Chapter 6 - Chemical Bonds Study Guide PDF

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A very concrete and detailed study guide of chapter 6 test that covers everything from the book and class' lecture....

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CHEM 1140: Chapter 6 Study Guide

Chemical Bonds Chemical bonds involve electrons. A bond forms between two atoms if it produces a more stable electron configuration for both atoms. Noble gases tend not to form bonds because they have their outer s + p subshells filled. This is a total of 8 electrons and is referred to as an octet of electrons. The octet rule states that atoms of the representative elements form bonds so as to have access to eight outer electrons. H, He, Li, and Be tend to form bonds to have access to 2 electrons (duet rule). Valence electrons: generally the s and p electrons in the outermost shell How can an atom alter its electron configuration to obtain an octet? 1. Metals may lose 1-4 electrons to form a cation with the electron configuration of the previous noble gas 2. Nonmetals may gain 1-4 electrons to form an anion with the electron configuration of the next noble gas 3. Atoms (usually 2 nonmetals) may share electrons with other atoms to obtain access to the number of electrons of the next noble gas Processes 1 and 2 are involved in the formation of ionic compounds. Process 3 produces molecular compounds.

Ions Representative metals (Group 1, 2, and Al) are willing to lose all of their valence electrons to acquire the octet of the previous noble gas. The loss of any additional electrons is expensive in terms of energy. Representative nonmetals (Nonmetals in Groups 13-17) are willing to gain electrons to fill up to the octet of the next noble gas. The gain of any additional electrons is expensive in terms of energy.

Formulas of Binary Ionic Compounds When forming ionic compounds, we now want to think in terms of electron movement. The metals give their electron to the nonmetal in order for both atoms to reach a more stable configuration and favorable configuration. Think about how this applies to how/why we use the cross/charge method to balance ionic compounds. It deals with the movement of electrons to make both the metal cation and the nonmetal anion “happiest” in their most stable configurations to reach an octet of outer electrons. Ionic compounds are generally solids at room temperature. They have high melting points and are often brittle. The 3D array of ions is called a lattice. The lattice is held together by strong and rigid electrostatic interactions known as ionic bonds, which lends to the solid nature of these compounds 1

(attractions between different units held together in the lattice). Polyatomic ions exist as ions because they have an imbalance of electrons compared to the total number of protons in their nuclei. Lattice energy tells us about the strength of the ionic bonds in a lattice crystal. This factor depends on charges, 3D lattice structure, and the distance between ions in the structure. Down a group, lattice energy decreases because the ions get bigger and the distance between ions gets bigger. Charge is important too because lattice energy is affected by charge, > the charge, > the lattice energy.

Covalent Bonds Covalent bonds involve the formation of bonds between nonmetals that share electrons. Molecular compounds resulting from covalent bonding are generally discrete molecular units (unlike the ionic lattice formations) and are only weakly attracted to one another. They can move past each other freely and therefore are generally liquids, gases, or soft solids with low melting points. Lewis Dot Structures: a way to represent the valence electrons available for an element as a series of dots around the written elemental symbol. For the representative elements, the number of dots (valence electrons) equals the group # (for Groups 1 and 2) or Group # - 10 (Groups 13-18) to form a neutral atom. For conventional sake, we fill one electron around a side first, and then pair them up. Ex/ B is in group 13 and would have 3 dots around the symbol B in a Lewis dot structure. In Sapling, cations that form to achieve noble gas configuration are shown to have “0” valence electrons. I’m not a huge fan of this, as in reality, you could say these cations have 8 valence electrons, similar to the preceding noble gas. I just wanted to give you a heads-up before you work on the Sapling problems. Lewis Structures: a representation of the order and arrangement of atoms in a molecule as well as all of the valence electrons for the atoms involved. Generally each unbound electron is represented by a dot and each pair of bonded electrons is represented by a dash (-). Either understand your book’s rules or print out and understand the rules I supplied with this study guide for writing Lewis Structures. Be comfortable with determining the number of valence electrons each atom has based on its electron configuration, so that you can determine the best structures. The formulas for many simple binary molecular compounds can be justified by the octet rule. The two types of atoms involved in the compound formation are trying to fill their valence electrons to a full 8 electrons. (For hydrogen (and He, Li, Be), these elements follow the duet rule.) Double bonds and triple bonds are instances where two or three pairs of electrons are shared between two atoms, respectively. Double bonds are represented by 2 dashes (=) and triple bonds are represented by 3 dashes (≡). The 6N + 2 rule is a handy method of determining if double or triple bonds are present in a compound. N = # of non-H atoms in the molecule. This rule only works when compounds form octet-based (or duet-based) Lewis structures. For example, sulfur is an atom that breaks the octet rule on a regular basis. It loves to have more than 8 electrons. Boron prefers to only have 6. For polyatomic ions, place the atoms in brackets with the charge of the ion in superscript to the right of the brackets. When determining the total number of electrons to assign within the polyatomic ion,

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figure out the number of valence electrons for each atom and add 1 electron for each negative charge or take away an electron for each positive charge.

Resonance Structures Resonance structures exist for molecules when equally correct Lewis structures can be written without changing the basic skeletal geometry or the position of any atoms. Each representation is connected by a double-headed arrow to indicate that all bonds are intermediate between the structures (a hybrid). Try to imagine resonance structures as a combination of two or more structures (not as oscillating back and forth).

Bond Length/Strength Bond length is the distance from one nucleus to the other in a covalent bond. Multiple bonds tend to be shorter than single bonds. Smaller atoms tend to have shorter bond lengths. Bond strength is measured experimentally as the amount of energy required to break the bond. Multiple bonds tend to be stronger than single bonds.

Geometry Electronegativity: the ability of an atom of an element to attract electrons to itself in a covalent bond. The most electronegative elements are: F > O > N ≥ Cl > S ≥ Br. Electronegativity increases from left to right and increases from bottom to top of a column. An electronegativity chart is shown in your text Fig. 6-15. Polar covalent bonds: occur when atoms of differing electronegativity bond. The shared electrons spend more time near the more electronegative atom vs. the other. The result is partial charges (δ- for the more electronegative element and δ+ for the less electronegative element) and the formation of a dipole bond (bond with 2 poles of differing charges). Exception: Even though C and H have differing electronegativity values (2.5 and 2.1 respectively), H- C single bonds are considered to be nonpolar (an important feature of organic chemistry). Nonpolar bonds: when electrons are shared equally between atoms. Occurs between the same types of atoms (Ex/ O2, Cl2). Occurs between two atoms with similar/same electronegativity (Ex/ C bonded to S). BOND SUMMARY: Ionic bonds: e- donated or accepted Covalent bonds: e- are shared Polar covalent bonds: e- are shared unequally Nonpolar covalent bonds: e- are shared eqully Polarity: Polar molecules have a net dipole, meaning the partial charges do not cancel out. To be a polar molecule, the molecule must contain polar bonds and the bond dipoles do not cancel out. So the molecule must by asymmetric. One side is more + or – charged than the other.

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Geometry: Electron groups repel each other to the maximum extent. (all – charges) The shape/geometry of a molecule can be predicted (VSEPR) and Lewis structures. Molecular geometry: shape described by bonded atoms and “ignores” unbound (lone pairs)

Study the VSEPR handout attached to this study guide and be able to tell me the molecular and/or electronic geometries for simple compounds. For those of you still wanting even more information/help with geometry, this is a good website with lots of pictures, tutorials, etc. about geometry: http://intro.chem.okstate.edu/1314F00/Lecture/Chapter10/VSEPR.html.

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Guidelines for Drawing Lewis Structures 1. Use the guidelines from your book if they are clearer for you. 2. An alternative approach to the same results is the following: a. Sum all valence electrons of all atoms. Be sure to account for any electrons gained or lost if you are drawing the Lewis Structure of an ion. For anions add the same number of electrons as specified by the charge (if an ion is “-2” add 2 electrons). For cations subtract the same number of electrons as specified by the charge (if an ion is “+1” subtract 1 electron). b. Draw a skeleton structure 1) LEAST electronegative element is usually in the middle (but never hydrogen) 2) Choose the most symmetrical structure c. Distribute the electrons 1) Put a pair of electrons between the atoms 2) Give each atom an octet of electrons (Hydrogen only gets a DUET) 3) Use multiple bonding if necessary d. Double check your structure (including a check of formal charge) Other guides: Electronegativity increases left to right across the periodic table and increases from bottom to top. The most electronegative elements: F > O > Cl ≥ N > S ≥ Br (For the rest of the elements, use the periodic trend) Carbon and hydrogen have essentially the same electronegativity Carbon always forms 4 bond, including double and triple bonds Nitrogen forms 3 bonds where possible including double and triple bonds Oxygen forms 2 bonds where possible including double and triple bonds Hydrogen forms 1 bond and never forms multiple bonds (The exceptions to the bonding of N and O are few, but we will see a couple) For the bonding of C, N, O, H, think of these Lewis representations: . . .. . . . . . . C N O .

4 dots, shares e- in 4 bonds

..

3 single dots, so shares e- in 3 bonds

H.

..

2 single dots, so shares e- in 2 bonds

1 dot, so can only make 1 bond

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Formal Charge The noble gas rule (or octet rule) is not always sufficient to determine the correct configuration of atoms in a molecule. In order to determine the favored configuration for a molecule, we employ formal charge values as a guide. The following formula is used to determine the formal charge (F.C.) of an atom in a molecule: F.C. = Ev – [Eu + ½ (Eb)] where Ev is the number of valence electrons in the neutral atom, Eu is the number of unshared valence electrons that the bonded atom has and Eb is the number of bonded electrons it shares. When deciding among possible Lewis structures for a molecule, one should choose the structure that yields: 1) F.C. values of zero or as near zero as possible, and 2) negative values associated with atoms that are the most electronegative, if they are not zero. Ex/ CH4O

a.

H I H-C-O-H I H

a. F.C. (C) = 4 – [0 + ½ (8)] = 0 F.C. (H) = 1 – [0 + ½ (2)] = 0 F.C. (O) = 6 – [4 + ½ (4)] = 0

b.

H-C-O-H I I H H

b. F.C. (C) = 4 – [2 + ½ (6)] = -1 F.C. (H) = 1 – [0 + ½ (2)] = 0 F.C. (O) = 6 – [2 + ½ (6)] = +1

Both structures have octets for the second period elements. From the above calculations we can see that structure (a) is the preferred structure since all F.C. values are zero while structure (b) has a positive value for the most electronegative atom. From this example you can now write the preferred Lewis structures for the following molecules and ions: a. H2O

b. NH4+

c. CCl2F2

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VSEPR SHAPES (Valence Shell Electron Pair Repulsion Model for molecular shapes)

Electronic Geometry (also known as Electron Pair Geometry or Electron Geometry) # of groups of electrons around the central atom (*bonding and **non-bonding): 2 3 4 Linear

Trigonal Planar

Tetrahedral

X

X

A

A

X–A–X Ex/ CO2

X

X

X

X X

Ex/ BF3

Ex/ CH4

A is the central atom, X represents other atoms (or pairs of non-bonding electrons)

Molecular Geometry If all groups of electrons around the central atom are involved in bonds (no lone pairs), the molecular geometry is the same as the electron geometry. If there are non-bonding pairs of electrons around the central atom, see below.

# of groups of electrons around the central atom (bonding and non-bonding):

3 One n

..

4 nding pair

One no

Two non

A

A X

..

4 ding pair

X

X

..

ing pairs

A X

X

X

..

X

Electron Geometry Trigonal Planar

Electron Geometry Tetrahedral

Electron Geometry Tetrahedral

Molecular Geometry Bent at 120o Ex/ SO2

Molecular Geometry Trigonal Pyramidal Ex/ NH3

Molecular Geometry Bent at 109o Ex/ H2O

* All bonds (single, double, or triple bonds) count as 1 group around the central atom ** Lone pairs (2 electrons) count as 1 group around the central atom 7

Summary Electronic Geometry: 2 groups around central atom = linear 3 groups around central atom = trigonal planar 4 groups around central atom = tetrahedral

Molecular Geometry: 2 groups around central atom (0 lone pairs) = linear 3 groups around central atom (0 lone pairs) = trigonal planar 3 groups around central atom (1 lone pair) = bent 4 groups around central atom (0 lone pairs) = tetrahedral 4 groups around central atom (1 lone pair) = trigonal pyramidal 4 groups around central atom (2 lone pairs) = bent

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