have 3-D shapes (called molecular geometry). In this course, you will need to be able to correctly identify the molecular geometry of molecules with one central atom.

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Remember that all atoms are a positive nucleus surrounded by a negative electron cloud. Also, all electrons in bonds and lone pairs are negatively charged. So consider that every bond and lone pair around an atom in a molecule will push against the others. All of the bonds and lone pairs repel each other and try to be as far apart from each other as they can. This is the basis of VSEPR theory.

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Everything is arranged around the inner atoms (central atoms) in a molecule so that is the focus. In this course, you will only be expected to determine the molecular geometry of molecules containing 2-5 atoms with one central atom at most. There are basically 3 steps to determining molecular geometry: 1)

Count the VSEPR groups around the central atom.

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Separate the VSEPR groups as far apart from each other as possible.

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Only consider all bonded atoms to determine the correct 3-D shape. This means understanding that lone pairs on a central atom are invisible, even though they repel just like bonds and electron clouds in atoms do.

Single bonds, double bonds, triple bonds and lone pairs around the central atom all count equally as one VSEPR group.

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A correct Lewis structure is critical for determining the correct 3-D structure of a compound. So make sure that you follow all steps outlined in creating the Lewis structure.

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Note how all of the VSEPR groups in the 3 molecules shown on this slide are treated equally, no matter the type.

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Once you have correctly created the Lewis structure and identified the VSEPR groups, you can start to determine the 3D structure. Think of this as a 2-step process: 1)

Molecules we study in this course will have a maximum of 1 central atom. Surrounding the central atom will be either 2, 3 or 4 VSEPR groups. Separate the VSEPR groups to get them as far apart as possible.

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Only consider atoms when determining the final structure. While lone pairs of electrons around the central atom are equal VSEPR groups and will separate equally just as bonds do, they are invisible and do not contribute visually to the final structure.

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Two VSEPR groups: Note how the 2 double bonds in CO2 are treated equally as 2 VSEPR groups, just as the single and triple bonds are in HCN. To separate 2 VSEPR groups, place them directly across from each other at 180o. Looking at the structures, you can see that the atoms are in a straight line, meaning that the structures are linear. When you have 2 VSEPR groups, the only option for 3-D structure is linear.

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Three VSEPR groups: Three groups can be separated equally by placing them 120o apart from each other all in the same flat plane. Imagine that if you lay each of the structures down on a flat table, all of the atoms will touch the table (this means that they are all in one plane). Formaldehyde – the structure on the left has 1 double bond and 2 single bonds as its 3 VSEPR groups. There are 4 atoms C, O, H, H – since you can see them all, the structure is considered Trigonal Planar (a triangular shape with all atoms in the same plane – think that if you draw a shape around all atoms, you will create a triangle). Sulfur dioxide – the structure on the right has 1 double bond, 1 single bond and 1 lone pair of non-bonding electrons as its 3 VSEPR groups. While the 3 groups separate into the same trigonal planar arrangement of 120o of separation, the lone pair above the Sulfur atom is invisible because it is just 2 tiny electrons, not an atom. This means that all that you see are the 3 actual atoms, S, O and O. If you trace around the atoms only and ignore the lone pair above the S atom, you will have an Angular structure (also called “Bent”). Therefore, when you have 3 VSEPR groups around a central atom, you have 2 choices for 3-D structure. If all VSEPR groups are bonded atoms, the structure will be Trigonal Planar. If one of those VSEPR groups is a lone pair of non-bonding electrons, the structure will be Angular. 10

Four VSEPR groups: For maximum separation of 4 groups, the structure formed is called Tetrahedral. Think of an orange as your central atom: If I handed you 4 toothpicks (acting as 4 VSEPR groups) and asked you to place those toothpicks in the orange as far apart from each other as possible, you would naturally create a Tetrahedral structure. However, the final visible structure you end up with has to do with how many atoms are present as VSEPR groups versus lone pairs. Molecules containing four VSEPR groups can be one of 3 possible molecular geometries: 1)

All 4 groups as bonded atoms – all groups visible – final 3-D shape remains Tetrahedral, just as how the VSEPR groups were originally separated. Example shown is methane (CH4). The separation angle between all bonded atoms is 109.5o.

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3 groups as bonded atoms and 1 group as an invisible lone pair – only bonded atoms are visible – results in a Trigonal Pyramidal structure – means “Triangular-shaped pyramid”. The example shown is Ammonia (NH3). The lone pair on top of Nitrogen pushes the H atoms in Ammonia away from each other but the lone pair itself is invisible. All you see is the N atom on top and the H atoms surrounding below it. The H atoms are separated from each other at 107o. The angle is not 109.5o since lone pairs push away a little stronger than regularly bonded atoms. 11

3)

2 groups as bonded atoms and 2 groups as invisible lone pairs – results in an Angular (“Bent”) structure. The example shown is water (H2O). The two H atoms are pushed away by the 2 lone pairs to an angle of 104.5o. Since you only see the bonded atoms, the structure appears angular.

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You do not have to worry about molecules with more than one central atom in this course. However, this information will be important next term.

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Covalent bonds of course involve sharing electrons. You might think that sharing means that all electrons evenly go back and forth between the two atoms that are sharing them. However, not all electrons are shared evenly. Some elements can pull on shared electrons more than others. Think of every bond as an electron tug-of-war between the 2 elements involved. The strength of each element in pulling on shared electrons is called electronegativity.

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This numerical scale was developed by Linus Pauling and gives elements their relative electronegativity strength. By comparing the electronegativity values between the two elements in a bond, you can determine whether the electrons are being pulled significantly toward one element more than the other. If they are, they will make that element more negatively charged (since electrons are negative) and the other element not seeing the electrons as much will be positively charged. The electronegativity difference determines whether the bond would be considered non-polar, polar-covalent or ionic.

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When considering any two elements that come together, to determine how they will interact regarding bonding you simply determine the difference between their electronegativity values. This slide is just introducing the concept of recognizing electronegativity differences by using the numerical scale.

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In chemistry, polarity is a measure of charge separation (the degree of positive and negative charge on atoms close together). When determining polarity, you first consider the electronegativity values of the atoms being compared. The closer the 2 electronegativity values are to each other, the more evenly the electrons are shared between the two atoms and the less polarity that is present. The more different the 2 values are, the more the electrons are unevenly shared and the higher the polarity.

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When comparing elements bonding together, there are three potential outcomes. This is a simplified version where we will only consider three basic categories. You may see 5 categories in the textbook where metalnonmetal interactions are separated from nonmetal-nonmetal interactions. We will keep it more straightforward and go by the electronegativity differences alone. 1)

Nonpolar covalent – This is where 2 elements bonding together have very similar or identical electronegativity values. If the difference between the two elements is 0.4 or less, the electrons are considered to be shared equally and no charge polarity develops. Charge polarity means that one element becomes significantly negative (due to having more electrons closer to it) and the other becomes significantly positive (due to having the electrons pulled away from it). In this case, since the electrons are evenly pulled toward each element, neither one of them becomes charged.

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2)

Polar covalent – This occurs when the two elements have a difference of electronegativity ranging between 0.4 and 2.0 (unequal sharing). This means that there is significant pull of the electrons toward one element, resulting in a permanent negative charge on that element and a permanent positive charge on the other. This results in a polar covalent bond meaning that the charges are present. These charges are considered dipoles, also called dipole moments.

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Ionic - This occurs when the difference in electronegativity between the 2 elements is so great that the electrons are not shared but instead are pulled away from one element to the other. This ionic bond is exactly what was covered in Chapter 4.

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Polar bonds can be shown by either drawing delta (that weird looking d) with a positive sign or negative sign over the atoms in the bond. This shows that permanent positive and negative regions have developed in the molecule (middle image). You can also show polar bonds by drawing an arrow pointing toward the atom with the higher electronegativity (the arrow points toward the negatively-charged atom). Then you add a plus sign to the arrow end over the positively charged atom in the bond.

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The bigger the difference in values, the more polar the interaction. A) C – 2.5 ; F – 4.0 = 1.5 difference : N – 3.0 ; F – 4.0 = 1.0 difference : O – 3.5 ; F – 4.0 = 0.5 difference B) Si – 1.8 ; F – 4.0 = 2.2 difference : C – 2.5 ; F – 4.0 = 1.5 difference : N – 3.0 ; O – 3.5 = 0.5 difference C) B – 2.0 ; Cl – 3.0 = 1.0 difference : S – 2.5 ; Cl – 3.0 = 0.5 difference : Cl – 3.0 ; Cl – 3.0 = 0 difference

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Least polar means least positive and negative charges developing on the two atoms – electronegativity values will be more similar between the two elements Remember: Electronegativity difference and bonding : 0 – 0.4 difference – nonpolar bond ; 0.5 – 2.0 difference – polar covalent bond ; > 2.0 difference – ionic bond Mg-O : Mg – 1.2 ; O – 3.5 = 2.3 difference – ionic bond C-O : C – 2.5 ; O – 3.5 = 1.0 difference – polar covalent bond O-O : O – 3.5 ; O – 3.5 = 0 difference – nonpolar bond Si-O : Si – 1.8 ; O – 3.5 = 1.7 difference – polar covalent bond N-O : N – 3.0 ; O – 3.5 = 0.5 difference – polar covalent bond N-O has the lowest difference in electronegativity between the two elements while still having enough of a difference to be polar covalent (0.5 difference)

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Most polar means highest positive and negative charges developing on the two atoms – electronegativity values will differ more between the two elements Remember: Electronegativity difference and bonding : 0 – 0.4 difference – nonpolar bond ; 0.5 – 2.0 difference – polar covalent bond ; > 2.0 difference – ionic bond Mg-O : Mg – 1.2 ; O – 3.5 = 2.3 difference – ionic bond C-O : C – 2.5 ; O – 3.5 = 1.0 difference – polar covalent bond O-O : O – 3.5 ; O – 3.5 = 0 difference – nonpolar bond Si-O : Si – 1.8 ; O – 3.5 = 1.7 difference – polar covalent bond N-O : N – 3.0 ; O – 3.5 = 0.5 difference – polar covalent bond Mg-O has the greatest electronegativity difference (2.3) but is ionic. Si-O has the greatest electronegativity difference while still being polar covalent (1.7).

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Once you determine the degree of polarity that exists within each of the separate bonds in a molecule, you need to compare all of them to see if the molecule itself will be polar or non-polar. This can be a very difficult concept so it can take time and practice to understand it really well. Polar molecules will result under several scenarios: 1)

The molecule contains polar bonds that cannot be canceled out, leaving some permanent charge within the molecule.

2)

Unbalanced lone pair electrons around the central atom create an unsymmetrical geometry and induce permanent charge within the molecule.

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Nonpolar molecules will result under several scenarios: 1)

The molecule contains nonpolar bonds and the molecule’s shape is symmetrical, leading to no permanent charge development.

2)

The molecule contains polar bonds of equal strength and the molecule’s shape is symmetrical, leading to the polar bonds cancelling each other out. No permanent charge develops.

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In both of these cases, the molecules have polar bonds that cannot be canceled out. Thus, the molecules are considered polar. In water, the two arrows are pointing toward one another, but both are pointing up as well. If the molecule was linear and the arrows pointed directly at each other, they could cancel out. However, since they are both angled and pointing up, there is a portion of the arrows that cannot be canceled out. If any portion of even one polar bond cannot be canceled out, the molecule is considered polar. In HCN (hydrogen cyanide), there is only one polar bond between the Carbon and Nitrogen atoms. Since there is nothing to cancel out this polar bond, the molecule is polar.

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You can have polar bonds in a molecule that is symmetrical. The polar bonds, if equal in strength and perfectly opposite in direction, can cancel out. This is what happens in Carbon dioxide, shown above.

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Binary Ionic Compounds – Named in Chapter 4 Binary Molecular Compounds – Compounds using covalent bonds only First element - If the first element in the formula is listed as singular, you do not use the ”Mono” prefix. Just list the element’s name as it appears on the periodic table. If the first element is plural, use the correct prefix from the table listed on the next slide. Second element – No matter whether it is singular or plural, always use the correct prefix from the table on the next slide.

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First element - If the first element in the formula is listed as singular, you do not use the ”Mono” prefix. Just list the element’s name as it appears on the periodic table. If the first element is plural, use the correct prefix from the table listed on the next slide. Second element – No matter whether it is singular or plural, always use the correct prefix from the table on the next slide.

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This was a prank put out by College students referring to water’s official chemical compound name, Dihydrogen monoxide. They listed all of these things as a reason to ban the use of water, making it sound incredibly dangerous.

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