Kmt, intermolecular forces, intro energy

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Information about Kmt, intermolecular forces, intro energy
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Published on March 18, 2014

Author: felipedelagarzam

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kinetic molecular theory,

Liquids and SolidsLiquids and Solids Quick brainstorm What are some basic differences between gases and liquids? between liquids and solids?

A look back at the three statesA look back at the three states Particle arrangement is the key!

The Kinetic Molecular TheoryThe Kinetic Molecular Theory The purpose of the KMT is to help explain particle behavior in the three states of matter. We will re-visit this theory when we look at gases which is the next unit. Now: Let’s look at liquids and solids.

Liquids and the KMTLiquids and the KMT • Liquids have a higher density than gases due to the closer arrangement of the particles. • Liquids are much less compressible than gases. • Particles are not bound to fixed positions, rather they move constantly. This motion is fluidity. Fluidity depends on forces and temperature. viscosity- resistance to flow

Liquids and the KMTLiquids and the KMT • The speed of motion of liquid particles depends on the temperature. Liquids under extremely cold temps. Would be very viscous- resist flow. • Liquids diffuse and mix with other liquids. Speed of diffusion is governed by temperature of the liquid and the types of forces between the liquid particles. • Liquids exhibit surface tension and capillary action.

Surface TensionSurface Tension The forces felt by particles on the surface are primarly down and to the sides. The surface tension is directly related to strength of forces between particles.

Solids and the KMTSolids and the KMT • Particles are closely packed, forces play a strong role. • Particle motion is limited to vibrations primarily due to the strength of particle attractions. • Most substances are most dense as a solid.

Solids and the KMTSolids and the KMT • Solids are less compressible than liquids due to close particle arrangement. • Diffusion does occur, but the rate is millions of times slower than in liquids.

Two types of solid structuresTwo types of solid structures crystalline solids- ordered particle arrangement. amorphous solids- random arrangement of particles.

Ionic CrystalsIonic Crystals Attractive forces between cations and anions. Very strong attractive forces.

Covalent network crystalsCovalent network crystals Each atom is covalently bonded to a neighboring atom. Examples are carbon in the form of diamond, graphite, nanotube and buckyball These are all Allotropes

Metallic CrystalsMetallic Crystals Metal atoms surrounded by a sea of electrons

The Forces between particlesThe Forces between particles We have been focusing on the fact that particles in a liquid and solid are closer together, and must experience some form of attraction (strong or weak) to neighboring particles. Now, we will define (describe) those forces.

List 2 ways that people can cause weathering and erosion to take place. Forces http://en.wikipedia.org/wiki/Sodium_Chloride

Intermolecular Forces • Polar Molecules have • Dipole-dipole Force: • + end of one molecule attracts – end of another (based on electronegativity difference IN polar molecules) – PH3

PolarityPolarity How do you determine polar vs. nonpolar? Draw the molecule: Perfectly symmetric – nonpolar = EQUAL sharing of electrons Asymmetric and has polar bonds – polar = UNEQUAL sharing of electrons

Dipole-DipoleDipole-Dipole The force of attraction between two polar molecules

Hydrogen BondsHydrogen Bonds (not really bonds)(not really bonds) STRONG DIPOLE-DIPOLESTRONG DIPOLE-DIPOLE The force of attraction between hydrogen (H) and the lone pairs of electrons on a highly electronegative atom (F, O, or N)

Non-Polar molecules have • London Dispersion Forces: nonpolar molecules have temporary dipoles (due to electron motion in atoms) – I2 A second atom or molecule, in turn, can be distorted by the appearance of the dipole in the first atom or molecule (because electrons repel one another) which leads to an electrostatic attraction between the two atoms or molecules http://www.chem.purdue.edu/gchelp/liquids/disperse.html

Intermolecular Forces • Hydrogen bonds: special type of dipole- dipole force that results from large electronegativity difference between hydrogen and nitrogen, oxygen, or fluorine • H and FON • Not a true bond

London dispersion forcesLondon dispersion forces The movement of electrons causes an instantaneous dipole, which can then induce a dipole in a neighboring atom, generating a temporary force.

Intermolecular forces Force Type of Compound Relative strength Boiling Points of Compounds with this force Melting Points of Compounds with this force Ionic Dipole- Dipole Hydrogen Bonds London Dispersion Generalization: As force strength _____________, energy required to separate molecules/ ions ___________, therefore MP and BP ______________.

Types of intermolecular forcesTypes of intermolecular forces Is the molecule polar or nonpolar? Is H bonded to either F, O, or N? Hydrogen Bond Dipole – Dipole force London Dispersion Force NO YES POLAR NONPOLAR

ExamplesExamples CH4 NH3 H2O Cl2 HBr

Strength of the forcesStrength of the forces Which of these forces would you consider the strongest? How do ionic bonds (aka ionic forces) fit into the picture? Hydrogen BondsHydrogen Bonds Ionic Bonds/Forces are the strongestIonic Bonds/Forces are the strongest forces between two particles.forces between two particles.

Quick Questions: • Cl2 and Br2 have approximately the same shape and neither is polar. – Which intermolecular force affects Cl2 and Br2? – Upon cooling, both Cl2 and Br2 form solids. Why? – At 25o C, chlorine (Cl2 ) is a gas whereas bromine (Br2 ) is a liquid. Why?

What about energy changes with physical changes? -Energy changes during phase transitions: Heating/ Cooling Curves - Phase Diagrams: Graphical depictions of the relationship between the phases of matter

How do we predict how much energy is required to change a substance’s phase? • Molecular Mass: as mass increases, the interaction between molecules increases (due to more electrons) • Molecular Geometry: polar or nonpolar molecules • Intermolecular Forces – Forces between two ions or molecules. – As opposed to intraparticlar forces • Covalent bonding

Thermochemistry The study of energy changes during physical and chemical changes

What is Energy? Definition of Energy: the capacity to do work or produce heat Equation: E = q (heat) + w (work) Forms that energy may take: Kinetic, potential, light, heat, work Units for energy: Joules (J) or Kilojoules (kJ) Calories (cal) = 4.184J = 1 cal

Thermochemistry in two big categories Energy associated with Chemical Changes Exothermic reactions Endothermic reactions Energy associated with Physical Changes Melting Freezing Boiling Vaporizing Sublimation Deposition

How is energy involved in phase changes? Endothermic Exothermic Melting Freezing Boiling (Evaporating) Condensing Sublimating Deposition

Physical Change What is the Kinetic Theory? - all particles are in constant random motion http://www.usatoday.com/weather/tg/wevapcon/wevapcon.htm http://www.nyu.edu/pages/mathmol/modules/water/water_concepts.html

Energy Causes Phase Changes • Adding energy to a substance can cause two things to occur – The intermolecular forces within the substance can weaken due to added heat causing molecular vibration and increased temperature. – The intermolecular forces within the substance can be reduced to a point at which the substance changes phase

Temperature Temperature - When samples of different temperatures are in contact, ENERGY is transferred.

Heat and Temperature Temperature Temperature (T) – a measure of heat flow An average of the kinetic energy of all the atoms/particles of a substance - Celsius, Kelvin - Kelvin = Celsius + 273 Heat Heat (q) – the total energy of a substance. A sum of the kinetic and potential energies of a substance (E=q constant pressure) The first law of thermodynamics: the law of the conservation of energy.

Enthalpy A substance’s energy can also be measured by enthalpy. (H) H = q (heat) + w (work) At constant pressure, no work can be done on or by the system. So, w = 0 H = q (at constant pressure) Enthalpy is equivalent to heat if the system is at constant pressure.

What is a heating/ cooling curve? Shows the energy changes for a substance during a phase change Movie Time Temperature(°C)

What if there is a phase transition (freezing/ melting)? Q = mHfus • Hfus = enthalpy of fusion – Energy required to change 1 gram of a substance from a solid to liquid (no temperature change) – Units: J/g • Why is Hfus positive when you’re melting a substance?

What if there is a phase transition (evaporation/condensing)? Q = mHvap • Hvap = enthalpy of vaporization – Energy required to change 1 gram of a substance from a liquid to a gas (no temperature change) – Units: J/g • Why is Hvap positive when you’re evaporating a substance?

How are energy changes within a phase measured? Q = mCp∆T (like the calorimetry equation) There is a temperature change!!! What is specific heat?

What is a Phase Diagram? http://www.chemistrycoach.com/Phase_diagram.htm http://jchemed.chem.wisc.edu/JCESoft/CCA/CCA2/MAIN/BENZENE/CD2R1.HTM Pressure(atm) Temperature (°C) 1.0 0.5 1.5 2.0 -10 57 60 B C A D E

Phase Diagram Terms • Triple Point: the temperature and pressure at which the solid, liquid, and gas phases coexist at equilibrium • Critical Point: the temperature and pressure at which the gas and liquid states of a substance become identical and form one phase

• Normal boiling point: The temperature at which the liquid and gas phases are in equilibrium. • Normal melting point: The temperature at which the solid and liquid phases are in equilibrium.

What is Vapor Pressure? • Gas pressure exerted by particles that have escaped from a liquid • As temperature increases, more molecules can escape into the gas phase, increasing the pressure above the liquid

Vapor Pressure (Pvap)Pvap(atm) Temperature (°C) Diethyl ether Ethanol Water

Physical Change and Phase Diagrams http://wps.prenhall.com/wps/media/objects/602/616516/Media_Assets/Chapter10/Text_Images/FG10_28.JPG

More Phase Diagrams http://www.chem.neu.edu/Courses/1131Tom/Lecture25/img007.GIF

More Phase Diagrams Temp (°C) Pressure(atm) A B C D E 1 5 15 32

Phase Diagrams Think ABOUT Based on the phase diagram for water explain to the person next to you how you think ice skating works… (Hint: think pressure and the solid liquid phase boundary)

Phase diagrams • What can we determine about density from a phase diagram? • Based on the relative density of the liquid and solid phases of water, why do you think ice skating works?

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