Liquids & Solids Advanced Chemistry 30S M. Patenaude GPHS Science Dept A. Three Types of Molecular Motion Gases Liquids Solids Translational Motion Free Hindered None Rotational Motion Free Hindered Hindered Vibrational Motion Free Free Free Molecules in solids and liquids are very close together. The average distance between particles is less than one molecular diameter. Molecules in gases, on the other hand, are very far apart. The average distance between particles is about 10 molecular diameters. B. Intermolecular Forces forces of attraction between molecules or atoms generally much weaker than covalent bonds Only 16 kJ/mol of energy is required to overcome the intermolecular attraction between HCl molecules in the liquid state (i.e. the energy required to vaporize the sample) However, 431 kJ/mol of energy is required to break the covalent bond between the H and Cl atoms in the HCl molecule Thus, when a molecular substance changes states the atoms within the molecule are unchanged The temperature at which a liquid boils reflects the kinetic energy needed to overcome the attractive intermolecular forces (likewise, the temperature at which a solid melts). Thus, the strength of the intermolecular forces determines the physical properties of the substance Types of Intermolecular Forces: Ion-Dipole Interactions – between an ion and a polar molecule – the forces depend on three factors: • The magnitude of the ion’s charge • The magnitude of the dipole • The distance between the ion and the dipole – An example occurs in the hydration of an ion in water. For example, consider the dissolving of table salt in water: NaCl (s) Na+ (aq) + Cl- (aq) – Each ion interacts with water and becomes “hydrated” - surrounded by polar water molecules. – As the ions are surrounded, energy is released. (Exothermic process) Dipole-Dipole Interactions – Forces of attraction between polar molecules. – Weaker than ion-dipole forces – require close proximity of molecules – More polar molecules = stronger dipole- dipole interactions Hydrogen Bonding Hydrogen bonds are considered to be dipole- dipole type interactions A bond between hydrogen and an electronegative atom such as F, O or N is quite polar: The hydrogen atom has no inner core of electrons, so the side of the atom facing away from the bond represents a virtually naked nucleus This positive charge is attracted to the negative charge of an electronegative atom in a nearby molecule this side of the hydrogen atom can get quite close to a neighboring electronegative atom (with a partial negative charge) and interact strongly with it Hydrogen bonds vary from about 4 kJ/mol to 25 kJ/mol (so they are still weaker than typical covalent bonds. But they are stronger than dipole-dipole and or dispersion forces. They are very important in the organization of biological molecules, especially in influencing the structure of proteins Water is unusual in its ability to form an extensive hydrogen bonding network Each water molecule can participate in four hydrogen bonds One with each non-bonding pair of electrons One with each H atom London Dispersion Forces Non-polar molecules would not seem to have any basis for attractive interactions. However, gases of non-polar molecules can be liquefied and solidified, at low temperatures - indicating that if the kinetic energy is reduced, some type of attractive force can predominate. Fritz London (1930) suggested that the motion of electrons within an atom or non-polar molecule can result in a temporary dipole moment As an atom or molecule becomes polarized (it forms an instantaneous dipole), it can induce an opposite dipole in a neighboring atom or molecule. These temporary dipole-dipole attractions form and dissolve over and over again The larger the atom/molecule, the more polarizable it will be - and the greater the London Dispersion Forces! (e.g. Consider the Halogens) C) The Liquid State Properties of Liquids: – very low compressibility – take the shape of their container – high density (comp. to gases) Many properties are related to forces between liquid molecules: • Interior molecules experience attractions all around them • Surface molecules experience attractions from their sides and below As a result of uneven forces, molecules at the surface are pulled inward, creating a “sphere” - a bead, or drop, of liquid COHESIVE forces are the forces between molecules within the liquid – Strong cohesive forces are the result of strong intermolecular forces – They result in high SURFACE TENSION - the resistance of a liquid to an increase in its surface area – They also result in high VISCOSITY - the resistance of a liquid to flow. (Another factor is molecular “complexity”) – Note that viscosity decreases with increased temperature (why?) ADHESIVE forces are the forces between liquid molecules and their container – These forces are strong if the liquid is polar and the surface is polar (or if the liquid is non-polar and the surface is non-polar) – A liquid will WET a surface if there are strong adhesive forces – CAPILLARY ACTION - the spontaneous rising of a liquid up a tube - is explained by strong adhesive forces between the liquid and the tube’s surface, as well as the strong cohesive forces between liquid molecules D) Structure & Bonding in Solids Crystalline solids have a characteristic regular arrangement of particles. – The positions of the particles can be represented by a LATTICE – The smallest repeating unit of a crystal lattice is called a UNIT CELL – We will study three common types of unit cells: simple cubic (sc), body centered cubic (bcc) and face centered cubic (fcc) Simple Cubic Unit Cells 8 atoms at the vertices of a cube Two atoms touch along each edge of the cube: 2r = s where “r” = atomic radius How many atoms and “s” = length of side are inside the cell? Body Centered Cubic Similar to simple cubic…with an extra atom in the center of the cube Atoms are forced apart along the edges 3 atoms touch along the diagonal of the cube: How many atoms are inside the cell? 4r = 3 s Face Centered Cubic 8 atoms at the vertices of a cube, with an additional atom in the center of each face of the cube 3 atoms touch along the diagonal of a face: How many atoms 4r = 2 s are inside the cell? Holes inside Unit Cells A hole formed by 3 atoms in a plane is called a trigonal hole: A hole when a fourth atom is placed on TOP of these three is called a tetrahedral hole: A hole formed by SIX atoms, four in a square plane, with one above and one below, is called an octahedral hole! Holes in the Unit Cells, cont’d Prove to yourself that a FCC cell has 8 tetrahedral holes Prove to yourself that a FCC cell has an octahedral hole in its center Prove to yourself that there is an octahedral hole in the center of each edge of the FCC unit cell What kind of holes are present in BCC unit cells? Closest Packing The first two crystal lattices (sc & bcc) begin with a square layer of atoms There is a more efficient way of packing atoms - beginning with a hexagonal array of atoms Square packed - simple cubic Square packed - body centered Body centered cubic Body Centered Cubic Closest packing In a closest packed crystal, the second layer of atoms will be placed over the holes in the first layer: Notice that there are now TWO types of holes in the blue layer of atoms! Closest Packing If atoms are now placed on top of holes in the second layer, above atoms in the first layer, an arrangement of atoms “A-B-A-B-A…” will result: This arrangement leads to a HEXAGONAL CLOSEST PACKED (HCP) Structure Hexagonal closest packing Be Co Mg Zn If the third layer is placed over top of holes in BOTH the first and second layers, an arrangement of atoms “A-B-C-A-B-C-…” will result: This arrangement leads to a CUBIC CLOSEST PACKED (ccp) Structure, with a Face-Centered Note that both closest packed Cubic Unit Cell! structures have a coordination number of “12” Cubic closest packing Ag Ni Al Pb Au Pt Ca Cu Coordination Numbers Coordination Structure Stacking Pattern Number Simple cubic 6 AAAAAAA Body Centered 8 ABABABAB cubic Hexagonal Closest 12 ABABABAB Packed Cubic Closest 12 ABCABCABC Packed Bonding in Metals The valence electrons in metals are delocalized - they are free to move from atom to atom. This is why metals conduct electricity Metal atoms can be thought of as cations floating in a sea of electrons” - + - + - + - + - + - + - + - + - + - + - The valence electrons are “shared” among all the atoms in the metal - DELOCALIZED COVALENT BONDING Metal Alloys An alloy is a metallic solid made up of a mixture of elements. – Substitutional Alloys • some metal atoms are replaced by other metal atoms in a crystal lattice • Brass (67% Cu, 33% Zn) • Bronze (~93% Cu, 7% Sn) • Sterling Silver (93% Ag, 7% Cu) – Interstitial Alloys • Small atoms fill some of the holes (interstices) between atoms in the crystal lattice • For example, carbon is added to steel (0.2% up to 1.5%) • The added carbon atoms form strong directional covalent bonds with the iron atoms, making the metal stronger Network Solids A network solid is an atomic solid where the atoms are bonded with strong “directional covalent” bonds. The network solid is really one giant covalent “molecule” The two best examples are Diamond and Graphite (two allotropes of carbon) A great resource for viewing solid structures (with Chime): http://learn.chem.vt.edu/archive/apache/htdocs/105a/crystal/crystal.html Graphite A network solid where carbon atoms are covalently bonded together in planes. Each carbon atom is bonded to three other atoms The planes are held together by weak London Dispersion Forces, and slide off easily Diamond Diamond is also a network solid, but in 3-dimensions. Carbon atoms are bonded to four other atoms, in a tetrahedral geometry Because of this 3-d network of strong, directional covalent bonding, diamond is the hardest natural substance Another look at NaCl E. Heating Curves & Changing State If a substance is heated slowly, a plot of temperature vs time can be created. Such a plot is called a heating curve. Plateaus will occur wherever a change of state occurs. The first plateau represents melting (fusion) and the second plateau boiling (vaporization). The energy (heat) required to melt one mole of a substance is called the molar enthalpy of fusion for the substance, DHfus. For water, DHfus is 6.0 kJ/mol. 6.0 kJ of heat must be absorbed to melt one mole of ice. The amount of energy (heat) needed to boil one mole of water is called the molar enthalpy of vaporization of water, DHvap. For water, DHvap is 40.7 kJ/mol. Recall: 1 mol water = 18 g = 18 mL The Heating Curve of Water 140 120 Temperature/°C 100 80 60 40 20 0 -20 0 100 200 300 400 500 600 Time/min A positive enthalpy change, DH, means that energy is being absorbed by the solid or liquid during melting and boiling. When a solid melts, the molecules in the liquid state have more energy than those in the solid state. When a liquid boils, the molecules in the gaseous state have more energy than those in the liquid state. The conversion of a solid directly to a gas (without melting) is called sublimation. The molar enthalpy of sublimation, DHsub, is the energy needed to cause one mole of a solid to sublime. Dry ice (solid CO2), iodine, naphthalene all sublime at room temperature and 1 atm pressure. F. Vapor Pressure A liquid in a closed container will start to evaporate. As more vapor is introduced into the space above the liquid, some of the vapor condenses back to the liquid state. At first, the rate of evaporation is greater than the rate of condensation. The rate of evaporation remains constant, and the rate of condensation increases over time. Eventually, the two rates are equal. At this point, the system is in a state of “equilibrium” or balance. The pressure of the vapor above the liquid at this point of equilibrium is called the vapor pressure of the liquid. A liquid that has a very HIGH vapor pressure is said to be very VOLATILE. This means it will evaporate very easily. Volatility is determined by two main factors: Molar Mass & Intermolecular Forces – Large molar mass = less volatile – Strong intermolecular forces = less volatile Vapor pressure increases with temperature. A sketch of a vapor pressure curve looks like: A Vapor Pressure B Temperature The equation that relates the vapor pressure of a liquid to its temperature: DHvap 1 ln( Pvap ) C R T Where DHvap is the molar enthalpy of vaporization for the liquid. This equation is can be interpreted as a straight line equation: y = m x + b DHvap 1 ln( Pvap ) C R T So a graph of ln(Pvap) vs (1/T) should be linear. The slope of the line would allow the calculation of DHvap. The Clausius-Clapeyron Equation: A more useful form of the first equation that allows the calculation of DHvap if only TWO vapor pressures and temperatures are measured. P1 DHvap 1 1 ln P2 R T2 T1 Normal Melting Point Is defined as the temperature where the solid and liquid states have the same vapor pressure under conditions where the total pressure is 1 atmosphere (standard pressure). Refer to apparatus in Zumdahl (Fig 10.44) Normal Boiling Point Is defined as the temperature where the vapor pressure of a liquid is exactly one atmosphere (standard pressure). When the vapor pressure reaches 1 atm, bubbles of vapor are able to form inside the liquid (which is what we see when the liquid “boils”) Supercooling occurs when a liquid remains “liquid” when it is slowly cooled below its freezing point. As the water cools, its molecules must become “organized” into the solid crystal shape. The temperature may drop below 0°C before this organization occurs. As the crystal forms, heat is released and the temperature rises back to 0°C as the liquid freezes. Superheating occurs if a liquid’s temperature rises above its boiling point before boiling starts. This may occur if the liquid is heated very rapidly. To form vapor bubbles within the liquid, many high-energy molecules must congregate. This may not happen immediately if the liquid is heated very quickly. Once enough molecules DO congregate, a bubble will form. Because the temperature is higher than the boiling point, the pressure of vapor will be higher than 1 atm. This may cause the bubbles of vapor to “burst” as the liquid starts to boil. (Called “bumping”) Boiling chips are added to a liquid as it is heated. The boiling chips contain trapped air that is released as the temperature rises. This creates small bubbles of gas in the liquid that act as “starters” for vapor bubbles to form. This prevents superheating from occurring. Phase Diagrams label axes label phase regions label: triple point critical point melting point boiling point sublimation point Critical Point The temperature and pressure at which the liquid and gaseous phases of a pure stable substance become identical. The critical temperature of a gas is the maximum temperature at which the gas can be liquefied; the critical pressure is the pressure necessary to liquefy the gas at the critical temperature.