1. Chapter 4 CHEMICAL BONDING AND MOLECULAR GEOMETRY
CHEMISTRY: Atoms First
Chapter 4 CHEMICAL BONDING AND MOLECULAR GEOMETRY
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2. Bond – an attractive force that holds two atoms together.
Atoms bond to obtain a more stable electronic configuration.

3. Types of Chemical Bonds:
Ionic bond: Electrostatic attraction between cations and anions. Occurs between metals and non-metals.
Covalent bond: Sharing of outermost electrons 'ties' atoms together. Occurs between non-metals.
Metallic bond: Outer electrons are 'shared' by all atoms. Occurs between metals.

4. Figure 4.2
Sodium is a soft metal that must be stored in mineral oil to prevent reaction with air or water.
Chlorine is a pale yellow-green gas.
When combined, they form white crystals of sodium chloride (table salt). (credit a: modification of work by “Jurii”/Wikimedia Commons)

5. Naming Ionic Compounds
Binary Compounds – composed of only two elements, though may have more than two atoms: NaCl, KF, CaCl2, Fe2O3 …
Binary ionic compounds consist of cations (usually metals) and anions (usually nonmetals), e.g., MgCl2
Charges must cancel out to give the compound a neutral charge overall

6. Figure 4.3
The atoms in sodium chloride (common table salt) are arranged to (a) maximize opposite charges interacting. The smaller spheres represent sodium ions, the larger ones represent chloride ions. In the expanded view (b), the geometry can be seen more clearly. Note that each ion is “bonded” to all of the surrounding ions—six in this case.

7. Li+1 → lithium ion, Ca+2 → calcium ion
●Cation (metal) – name is the same as the element, + 'ion'
●Fixed charge cations – metals that only form one cation (such as Group 1 and 2 metals):
Li+1 → lithium ion, Ca+2 → calcium ion
●Variable charged cations – metals that may form different cations (most transition metals). Use Roman numerals to show the charge:
Fe+2 → iron (II) ion
Fe+3 → iron (III) ion

8. ●Anion (non-metal) – use the root of the element name, change the ending to 'ide', + 'ion':
S → S-2
sulfur → sulfide ion
N → N-3
nitrogen → nitride ion
O → O-2
oxygen → oxide ion

9. ●List the cation first, then the anion
●Do not include 'ion' in the name
●Names must be distinctive, in order to distinguish between similar compounds, such as with variable-charged metals
NaCl – sodium chloride
CaF2 – calcium fluoride
FeI2 – iron (II) iodide
FeI3 – iron (III) iodide

10. For ionic compounds with variable charged cations:
●Basically, all metals are variable charged, except for: Group 1, Group 2, Ag+1, Zn+2, Cd+2, Al+3, Ga+3
●For all other metals, the Stock System (Roman Numerals) must be used:
Cu2O – copper (I) oxide
CuO – copper (II) oxide

12. Group
Ion usually formed
1 (1A) +1
H can form -1 ion
2 (2A) +2
13 (3A) +3
14 (4A)
+/-4
15 (5A) -3
16 (6A) -2
17 (7A) -1
18 (8A)

13. To determine the charge on a variable charge cation, treat the formula as an algebraic expression:
To determine the iron charge in Fe2O3
●let Fe = x and O = y (x and y are ionic charges)
●the charges of the ions must add up to the overall charge, which is 0 in this case, so
2x + 3y = 0
●we know that y = -2 (oxide ion)
2x + 3 (-2) = 0
x = +3
●so Fe2O3 is named iron (III) oxide

14. ●The formula shows a ratio of one ion to the other.
●The ionic charges must cancel out so that the overall charge is neutral
●Always list the metal first, then the non-metal
●Select subscripts to balance charges
●Reduce subscripts if needed to obtain the lowest whole number ratio between ions

15. Naming Molecular Compounds
Binary Molecular Compounds (e.g., SO3)
Compounds consisting of two nonmetals
First element in the formula is named first.
S = sulfur
Second element name is changed by adding suffix -ide.
O = oxygen → oxide
Add prefixes to identify quantity of atoms (see Table 4.4).
SO3 = sulfur trioxide

16. Figure 4.4
The potential energy of two separate hydrogen atoms (right) decreases as they approach each other, and the single electrons on each atom are shared to form a covalent bond. The bond length is the internuclear distance at which the lowest potential energy is achieved.

17. ●The formula name must indicate the subscripts
●use prefixes to show subscripts
●The prefix 'mono' is not used on the first element listed
●Remember that molecules have fixed numbers of atoms linked together, so DO NOT reduce subscripts to lower ratios

18. Naming Binary Compounds
1. Prefixes are used to designate the number of each type of element:
number of atoms prefix
1 mono
2 di
3 tri
4 tetra
5 penta
6 hexa
7 hepta
8 octa
9 nona
10 deca

19. For elements in standard state, use chem. symbols
Except for diatomic elements:
H2, N2, O2, F2, Cl2, Br2, and I2 exist as diatomic molecules in their standard states
These seven elements are H2 + “7”

20. Polyatomic Ions
These are covalently bonded atoms with an overall charge (an ionic 'molecule'):
NO3-1 – nitrate ion
ClO3-1 – chlorate ion
C2H3O2-1 – acetate ion
OH-1 – hydroxide ion
SO4-2 – sulfate ion
CO3-2 – carbonate ion
PO4-3 – phosphate ion
H3O+1 – hydronium ion
NH4+1 – ammonium ion (NH3 – ammonia)

21. Polyatomic ions containing oxygen and another non-metal
Oxyions (oxoanions)
Polyatomic ions containing oxygen and another non-metal
Most common forms end in 'ate'
One less oxygen ends in 'ite'
Two less oxygens, 'hypo' prefix and 'ite' suffix
One more oxygen, 'per' prefix and 'ate' suffix
ClO-1 – hypochlorite ion
ClO2-1 – chlorite ion
ClO3-1 – chlorate ion
ClO4-1 – perchlorate ion

22. Electronegativity (EN) – attraction for bonding (shared) electrons
The difference in EN between the bonding atoms determines if the bond is ionic or covalent
Not all covalent bonds are identical

23. Figure 4.6
The electronegativity values derived by Pauling follow predictable periodic trends with the higher electronegativities toward the upper right of the periodic table.

24. With a large EN difference (such as between metals and non-metals), the more electronegative atom takes the electrons from the other atom and forms ions.
If the ENs are identical, both have the same attraction, and the bonding electrons are shared equally. (non-polar covalent bonds)
For small EN differences, the electrons are shared, but not equally (polar covalent bonds)

25. Polar (covalent) bonds have a slight positive charge on one end, and a slight negative charge on the other end.
The more EN atom pulls the bonding electrons closer to itself, which creates the slight negative charge.
The less EN atom then has a slight positive charge.
This is called a dipole (two poles).

26. Figure 4.5
The distribution of electron density in the HCl molecule is uneven. The electron density is greater around the chlorine nucleus. The small, black dots indicate the location of the hydrogen and chlorine nuclei in the molecule.
Symbols δ+ and δ– indicate the polarity of the H–Cl bond.

27. Figure 4.8
As the electronegativity difference increases between two atoms, the bond becomes more ionic.

28. Figure 4.6
The electronegativity values derived by Pauling follow predictable periodic trends with the higher electronegativities toward the upper right of the periodic table.

29. Figure 4.26
There is a small difference in electronegativity between C and H, represented as a short vector.
The electronegativity difference between B and F is much larger, so the vector representing the bond moment is much longer.

30. Lewis Structures (Dot Structures)
Used to determine structures and shapes of molecules
Dots are used to represent valence (outer) electrons
Electrons are placed above, below, left and right of each chemical symbol
First four electrons placed unpaired
Then paired
Electrons placed between two atoms are shared by both atoms

31. Only outer electrons participate in bonding
Inner electrons are in stable Noble Gas configurations
Valence Shell – outermost occupied ground state shell
Valence electrons (v.e.) – electrons in valence shell

32. Valence electrons (ve)
The outer electrons are called “valence electrons” or “bonding electrons”
Most atoms want to get 8 v.e. (octet rule), which they do by gaining, losing, or sharing electrons.
Main exception: Hydrogen wants 2 v.e.
Noble gases already have 8 v.e. (except helium which has 2), which is why they almost never react.

33. Figure 4.10
Lewis symbols illustrating the number of valence electrons for each element in the third period of the periodic table.

34. Figure 4.11
Cations are formed when atoms lose electrons, represented by fewer Lewis dots, whereas anions are formed by atoms gaining electrons. The total number of electrons does not change.

36. To draw Lewis Structures for molecules:
Add up valence electrons for all atoms
Take into account the charge on an ion
Draw backbone ('skeletal structure') for the molecule
Link atoms with single bonds (two shared electrons)
Usually elements down and to left on PT are central (carbon is usually central)
H and Group 17 elements almost always only form single bonds

37. Add electrons to satisfy the octet rule for all atoms:
Add electrons to outer atoms first
Electrons usually placed in pairs above, below, right, or left of symbol
Electrons between two atoms are bonding electrons, and count toward valence electrons for both atoms
2 shared e- - one line – single bond
4 shared e- - two lines – double bond
6 shared e- - three lines – triple bond

38. Use/move unshared pairs of electrons to form double and triple bonds, if needed, to satisfy octets
Double check:
Final structure must have the exact number of total valence electrons
Check that all atoms obey the octet rule, not counting the few exceptions

42. Following the above procedure may result in several different Lewis Structures being drawn
Formal Charges will often show which is the more stable, preferred, structure
Formal charges for each atom should add up to the overall charge for the atom/molecule

43. Formal charge (FC) shows the effective charge on each atom
FC = valence electrons – unshared electrons – 1/2 bonding electrons
FC = v.e. – u.e. – 1/2 b.e.
Structures with FCs closer to 0 are preferred
Negative FCs should be on more electronegative atoms

44. The structure on the left is preferred due to the formal charges.

45. Exception to the Octet rule
Although nitrogen does not follow the octet rule, the formal charges of both atoms are zero.

46. Exception to the Octet rule
In both of these cases, the formal charges of all atoms are zero.

47. FIGURE 4.12
In PCl5, the central atom phosphorus shares five pairs of electrons. In SF6, sulfur shares six pairs of electrons.

48. Not all molecules can be accurately represented by a single Lewis Structures
Sometimes two or more structures may be drawn with identical formal charges
In this case the actual structure of the molecule is an average of all of the similar structures
These are called Resonance Structures

49. Resonance structures
All structures are equal according to formal charges. The actual structures is an ‘average’ of these three Lewis Structures

50. Molecular Geometry
The geometric shape of a molecule may be determined from the Lewis Structure
Groups of valence electrons on an atom repel each other, and move to maximize their angles of separation
Groups of electrons are:
Unshared pairs of electrons
single bonds
double bonds
triple bonds
single unshared electron

51. Valence Shell Electron Pair Repulsion
also called:
VSEPR Theory
Valence Shell Electron Pair Repulsion

52. With two groups – 180° between groups
Linear
With three groups – 120° between groups
Trigonal
With four groups – 109.5° between groups
Tetrahedral
The actual shape of the molecule is dependent on the actual bonds.

53. Electronic Geometry
Use all groups of electrons to determine the angles between groups
Molecular Geometry
Look only at bonds to determine the shape of the molecule
Angles between bonds determined by electronic geometry
For central atoms with no unshared electrons, el. geom. = mol. geom.

54. Figure 4.15
The BeF2 molecule adopts a linear structure in which the two bonds are as far apart as possible, on opposite sides of the Be atom.

55. Figure 4.16
The basic electron-pair geometries predicted by VSEPR theory maximize the space around any region of electron density (bonds or lone pairs).

56. Example
Lewis structures often do not represent the true bond angles. This molecule is trigonal, and all bond angles are 120°.

57. Figure 4.21

58. Figure 4.17
The molecular structure of the methane molecule, CH4, is shown with a tetrahedral arrangement of the hydrogen atoms. VSEPR structures like this one are often drawn using the wedge and dash notation, in which solid lines represent bonds in the plane of the page, solid wedges represent bonds coming up out of the plane, and dashed lines represent bonds going down into the plane.

59. Figure 4.18
The electron-pair geometry for the ammonia molecule is tetrahedral with one lone pair and three single bonds.
The trigonal pyramidal molecular structure is determined from the electron-pair geometry.
The actual bond angles deviate slightly from the idealized angles because the lone pair takes up a larger region of space than do the single bonds, causing the HNH angle to be slightly smaller than 109.5°.

60. Figure 4.19
The molecular structures are identical to the electron-pair geometries when there are no lone pairs present (first column). For a particular number of electron pairs (row), the molecular structures for one or more lone pairs are determined based on modifications of the corresponding electron-pair geometry.

61. Example
For small molecules, the molecular geometry is referring to the central atom. For larger molecules like the one above, the geometry of each central atom needs to be specified.

62. Polar Molecules
Molecules that have a dipole moment – a difference in polarity from one end to the other
Molecules must have polar bonds or asymmetric unpaired electrons.
Molecule must also be asymmetric (not symmetric) to be a polar molecule
Polar molecules have an attraction to other polar molecules, like a weak ionic bond.

63. Figure 4.26
There is a small difference in electronegativity between C and H, represented as a short vector.
The electronegativity difference between B and F is much larger, so the vector representing the bond moment is much longer.

64. Figure 4.6
The electronegativity values derived by Pauling follow predictable periodic trends with the higher electronegativities toward the upper right of the periodic table.

65. Figure 4.27
The overall dipole moment of a molecule depends on the individual bond dipole moments and how they are arranged.
Each CO bond has a bond dipole moment, but they point in opposite directions so that the net CO2 molecule is nonpolar.
In contrast, water is polar because the OH bond moments do not cancel out.

67. Slightly polar bonds, but non-polar molecule since the molecular geometry is exactly symmetric

69. Figure 4.28
Molecules are always randomly distributed in the liquid state in the absence of an electric field.
When an electric field is applied, polar molecules like HF will align to the dipoles with the field direction.

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