Ch. 13 Notes:

 

Physical Properties of Liquids-

Surface tension- an increase in the attractive forces between molecules at the surface of a liquid; thereby decreasing the surface area.

            *The total attractive force is divided between fewer adjacent molecules (no molecules on top of liquid) so the attraction is greater.

 

Capillary Action- spontaneous rising of a liquid in a tube.

Cohesive forces- attractions between identical molecules in the liquid.

 

Adhesive forces- attractions between different molecules.

 

Cohesive>adhesive = beading

Cohesive<adhesive = no beading

 

e.g., Explain the meniscus in a graduated cylinder.

 

e.g., Explain the use of surfactants in dishwashers.

 

 

Viscosity- a liquid’s resistance to flow.

            e.g., water has a low viscosity à flows easily. Syrup has a high viscosity à flows slowly.

            *Why? Attractive forces in syrup are strong holding the substance together.

 

Viscosity decreases with high temperatures due to increased K.E. of the molecules.

 

Evaporation-

            *depending on the attractive forces in a substance a certain amount of K. E. is needed to change a liquid to a gas.

 

At a constant T, the proportion of molecules w/enough K.E. to escape will remain constant and the liquid will evaporate at a uniform rate.

            *The greater the surface area, the greater the evaporation rate

            *The greater the T, the greater the K.E. and therefore the # of molecules with enough energy to evaporate.

 

E.g., The temperature of a beaker of liquid evaporating on a lab bench is usually constant b/c it can absorb heat from its surroundings easily and quickly. If the beaker is insulated, however, the liquid cools and the rate of evaporation decreases. Explain what is occurring?

 

Answer: when a molecule escapes from the liquid, the average K.E. of the remaining molecules decreases. If the average K.E. decreases then the Temperature decreases explaining the cooling. Also, as the average K.E. decreases, the # of molecules that have K.E. above the escape energy decreases. So, because fewer molecules have enough energy to escape, the rate of evaporation decreases.

 

Vapor Pressure- the pressure of a gas above a liquid in a closed container.

            *molecules still evaporate but some bounce off the walls of the container and those hitting the liquid are condensed. This continues w=until equilibrium is reached.

            *The # of molecules evaporating depends on T and the attractive forces of the liquid.

            *As T increases – V.P. increases

 

Boiling- when the V.P. = Atmospheric Pressure

 

Since B.P. varies with pressure – normal boiling point refers to 1.00 atm pressure.

            e.g., Denver, CO = is less than 760 mmHg. How does this change the boiling temperature? Answer- It decreases. It takes longer to cook food properly b/c at a lower temperature.

 

Heat of Vaporization (∆Hvap) – energy needed to convert 1 gram of liquid into 1 gram of gas at a liquid’s normal B.P.

            *J/g or J/mol

            * positive because you always add energy to a liquid to vaporize it.

            *value depends on intermolecular forces.

                        e.g., Of same-size molecules, which forces (London dispersion, dipole-dipole or H-bonding) have the highest Hvap values?

 

Predict which compound will have 1-the lower B.P., 2 – the higher ∆Hvap , 3 – the higher evaporation rate, 4 – the lower V.P.

            a. C6H14  or  C8H18                   b. HF or HCl               c. C6H13OH or C3H7OC3H7

 

 

Clausius-Clapeyron Eqn.- relates V.P. and ∆Hvap

            lnP = ∆Hvap  + C                       P = Vapor Pressure

                        RT                               C = constant

 

*lnP vs. 1/T on a graph gives a slope of  - ∆Hvap/R

 

Example- What are the heat of vap (kJ/mol) and the normal boiling point of a liquid that has a vapor pressure of 254 mmHg at 25oC and a vapor pressure of 648 mmHg at 45oC? (R= 8.314 J/mol K)  Answer: 36.9 kJ/mol, 322 K

 

 

Solids-

Metallic crystals – rigid structure of metal nuclei with a sea of electrons.

            *mobile e- - explains conductivity and malleability of metals.

 

The atoms in metallic crystals are arranged in the most compact form so they crystallize as

·        face-centered cubic structure   OR

·        body-centered cubic structure

 

alloy – mixture of elements with metallic properties

            *substitutional – when metal atoms are replaced by metal atoms of similar size

                        e.g., Brass = Cu and Zn

            *interstitial – when smaller atoms “fill” the empty spaces in the metals’ crystal structure

                        e.g., steel = Fe and C

 

alloy steels – interstitial alloy with substitutional alloy

 

Network Atomic Solids – contain strong directional covalent bonds.

            *typically brittle and poor conductors.

            e.g., C – two forms

                        *diamond – tetrahedral (sp3) , no conductivity

                        *graphite – layers (w/weak bonds) and trigonal planar (sp2), other p orbital accounts for conductivity b/c it has delocalized electrons that can pass current – “conductive bands”

 

e.g., Si – bonds mostly with Oxygen = silica – SO2, found in quartz or sand

            *not like CO2. Why? Si cannot form pi bonds w/O b/c of its large radius so it makes 4 sigma bonds instead.

            *empirical formula for silica is SiO2 but it is really SiO4 as a tetrahedral.

 

Semiconductors – Silicon

            *The gap between a filled and empty molecular orbitals is small so a few electrons are able to (with increasing T) gain enough energy to jump up to the empty conductive band.

 

Molecular Solids – forces between depend on nature of the molecule

            *Dispersion and dipole-dipole

 

Phase Changes –

Heating Curve – a plot of T vs. time showing the changes of state of a substance goes through. (considering P constant)

 

 

When a substance reaches its melting point – what happens to the curve? The molecules?

 

Heating a solid below its melting point at a constant rate undergoes the following:

  1. T of solid increases at a constant rate until it melts.
  2. When melting, T stops rising and stays constant until s à l
  3. T of liquid increases at a constant rate until reach B.P.
  4. When reach boiling, T stops increasing and stays constant until l à g
  5. T of gas increases at a constant rate.

 

 

Heat capacity – (J/oC) reciprocal of slope of curve where T increases as heat is added.

 

Specific heat – (J/goC) heat capacity divided by # g of sample.

 

Enthalpy of fusion (melting) – length of solid-melting plateau

            Amount of heat added divided by moles of sample

 

Enthalpy of vaporization(boiling) – length of liquid-boiling plateau

            Amount of heat/# moles of sample

 

Supercooling – a metastable liquid that is cooled to a T below its M.P.

 

Metastable – unstable, if disturbed liquid will crystallize rapidly.

 

How much energy does it take to convert 130. grams of ice at  - 40oC to steam at 160oC.

            (7.22 mol = 130. g)                  (2.1 J/goC)

1. 40oC x 130. g x 2.1 J/goC (specific heat cap of water(s) = 10.92 kJ

2. 7.22 mol x 6.0kJ/mol (Hfus of water) = 43.4 kJ

3. 100oC x 130. g x 4.2 J/goC (specific heat cap of water(l)) = 54.6 kJ

4. 7.22 mol x 43.9 kJ/mol (Hvap of water) = 317kJ

5. 6ooC x 130. g x 1.8J/goC (specific heat cap of water(g)) = 14.0 kJ

 

Total = 440. kJ

 

Phase Diagrams – shows the relationship between pressure and temperature and the three states of matter.

            Triple point- when all three phases are in equilibrium

            Equilibrium – mixture of two phases on that line (exist as same time)

 

Know the standard form of a phase diagrams- what is on the x-axis? Y-axis? Where are your phases?

 

Critical point – the max. temperature at which any liquid can exist.

 

Supercritical fluid – point above the critical point.