Liquids & Solids

Chapter 12

Properties of LiquidsO Have definite volume & take shape

of containerO Lower kinetic energy than gasesO Particles move about constantly,

more orderly, more attractive forces between particles

O Liquids are fluids (they flow)

Properties of LiquidsO Higher density (x 1000s) than gasesO Most have lower density than solids,

(H2O exception)

Properties of LiquidsO Transmit pressure equally in all

directions, because they are relatively incompressible.

Properties of LiquidsO Ability to diffuse more slowly than

gases; will diffuse more rapidly with increase in temp.

Properties of LiquidsO Surface tension increases with more

attraction between molecules

Cohesion & AdhesionO Cohesion – force of attraction

between particles of the same substance

O Adhesion –force of attraction between particles of different substances

Properties of LiquidsO Capillary action – the attraction of

the surface of a liquid to the surface of a solid.

VaporizationO The process by which a liquid or solid

changes to a gas.O Evaporation – particles escape from

the surface of a nonboiling liquid and enter the gas state.

VaporizationO Boiling – change of a liquid to

bubbles of vapor that appear throughout the liquid; not just at the surface. It occurs when the equilibrium vapor pressure of the liquid equals the atmospheric pressure.

FreezingO When a liquid cools,

average kinetic energy of the particles decreases. If low enough, the attractive forces between particles will pull the particles into an even more orderly arrangement. (solid)

Properties of SolidsO Definite shape and volumeO Intermolecular forces between particles are

stronger, more closely packed.O Exhibit relatively fixed positionsO More ordered arrangement than liquids or gases

Properties of solidsO 2 types of solids

O Crystalline – orderly, geometric arrangementO Definite melting point (temperature at

which a solid becomes a liquid)O Amorphous – random arrangement

O No definite melting point (glass, plastic, rubber)

O Sometimes called supercooled liquids

O Higher density than gases and most liquids

O Higher incompressibility than liquids

Properties of SolidsO Low rate of diffusion – while diffusion

does occur in solids, it is millions of times slower than in liquids.

Crystalline SolidsO Exist as

O Single crystalsO Groups of crystals fused together

O Crystal structure – total 3-dimensional arrangement of particles of a crystal

O Crystal lattice – a coordinate system of particles of a solid, consisting of multiple unit cells.

Crystalline solids O 7 types of symmetry

4 Binding Forces in Crystal Structures

O Ionic crystals – form hard, brittle crystals with high melting points and make good insulators – form between pos/neg ions.

O Covalent network crystals – form large bonds between atoms that are covalently bonded to their nearest neighboring atoms – (Cx, SiO2, SiCx) – hard, brittle, rather high melting pts., non conductors or semi-conductors.

4 Binding Forces in Crystal Structures

O Metallic crystals – metal atoms are surrounded by a sea of valence electronsO Melting points vary greatly

O Covalent molecular crystals – held together by weak intermolecular forcesO Non-polar – H2, CH4, C6H6 – held together by

weak London dispersion forcesO Polar – H2O, NH3 – can be held together by

dispersion forces, dipole-dipole forces, or even hydrogen bonding

O Both have low melting points, easily vaporized, soft, and are good insulators

Amorphous SolidsO - from the Greek word, “without

shape”O Do not have a regular shapeO Hold their shape for a long time, but

DO flow!

Changes of StateO Equilibrium – a dynamic (changing)

condition in which two opposing changes occur at equal rates in a closed system. A closed system is one in which matter cannot enter or leave, but energy can.Equilibrium- individuals

constantly coming &going, but total # peopleStay the same.Not in equilibrium- in the morning more people come, as time passes, pool gets crowded.

Le Châtelier’s PrincipleO When a system at equilibrium is disturbed

by application of a stress, it attains a new equilibrium position that minimizes the stress. A stress is typically a change in concentration, pressure, or temperature.O Ex. Suppose the temperature to a closed

system is raised from 25oC to 50oC. O Liquid + heat energy vapor

O Endothermic, shifts to the right in favor of the forward reaction.(vapor produced at higher rate, until equilibrium is attained.)

O Ex. Holding mass and temperature constant in a closed system, yet increasing the volume causes a stress to the equilibrium of the system. What happens?O Concentration of the molecules in the

system decreases because volume has increased. Because of this, fewer molecules strike the liquid surface and condense. Rate of evaporation is now higher than rate of condensation.O Liquid + heat energy

vaporO Shifts to the right because of more

evaporation

Equilibrium ShiftsShifts in the Equilibrium for the ReactionLiquid + Heat Energy VaporChange Shift

Addition of liquid Right

Removal of liquid Left

Addition of vapor Left

Removal of vapor Right

Decrease in container volume

Left

Increase in container volume

Right

Decrease in temperature Left

Increase in temperature right

Chemical Equilibrium(Ch.18)

O A chemical reaction in which the products can react to re-form the reactants is called a reversible reaction.

O A reversible reaction is in chemical equilibrium when the rate of its forward reaction equals the rate of its reverse reaction and the concentrations of its products and reactants remain unchanged.

Chemical reaction convention

O A chemical reaction is written left right (forward) and right left (reverse).

Predicting the Direction of Shift

O Changes in PressureO A change in pressure affects only

equilibrium systems in which gases are involved.

O For changes in pressure to affect the system, the total number of moles of gas on the left side of the equation must be different from the total number of moles of gas on the right side of the equation.

Change in PressureO Example

O Let’s consider the synthesis of ammonia:O N2(g) + 3H2(g) 2NH3

O If pressure is increased on this system, equilibrium shifts to the right. This is because we have 4 moles of total gas on the left, and only 2 moles on the right. The system can reduce the number of molecules, and hence, the total pressure, by shifting the equilibrium in the direction of the lesser number of moles.

Remember:O Changes in partial pressure of a gas

within a reaction WILL cause a shift in equilibrium.

O Changes in pressure cause changes in concentration of gases within a reaction.

O The addition or removal of some other gas has no effect even though it changes the total pressure.

Changes in Temperature

O Reversible reactions are exothermic in one direction and endothermic in the other.

O According to Le Châtelier’s principle, the addition of heat shifts the equilibrium so that heat is absorbed, which favors the endothermic reaction. The reverse is true for exothermic reaction.

O HINT: Treat heat energy absorbed as a reactant, and heat energy released as a product, to determine shift.

Changes in Temperature

O ExampleO 556 kJ + CaCO3(s) CaO(s) +

CO2(g)

O An increase in temperature caused by adding heat to the system would shift the equilibrium to the right. A decrease in temperature would shift it in the reverse direction.

Changes in Concentration

O An increase in the concentration of a reactant is a stress on the equilibrium system. It causes an increase in collision frequency and generally an increase in reaction rate.O Ex. A + B C + DO An increase in the concentration of A

shifts the equilibrium to the right. O A decrease in the concentration of A

shifts to the left.

Possible changes of state

O Phase – any part of a system that has uniform composition and properties.

O Evaporation can be expressed as:O Liquid + heat energy vaporO When a liquid changes to a vapor, it

absorbs heat energy from its surroundings.

O Condensation can be expressed as:O Vapor liquid + heat energyO When a vapor condenses, it gives off

energy to its surroundings.

O Liquid-Vapor equilibrium is expressed:O Liquid + heat energy vapor

O Double arrow indicates a reversible reaction

Phases of MatterO Phase- any part of a system that

has uniform composition and properties

O Equilibrium vapor pressure – the pressure exerted by a vapor in equilibrium with its corresponding liquid at a given temperature.

O Volatile liquids – liquids that evaporate readily, have weak forces of attraction between their particles.

Which liquid is more volatile?

O Oil, water, alcohol

Oil, water, alcoholO Which liquid is more evaporates

most readily?O Which liquid exhibits a higher

amount of kinetic energy at room temperature?

O Which liquid has the weakest attractive forces between molecules?

O Which liquid is least volatile?

BoilingO A liquid boils when it has absorbed enough

energy to evaporate. O Boiling – the conversion of a liquid to a

vapor within the liquid as well as at its surface.

O Boiling point – the temperature at which the equilibrium vapor pressure of the liquid equals the atmospheric pressure.

O The lower the atmospheric pressure is, the lower the boiling point. (Remember, boiling water in the bell jar?)

The normal boiling point of each liquid shown occurs when its equilibrium vapor pressure equals 760 torr. Energy must be added continuously in order to keep a liquid boiling. The temperature of the liquid and its vapor at the boiling point remains constant despite the continuous addition of energy. The added energy is used to overcome the attractive forces between the molecules of the liquid during the phase change, and is stored in the vapor as PE.

O s

Molar Enthalpy of Vaporization

O Symbol - ΔHv

O The amount of energy as heat that is needed to vaporize one mole of liquid at the liquid’s boiling point at constant pressure is called the liquid’s molar enthalpy of vaporization.

O Measures the attraction between particles of a liquid.

O The stronger the attraction is, the more energy is required to overcome it, resulting in a higher molar enthalpy of vaporization.

FreezingO Occurs when a substance loses enough heat

energy to solidify.O Liquid solid + energy (results in loss

of energy)O The normal freezing point – the temperature

at which the solid and liquid are in equilibrium at 1 atm, 760 torr, or 101.3 kPa, pressure. (for pure crystalline substances)O Freezing occurs at constant temperature, and

the energy lost during freezing is the PE that was present in the liquid.

MeltingO Also occurs at constant temperatureO Solid + energy liquidO For pure crystalline solids, the

melting point and freezing point are the same. At equilibrium, melting rate = freezing rate

Molar Enthalpy of Fusion

O The amount of energy as heat required to melt one mole of solid at the solid’s melting point is the solid’s molar enthalpy of fusion.

O Symbol – ΔHf

O Energy absorbed increases the solid’s PE as its particles are pulled apart, overcoming the attractive forces holding them together.

Sublimation & Deposition

O At sufficiently low temperature and pressure conditions, a liquid cannot exist.

O A solid substance exists, in this case, in equilibrium with its vapor instead of its liquid.

O Solid + energy vaporO Change of state from solid to gas-

sublimationO Reverse process - deposition

Phase Diagram of Water

O A graph of pressure vs. temperature that shows the conditions under which the phases of water exist. It also reveals how the states of a system change with changing temperature or pressure.

Triple point – the temp. and pressure at which the solid, liquid, and vapor of the substance can coexist at equilibrium.Critical point- the critical temperature & pressure of the substance.Critical temperature- (tc), the temp above which the substance cannot exist in the liquid state.Critical pressure-(pc), the lowest pressure at which the substance can exist as a liquid at the critical temperature.

Water – Most Abundant Liquid

O 75% of Earth’s surface – oceans, rivers, lakes

O Frozen in glaciersO 70%-90% of the

mass of all living things

O Chemical reactions of life take place in water

Structure of WaterO Bent shape, polar moleculeO Intermolecular – between molecules, hydrogen

bondingO Intramolecular – between atoms within a

molecule – covalent bonding

O Rigid, open structure of solid ice crystals, has much space between them making water less dense as a solid.

O Water molecules are as tightly packed as possible at 3.980C.

O Water’s relatively high boiling pt. is due to strong hydrogen bonding between molecules.

Properties of WaterO Room Temperature-transparent,

odorless, tasteless, almost colorless.O Freezes at a pressure of 1 atmO Melts at a temperature of 00CO Molar heat of fusion of ice = 6.009

kJ/molO Density of ice = 0.917 g/cm3 at 00CO Density of liquid water = .999 84

g/cm3 at 00C

Properties of WaterO Both the boiling point and the

molar heat of vaporization of water are high compared to those of nonpolar substances of comparable molecular mass, like CH4, methane.

O This is because of the strong hydrogen bonding that must be overcome for boiling to occur.

O Molar heat of vaporization at 1000C is 40.79 kJ/mol, relatively high making steam a good source for cleaning, use in radiators or steam-heating systems.

Floating Ice

The lower density of ice acts as an insulator in Earth’s oceans,allowing living organisms to thrive. If ice were denser than water, it would sink and would cause theoceans to freeze killing all life within them!

Quantity of Heat Energy

O How much heat energy is absorbed when 47.0g of ice melts at STP? How much heat energy is absorbed when the same mass of liquid water boils?

O Use:O Amount of substance(mol) x molar

heat of fusion or vaporization (kJ/mol) = heat energy (kJ)

ExamplesO What quantity of heat energy is

released when 506g of liquid water freezes?

O Answer: 169 kJ

Example 2O What mass of steam is required to

release 4.97 X 105kJ of heat energy on condensation?

O Answer: 2.19 X 105g