chapter 14: liquids and solids - seattle...

CHEM 139: Zumdahl Chapter 14 page 1 of 16

CHAPTER 14: LIQUIDS AND SOLIDS Active Learning Problems: 1-13, 15-19, 21-22

End-of-Chapter Problems: 1-3, 5-11, 19, 21-37, 39-61, 63-65, 69, 71, 74-90

PROPERTIES OF LIQUIDS AND SOLIDS Properties of Liquids 1. Liquids have an indefinite shape. → They take the shape of their container.

2. Liquids have a fixed volume. → They cannot be compressed; nor can they be

expanded very much.

3. Liquids are much denser compared to gases. – Liquids are about 1000 times denser than gases.

4. Liquids usually flow readily, so liquids that mix will eventually form a homogeneous mixture.

– This is because molecules in a liquid are constantly in motion at the molecular level.

Properties of Solids 1. Solids have definite shape. – Particles are fixed in place and vibrate but do not move.

2. Solids have a fixed volume. → Like liquids, they cannot be compressed; nor can they be expanded very much.

3. Solids are either crystalline or noncrystalline (amorphous). – A crystalline solid contains particles in a regular, repeating pattern. – A noncrystalline or amorphous solid is disordered.

4. Like liquids, solids have much greater densities than gases.

5. Solids do not mix by diffusion because particles in a solid do not move and mix. 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 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


14.1 WATER AND ITS PHASE CHANGES

phase change: change in physical state

Water is virtually everywhere, covering 70% of the earth’s surface. • About 97% of the earth’s water is in the oceans, which help regulate the earth’s temperature. • Water continually vaporizes from lakes, rivers, and oceans to become water vapor. • At high concentrations water vapor becomes saturated and when the temperature drops,

gaseous water molecules move more slowly and become attracted to one another. → The water molecules nucleate (form clusters and water droplets that become clouds).

→ When the water droplets become large enough, gravity pulls them towards the Earth’s surface, which results in rain. • At even lower temperatures, liquid water loses more kinetic energy until the molecules move

so slowly they cannot overcome the attractive forces with nearby molecules → The liquid molecules are fixed in place and become ice.

condensed states: solid and liquid phases – In solids and liquids, molecules are in close contact and attracted to one another.

→ These attractive forces holding particles together are called intermolecular forces.


HEATING/COOLING CURVES Consider the changes that H2O undergoes when a block of ice is taken from a freezer and heated in a pan until it is completed converted into steam. – A heating/cooling curve shows the phase changes with temperature and heat added to or

removed from any system. Draw a heating curve indicating the following: 1. Regions for solid only, liquid only, gas only, solid-liquid, liquid-gas 2. Know the relationship between melting point and the phases present. 3. Know the relationship between boiling point and the phases present. 4. Know where the slope is zero, where the slope is positive, and why. Example: Assume this is a heating-cooling curve for water. Indicate the initial and final

conditions described below on the heating-cooling curve. a. A sample of water is initially at 25°C. Indicate this on the curve with an xi. b. The sample is heated until it is at 110°C. Indicate this on the curve with an xf. 14.2 ENERGY REQUIREMENTS FOR THE CHANGES OF STATE – To undergo a change in physical state (e.g. solid → liquid or liquid → gas), particles in

solids or liquids must overcome the intermolecular forces with surrounding molecules. – Thus, for a liquid or solid to become a gas, the thermal energy (associated with motion) of

particles in the liquid/solid must be great enough to overcome the attractive forces with surrounding particles. The attractive forces are ionic bonds, covalent bonds, metallic bonds, or intermolecular forces.

Heat Added

Tem

pera

ture

(°

C)


14.3 INTERMOLECULAR FORCES (IMF’s) Dispersion (or London) Forces (also called Induced-Dipole Forces) – In nonpolar molecules (shown as green) the electrons can shift and concentrate on one end → temporary dipole (red = − end; blue = + end)

– The partial positive charge caused by the temporary dipole in one molecule causes the electrons in an adjacent molecule to concentrate around the atom closest to the first molecule. → A temporary dipole results in the second molecule. → The temporary dipoles in both molecules cause them to be attracted to one another. – But that attraction lasts only until the electrons shift again, and the temporary dipoles go away.

– Every molecule experiences London forces. – They are the ONLY type of IMFs between

nonpolar molecules

– Polar molecules also experience London forces they usually have other types of IMFs that are stronger than London forces. The strength of London forces is determined by the number of electrons. – The bigger the molecule → more electrons

→ the greater its polarizability (ability to distort electron clouds to get a temporary dipole) → the stronger its London/dispersion forces

– Thus, the more electrons in an atom or molecule → The stronger the London/Dispersion forces Dipole-Dipole Forces: Attraction between polar molecules – generally stronger than dispersion forces because attraction is due to permanent dipoles. – The permanent dipoles mean the molecules are always attracted to nearby molecules. → Dipole-dipole forces are stronger than London forces for molecules of similar size (or number of electrons).


The Hydrogen Bond: – Especially strong type of dipole-dipole force – Exist between molecules with the following bonds: H–F, H–O, H–N – because these are small atoms with large electronegativity differences → very strong dipole in molecules – Strongest type of intermolecular force – Responsible for the relatively high melting and boiling point for water compared to molecules of

similar size.

Hydrogen bonds are also responsible for the bending and twisting in proteins, DNA, and other important biological molecules.

Image from http://blog.targethealth.com/?p=6846

Note: Hydrogen bonds are the strongest type of intermolecular forces between different molecules, BUT ionic and covalent bonds (holding ions or atoms together in compounds) are stronger than hydrogen bonds!


Ion-Dipole Forces – Attraction between an ion and the oppositely charged end of a polar molecule – e.g., between Na+ and the negative end (O atom) of a H2O molecule or between Cl– and the positive end (H atoms) of a H2O molecule.

– Note that when an ionic compound like NaCl dissolves in water, the formation of ion-dipole

forces between the Na+ (or Cl−) ions with water molecules results in the ionic bonds breaking. How to determine type of intermolecular forces involved: Ex. 1 Indicate the type(s) of intermolecular forces for each molecule below then circle the

molecule in each pair that experiences the stronger intermolecular forces. a. N2 or NO c. Cl2 or Br2 b. H2S or H2O d. PH3 or CH4

Is the molecule polar

or nonpolar?

polar

nonpolar

London Forces

Any H-F, H-O, or H-N bonds in the molecule?

no

yes

dipole-dipole forces and London Forces

hydrogen bonds and London Forces


Ex. 2: For each of the following, i. Identify the bond between atoms as ionic, polar covalent, or nonpolar covalent. ii. Identify the intermolecular forces as London dispersion forces, dipole-dipole

forces, and/or hydrogen bonding. Water (H2O)

A: _____________________ B: ______________________

Oxygen (O2)

A: _____________________ B: ______________________

HCN (H=white, C=charcoal, N=blue)

A: _____________________ B: ______________________

A B


Ex. 3: Indicate the bond or intermolecular forces described for each below: A. ionic bond D. London (dispersion) forces G. metallic bond B. polar covalent bond E. dipole-dipole forces H. ion-dipole forces C. nonpolar covalent bond F. hydrogen bond _________ i. The bonds holding the atoms together in a HF molecule _________ ii. The bonds holding two H2S molecules together in liquid H2S _________ iii. The bonds holding two Br2 molecules together in liquid Br2 _________ iv. The bonds holding the atoms together in a chlorine molecule _________ v. The bonds holding the CO2 molecules together in dry ice, solid CO2 _________ vi. The bonds in a sample of CuO _________ vii. The bonds holding atoms together in magnesium _________ viii. The bonds holding atoms together in water _________ ix. The bonds broken when NaCl dissolves in water _________ x. The bonds made when NaCl dissolves in water


14.4 EVAPORATION AND VAPOR PRESSURE vaporization: liquid → gas – From a molecular viewpoint, molecules “escapes” from the liquid to the gaseous state.

As the liquid evaporates, more molecules go into the gas phase. → vapor: The gas above a liquid when the liquid and gaseous states are both present

vaporization: liquid + heat → vapor condensation: vapor → liquid + heat

Liquid-Gas Equilibrium: liquid + heat ⎯⎯⎯⎯⎯ ⎯←⎯⎯⎯⎯⎯ →⎯

oncondensati

onvaporizati vapor

When the molecules in the liquid have enough energy, they escape to the gas phase. – In a closed system, when enough vapor exists above the liquid, some gaseous molecules condense back to the liquid. – Ultimately, the rate of vaporization = the rate of condensation. → The system has reached a state of dynamic equilibrium in which the forward

process occurs at the same rate as the reverse process. In an open system, molecules in the liquid have enough energy to escape to the gas phase and continue to escape in a process called evaporation. – The vaporized molecules continue to escape → little or no condensation occurs. → Ultimately, all of the liquid is converted into a gas. – Since vaporization requires energy, the liquid molecules take energy from the

surroundings, so the temperature of the surroundings often decreases. → Why evaporation is a “cooling process” and is used to reduce body temperature


Activity: Cap the end of the syringe with your finger to make it airtight then pull the plunger out at least half way. Release the plunger.

Record your observations: Explain why this occurs. 13.1 PRESSURE gas pressure: force per unit area exerted by gas molecules colliding against the inside walls of their container – a measure of how often gas particles hit the container walls In the 1600’s, Evangelista Torricelli published the first explanation for a vacuum. – Ancient Greeks observed that a wine barrel empties slowly if only one hole is drilled. – Torricelli explained that a “sea of air” surrounding the Earth slows the flow of wine out of the barrel. – If you create a second hole on the top of the barrel, air molecules can rush in and push the liquid out. → The wine flows out faster. vacuum: empty space with no gas molecules present → gas pressure equals zero – Don’t think of a vacuum cleaner because there’s no

suction, just empty space! atmospheric pressure: – pressure exerted by air molecules colliding with surfaces in the environment → Why anything entering the Earth's atmosphere burns up! – At sea level, atmospheric pressure can hold up a column of

mercury about 760 mm in height = 760 mmHg. – decreases as altitude increases – At higher altitudes like Denver and Mt. Everest, air becomes thinner, so atmospheric pressure is lower. → Less O2 present makes it harder to breathe. barometer: instrument invented by Torricelli to measure atmospheric pressure


standard atmospheric pressure: a column of mercury measuring 760 mmHg.

1 atm ≡ 760 torr ≡ 760 mm Hg = 14.7 psi (approx.) Ex. 1: The tire pressure for tires used on most automobiles is about 32 psi. Express this

pressure in units of atm, torr, and mmHg. Vapor Pressure (v.p.): pressure exerted by gas molecules above a liquid – For a molecule to go from liquid to gas, it has to break the intermolecular forces with other liquid molecules around it → Weaker the intermolecular forces are easier to break → more gas molecules → higher vapor pressure → Stronger the intermolecular forces are harder to break → fewer molecules go from liquid to gas → lower vapor pressure.

Example: Consider the two closed systems above. In which container, (a) or (b), does the

liquid have stronger intermolecular forces? Explain why.


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

For a liquid to boil, its vapor pressure must equal the atmospheric pressure. – Note that atmospheric pressure is the pressure exerted by air molecules (i.e., the gas

molecules in the atmosphere). → Since intermolecular forces influence vapor pressure, they also influence boiling point.

→ Weaker the intermolecular forces → more gas molecules → higher vapor pressure → Less energy is needed to get vapor P = atmospheric P → lower boiling point

→ Stronger the intermolecular forces → fewer gas molecules → lower vapor pressure → More energy is needed to get vapor pressure = atmospheric pressure

→ higher boiling point – normal boiling point (b.p.) is the b.p. when atmospheric pressure is 1 atm (760 mmHg) – e.g. Water boils at 100°C at sea level (where atmospheric pressure is ~1 atm) but at ~95°C in Denver where atmospheric pressure is much lower (~.85 atm). Ex. 1: If hexane (C6H14) molecules are nonpolar, fill in the blanks for the following statements

about water and hexane:

a. Hexane’s intermolecular forces are _______________________________. b. Water’s intermolecular forces are _______________________________. c. Water's intermolecular forces are _________ than hexane's. stronger weaker d. Water has a __________ vapor pressure compared with hexane. higher lower e. Water has a ____________ boiling point than hexane. higher lower

Ex. 2: Circle the compound in each pair below with the higher boiling point, and explain why. a. CO or O2 b. HF or HCl c. Cl2 or F2 Ex. 3: Which of the following compounds has the higher boiling point: CH2O or COF2? Explain why.


14.5 THE SOLID STATE: TYPES OF SOLIDS (MOLECULAR , IONIC, AND ATOMIC) Solids can be crystalline or amorphous. Crystalline Solids: Have an ordered arrangement extending over a long range – different types of crystalline solids: molecular, ionic, covalent network, and metallic.

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.

– Because of hydrogen bonds, the arrangement of H2O molecules in ice leaves "holes" or empty space. – Note the hexagonal holes in the molecular-level image for ice at the right. → Snowflakes have hexagonal symmetry because of

the hexagonal holes in the molecular-level arrangement/structure of water molecules in ice!


The density of ice (d=0.917 g/cm3) is lower than the density of liquid water (d=1.00 g/cm3). With all other substances, the solid is more dense than its liquid. – When ice melts, the H2O molecules fill in the holes, so liquid H2O is denser than ice.

ATOMIC SOLIDS: consist of metal or nonmetal atoms, where nonmetal atoms in a covalent network solid are held together by covalent bonds, and metal atoms are held together by metallic bonds. Covalent Network Solids: covalently bonded atoms forming a large network of indefinite size. (a) Graphite consists of covalently bonded carbon atoms that form layers of carbon atoms. (b) Diamond consists of covalently bonded carbon atoms that form such a network of carbon

atoms in 3D tetrahedral structure. → Diamond is so hard because the covalent bonds between the carbon atoms must be

broken to cleave or melt the diamond crystal.


Metallic Solids: exist as positive ions surrounded by a “sea of electrons” – electrons move freely throughout the metal → metals having high electrical conductivity → metals conduct heat, which is why most pots

and pans are made out of metal – electrons can act as a “glue” to hold the metal

atoms together → metals are malleable (can hammered into a

thin foil) and ductible (can stretched into a fine wire)

IONIC CRYSTALS: lattice of metal & nonmetal ions – e.g. NaCl, Al2O3, CaBr2 – 3D network of ions held together by electrostatic attraction = ionic bonds → high melting points, hard and brittle – conduct electricity only when melted or dissolved in solution

AMORPHOUS: solids lacking 3D arrangement of atoms – similar in appearance to liquids since they lack the ordered structure of crystalline solids

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

silica (SiO2) glass


Ex. 1: Indicate the type of bonds or intermolecular forces that must be broken to melt the following substances: a. ice b. diamond c. sodium chloride (NaCl) d. aluminum Ex. 2: If sodium chloride’s melting point is 801˚C while diamond’s melting point is about

3550˚C, compare the relative strength of the bonds in sodium chloride with the bonds in diamond.

Ex. 3: Copper’s melting point is 1083˚C while NaCl’s melting point is about 801˚C. a. Compare the relative strength of the bonds in copper with the bonds in NaCl. b. Can you assume metallic bonds are always stronger than ionic bonds? Yes No Explain. Ex. 4: Rank the following substances in terms of increasing melting point: ice, diamond, copper, and table salt (NaCl).

_______________ < _______________ <_______________ <_______________ lowest m.p. highest m.p.

chapter 14: liquids and solids - seattle...

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