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Chem181-Chemistry for Engineers 1A 1

CHEMICAL BONDS

INTERMOLECULAR FORCES are attractive forces between molecules.

Generally, intermolecular forces are weaker than intramolecular forces. While intramolecular forces stabilize individual molecules, intermolecular forces are responsible for the bulk properties of matter e.g. melting and boiling points. Examples: 41 kJ of energy is required to vapourise 1 mole of H2O at its boiling point. While 930 kJ of energy is required to break 2 x O-H bonds in H2

• Dipole-dipole interactions

O These intermolecular forces as a group are referred to as van der Waals forces. Examples of van der Waals forces are:

• Hydrogen bonding • London dispersion forces

Ion-Dipole Interactions

• A fourth type of force, ion-dipole interactions are an important force in solutions of ions.

• The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.

The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together. They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.


Dipole-Dipole Interactions Molecules that have permanent dipoles are attracted to each other.

The positive end of one is attracted to the negative end of the other and vice-versa.

These forces are only important when the molecules are close to each other.

The more polar the molecule, the higher is its boiling point. London Dispersion Forces While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom.


At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side

Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1.

London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.

• These forces are present in all molecules, whether they are polar or nonpolar. • The tendency of an electron cloud to distort in this way is called polarizability.

Factors Affecting London Forces

• The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane).

• This is due to the increased surface area in n-pentane.


• The strength of dispersion forces tends to increase with increased molecular weight.

• Larger atoms have larger electron clouds, which are easier to polarize. WHICH HAVE A GREATER EFFECT:

dipole-dipole interactions or dispersion forces

• if two molecules are of comparable size and shape, dipole-dipole interactions will likely be the dominating force.

• if one molecule is much larger than another, dispersion forces will likely determine its physical properties.

HYDROGEN BONDING

• The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong.

• We call these interactions hydrogen bonds. • Hydrogen bonding arises in part from the high electronegativity of nitrogen,

oxygen, and fluorine. • Also, when hydrogen is bonded to one of those very electronegative elements, the

hydrogen nucleus is exposed. •


VISCOSITY

• Resistance of a liquid to flow is called viscosity. • It is related to the ease with which molecules can move past each other. • Viscosity increases with stronger intermolecular forces and decreases with higher

temperature.


SURFACE TENSION Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

INTRAMOLECULAR FORCES hold atoms together in a molecule. While intramolecular forces stabilize individual molecules, intermolecular forces are responsible for the bulk properties of matter e.g. melting and boiling points. Examples: 41 kJ of energy is required to vapourise 1 mole of H2O at its boiling point. While 930 kJ of energy is required to break 2 x O-H bonds in H2O Intramolecular forces include: (A) Ionic bonding (B) Covalent Bonding (A) IONIC BONDING: is the electrostatic interaction of oppositely charged ions in an ionic compound e.g. NaCl (m.pt. 104oC), LiF, CaO, MgO (1100oC), KCl, KBr. Takes place between cations and anions: e.g. Na+ + Cl- → NaCl Li+ + Cl- → LiCl Ca2+ + O2- → CaO Elements most likely to form cations in ionic compounds are alkali metals (group I) and alkaline earth metals (group II). Elements most likely to form anions in ionic compounds are halogens (Group VII) and oxygen. Elements with low ionization energy (IE) – from cations Elements with high electron affinity (EA) – from anions


Recall: Electronegativity: is the ability of an atom to attract electrons toward itself in a chemical bond. e.g. in H-F, EN of F = 4.0 Generally the halogens have high EN values (F = 4.0, Cl = 3.0, Br = 2.8) On the periodic chart, electronegativity increases as you go…

…from left to right across a row. …from the bottom to the top of a column.

As a general rule an ionic bond forms when the EN difference between the 2 bonding atoms is 2 or more. Electron Affinity: is the energy change that occurs when an electron is accepted by an atom in the gaseous state.

e.g. X (g) + e- → X- (g)

Cl2(g) + 2e- → 2Cl-

Elements most likely to form cations in ionic compounds are alkali metals (group I) and alkaline earth metals (group II). Elements most likely to form anions in ionic compounds are halogens (Group VII) and oxygen. Elements with low ionization energy (IE) – from cations Elements with high electron affinity (EA) – from anions


(B) Covalent Bonding Covalent bond: a bond in which two electrons are shared by two atoms Covalent compounds are compounds that contain only covalent bonds e.g. HCl, NH3, CH4, CCl4, H2

Attractions between electrons and nuclei

O.

There are several electrostatic interactions in these bonds:

Repulsions between electrons Repulsions between nuclei

Polar Covalent Bonds

• Fluorine pulls harder on the electrons it shares with hydrogen than hydrogen does.

• Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.


Although atoms often form compounds by sharing electrons, the electrons are not always shared equally

• When two atoms share electrons unequally, a bond dipole results.

• • The dipole moment, µ, produced by two equal but opposite charges separated by a

distance, r, is calculated: µ = Qr

• It is measured in debyes (D). Covalent bonding between many-electron atoms involves only the valence electrons. Let us consider the F2 molecule: It involves the covalent bonding of 2 F atoms. The electronic configuration of F is 1s22s22p5 The 1s electrons are low in energy and stay near the nucleus, hence do not participate in bonding. Each F has 7 valence electrons: 2s22p5

+ F F F F

= 2 + 5= 7

A LEWIS structure is a representation of covalent bonding where shared electrons are shown as lines or dots as shown in the F2 molecule above.

Example: CO2 O C O O C O

Other non-bonding electrons are called lone pairs

H O H

Lone pair of electrons


The Octet Rule: An atom other that hydrogen tends to form bonds until it is surrounded by 8 valence electrons. The Octet rule works mainly for elements in period 2 . Why? The following molecules illustrates the Octet Rule The F2

F F molecule

The N2

N N molecule

Each atom is surrounded by 8 electrons in each case: Octet Rule Exercise 1.: Draw the Lewis Dot Structure for the CCl4 molecule Solution: Step 1. Count up all the valence electrons first

C = group 4 = 4 e- = 4 e- Cl = group 7 = 4 x 7e- = 28 e- Total electrons 32 e

C

Cl

Cl

Cl

ClCl

- Step 2. Insert the bonding electrons

A total of 10 electrons are used in the 5 bonds Step 3. Distribute the remaining electrons around substituent atoms i.e. complete the Octet around substituent groups

CCl

Cl

Cl

ClCl


Step 4. Place any remaing electrons around the central atom Exercise 2. Draw the Lewis Dot Structure for the H2SO4

C C

molecule (S is the central atom) Atoms can form different types of covalent bonds (a) A Single bond: 2 atoms held together by one electron pair e.g C:C or C-C. (b) A Double bond: two atoms are held together or share 2 pairs of electrons e.g. C=C. (c) Triple bond: arises when two atoms share 3 pairs of electrons e.g.

Double and triple bonds are Multiple bonds which are shorter than single bonds. Bond length: is the distance between the nuclei of two covalently bonded atoms in a molecule. e.g. C-H 107 pm (picometers) C=O 121 pm C-O 143 pm C-C 154 pm C≡C 120 pm

Comparison between Ionic and Covalent Compounds

They differ markedly in physical properties because of differences in the nature of their bonds. There are two types of attractive forces in covalent compounds i.e. intermolecular and intramolecular forces. Covalent compounds are usually gases, liquids or low melting solids at room temp. Ionic compounds are solids at room with high melting points. Most ionic compounds are soluble in water, conduct electricity hence are strong electrolytes. Most covalent compounds are insoluble in water or if they do dissolve, they do not conduct electricity. Liquid or molten covalent compounds do not conduct electricity because no ions are present. Molten ionic compounds conduct electricity because they contain mobile cations and anions.


Example. Bond type Ionic Covalent Compound NaCl CCl4 Appearance White solid Colorless liquid Melting point (M.pt.) 801o -23C oC Boiling point (B.pt.) 1413o 76C oC Solubility in H2 High O Very low Electrical conductivity of: (a) Solid Poor Poor (b) Liquid Good Poor

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