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CHM 150 Chapter 10 McMurry-Fay Notes page 1 of 14

Chapter 10: Liquids, Solids, and Phase Changes

In-chapter exercises: 10.1–10.6, 10.11; End-of-chapter Problems: 10.26, 10.31, 10.32, 10.33, 10.34, 10.35, 10.36, 10.39, 10.40, 10.42, 10.44, 10.45, 10.66, 10.82, 10.84, 10.86 Kinetic Molecular Theory of Liquids and Solids phase (=physical state): solid, liquid, or gas Solids have the lowest kinetic energy (KE)—i.e. do not move very much – Highest attraction between particles → particles are stuck in specific sites = very confined Liquids have slightly higher KE—i.e. particles moving more than in solid – Particles are still attracted and maintain contact with one another but can move past one another → particles are less confined Gases have greatest KE—i.e. particles move quickly and randomly – Attractive forces almost (if not) completely overcome, so particles can fly freely within container → particles are far away from each other = unrestricted 10.2 INTERMOLECULAR FORCES intermolecular forces: Attractive forces between 2 molecules – e.g. between 2 water molecules Ion-Dipole Forces – Attraction between an ion and the oppositely charged end of a polar molecule

– e.g. between Na+ and the negative end of a H2O molecule (O in H2O) or between Cl– and the positive end of a H2O molecule (H atoms in H2O)


Dispersion (London or Induced-Dipole) Forces – Attraction between temporary or induced dipoles in adjacent molecules – Electrons constantly shift and can sometimes concentrate in one region → instantaneous dipole that goes away once electrons shift again – Chocolate-chip cookie dough analogy – Most common type and weakest of intermolecular forces, found between all types of molecules – Only type of intermolecular force between nonpolar molecules – In general, larger molecules w/ more electrons are more “polarizable” – polarizable = tendency to experience electron shifts that result in charges → the larger the molecule, the stronger the dispersion forces Dispersion (Induced-dipole) Forces Electrons shift in one molecule and concentrate on one side → temporary dipole (light area = +ve; dark area = –ve) The temporary dipole causes electrons to shift in adjacent molecules → another temporary dipole When electrons shift again, the temporary dipoles go away. Dipole-Dipole Forces: Attraction between polar molecules – generally stronger than

dispersion forces because attraction is due to

permanent dipoles rather than temporary dipoles

Note: Van der Waals forces refer to intermolecular forces due to either London dispersion forces or dipole-dipole forces. Hydrogen Bonds:


AB

– Exist between molecules with following bonds: H–F, H–O, H–N – Special type of dipole-dipole force caused by small radii and large electronegativity differences between H and O, N, and F atoms – strongest type of intermolecular force – Responsible for the relatively high melting and boiling points for water, bending

and twisting in proteins, DNA, and other important biological molecules

Note: Hydrogen bonds are the strongest type of intermolecular force but ionic and covalent bonds are stronger than hydrogen bonds.

How to determine type of intermolecular forces involved:

Is the molecule polar or nonpolar

nonpolar

polar

dispersion (London) forces

Are there H–F, H–O, or H–N bonds

hydrogen bonding

dipole-dipole forces

yes

no

Example: For each of the following, identify the type of bond holding atoms

together in the molecules and the type of intermolecular forces between the molecules.

A: _____________________ B: ______________________


BA

A: _____________________ B: ______________________

Ex. 2: Consider the following six choices:

A. ionic bond D. dispersion (induced-dipole) forces B. polar covalent bond E. dipole-dipole forces C. nonpolar covalent bond F. hydrogen bond

Identify the type of bond or intermolecular force described for each below: _____ a. The C–C bond in C2H6. _____ b. The bonds in NaCl. _____ c. The bonds holding HBr molecules together. _____ d. The bonds holding atoms together in a water molecule. _____ e. The bonds holding two NH3 molecules together. _____ f. The bonds holding atoms together in a HF molecule. _____ g. The bonds broken when KBr dissolves in water. _____ h. The bonds formed when KBr dissolves in water. Ex. 3 Circle the molecule in each pair which experiences the stronger

intermolecular forces: a. N2 or NO b. H2S or H2O c. Cl2 or Br2


10.3 SOME PROPERTIES OF LIQUIDS 10.5 EVAPORATION, VAPOR PRESSURE, AND BOILING POINT evaporation: for a liquid to vaporize, the surface molecules must break the

intermolecular forces with other molecules in the liquid – at the boiling point, molecules have enough energy to break the

intermolecular forces with other molecules and become gas

Boiling Point: temperature where vapor pressure of liquid is equal to external pressure (usually atmospheric pressure)

– stronger intermolecular forces → more energy is required to break intermolecular bonds between molecules in the liquid → higher boiling point – normal boiling point is the boiling point at a pressure of 1 atm – e.g. water boils at 100°C at 1 atm, but it boils at ~95°C in

Denver where the atmospheric pressure is ~0.85 atm

Vapor Pressure: partial pressure exerted by gas molecules above the liquid – varies for different liquids, varies for different temperatures

– more gas molecules → higher vapor pressure

– The weaker the intermolecular forces → more molecules can go from liquid to vapor → higher vapor pressure

– In the examples below, liquid A has weaker intermolecular forces than B

Viscosity: resistance of a liquid to flow – for example, honey has high viscosity; water has low viscosity – stronger intermolecular forces = stronger attraction → higher viscosity

A B


Surface Tension: attraction between surface molecules in a liquid – stronger intermolecular forces = stronger attraction → surface molecules are held together more strongly → higher surface tension

Ex. 1: Water molecules experience hydrogen bonding and hexane molecules

experience induced-dipole (or dispersion) forces. Which of the following statements are true/false?

a. Water's intermolecular forces are weaker than hexane's. T F b. Hexane’s vapor pressure is higher than water’s. T F c. Hexane’s boiling point is lower than water’s. T F d. Water’s surface tension is higher than hexane’s. T F e. Water’s viscosity is lower than hexane’s. T F

Ex. 2: Explain each the following in terms of intermolecular forces: a. Why O2's boiling point is -183°C while NO's boiling point is -151°C. b. Why N2's boiling point is -196°C while Br2's boiling point is 59°C. c. Why H2S’s boiling point is –61°C and H2O’s boiling point is 100˚C.


10.6 KINDS OF SOLIDS Crystalline Solids: Have an ordered arrangement extending over a long range – different types of crystalline solids: molecular solids, covalent network solids,

ionic and metallic solids. Molecular Solids: consist of molecules held together by intermolecular forces The Structure and Properties of Ice – Ice is an example of a molecular solid. – The hydrogen bonds between water

molecules are responsible for many unusual properties of ice and water.

The density of ice (d=0.917 g/cm3) is less than the density of liquid water (d=1.00 g/cm3). With all other substances, the solid is more dense than its liquid. – The density of ice (d=0.917 g/cm3) is less

than the density of liquid water (d=1.000 g/cm3) whereas for all other substances, the solid is more dense than its liquid.

– Because of the hydrogen bonds, the arrangement of water molecules in ice crystal has "holes" or empty space. – When ice melts, the water molecules fill in the holes, so liquid water is more dense than ice.

– Note the hexagonal holes in the molecular-level image above. – Snowflakes have hexagonal symmetry because of the hexagonal holes in the molecular-level arrangement of water molecules in ice!


Ionic Crystals: lattice of metal & nonmetal ions – e.g. NaCl, MgO, CaBr2 – 3D network of ions held together by electrostatic attraction → high melting points, hard and brittle – conduct electricity when melted or dissolved in solution

10.10 STRUCTURE OF SOME COVALENT NETWORK SOLIDS

Covalent Network Solids are covalently bonded atoms that form a large network of indefinite size. (a) Graphite is made up of covalently bonded carbon atoms that form layers of

sp2 hybridized carbon atoms. (b) Diamond is made up of covalently bonded carbon atoms that form such a

network of sp3 hybridized carbon atoms in 3D tetrahedral structure. → Diamond is so hard because so many covalent bonds must be broken to

break up the diamond crystal.


Fullerenes are a class of covalently bonded carbon atoms, similar to graphite. (a) C60 has a shape similar to a soccer ball and is often called a Buckyball. (b) Nanotubes consist of sheets of graphite rolled into tubes. – up to ten times as strong as steel

Amorphous Solids: solids lacking 3D arrangement of atoms

Silica (SiO2) makes up sand and quartz glass: optically transparent solid of inorganic materials cooled to a rigid but

non-crystalline arrangement of Si-O bonded atoms called quartz glass


10.7 Probing the Structure of Solids: X-Ray Crystallography

X-ray diffraction: scattering of X-rays by units of a solid crystal – we can construct electron density contour map, where maximum

densities are near center of each atom can be used to determine the positions of nuclei → bond lengths and bond angles in crystals

Metallic Crystals: positive metal ions surrounded by a sea of electrons – electrons are “delocalized”—i.e. free to

move around the entire metal → electrons move freely throughout the

metal resulting in good heat and electrical conductivity

→ electrons act as a glue holding nuclei together, so shape of metal can be easily

manipulated → metals are malleable and ductile – some metals will react with water but are

never soluble in water or other solvents 10.4 Phase Changes: change from one physical state to another

fusionfreezingSOLID LIQUID vaporization

condensationGAS

sublimation

deposition The Equilibrium Nature of Phase Changes dynamic equilibrium: rate of forward process is exactly equal to the rate of

reverse process


Liquid-Solid Equilibrium

freezing: liquid → solid melting (or fusion): solid → liquid melting point: temperature at which solid and liquid phases coexist in equilibrium – normal melting point is melting point at 1 atm

Consider solid-liquid equilibrium of water and ice (at 0°C and 1 atm):

ice water

When ice cubes are placed into a glass of water, the ice cubes begin to melt, but some water between the ice cubes freezes, causing ice cubes to fuse molar heat of fusion (∆Hfus): energy required to melt one mole of solid supercooling: a substance remains liquid even below its freezing point – results when a liquid is cooled so rapidly that molecules don't have time to

arrange themselves properly – unstable condition; stirring or adding "seed" crystal causes solidification

– For example, supercooling can happen in your refrigerator. – A bottle of carbonated soda can be supercooled, so when you pick it up, you give it enough energy to freezes instantly → ice crystals form Liquid-Gas Equilibrium: liquid vapor

evaporation (or vaporization): liquid → vapor condensation: vapor → liquid – the process of a gas liquefying – can result from two ways: 1. cooling sample of gas

→ lower KE, and molecules start to aggregate to form small drops of liquid, eventually causing condensation

2. applying pressure to gas → minimize space between molecules, so molecules are attracted to each other to form droplets, eventually causing condensation


Solid-Gas Equilibrium: Consider the dynamic equilibrium: solid vapor

For example, dry ice, CO2(s), sublimes at room temperature, completely skipping a liquid phase.

sublimation: solid → gas (with no liquid phase) deposition: gas → solid (with no liquid phase) molar heat of sublimation (∆Hsub): – energy required to sublime one mole of solid

Heat of Phase Transition

heating-cooling curve: Shows the phase changes that occur when heat is added or removed from a sample Draw a heating curve indicating the following: 1. Regions for solid only, liquid only, gas only, solid-liquid, liquid-gas 2. The relationship between melting point and phases present 3. The relationship between boiling point and phases present 4. Where the curve is flat, where the slope is positive

Heat Added

Temperature (°C)


10.11 Phase Diagrams – summarize the conditions at which a substance exists as solid, liquid, or gas – allow us to determine melting and boiling points at different external pressures Water’s phase diagram is shown in Fig. 10.28 – graph divided into three regions corresponding to each phase – lines separating two regions indicate conditions when both phases exist – triple point: point at which all three curves meet – when all three phases exist in equilibrium with one another – for water, at 0.0098°C and about 6.0×10-3 atm

Phase Diagram for Water (Fig. 10.28)

Phase Diagram for CO2 (Fig. 10.29)


Phase Diagram for Water versus that for CO2: – Note that line separating solid and liquid phases has positive slope for CO2

but negative slope for H2O – for CO2, the solid is more dense than the liquid, so increasing pressure

converts the liquid to a solid – for H2O, the solid is less dense than the liquid, so increasing pressure

converts the solid to a liquid Critical Temperature and Pressure:

critical temperature (Tc): above which its gas form cannot be made to liquefy, no matter how great the applied pressure – highest temperature at which a substance can exist as a liquid – intermolecular attraction is a finite quantity for given substance – below Tc, molecules are moving slowly enough to maintain contact – above Tc, molecular motion so energetic that molecules will always

break away from attraction critical pressure (Pc): minimum pressure that must be applied to liquefy sample at the critical temperature

Given a Phase Diagram, be able to do the following: – Determine what phase(s) is/are present at a given temperature and pressure – Indicate the melting point or boiling point at a given pressure – e.g. the normal melting or boiling point – Describe what phase change occurs when temperature is changed at a constant pressure – Describe what phase change occurs when pressure is changed at a constant

temperature.

chapter 10: liquids, solids, and phase changesdiscoverchemistry.yolasite.com/resources/solids...

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