Viscosity and intermolecular forces relationship goals

Effect of Intermolecular Forces on Physical Properties |

viscosity and intermolecular forces relationship goals

The concepts of cohesion, adhesion, and viscosity are defined. Liquids flow because the intermolecular forces between molecules are weak enough to . and so have relatively weak intermolecular forces; the difference is the size. . Our Content · NGSS · Mission & Goals · Sponsorship · Recognition & Awards · Terms. LEARNING GOALS , Deduce the relative vapor pressure of substances when given their structures or formulas. The relative strength of intermolecular forces can often be determined by observing their Evaporation vs Boiling Point - They BOTH represent liquids turning to gases but there is a distinct difference. LEARNING OBJECTIVES. By the end of this section, you will be able to: Distinguish between adhesive and cohesive forces; Define viscosity, surface tension, roles of intermolecular attractive forces in each of these properties/ phenomena . In this equation, h is the height of the liquid inside the capillary tube relative to the.

So, in van der Waals dispersion forces: Boiling point is higher for larger compounds Melting point is higher for larger compounds Freezing point is lower for smaller compounds Vapor pressure is higher for smaller compounds Van der Waals Dipole-Dipole Interactions A partial positive charge and a partial negative charge can be created between two atoms when there is a difference in electronegativity.

These interactions are called van der Waals dipole-dipole interactions. For example, carbon is less electronegative than oxygen, creating a partial positive on carbon and a partial negative on oxygen.

viscosity and intermolecular forces relationship goals

The dipole interactions are stronger than the dispersion forces because the oxygen will almost always have slightly more electrons than the carbon, instead of constantly changing.

There still isn't a full negative charge on the oxygen, or a full positive charge on the carbon. But the partial positive and negative charges are still enough to attract opposite charges together. The higher the difference in electronegativity, the strong the dipole-dipole interactions will be.

So compounds with a higher electronegativity difference will have strong intermolecular forces. Hydrogen Bonding Under certain circumstances the dipole bonds created from hydrogen can create very strong bonds. Let's say that you have an atom that is attached to several hydrogen atoms: You may have even seen these attractive adhesive forces in action in lab. When water is in a glass graduated cylinder, for example, the water creeps up the sides of the glass, creating a concave curve at the top called a meniscus, as shown in the figure below.

Chapter 10 – Liquids and Solids AP Chemistry Goals 1. Differentiate

Water in graduated cylinders made out of some types of non-polar plastic, on the other hand, forms a flat meniscus because there are neither attractive nor repellant cohesive forces between the water and the plastic.

See Figure 8 for a comparison of polar and non-polar graduated cylinders. In graduated cylinder A, made of glass, the meniscus is concave; in cylinder B, made of plastic, the meniscus is flat. Viscosity At the beginning of the module, we said that one of the defining features of liquids is their ability to flow. But among liquids there is a huge range in how easily this happens. Consider the ease with which you can pour yourself a glass of water, as compared to the relative challenge of pouring thick, slow-moving motor oil into an engine.

The difference is their viscosityor resistance to flow. Motor oil is quite viscous ; water, not so much. Remember, water molecules form strong hydrogen bonds with each other. Pentane, on the other hand, made up of just hydrogen and carbon atoms, is nonpolar, so the only types of intermolecular forces it can form are the relatively weak London dispersion forces.

The weaker intermolecular forces mean that the molecules can more easily move past each other, or flow — hence, lower viscosity. But both water and pentane are relatively small molecules. For example, compare pentane to motor oil, which is a complex mixture of large hydrocarbons much larger than little pentane, and some with dozens or even hundreds of carbons in a chain.

Both liquids are nonpolar, and so have relatively weak intermolecular forces ; the difference is the size. The big, bendy motor oil hydrocarbons can literally get tangled with their neighbors, which slows the flow.

Group A consists of large molecules in a tangled blob a viscous liquid and Group B consists of smaller molecules with fewer entanglements a less viscous liquid. Returning to our original comparison of motor oil versus water, even though water has such strong intermolecular forcesthe much larger size of the molecules in the motor oil makes the oil more viscous.

This is the case because temperature affects both of the factors that determine viscosity in the first place. It also makes the molecules move around more, so those big molecules that got tangled up when they were cold become more dynamic and are more able to slide past each other, allowing the liquid to flow more easily.

viscosity and intermolecular forces relationship goals

Comprehension Checkpoint Motor oil pours more slowly than the solvent pentane because motor oil is made up of a. Complex liquids When you think of water, you might think of its chemical formulaH2O. This formula describes a pure liquid composed only of H2O molecules, with absolutely no other components.

The reality, though, is that the vast majority of liquids we encounter are complex mixtures of many compounds.

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Solutions are made of a liquid solvent in which one or more solutes are dissolved. Solutes can be solidsliquidsand gases. There are many, many common solutions that use water as the solvent, including salt water and pretty much any type of flavored drink.

This is because the vapor pressure is below the applied pressure, and the piston moves in against the gas until it all condenses into the liquid. For all temperature and pressure combinations to the left of the curve, only liquid exists. What if we start at a temperature-pressure combination on the curve and elevate the applied pressure without raising the temperature? The applied pressure is now greater than the vapor pressure, and as before all of the gas will condense into the liquid.

Just as before, for all points to the left of or above the curve, only liquid exists. The opposite reasoning applies if we decrease the applied pressure.

Figure 1 thus actually reveals to us what phase or phases are present at each combination of temperature and pressure: We know that, if the temperature is low enough, we expect that the water will freeze into solid. To complete the phase diagram, we need additional observations. We go back to our apparatus we used before, with a piston in a cylinder trapping liquid water and vapor in phase equilibrium.

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If we slowly lower the temperature, the vapor pressure decreases slowly as well, as shown in Figure 1. However, if we continue to lower the temperature, we observe an interesting transition, as shown in the more detailed Figure 2. Below this temperature, the pressure continues to vary smoothly, but along a slightly different curve.

To understand what we have observed, we examine the contents of the container. As with the liquid-vapor curve, we can interpret this new curve in two ways. The solid-gas curve gives the vapor pressure of the solid water as a function of temperature, and also gives the sublimation temperature as a function of applied pressure.

Figure 2 is still not a complete phase diagram, because we have not included the combinations of temperature and pressure at which solid and liquid are at equilibrium. Very careful measurements reveal that the solid-gas line and the liquid-gas line intersect in Figure 2 where the temperature is 0. Under these conditions, we observe inside the container that solid, liquid, and gas are all at equilibrium inside the container.

If we raise the applied pressure slightly above the triple point, the vapor must disappear. We can observe that, by only slightly varying the temperature, the solid and liquid remain in equilibrium.

We can further observe that the temperature at which the solid and liquid are in equilibrium varies almost imperceptibly as we increase the pressure. If we include the solid-liquid equilibrium conditions on the previous phase diagram, we get Figure 3, where the solid-liquid line is very nearly vertical. Figure 3 Figure 3 is an example of a complete phase diagram. This diagram shows for each temperature and pressure which phase or phases are present at equilibrium.

Figure 3 is for water, but each substance has its own unique phase diagram, similar in appearance. Boiling Points and Intermolecular Forces Earlier in this study, we determined that the boiling point of a liquid is the temperature at which the vapor pressure of the liquid equals the pressure applied externally, e.

From our work on dynamic equilibrium, we know that this higher temperature is required to provide sufficient kinetic energy for the molecules in the liquid to overcome stronger attractions between the molecules.