Chapter 11 – Intermolecular Forces, Liquids and Solids Homework: 13, 16, 18, 19, 23, 43, 45, 47, 48, 49, 50, 51, 54, 55, 56

11.2 – Intermolecular Forces Strengths of intermolecular forces of different substances varies Generally weaker than ionic or covalent bonds Less energy needed to vaporize (or evaporate) a liquid or melt a solid than to break the covalent bonds in molecules In other words: Compound stay intact when melting or boiling, just breaking intermolecular forces

Many properties of liquids reflect strength of the intermolecular forces Such as boiling point Example: Because the forces between HCl molecules are so weak, HCL boils at very low temperature -85ºC at atmospheric pressure

Boiling/Melting A liquid boils when bubbles of its vapor form within the liquid Molecules in the liquid must overcome their attractive forces to do vaporize The stronger the forces, the higher the temperature at which the liquid boils Same general principle applies to melting

Types of Intermolecular Forces Three types exist between neutral molecules dipole-dipole forces London dispersion forces hydrogen-bonding forces a.k.a. van der Waals forces There is one other force that mostly applies to solutions ion-dipole force All of these forces tend to be less than 15% as strong as covalent or ionic bonds

Dipole-Dipole Forces Neutral POLAR molecules attract each other when the positive end of one molecule is near the negative end of another These dipole-dipole forces only work when the polar molecules are very close together Weaker than ion-dipole forces

In liquids polar molecules are free to move with respect to one another Different configurations create orientations that are attractive, and orientations that are repulsive.

Ion-Dipole Forces Pretty simple Exists between an ion and a partial charge on a polar molecule. Increases as either the charge of the ion increases, or the magnitude of the dipole moment increases

Rules for Dispersion Forces  When the molecules of two substances have similar weight and shape, dispersion forces are approximately equal Differences in the attraction forces are due to differences in strengths of the dipole-dipole forces More polar molecules have the stronger attractions

 When the molecules of two substances differ in molecular weights, dispersion forces tend to be decisive in finding which substances has the stronger attractions Differences in the magnitudes of the forces usually because of differences in atomic weight More massive molecules having stronger attractions

Hydrogen Bonding Hydrogen bonding is a special type of intermolecular force Always between a H atom in a polar bond and an unshared electron pair on a nearby, small, electronegative ion or atom Usually a H-F, H-O or H-N bond Usually a F, O or N atom in another molecule

Example A hydrogen bond exists between the H atom in HF molecule and the F atom of a nearby HF molecule F-HÅÅÅÅF-H dots represent the hydrogen bond

The hydrogen bond is a unique form of a dipole-dipole attraction Because F, N and O are so electronegative, bonds are VERY polar H on positive end H has no inner core of electrons Positive side has partially exposed nucleus Gives a large dipole effect (since not just electron density, but actual nuclear charge) Also, since H is so small, it can approach an electronegative atom very closely

Hydrogen bonds still weaker than ordinary bonds Stronger than dipole-dipole or dispersion forces

Comparing Intermolecular Forces Dispersion forces are found in all substances Strengths increase with increasing molecular weight Strength increases with longer molecules

Dipole-dipole forces adds to dispersion forces Found only in polar molecules Hydrogen bonds require H atoms bonded to F, O or N Also adds to dispersion forces

11.5 – Vapor Pressure Molecules can escape from the surface of a liquid to the gas phase by evaporation. Suppose we place an amount of ethanol (C 2 H 5 OH) in an evacuated, closed container. Ethanol will quickly begin to evaporate So pressure exerted by the vapor above the liquid will increase Eventually, the pressure of the vapor will become a constant value This is called the vapor pressure of the substance

Explaining Vapor Pressure Molecules of a liquid move at various speeds At any instant, some of the molecules at the surface of the liquid get enough kinetic energy to overcome the attractive forces of their neighbors Thus escaping into the gas phase The weaker the attractive forces, the more particles that can escape, and therefore, more vapor pressure

At any given temperature the movement of molecules from liquid to gas goes on continuously. As the number of gaseous molecules increase, the probability increases that a molecule in the gas phase will hit the liquid surface and be recaptured by the liquid Eventually, rate at which molecules return to liquid equals the rate at which they escape So we get a steady number of molecules in the gas phase

Dynamic Equilibrium The condition when two opposing processes are occurring at the same time, and at the same rate, is called dynamic equilibrium Often referred to simply as equilibrium A liquid and vapor are in equilibrium when evaporation and condensation occurs at equal rates Often appears like nothing is happening, but no net change But particles are constantly changing from liquid  gas and from gas  liquid The vapor pressure of a liquid is the pressure exerted by its vapor when vapor and liquid are in equilibrium.

Volatility, Vapor Pressure and Temperature Substances with high vapor pressure evaporate more quickly than substances with low vapor pressure Liquids that evaporate easily are said to be volatile As temperature increases, particles move more, and will evaporate more. Vapor pressure will increase as temperature increases Non-linear progression

Vapor Pressure and Boiling Point A liquid boils when vapor pressure equals the external pressure acting on the surface of the liquid The temperature at which a liquid boils increases with increasing external pressure The boiling point of a liquid at 1 atm pressure is called its normal boiling point

11.6 – Phase Diagrams The equilibrium between liquid and vapor is not the only dynamic equilibrium we deal with. Under the right conditions (of pressure and temperature) we can have other dynamic equilibriums solid / liquid solid / vapor

A phase diagram is a graphical way to summarize the conditions which cause the various equilibrium to exist. Allow us to predict the phase of matter given a temperature and pressure The pressure shown in the diagram is either the pressure applied to the system or the pressure generated by the substance itself

Three main curves on the graph  The line from A to B Equilibrium between gas and liquid phase The pressure at 1 atm on this graph represents the normal boiling point of the substance The curve ends at the critical point

Critical Point? The critical point is the critical temperature and critical pressure of the substance Critical temperature is the temperature where beyond it, the substance will never again be a liquid Intermolecular forces too weak, no matter what the pressure Critical pressure is the pressure that is needed to create a liquid at the critical temperature

 The line AC represents the variation in the vapor pressure of the solid as it sublimes at different temperatures

 The line AD represents the change in melting point of the solid with increasing temperature Usually slopes right as pressure increases Because solid is usually denser than liquid An increase in pressure usually favors the more compact solid phase, so higher temperature needed Melting and freezing point are the same The melting point at 1 atm is normal melting point

Point A, where the lines meet, is known as the triple point All three phases are in equilibrium at this temperature and pressure

Practice Interpreting Phase Diagrams - Water

Carbon Dioxide