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Liquids and solids 10.1 Intermolecular forces u Inside molecules (intramolecular) the atoms are bonded to each other. u Intermolecular refers to the.

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Presentation on theme: "Liquids and solids 10.1 Intermolecular forces u Inside molecules (intramolecular) the atoms are bonded to each other. u Intermolecular refers to the."— Presentation transcript:


2 Liquids and solids

3 10.1 Intermolecular forces u Inside molecules (intramolecular) the atoms are bonded to each other. u Intermolecular refers to the forces between the molecules. u These are what hold the molecules together in the condensed states.

4 Intermolecular forces u Strong covalent bonding ionic bonding u Weak Dipole dipole London dispersion forces Hydrogen bonding u During phase changes the molecules stay intact. u Energy used to overcome forces.

5 Dipole - Dipole u Remember where the polar definition came from? u Molecules line up in the presence of a electric field. The opposite ends of the dipole can attract each other so the molecules stay close together. u 1% as strong as covalent or ionic bonds. u Weaker with greater distance. u Small role in gases.


7 Hydrogen Bonding u Especially strong dipole-dipole forces when H is attached to F, O, or N u These three because- They have high electronegativity. They are small enough to get close. u Affects boiling point.

8 Water Hydrogen Bonding Clip

9 CH 4 SiH 4 GeH 4 SnH 4 PH 3 NH 3 SbH 3 AsH 3 H2OH2O H2SH2S H 2 Se H 2 Te HF HI HBr HCl Boiling Points 0ºC

10 London Dispersion Forces u Non - polar molecules also exert forces on each other. u Otherwise, no solids or liquids. u Electrons are not evenly distributed at every instant in time. u Have an instantaneous dipole. u Induces a dipole in the atom next to it. u Induced dipole - induced dipole interaction.

11 Example HH HH HH HH + HH HH + + LD Video

12 London Dispersion Forces u Weak, short lived. u Lasts longer at low temperature. u Eventually long enough to make liquids. u More electrons, more polarizable. u Bigger molecules, higher melting and boiling points. u Exist in all molecules, much weaker than other forces. u Also called Van der Waals forces.

13 #36

14 #38

15 #39

16 #40 (In Webassign)

17 10.2 Liquids u Many of the properties due to internal attraction of atoms. Beading Surface tension Capillary action u Stronger intermolecular forces cause each of these to increase.

18 Surface Tension u Molecules in the middle are attracted in all directions. Molecules at the top are only pulled inside. u Minimizes surface area.

19 Capillary Action u Liquids spontaneously rise in a narrow tube. u Intermolecular forces are cohesive, connecting like things. u Adhesive forces connect to something else. u Glass is polar. u It attracts water molecules.


21 Beading u If a polar substance is placed on a non-polar surface. There are cohesive, But no adhesive forces. u And Visa Versa

22 Viscosity u How much a liquid resists flowing. u Large forces, more viscous. u Large molecules can get tangled up.

23 Model for Liquids u Cant see molecules so picture them in motion but attracted to each other. u With regions arranged like solids but with higher disorder. with fewer holes than a gas. Highly dynamic

24 Phases u The phase of a substance is determined by three things. u The temperature. u The pressure. u The strength of intermolecular forces. u Changes of State Video

25 10.3 Solids u Two major types. u Crystalline - have a regular arrangement of components in their structure. (table salt) u Amorphous - those with much disorder in their structure. Said to be frozen in place. (glass, plastic)

26 10.4 Bonding Models for Metals u Why do metal atoms stay together? How does their bonding effect their properties? u Two Models: u Electron Sea Model: A regular array of metals in a sea of electrons. u Band (Molecular Orbital) Model: Electrons assumed to travel around metal crystal in MOs formed from valence atomic orbitals of metal atoms.

27 Electron Sea Model

28 10.5 C & Si - Atomic Network Solids u Composed of strong directional covalent bonds that are best viewed as a giant molecule. 4 brittle 4 do not conduct heat or electricity 4 carbon, silicon-based 4 graphite, diamond, ceramics, glass u Diamond - hardest natural substance on earth, insulates both heat and electricity. u Graphite - slippery, conducts electricity.

29 Diamond - each Carbon is sp3 hybridized, connected to four other carbons. u Carbon atoms are locked into tetrahedral shape. Strong bonds give the huge molecule its hardness.

30 u Each carbon is connected to three other carbons and sp 2 hybridized. u The molecule is flat with 120º angles in fused 6 member rings. The bonds extend above and below the plane. Graphite is different

31 This bond overlap forms a huge bonding network. u Electrons are free to move through out these delocalized orbitals. u The layers slide by each other.

32 10.6 Molecular solids. u S 8, P 4, CO 2, H 2 O u Different molecules have different forces between them. u These forces depend on the size of the molecule. u They also depend on the strength and nature of dipole moments. u Non-Polar: Large molecules (such as I 2 ) can be solids even without dipoles. u Polar: Dipole-dipole forces are generally stronger than L.D.F. u Hydrogen bonding is stronger than Dipole-dipole forces. u No matter how strong the intermolecular force, it is always much, much weaker than the forces in bonds. u Stronger forces lead to higher melting and freezing points.

33 10.7 Ionic Solids u The extremes in dipole dipole forces- atoms are actually held together by opposite charges. u Huge melting and boiling points. u Atoms are locked in lattice so they are hard and brittle. u Every electron is accounted for so they are poor conductors - good insulators.

34 What type of solid will each substance form? #68 u Diamond Quartz u PH 3 NH 4 NO 3 u H 2 Ar u Mg Cu u KCl C 6 H 12 O 6 u SF 2

35 10.8 Vapor Pressure u Vaporization - change from liquid to gas at boiling point. u Evaporation - change from liquid to gas below boiling point. Heat (or Enthalpy) of Vaporization ( H vap ) - the energy required to vaporize 1 mol at 1 atm.

36 u Vaporization is an endothermic process - it requires heat. u Energy is required to overcome intermolecular forces. u Responsible for cool earth. u Why we sweat.

37 Condensation u Change from gas to liquid. u Achieves a dynamic equilibrium with vaporization in a closed system. u What is a closed system? u A closed system means matter cant go in or out. u What the heck is a dynamic equilibrium?

38 Dynamic equilibrium u When first sealed the molecules gradually escape the surface of the liquid.

39 u Rate of Vaporization = Rate of Condensation u Molecules are constantly changing phase Dynamic u The total amount of liquid and vapor remains constant Equilibrium Dynamic equilibrium

40 Vapor pressure u The pressure above the liquid at equilibrium. u Liquids with high vapor pressures evaporate easily. They are called volatile. u Increases with increasing temperature. u Decreases with increasing intermolecular forces. Bigger molecules (bigger LDF) More polar molecules (dipole-dipole)

41 Changes of state u The graph of temperature versus heat applied is called a heating curve. u The temperature a solid turns to a liquid is the melting point. The energy required to accomplish this change is called the Heat (or Enthalpy) of Fusion H fus

42 Heating Curve for Water Ice Water and Ice Water Water and Steam Steam

43 Heating Curve for Water Heat of Fusion Heat of Vaporization Slope is Heat Capacity

44 Melting Point u Melting point is determined by the vapor pressure of the solid and the liquid. u At the melting point the vapor pressure of the solid = vapor pressure of the liquid at 1 atm

45 Boiling Point u Reached when the vapor pressure equals the pressure of the surrounding atmosphere. u Normal boiling point is the boiling point at 1 atm pressure. u Super heating - Rapid heating above the boiling point allows liquid state to exist above normal boiling point. u Supercooling - Rapid cooling below the freezing point allows liquid state to exist below the normal freezing point.

46 Practice


48 10.9 Phase Diagrams u A plot of temperature versus pressure for a closed system, with lines to indicate where there is a phase change. u Critical temperature: temperature above which the vapor can not be liquefied. u Critical pressure: pressure required to liquefy at the critical temperature. u Critical point: critical temperature and pressure (for water, T c = 374°C and 218 atm). Liquid & solid are indistinguishable.

49 Solid Liquid Gas Triple Point Critical Point Temperature Pressure

50 Temperature Solid Liquid Gas 1 Atm A A B B C C D D D Pressure D

51 Solid Liquid Gas u This is the phase diagram for water. u The density of liquid water is higher than solid water. Temperature Pressure

52 Solid Liquid Gas 1 Atm u This is the phase diagram for CO 2 u The solid is more dense than the liquid u The solid sublimes at 1 atm. Temperature Pressure


54 u Like most substances, bromine exists in one of the three typical phases. Br 2 has a normal melting point of -7.2°C and a normal boiling point of 59°C. The triple point for Br 2 is -7.3°C and 40 torr, and the critical point is 320°C and 100 atm. Using this information, sketch a phase diagram for bromine indicating the points described above. –Based on your phase diagram, order the three phases from least dense to most dense. –What is the stable phase of Br 2 at room temperature and 1 atm? –Under what temperature conditions can liquid bromine never exist? –What phase changes occur as the temperature of a sample of bromine at 0.10 atm is increased from -50°C to 200°C?

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