KINETIC THEORY OF MATTER - molecules in matter are always in motion - speed of molecules is proportional to the temperature

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1 1 KINETIC TERY F MATTER - molecules in matter are always in motion - speed of molecules is proportional to the temperature TE STATES F MATTER 1. Gas a) ideal gas - molecules move freely - molecules have no interactions with each other - no attraction or repulsion - like ping-pong balls b) real gas - molecules move freely - molecules have small attractions for each other - like sticky ping-pong balls - attractions decrease pressure of gas - molecules have a volume - inclusion of the molecular volume increases the pressure of gas (more crowding) 2. Liquid - molecules move but movement is heavily influenced by other molecules - molecules must move through a crowd of molecule - ideal gas can never condense to be a liquid. 3. Solid - molecules (or ions) vibrate about fixed position in crystal lattice - motion is not great enough to overcome attractions - molecules (or ions) stay in a specific arrangement INTERMLECULAR FRCES verriding Principles 1. pposite charges attract. 2. The greater the charge, the greater the attraction. - These principles are a restatement of Coulomb s Law.

2 2 CARGE DISTRIBUTINS F MATTER 1. Ion 2. Dipole - A polar molecule has a dipole. - The value of a dipole depends on a) polarity of bonds b) molecular geometry C N 3. Induced Dipole - A molecule can have its electron cloud polarized by neighboring charge. - Consider F 2. F F - Fluorine molecule has no dipole. - Consider F 2 next to sodium ion, Na +. F F Na + F F δ - δ + - Negative electron cloud adjusts to move closer to positive ion. - Fluorine molecule now has a dipole! - Sodium ion induces dipole in fluorine. - Thus fluorine may have an induced dipole. Some atoms in molecules have electron clouds that adjust easier than others. Na + F F Na + I I - Sodium ion induces larger dipole in iodine molecule than in fluorine.

3 3 The ability of an atom or molecule to adjust its electron cloud in response to an outside charge is called its polarizability. The less tightly bound to the nucleus the electrons are, the more polarizable they are. - As a periodic trend, polarizability increases as row increases. Each electron in an atom, molecule or ion contributes to the overall polarizability. - Thus, the more electrons a molecule has, the greater the strength of its induced dipole. Example: Arrange the following atoms in order of increasing polarizability: Ar, e, Kr, Ne and Xe. Answer: e, Ne, Ar, Kr, Xe Xe very polarizable e barely polarizable TYPES F INTERMLECULAR FRCES 1. Ion Dipole Example: salt water strongest type of intermolecular force - as the charge increases, the strength of the force increases Ca forces are stronger than Na forces

4 4 2. Dipole Dipole Example: Formaldehyde δ + δ - C C dipole dipole forces C Molecules with greater dipole moments have greater attractions to each other. Intermolecular forces (such dipole-dipole forces) can occur between different molecules. δ - δ + δ + C C N C 3. ydrogen bonding - A hydrogen bond is a very strong dipole dipole bond. - Molecules with N,, F covalently bonded to can form hydrogen bonds. - N,, F atoms have high electronegativity - Electrons in sigma bond are greatly polarized. δ - δδ - δδ + - ydrogen is bare naked. - A hydrogen bond is formed when polarized is attracted to nonbonding pair of electrons.

5 5 Example: 2 hydrogen bond - ydrogen bond is much stronger than dipole dipole bond. - ydrogen bonding gives water many important and distinctive properties. - ydrogen bonding holds together DNA helix and proteins. 4. Induction forces (Dipole Induced Dipole) Example: Cl in I 2. δ + δ - I I Cl I I I I I I - Dipole can induced by other dipoles. - Induction forces are weaker than dipole dipole forces. 5. Dispersion forces (Induced dipole induced dipole) - also known as London (dispersion) forces Charge distribution of molecules fluctuates over time. - i. e., charge distribution is unequal in nonpolar molecules at times - i. e., nonpolar molecules become polar for a short time, though average dipole is zero.

6 6 Spontaneous fluctuation of electron cloud. Consider a single bromine molecule. δ - δ + - dipole appears, then disappears, then appears reversed, then disappears, etc Fluctuating dipole induces dipole in another molecule so that they fluctuate together. Consider two bromine molecules δ + δ - δ + δ - δ + δ - Strength of dispersion forces depends on strength of induced dipole. - Molecules with greater polarizability have greater dispersion forces. Illustration: stronger bond Cl Cl Cl Cl weaker bond Bond between bromine molecules is stronger than bond between chlorine molecules. - degree of contact with other molecules is also important

7 7 Aside: abbreviated carbon chain structure C C C C Example: Pentane, C 5 12 has stronger dispersion forces than butane, C 4 10 butane pentane Bonding between butane molecules is weaker than bonding between pentane molecules. 2-methylbutane pentane Bonding between 2-methylbutane molecules is weaker than bonding between pentane molecules. Boiling Point Butane -0.5 ºC 2-methylbutane 27.8 ºC Pentane 36.1 ºC Increasing Intermolecular Forces

8 8 Summary of Intermolecular Forces Molecule A Molecule B ion ionic bonding ion - dipole ion dipole hydrogen bonding dipole - dipole dipole induction induced dipole dis persion induced dipole Principles of Intermolecular Force Strength 1. pposite charges attract. 2. The greater the charge, the greater attraction. 3. Intermolecular forces are additive. 4. Nonpolar molecules have only dispersion forces. 5. Polar molecules have dipole-dipole, induction forces and dispersion forces (and possibly hydrogen bonding). 6. Induced dipoles are often stronger than permanent dipoles. - Dispersion forces can be a significant reason why polar molecules are attracted to each other. Example: Fluoromethane, C 3 F (18 e - ) has stronger total forces between molecules than the force between fluorine gas molecules, F 2 (18 e - ). Example: Iodine, I 2 (nonpolar, 106 e-) has stronger forces in its molecules than does hydrogen sulfide (polar, 18 e-). Why? - Because the induced dipole of iodine is much greater than induced dipole and permanent dipole of hydrogen sulfide since iodine has so many more electrons.

9 9 INTERMLECULAR FRCES AND PASE TRANSITINS Solid Gas condensation evaporation Liquid Phase changes occur when intermolecular bonds are affected by an increase or decrease in kinetic energy. Temperature during a phase transition is always constant. - All energy input goes into phase transition, not temperature increase. eating Curve of Water Temperature (C) liquid water is heated Time (sec) water is boiling steam is getting hotter PRPERTIES F LIQUIDS Viscosity - resistance to flow - thickness of a fluid - corn syrup has a higher viscosity than water - viscosity is dependent on dynamic interaction between liquid molecules - intermolecular forces and geometry of molecule are important. Surface Tension - The ability of a liquid to form a film on its surface that can resist puncture. - Water and g are liquids with high surface tension, dish soap has low surface tension. - Bugs walk on water because water has high surface tension. - Meniscus forms in graduated cylinder because of surface tension.

10 10 PASE DIAGRAMS - indicates phase of substances at different temperatures and pressures - indicates under what conditions phase changes occur Example: C 2 critical point (31.1 C, 73 atm) s l g triple point (-56.4 C, 5.11 atm) Temperature - Lines indicate phase boundaries where both phases exist at the same time. - At the triple point, all three phases exist at the same time. - Each substance has only one triple point and it is unique to that substance. - The critical point is where distinction between gas and liquid is lost. - Substance is called a supercritical fluid. - Note how the phases diagram illustrates how the boiling, freezing and sublimation points change with pressure. - Slope of the phase boundary between liquid and solid is a consequence of the solid having a higher density than the liquid. - The boundary in most substances has a positive slope. Clausius Clapeyron Equation The line between the liquid and gas phase is described by the Clausius Clapeyron equation. P 2 vap 1 1 ln Δ P = 1 R T 1 T 2 - ln is the natural logarithm (more in 1190) - With the heat of vaporization, the boiling point of a substance can be predicted.

11 11 Example: 2 critical point (374 C, 224 atm) s l g triple point ( C, atm) Temperature - Note triple point is much different than C 2. - Note negative slope of solid-liquid boundary - Indicates solid water has lower density than liquid water. - Thus ice floats! Estimating the relative values of boiling points and freezing points Boiling point of substance is temperature where vapor pressure of liquid equals atmospheric pressure. The boiling point and freezing point of substances depends on 1. Intermolecular forces. - The stronger the forces between molecules, the harder it is to separate them - Therefore the stronger the forces, the higher the boiling points and melting points of substances. 2. Kinetic energy of molecules - The greater the mass of the molecules, the greater the boiling and freezing points. - The higher the temperature, the faster the molecules move and the more easily the molecules can break away from each other.

12 12 Increasing boiling points and melting points imply increasing heats of vaporization and fusion. Example: rder the following according to increasing boiling point: 2, Cl 2, F 2, I 2. Recall: Dispersion forces increase as the total polarizability increases. Polarizability increases as row of element increases and the number of electrons increases. Therefore answer is: F 2, Cl 2, 2, I 2 Example: rder the following according to decreasing boiling point. μ Acetaldehyde C 3 C 2.7 D Acetonitrile C 3 CN 3.9 D Methyl Ether C 3 C D Methyl Chloride C 3 Cl 1.9 D μ - strength of dipole D Debye unit The strength of the intermolecular force is proportional to the strength of the dipole. As the value of the dipole increases, the strength of the intermolecular force increases. Therefore answer is: C 3 CN, C 3 C, C 3 Cl, C 3 C 3