Intermolecular Forces • Forces of attraction between neighboring particles • Much weaker than bonding forces • Responsible for state of matter and some physical properties • e.g., The stronger the attractive forces, the higher the melting and boiling points • Also involved in change of state
Three Types • London Dispersion forces • Dipole-dipole forces • Hydrogen bonds
London Dispersion Forces • The motion of electrons can create an instantaneous dipole moment on an atom • For example, if at any one time both of a helium atom’s electrons are on the same side of the atom at the same time • A temporary dipole on one atom can cause, or induce, a temporary dipole on an adjacent atom
London Dispersion Forces • These forces are significant only when molecules are very close together, as in a compressed gas • These forces are found only in nonpolar compounds • Molecules and atoms will lose their spherical shape
More compact molecules have smaller surface areas, weaker London dispersion forces, and lower boiling points. • Flatter, less compact molecules have larger surface areas, stronger London dispersion forces, and higher boiling points.
Dipole-Dipole Forces • Polar molecules have a positive end and a negative end • Dipole-dipole forces occur when the positive end of one molecule is attracted to the negative end of another • Only effective when polar molecules are very close together • For molecules of about the same size, dipole forces increase with increasing polarity
If two neutral molecules, each having a permanent dipole moment, come together such that their oppositely charged ends align, they will be attracted to each other.
Hydrogen Bonds • Type of dipole-dipole force • Not a true bond! • Occurs between molecules containing a hydrogen atom bonded to a small, highly electronegative atom with at least one lone pair of electrons (e.g., N, O & F) • The hydrogen in one molecule will be attracted to the electronegative atom in another molecule
Hydrogen Bonds • Hydrogen has no inner core of electrons, so a dipole will expose its concentrated charge on the proton, its nucleus. • Hydrogen can approach an electronegative atom very closely and interact strongly with it.
Electron shell around a hydrogen atom is rather thin, giving the hydrogen atom a small positive charge. • Electron shell round an oxygen atom is quite thick, and so oxygen carries an extra bit of negative charge. • These opposite charges attract, although quite weakly. • This weak force is called a hydrogen bond. The hydrogen atoms of one water molecule stick to the oxygen atoms of nearby water molecules.
Properties of Liquids • Have much greater densities than their vapors • Only slightly compressible; not a discernable difference when compressed • Fluidity: ability to flow • Liquids can diffuse through one another, but at a much slower rate than gases
Properties of Liquids • Viscosity: resistance to flow • Determined by the type of intermolecular forces involved, the shape of the particle, and the temperature • The stronger the attractive forces, the higher the viscosity • The larger the particles, the higher the viscosity • Increases as temp decreases
Properties of Liquids • Surface Tension: the imbalance of forces at the surface of a liquid • The uneven forces make the surface behave as if it has a tight film stretched across it • The stronger the intermolecular forces, the higher the surface tension
Properties of Liquids • Surfactants: compounds that lower the surface tension of water • Frequently added to detergents • Capillary action: movement of a liquid through narrow spaces
Properties of Solids • Have extremely strong intermolecular forces in order for solids to have definite shape and volume • Particle arrangement causes solids to almost always have higher densities than liquids • Ice is an exception: it expands when it freezes because of the way the particles arrange themselves during the freezing process
Properties of Solids • Particle arrangements cause different types of solids: • Crystalline solids • Molecular solids • Covalent network solids • Ionic solids • Metallic solids • Amorphous solids
Crystalline Solids • Has atoms, ions, or molecules arranged in an orderly, geometric, 3-D structure • Individual pieces of a crystalline solid are called crystals • Smallest arrangement of connected points that can be repeated in 3 directions to form a lattice is called a unit cell • There are 7 different crystal systems based on shape
Molecular Solids • Held together by dispersion forces, dipole-dipole forces or hydrogen bonds • NOT held together by genuine bonds (ionic and covalent) • Most are NOT solids at room temperature • Poor conductors of heat and electricity (don’t contain ions) • Examples are sucrose and ice
–Molecular such as sucrose or ice whose constituent particles are molecules held together by the intermolecular forces.
Arrangement of molecules in liquid water Arrangement of molecules in ice
Covalent Network Solids • Atoms that can form multiple covalent bonds • Form a network of atoms that do not have a unit cell • Most allotropes exist in this form • Allotropes are forms of the same element that have different bonding patterns of arrangement • Examples include diamonds and graphite, quartz
Covalent network solids such as quartz where atoms are held together by 3-D networks of covalent bonds. Here the hexagonal pattern of Si (violet) and O (red) atoms in structure matches the hexagonal crystal shape
Ionic Solids • Type of crystalline solid • Type and ratio of ions determine the structure of the lattice and the shape of the structure • The network of attractions that extend through an ionic compound gives these compounds their high melting points and hardness
Ionic Solids • Strong but brittle • When struck, cations and anions are shifted, which causes repulsion that in turn shatter the crystal • Poor conductors of heat and electricity in solid form
Ionic solids are an orderly pattern of one ion, generally the anion, with cations positioned in 'holes' between the anions • The occupation of these 'holes' depends on the formula of the ionic compound
Sodium chloride Cupric chloride
Metallic Solids • Consist of positive metal ions surrounded by a sea of mobile electrons • Mobile electrons make metals malleable, ductile, and good conductors of heat and electricity
A series of metals atoms that have all donated their valence electrons to an electron cloud that permeates the structure • This electron cloud is referred to as an electron sea • Visualize the electron sea • model as if it were a box of marbles that are surrounded by water. The marbles are the metal atoms and the water represents the electron sea.
The marbles can be pushed anywhere within the box and the water will follow them, always surrounding the marbles. • This unique property, allows metallic bonds to be maintained when pushed and pulled in all sorts of ways. • As a result, they are malleable and ductile.
Gold Copper Silver
Amorphous Solids • Solid in which the particles are not arranged in a regular, repeating pattern, but still retain rigidity • Examples include glass, rubber, many plastics, tar and wax • Particles are trapped in a disordered arrangement thatis characteristic of liquids
Phase Changes • Always involve a change in energy • Energy is needed either to overcome or form attractive forces between particles
Melting and Freezing • Melting point/freezing point: temp at which solid and liquid forms exist in equilibrium • Melting is endothermic • Freezing is exothermic
Vaporization • The change of state from a liquid to a gas • Endothermic process • Two methods of vaporization: • Evaporation • Boiling
Evaporation • Occurs at the surface of a liquid • Occurs because molecules close to the surface have enough energy to overcome the attractions of neighboring molecules and escape • Slower molecules stay in the liquid state • Rate of evaporation increases as temp increases
Boiling • Occurs within the liquid • Boiling point: temp at which vapor pressure equals atmospheric pressure • If vapor pressure is less than atmospheric pressure, bubbles do not form
Condensation • Change of a gas to a liquid • Exothermic process • Molecules of vapor can return to the liquid state by colliding with the liquid surface • The particles become trapped by the intermolecular attractions of the liquid