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Flashcards in chem_bonding Deck (27):
1

Define valence electrons.

Valence electrons are those electrons furthest from the nucleus (i.e. in the outer shell) - which are typically involved in covalent bonding.

2

Describe a reliable method to determine the number of valence electrons for an element.

The number of valence electrons can be determined by refering to which main-group column the element is found in. For example nitrogen is found in main-group 5 and therefore nitrogen has 5 valence electrons. These are the electrons described by the 2s22p3 electron configuration.

3

List the steps in drawing a Lewis structure.

The steps are

  1. Count the total number of valence electrons (add 1 for a negative charge, subtract 1 for a postive charge).
  2. Draw a molecular skeleton generally with the least electronegative element in the center and joining each atom with a dashed line.
  3. Place enough electrons to give each element an octet (hydrogen only needs a duplet).
  4. If Step 3 uses too many electrons- take away the excess electrons and make the appropriate double bond or triple bonds. If Step 3 uses too few electrons, add the required number of electrons to the central atom.
  5. Assign formal charges to all atoms.
  6. Draw all necessary resonance forms.

4

Define formal charge.

Formal charge is an apparent charge which results if an atom "owns" more or less electrons than its normal count of valence electrons. All lone pairs plus 1/2 of the shared pairs are "owned". Formal charge is therefore calculated by subtracting the number of electrons "owned" from the normal count of valence electrons.

5

Define resonance (structure) form.

Resonance forms are two or more Lewis structures drawn (seperated by double headed arrows) to indicate the existance of delocalized electrons. Delocalized electrons are the result of pi electrons being shared between three or more atoms (instead of just two atoms).

6

Explain VSEPR theory.

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts that electron pairs about a central atom will arrange themselves so as to minimize repulsion. Electron pairs include both bonding and lone pairs.

7

Define steric number (SN).

Steric number (SN) is a parameter about a central atom which accounts for the number of regions of electron density. It is calculated by: number of sigma bonds + number of lone pairs.

8

Compare and contrast electronic versus molecular geometry.

Electronic geometry describes the arrangement of the electron pairs about the central atom. There are five basic geometries: linear, trigonal planar, tetrahedral, trigonal bipyramidal and octahedral. The electronic geometry is essentially determined by the central atom's steric number (SN).

Molecular geometry describes the arrangement of atoms around a central atom. Molecular geometry is essentially determined by the electronic geometry.

Molecular geometry is identical to electronic geometry when there are no lone pairs of electrons. For example the electronic and molecular geometry of CH4 is tetrahedral (SN = 4, AX4-type molecule). By contrast the electronic geometry of NH3 is tetrahedral (SN = 4), but the electronic geometry is trigonal pyramidal (AX3E-type molecule). In this case NH3 has one lone pair of electrons.

9

Give the name of the electronic and molecular geometries for an AX3E2-type molecule.

AX3E2-type molecule has an electronic geometry essentially the same as an AX5-type molecule - that is trigonal bipyramidal. Its molecular geometry is T-shaped.

10

Give the name of the electronic and molecular geometries for an AX4E2-type molecule.

AX4E2-type molecule has an electronic geometry essentially the same as an AX6-type molecule - that is octahedral. Its molecular geometry is square planar.

11

Give the name of the electronic and molecular geometries for an AX4E-type molecule.

AX4E-type molecule has an electronic geometry essentially the same as an AX5-type molecule - that is trigonal bipyramidal. Its molecular geometry is see saw.

12

Give the name of the electronic and molecular geometries for an AX5E-type molecule.

AX5E-type molecule has an electronic geometry essentially the same as an AX6-type molecule - that is Octahedral. Its molecular geometry is square pyramidal.

13

Give the name of the electronic and molecular geometries for an AX2E3-type molecule.

AX2E3-type molecule has an electronic geometry essentially the same as an AX5-type molecule - that is trigonal bipyramidal. Its molecular geometry is linear.

14

When does electronic geometry = molecular geometry?

Molecular geometry is identical to electronic geometry when there are no lone pairs of electrons. For example the electronic and molecular geometry of CH4 is tetrahedral (SN = 4, AX4-type molecule). By contrast the electronic geometry of NH3 is tetrahedral (SN = 4), but the electronic geometry is trigonal pyramidal (AX3E-type molecule). In this case NH3 has one lone pair of electrons.

15

Define electronegativity.

Electronegativity is the attraction the nucleus has for a shared pair of electrons.

16

Define bond polarity.

Bond polarity refers to the degree to which there is an unequal attraction for a shared pair of electrons in a covalent bond.

A equal attraction occurs when both atoms involved in a bond are identical or if both atoms of similar electronegativities. This would result in a non-polar covalent bond.

An unequal attraction occurs when the two atoms involved in a bond have different electronegativities. The greater the difference in electronegativity, the more polar is the bond.

17

Compare and contrast bond versus molecular polarity.

Bond polarity occurs when electrons are "shared" between atoms of different electronegativities.

Molecular polarity occurs as a result of bond polarity and when there is an unequal distribution of electron density throughout the molecule (or at least in a portion of the molecule). In very simple cases (e.g. H-F), bond polarity is the same as molecular polarity. In other cases, molecular polarity can only be determined by considering bond polarity and the molecular geometry.

It is possible for bonds to be polar, but because of the molecular geometry this effect is cancelled (e.g. CO2). In other cases the bonds are polar and their effect is reinforced and as a result the molecule is polar (e.g. SO2).

18

Describe atomic orbital hybridization.

Atomic orbital hybridization is the result of the original atomic orbitals "fusing" to produce unique hybrid orbitals. These hybrid orbitals have different energies and "shapes" as compared to the original atomic orbitals.

For example, carbon's valence electrons sit in the 2s and 2p atomic orbitals according to the electon configuration (2s22p2). When carbon forms four single bonds, these atomic orbitals hybridize into four unique sp3 orbitals. Collectively the four sp3 orbtitals produce a tetrahedral orientation. All four sp3 hybrid orbitals are degenerate and their energies are distinct from the energies of the 2s and 2p orbitals.

19

Compare and contrast sigma versus pi bonds.

A sigma bond is the result of a head to head overlap of orbitals. This results in what we normally label as a single bond.

A pi bond is the result of a side by side overlap of p orbitals. This side by side overlap occurs both above and below the plane of the sigma bond.

A sigma bond and a pi bond between two atoms would constitute a double bond. A sigma bond and two pi bonds would constitute a triple bond.

20

Describe the relationship of steric number (SN) and hybridization.

For the purpose of the Chem 112 course, the steric number (which is easily determined by considering the number of sigma bonds and lone pair of electrons about an atom) is a good predictor of the atom's hybridization. SN = 2 = sp hybridization; SN = 3 = sp2 hybridization; SN = 4 = sp3; SN = 5 = sp3d hybridization; SN = 6 = sp3d2 hybridization.

21

Describe the relationship of electronic geometry and hybridization.

Given that electronic geometry is essentially determined by an atom's steric number (SN) and SN is also a good predictor of hybridization, there is a natural link between electronic geometry and hybridization. For example, all central atoms with SN = 4, will be sp3 hybridized and will have a tetrahedral electronic geometry.

22

Those electrons furthest from the nucleus (i.e. in the outer shell) - which are typically involved in covalent bonding.

Valence electrons

23

An apparent charge which results if an atom "owns" more or less electrons than its normal count of valence electrons. All lone pairs plus 1/2 of the shared pairs are "owned". This parameter is therefore calculated by subtracting the number of electrons "owned" from the normal count of valence electrons.

Formal charge

24

These are two or more Lewis structures drawn (seperated by double headed arrows) to indicate the existance of delocalized electrons. Delocalized electrons are the result of pi electrons being shared between three or more atoms (instead of just two atoms).

Resonance (structure) form.

25

This parameter is the attraction the nucleus has for a shared pair of electrons.

Electronegativity

26

This parameter refers to the degree to which there is an unequal attraction for a shared pair of electrons in a covalent bond.

A equal attraction occurs when both atoms involved in a bond are identical or if both atoms of similar electronegativities. This would result in a non-polar covalent bond.

An unequal attraction occurs when the two atoms involved in a bond have different electronegativities. The greater the difference in electronegativity, the more polar is the bond.

Bond polarity

27

This is a parameter about a central atom which accounts for the number of regions of electron density. It is calculated by: number of sigma bonds + number of lone pairs.

Steric number (SN)