Ap Chemistry Chapter 10 Liquids and Solids

Ap Chemistry Chapter 10 Liquids and Solids

Liquids and solids are unique from gases in that they have very similar compressibility, density and behavior of particles. In this chapter, we will study how liquids and solids behave. AP CHEMISTRY CHAPTER 10 LIQUIDS AND SOLIDS 10.1 Intermolecular Forces Molecules interact in many ways. Forces inside of molecules are different than those outside. Intramolecular bonding- bonding that

occurs inside of covalently bonded compounds. Intermolecular bonding- interactions between particles (atoms, molecules or ions) Intermolecular forces act on the condensed states of matter (liquids and solids), but less often on gases Changes in state are due to changes in intermolecular bonding, not

intramolecular bonding. Dipole-Dipole Interactions Dipole-dipole attraction- attraction of molecules having dipole moments for each other. (negative-positive) About 1% as strong as covalent or ionic bonds. Molecules

orient themselves to Hydrogen Bonding A particularly strong form of dipole-dipole attraction is hydrogen bonding- unusually strong dipole-dipole attractions involving hydrogen atoms which are covalently bonded to a very electronegative atom (F,O,N) with unshared electrons. There are two reasons for strength of hydrogen bonds:

1. Small size of H atom allows closeness 2. Large variation in polarity Substances with large amounts of H-bonding have high boiling points compared to similar substances. Ex. H2O, NH3, HF The boiling point of covalent hydrides increases with molecular weight in Group 4. In other groups, the first hydride has a high boiling point because of

hydrogen bonding. London Dispersion Forces (LDFs) London dispersion forces (LDFs) -relatively weak forces (usually) that exist between noble gas atoms and nonpolar molecules. LDFs also exist in compounds that have dipole-dipole and/or hydrogen bonding. LDFs may be the most important force in large molecules of these types. LDFs occur because of

momentary electron imbalance (temporary dipole) which can induce the same in adjacent molecules. Strength of LDFs This force is often very weak, thus the low freezing point of noble gases. The freezing point of noble gases gets higher as we go down the group because heavier atoms have more electrons and are thus more polarizable because they can develop temporary dipoles. This causes London dispersion forces to increase down a group. When writing AP essays, be sure to note the number of electrons that an atom or molecule has. Refer to a molecules polarizability!

Boiling points reflect LDFs in nonpolar covalent compounds. The boiling point nonpolar covalent compounds goes up as the size of the molecule increases. More electrons means that the molecule is more polarizable, thus higher LDFs and more energy required for the state change. LDF < Dipole-dipole < H-bonding < Metallic Bonding < Ionic Bonding

General trends in strength of attractionMemorize These! Increasing Strength LDF dipole-dipole H-bonding metallic bonding ionic bonding covalent network solid 1. Dipole-dipole attractionattraction of molecules having dipole moments for each other. (negative-positive) -about 1% as strong as covalent or ionic bonds. -molecules orient themselves to

minimize repulsion and maximize attractions Practice! Determine the type(s) of intermolecular force each of the follow substances would experience: H2 LDF-very weak HCl Dipole-dipole, LDF Fe Metallic, LDF

H 2S Dipole-dipole, LDF SiC Covalent Network CO2 LDF Rank these substances from lowest boiling point to highest boiling point. Justify your answers. NaCl, SiO2, NH3, H2O, F2, SO2 F2SO2 NH3 H2ONaCl Ge

TAMUs Summer Research Opportunities for High School Students TAMUs Summer Research Opportunities for High School Students Hosted by Drs Umesh Bageshwar and S. M. Musser Many faculty who work in different areas of research are interested in taking SERIOUS high school students to work in laboratory setting over the summer as research assistants.

Points to Note: The most appropriates students are high school juniors and seniors. Research areas include the following, but may be expanded due to student interest. Protein transport across cell membranes Aniogenesis

Cancer Protein stability Antibody engineering Students will work throughout the summer full time (June 1-August 26, when school opens). They should be fully committed to the lab schedule and they themselves are responsible for their transport.

Each lab will provide a choice of projects. Most projects are designed to the skills and understanding of high school students. They will be 100% supervised. Projects are usually designed by a researcher. There is a questionnaire for interested students. This will help in shortlisting students and speed along the process of determining how many CSHS students will be able to participate. We will need a lits of interested students by January. The aim is to begin the project the same day summer classes start. 10.2 The Liquid State Liquids exhibit certain

traits that are related to their intermolecular attractions. Surface tensionresistance of a liquid to an increase in its surface area. Liquids with relatively large intermolecular forces have high surface tensions. Ex. H2O. Polar

molecules have more surface tension than nonpolar molecules. Capillary action- spontaneous rising of a polar liquid in a narrow tube. Caused by 2 forces: Cohesive forces: intermolecular forces among the liquid molecules Adhesive

forces: forces between the liquid molecules and their container. The container must be made of polar material such as glass. A concave meniscus forms because the waters adhesive forces toward the glass are stronger than its cohesive forces. A nonpolar liquid can produce a convex meniscus (cohesive > adhesive) Viscosity- resistance of a liquid to flow.

Liquids with large intermolecular forces tend to be highly viscous. Ex. glycerol (AKA glycerine) More complex molecules are more viscous because they tangle up. (gasoline vs. grease) 10.3 An Introduction to Structures and Types of Solids Amorphous solids- very disordered, usually long chain-like molecules twisted up like spaghetti. (plastics, asphalt, rubber)

Crystalline solids- highly regular arrangement of components lattice- a 3-D system of points designating the centers of the components (atoms, ions or molecules) unit cell - smallest repeating unit of the lattice. (Learn the 3 common cubic unit cells pg 446!)

coordination number- number of nearest neighbors surrounding a particle in a crystal Crystal Lattice Patterns X-ray diffraction- method of determining crystal structure. -X-rays of a single wavelength are directed at a crystal and are scattered by it, producing a diffraction pattern which can be used to determine the crystal structure. Types of Crystal Arrangements simple cubic or primitive 1 net sphere (8 -1/8 spheres) The

coordination number of each particle is six. Body-centered cubic 2 net spheres (1 complete and 8 1/8 spheres) The coordination number is eight. Face-centered cubic 4 net spheres (6 -1/2 spheres and 8 -1/8 spheres). The coordination number is twelve. NaCl and other alkali halides are facecentered cubic (fcc). Types of crystalline solids

Ionic solids (NaCl)have ions at lattice points, held together by strong ionic forces. Molecular solids (sucrose)- have molecules at lattice points, held together by LDF, dipoledipole, &/or hydrogen bonding (Ice has H2O molecules at each point of the lattice) Types of crystalline

solids Metallic solids (gold)- have a metal atom at lattice points, held together by metallic bonds. Atomic solids (argon)- have a noble gas atom at lattice points, held together by LDF Types of crystalline

solids Covalent network solid (diamond and silicon compounds) (essentially one giant molecule) covalently bonded , have an atom at each lattice point, held together by very strong covalent bonds. The properties of a solid depend on the nature of the forces that hold the solid together. Diamond has very strong, covalent network forces.

10.4 Structure and Bonding in Metals Metallic crystals Have nondirectional covalent bonding that leads to properties like conductivity, malleability and ductility Spherical atoms packed together and bonded to each other equally-called metallic bonding

Types of bonding include closest-packed (with hexagonal or face centered cubic unit cells) and body centered cubic unit cells Bonding in Metals Strong and nondirectional which means they are difficult to separate, but easy to move. Two models are used to describe metallic bonding. They are: Electron sea model- This is the simplest model. The metal cations are in a sea of valence

electrons. The mobile electrons can conduct heat and electricity. The cations can be easily moved around (lead to properties of ductility and malleability). The Band Model Band model (MO model)- This is a more detailed model. The electrons are assumed to travel around the metal crystal in molecular orbitals formed from the valence atomic orbitals of the metal ions. The MOs that result are very closely spaced in energy levels, thus they form a continuous band. Some MOs are empty. Mobile electrons are excited into these empty M0s. The

Metal alloy - substance that contains a mixture of elements and has metallic properties. There are two types of metal alloys. Substitutional alloy- some of the host metal atoms are replaced by other metal atoms of similar size. Ex. brass, sterling silver, pewter Interstitial

alloy -holes in closest packed metal structure are filled by small atoms. Ex. steel (carbon in iron) 10.5 Carbon and Silicon: Network Atomic Solids Network solids- atomic solids with strong directional covalent bonds

Brittle and dont conduct heat or electricity Ex. C and Si Allotropes- different forms of the same element Ex. diamond, graphite, and buckminsterfullerene are allotropes of carbon. Sulfur has several allotropes Strongest type of bonding

Allotropes of Carbon Bonding of Carbon Diamond has a tetrahedral sp3 arrangement. Its MOs are far apart in energy, thus no conduction of electricity. Graphite has layers of six member rings (sp2). The pi molecular orbitals allow it to conduct electricity. It has strong bonding within its

layers. We can convert graphite to diamond at a pressure of 150,000 atm and a temperature of 2800oC. Partial Representation of the Molecular Orbital Energies in A) Diamond and B) a Typical Metal Silicon Compounds Silicon compounds are to geology as carbon compounds are to biology. Most important Si compounds contain Si and O. Silica, SiO2 (empirical formula) is the fundamental Si-O compound. Quartz and some types of sand are silica. Silica is a network of SiO4 tetrahedra. The bonds are single because the 3p orbitals are too big to bond strongly

with oxygen. Silicates are anions of Si and O. Ex. SiO44-, Si2O76-, Si3O96 Glass is formed when silica is melted and cooled rapidly. Other substances are added to glass to vary its properties. Glass is amorphous and can be called a supercooled liquid. Ceramics contain tiny crystals of silica in a glassy cement. Semiconductors Semiconductors- Elemental silicon has the same structure as diamond but the

energy gap between filled and empty MOs is much smaller. Thus, a few electrons can cross this gap at room temperature, making Si a semiconductor. At higher temperatures, where more energy is available to excite electrons into the conduction bands, the conductivity of Si increases. Most metals have decreased conductivity at higher temperatures. Doping Doping is a process which increases the conductivity of

silicon. A very few of the Si atoms are replaced by other atoms such as arsenic. Arsenic has one more valence electron than carbon. This gives more electrons for conduction and produces an n-type semiconductor (n stands for negative). If we instead replace a few Si atoms with atoms with one fewer electron (such as boron), we produce holes that electrons can travel into (producing a chain effect). This produces a p-type

Molecular Solids Dry ice has CO2 molecules at each point in the lattice. Molecular solids have molecules at lattice points instead of atoms or ions. Ex. ice, dry ice, sulfur (S8), phosphorus (P4) Strong covalent bonding within molecules but relatively weak forces between molecules.

The atoms within the molecule are closer to each other than atoms from adjacent molecules are. This indicates stronger bonding. CO , S8, I2, and P4 have no dipole-dipole forces. The last three are solids at room temperature because London dispersion forces are stronger in larger molecules. 2 Ionic solids have ions at lattice points. Stable, high melting point, held together by strong electrostatic forces

Structure of most binary ionic solids can be explained by the closest packing of spheres The smaller ions (usually cations) fit into the holes between the larger ions (anions). Attractions are maximized and repulsions are minimized. Atomic

Type of Solid Molecular Ionic Atom Molecule Ion Directional covalent bonding Nondirectional covalent

bonding; electrons are delocalized throughout the crystal LDF Polar molecules have dipoledipole interactions Nonpolar molecules have LDF Ionic bonding

Properties Hard with high mp, good insulators Range of hardness and mp, good conductors, malleable, ductile Low mp but increases with LDF Soft with low

mp, good insulators Hard with high mp and good insulator Ex. Diamond, Si cmpds Fe, Ag, Brass Ar Ice, CO2(s)

NaCl, CaF2 Structural Unit (at lattice point) Type of Bonding Covalent Network Metallic Group 8A Atom

Atom Vapor Pressure and Changes in State We will now consider how the properties of matter we have discussed thus far affect state changes in that same matter. Vaporization (evaporation)- endothermic process because we add energy to break intermolecular bonding. Heat (enthalpy) of vaporization- energy required to vaporize 1 mole of a liquid at a pressure of 1 atm

(Hvap) Evaporation is a cooling process and is the opposite process of condensation (uses the same value of H) Condensation- process by which vapor molecules reform a liquid dynamic equilibrium occurs when the rate of condensation = rate of vaporization vapor

pressure- pressure of the vapor present at equilibrium measured Pvapor by a barometer = Patmosphere - PHg volatile column -liquids which evaporate rapidly and have high vapor pressure

Water has most H-bonding. Ethanol has medium Hbonding and LDF Diethyl ether has no Hbonding but has LDF Vapor pressure is affected by two factors: Molecular weight- At a given temperature, heavy molecules have lower velocities than light molecules and thus have a lower tendency to

escape from the liquid surface. A liquid with a high molecular weight tends to have a small vapor pressure. Intermolecular forces- molecules with strong intermolecular forces also tend to have low vapor pressure because they need lots of energy to escape. Vapor pressure increases with temperature (higher KE) but it is not a direct relationship. You do not need to be able to solve VP-temp problems.

Below 0oC, The VP of ice is less than the VP of liquid water. Sublimation -direct change of a solid to a gas Ex. dry ice, freeze drying, iodine Heating curve - a plot of temperature versus time for a process where energy is added at a constant rate. 2.0J/gC 40.7 kJ/mol 4.18J/gC

6.02 kJ/mol 2.1J/gC Temperature remains constant during a phase change. Heat (enthalpy) of fusion- enthalpy change that occurs when a solid melts, Hfus This is an endothermic process as well

The Hfus of water is 6.02kJ/ mol Normal melting point- the temperature at which the solid and liquid have the same vapor pressure at 1 atm total pressure. Boiling -occurs when VP of the liquid = external pressure

Normal boiling point - the Ex. How many joules are needed to convert 5.0 g of ice at -15oC to steam at o 130 Phase C? changes and temp changes q = 5.0g (2.1J/goC)15oC = 160J = 0.16kJ 5.0g ice 1 mol ice 6.01 kJ 18.0 g 1 mol ice = 1.7 kJ q = 5.0g (4.18J/goC)100oC = 2100J = 2.1kJ

5.0g water 1 mol water 40.7 kJ 18.0 g water 1 mol water = 11 kJ q = 5.0g (2.0J/goC)30oC = 300J = 0.30kJ q = 0.16 + 1.7 + 2.1 + 11 + 0.30 = 15 kJ

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