Liquids and Solids

Liquids and Solids

Gas, Liquid, and SolidGas Liquid SolidHighly Compressible Slightly Compressible Very slightly compressible

Low Density High Density High Density

Fills container completely Does not expand to fill container

Rigidly retains its volume

Assumes shape of container

Assumes shape of container

Retains its own shape

Rapid diffusion Slow diffusion Extremely slow diffusion, surface only

High expansion on heating Low expansion on heating Low expansion on heating

Total Disorder Disordered Ordered arrangement

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

Diamond

Graphite

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

Silver

Copper

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 that is 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

Sublimation and Deposition Sublimation: solid goes directly to a gas without

passing through the liquid phase Deposition is the reverse process Sublimation is endothermic Deposition is exothermic

Heating Curves Graphic illustrations of phase changes Plot of temp of a sample as a function of time Notice temp remains constant during phase

changes while amount of energy varies

A: Rise in temperature as ice absorbs heat.B: Absorption of heat of fusion.C: Rise in temperature as liquid water absorbs heat.D: Water boils and absorbs heat of vaporization.E: Steam absorbs heat and thus increases its temperature.The above is an example of a heating curve. One could reverse the process, and obtain a cooling curve. The flat portions of such curves indicate the phase changes.

Heating Curve of Water

Phase Diagrams Diagram that relates the states of a substance to

temp and pressure State depends on temp and pressure 2 states can exist simultaneously at certain

temps and pressures Triple point: the temp and pressure when all

three states exist at the same time

• TRIPLE POINT - The temperature and pressure at which the solid, liquid, and gas phases exist simultaneously.

• CRITICAL POINT - The temperature above which a substance will always be a gas regardless of the pressure.

• FREEZING POINT - The temperature at which the solid and liquid phases of a substance are in equilibrium at atmospheric pressure.

• BOILING POINT - The temperature at which the vapor pressure of a liquid is equal to the pressure on the liquid.

• Normal (Standard) Boiling Point - The temperature at which the vapor pressure of a liquid is equal to standard pressure (1.00 atm = 760 mmHg = 760 torr = 101.325 kPa)

• NOTE – • The line between the solid and liquid phases is a curve of all

the freezing/melting points of the substance. • The line between the liquid and gas phases is a curve of all the

boiling points of the substance.