Transcript Slide 1
Liquids and Solids Properties of Liquids • Unlike gases, liquids do not respond dramatically to temperature and pressure changes. • We can study the liquid state and make 5 general observations. 1. Liquids have a variable shape, but a fixed volume. • Liquids take the shape of their container 2. Liquids usually flow readily. • However, not all liquids flow at the same rate. Properties of Liquids Continued 3. Liquids do not compress or expand significantly • The volume of a liquid varies very little as the temperature and pressure change. 4. Liquids have a high density compared to gases. • Liquids are about 1000 times more dense than gases. 5. Liquids that are soluble mix homogeneously. • Liquids diffuse more slowly than gases but eventually will form a homogeneous mixture. Intermolecular Bond Concept • An intermolecular bond is an attraction between molecules, whereas intra molecular bonds are between atoms in a molecule. • Some properties of liquids, such as vapor pressure, viscosity, and surface tension, are determined by the strength of attraction between molecules. • Intermolecular bonds are much weaker than intramolecular bonds. Intermolecular Bonds • Recall, that a polar molecule has positive and negative charges concentrated in different regions due to unequal sharing of electrons in bonds. • This uneven distribution of electrons in a molecule is called a dipole. • Intermolecular attractions result from temporary or permanent dipoles in molecules. • There are three intermolecular forces: dispersion forces, dipole forces, and hydrogen bonds. Dispersion Forces • Dispersion forces, or London forces, are the result of a temporary dipole. • Electrons are constantly shifting and a region may become temporarily electron poor and slightly positive while another region becomes slightly negative. • This creates a temporary dipole and two molecules with temporary dipoles are attracted to each other. Dispersion Forces Continued • Dispersion forces are the weakest intermolecular force. • Dispersion forces are present in all molecules. • The strength of the dispersion forces in a molecule is related to the number of electrons in the molecule: – The more electrons in a molecule, the stronger the dispersion forces. Dipole Forces • Polar molecules have a permanent dipole. • The oppositely charged ends of polar molecules are attracted to each other, this is the dipole force. • The strength of a dipole force is typically 10% of a covalent bond’s strength. • Dipole forces are stronger than dispersion forces. Hydrogen Bonds • Hydrogen bonds are a special type of dipole attraction. • Hydrogen bonds are present when a molecule has an N-H, O-H, or F-H bond. • Hydrogen bonds are especially important in water and living organisms. Physical Properties of Liquids • There are 4 physical properties of liquids that we can relate to the intermolecular attractions present in the molecules: – Vapor Pressure – Boiling Point – Viscosity – Surface tension Vapor Pressure • At the surface of a liquid, some molecules gain enough energy to escape the intermolecular attractions of neighboring molecules and enter the vapor state. This is evaporation. • The reverse process is condensation. • When the rates of evaporation and condensation are equal, the pressure exerted by the gas molecules above a liquid is called the vapor pressure. Vapor Pressure Continued • The stronger the intermolecular forces between the molecules in the liquid, the less molecules escape into the gas phase. • As the attractive force between molecules increases, vapor pressure decreases. Vapor Pressure Comparison • Lets compare water and ether. – Water has strong intermolecular attractions and ether has only weak intermolecular attractions. • At 0C, neither has a significant vapor pressure. • At 35C, ether has a significant vapor pressure and water does not. Vapor Pressure vs. Temperature • As the temperature increases, the vapor pressure of a liquid increases. • Again, the stronger the intermolecular attractions, the lower the vapor pressure at a given temperature. Boiling Point • The normal boiling point of a substance is the temperature where the vapor pressure is equal to the standard atmospheric pressure. • As we saw in the previous graph, the stronger the intermolecular attractions, the higher the boiling point of the liquid. • A liquid with a high boiling point has a low vapor pressure. Viscosity • The viscosity of a liquid is a liquid’s resistance to flow. • Viscosity is the result of an attraction between molecules. • The stronger the intermolecular forces, the higher the viscosity. Surface Tension • The attraction between molecules at the surface of a liquid is called surface tension. • For an object to sink in a liquid, it must first break through the surface. • The stronger the intermolecular attractions, the stronger the surface tension of a liquid. Properties of Solids • Unlike gases, solids do not respond dramatically to temperature and pressure changes. • We can study the solid state and make 5 general observations. 1. Solids have a fixed shape and volume. • Unlike liquids, solids are rigid. 2. Solids are either crystalline or noncrystalline. • Crystalline solids contain particles in a regular, repeating pattern. Properties of Solids Continued 3. Solids do not compress or expand to any degree • Assuming no change in physical state, temperature and pressure have a negligible effect on the volume of a solid. 4. Solids have a slightly higher density than their corresponding liquid • One important exception is water; ice is less dense than liquid water. 5. Solids do not mix by diffusion • The particles are not free to diffuse in a solid heterogeneous mixture. Crystalline Solids • There are three types of crystalline solids examples of which are shown below: – Ionic solids like NaCl, (a) – Molecular solids like H2O, (b) – Metallic solids like Cu, (c) Ionic Solids • A crystalline ionic solid is composed of positive and negative ions arranged in a regular, repeating pattern. • In table salt, NaCl, sodium ions and chloride ions are arranged in a regular threedimensional structure referred to as a crystal lattice. • Other ionic compounds will have different crystal lattices. Molecular Solids • A crystalline molecular solid has molecules arranged in a particular conformation. • In water, H2O, the molecules are arranged in a regular threedimensional structure. • Other examples of crystalline molecular solids are table sugar, C12H22O11, and sulfur, S8. Metallic Solids • A crystalline metallic solid is composed of metal atoms arranged in a definite pattern. • A metallic crystal is made up of positive metal ions surrounded by valance electrons. • Metals are good conductors of electricity because electrons are free to move about the crystal. • This is referred to as the “electron sea” model. General Properties of Solids Changes of Physical State • Heat is necessary to raise the temperature and change the physical state of a substance. • Specific heat is the amount of heat required to raise 1.00 g of a substance 1C. • Water is the reference and its specific heat is 1.00 cal/(g×C). • The specific heats of ice and steam are about half that of liquid water. Solid/Liquid Phase Changes • As a solid melts, the temperature is constant until all of the solid is changed to liquid. • The amount of heat required to melt 1.00 g of substance is called the heat of fusion (Hfusion). For water, the heat of fusion is 80.0 cal/g. • When a liquid changes to a solid, the heat change is the heat of solidification (Hsolid). • The value of Hfusion is the same as the value of Hsolid. Liquid/Gas Phase Changes • As a liquid vaporizes, the temperature is constant until all of the liquid is changed to gas. • The amount of heat required to vaporize 1.00 g of substance is called the heat of vaporization (Hvapor). For water, the value is 540 cal/g. • When a gas changes to a liquid, the heat change is the heat of condensation (Hcond). • The value of Hvapor is the same as the value of Hcond. Temperature/Energy Graphs • We can graph the amount of energy required to change the temperature and physical state of a substance. • The heating curve for water is shown here. • As energy is added the temperature increases and changes the physical states. Energy from Heating Curves • We can use the heating curve and heat values for water to calculate how much energy is required to change the temperature of a sample of water. • The amount of energy required to raise the temperature of a substance is calculated using the following formula: heat = (specific heat) × (change in temperature) × (mass of sample) • The amount of energy required to change the state of a substance is calculated as follows: heat = (Hxxx) × (mass) Chapter 13 29 Energy Calculation • How much energy is required to raise 25.5 g of ice at –5.0C to steam at 100.0C? • Looking at the heating curve for water, there are 4 regions: 1) Heating of solid ice from –5.0C to 0.0C 2) Melting of ice at 0.0C 3) Heating of liquid water from 0.0C to 100.0C 4) Vaporization of water at 100.0C Energy Calculation • The total energy is the sum of the energy in steps 1) through 4). Calculate the energy for each step. 1) (25.5 g) × [0.0 – (–5.0)]C × (0.50 cal/g×C) = 64 cal 2) (25.5 g) × (80.0 cal/g) = 2040 cal 3) (25.5 g) × [100.0 – 0.0)]C × (1.00 cal/g×C) = 2550 cal 4) (25.5 g) × (540 cal/g) = 13,800 cal • The total energy is: 64 cal + 2040 cal + 2550 cal + 13800 cal = 18,500 cal Structure of Water • Lets start with the electron dot formula for water: • Water has a bent molecular shape and the bond angle is 104.5. • Water is a polar molecule that exhibits strong hydrogen bonding. 32 Properties of Ice • Water is a one of the few substances that is less dense as a solid than as a liquid. • As water freezes, the hydrogen bonds organize the water molecules into a three-dimensional structure where the molecules are farther apart in the liquid. • Liquid water has a density of 1.00 g/mL while ice has a density of 0.917 g/mL. Physical Properties of Water • Water has unusually melting and boiling points, especially compared to the other hydrogen compounds of Group IVA/16. • This is due to hydrogen bonding which is present in water but not present in H2S, H2Se, or H2Te. Chemical Properties of Water • Water can undergo an electrolysis reaction to produce hydrogen and oxygen: 2 H2O(l) electricity → 2 H2(g) + O2(g) • Water reacts with active metals to produce hydrogen and a metal hydroxide: 2 K(s) + 2 H2O(l) → 2 KOH(aq) + H2(g) • Water reacts with metal oxides to produce a base: CaO(s) + H2O(l) → Ca(OH)2(aq) • Water reacts with nonmetal oxides to produce an acid: CO2(g) + H2O(l) → H2CO3(aq) 35 Hydrates • A hydrate is a crystalline ionic compound that contains water: CuSO45H2O • The dot indicates that water molecules are bonded directly to each unit of hydrate. • Heating a hydrate drives off the water and produces an anhydrous compound (without water). CuSO45H2O(s) → CuSO4(s) + 5 H2O(l) heat Conclusions • There are three types of intermolecular bonds: – Dispersion forces – Dipole forces – Hydrogen bonds • Dispersion forces are the weakest and hydrogen bonds are the strongest. • These intermolecular attractions affect the physical properties of substances. Conclusions Continued • There are 5 properties of liquids that are affected by intermolecular bonds: – Vapor pressure decreases as intermolecular forces increase – Boiling point increases as intermolecular forces increase – Viscosity increases as intermolecular forces increase – Surface tension increases as intermolecular forces increase Conclusions Continued • There are 3 types of crystalline solids: – Ionic solids – Molecular solids – Metallic solids • Energy is required to raise the temperature of a substance. • Water displays many unique properties due to the presence of strong hydrogen bonds.