Transcript Liquids - Derry Area School District
Liquids
Kinetic Theory
Explains states of matter in terms of movement of particles.
For liquids: Molecular forces stronger than gases molecules closer than gases.
Molecular forces weaker than solids molecules farther apart than solids.
Phase Changes
Solid
LOW E k Temperature (E k ) increases The
movement Collisions Distance
of particles between particles between particles , due to increased force of collisions.
Particles move apart sufficiently to enter another phase of matter -
Liquid.
For pure substances the melting/freezing phase change takes place at a definite temperature.
Vapor Equilibrium
Vapor Equilibrium
Represents a CLOSED SYSTEM Some particles lose Ek and change from a gas to a liquid. CONDENSATION •
Vapor – Ek
Liquid
Some particles have/gain sufficient Ek to escape from liquid phase to become gaseous.
EVAPORATION •
Liquid + Ek
Vapor
(Gas)
Rate evaporation = Rate condensation
the system is in a state of dynamic
equilibrium.
The number of gas particles changing to liquid equals the number of liquid particles changing to gas.
The vapor region of the system is saturated (contains as many particles possible at a specific temperature).
Equilibrium Vapor Pressure
The pressure exerted by the vapor portion of the system on the liquid portion.
Depends on: Type of liquid Temperature Generally, the LOWER the vapor pressure the STRONGER the intermolecular forces of the particles on the system.
LeChatelier’s Principle
If
stress
[a change in the conditions of the system] is applied to a system, the equilibrium moves in the direction that will reduce the stress.
Pressure increases --- System tries to relieve the stress by reducing pressure.
In the system initially: Vapor ↔ Liquid + E k
Pressure
↑ more liquid produced
Volume
↑
Temperature
↓ particles allowed to expand
Pressure
↓ more vapor prod. more vapor produced
Solid to Liquid
Solids become liquids at the
melting point
.
Temperature at which vapor pressure of the solid = vapor pressure of liquid Depends on: Size of the molecules Molecular forces • Weak intermolecular forces --- LOW m.p.
• High intermolecular forces --- HIGH m.p.
Ex:
H 2 O
Polar m.w. 18 m.p. 0o C b.p. 100 0C
CH 4
Nonpolar m.w. 16 m.p. –183 oC b.p. –162 oC
Standard Boiling Point
Temperature @ which the vapor pressure of the liquid equals standard atmospheric pressure.
[1 ATM, 101.325 kPa, 760 mm Hg, 760 Torr]
Boiling Point
Temperature @ which the vapor pressure of the liquid equals the existing atmospheric pressure.
If a liquid Boils quickly Evaporates quickly @ room temperature the liquid is said to be
volatile.
Volatile liquids usually have HIGH vapor pressures.
Ex: alcohols If the vapor pressure is extremely high the substance will change from a solid directly to a gas.
Sublimation
Pascal’s Law
Applies the idea of vapor pressure to BOILING… Pressure is exerted evenly on the surface of a confined liquid, and is transferred undiminished throughout.
As a liquid is heated…
Gases become less soluble @ higher temperatures.
bubbles form at the hottest point Bubbles collapse due to atmospheric pressure [Pascal’s Law] As the temperature of the liquid increases, the Ek of the molecules increases.
the vapor pressure of the bubbles increases When the vapor pressure of the bubbles equals the vapor pressure of the atmosphere the bubbles survive to the surface of the liquid – BOILING During Boiling the temperature of the liquid remains
constant.
Why can steam burn more severely than boiling water?
A: higher E k
Liquification of Gases
Michael Faraday
Discovered it was possible to liquefy gases by
simultaneously
cooling and compressing the gas.
Modern Liquification of Gases
Compress gas and increase temperature Use a coolant to remove the temperature increase Allow the gas to expand Temperature drops [Joule-Thomas Effect] Process repeated As temperature decreases Ek decreases A point is reached where the intermolecular forces of attraction [van der Waal / London forces] can cause the molecules to combine if they are close enough.
For every gas…
there is a temperature above which NO AMOUNT of pressure will cause the gas to liquefy –
critical temperature
[Tc].
Pressure that results in liquification at Tc –
critical pressure
.
In order to liquefy a gas Conditions must be at / below Tc Pressure at / above vapor pressure of liquid
Energy and Changes of State
Ep inc.
Ek inc Ep inc.
Ek inc Before a phase change – Kinetic Energy increases.
During a phase change – Potential Energy increases.
Enthalpy
Enthalpy of Fusion
[H fus ] – heat energy required to melt 1 gram of a substance.
Ex: ice – 334 J/g
Enthalpy of Vaporization
Ex: H2O – 2260 J/g [H vap ] – heat energy required to convert 1 gram of liquid to gas @ its boiling point.
Enthalpy
[q] – heat transferred q = m H fus q = m H vap q = m TC p [specific heat]
WATER
Most abundant liquid 75% of Earth’s surface covered by it.
10% of the land covered by water in the form of glaciers Water vapor is always present in the air.
70% -90% of all living things are water by weight. Large reserves of water underground
Physical Properties
Transparent Odorless Tasteless Almost colorless Exists as solid, liquid, gas @ Standard Pressure [1 Atm., 760 mm Hg, 760 Torr, 101.325 kPa] 0 o C ----- liquid water -------- 100 cm 3 solid water + Energy 111.111 cm 3 0 o C 4 o C -------- 1/9th greater volume Water contracts > 4 o C ---------------- Water begins to expand Water
most dense @ 4 o C.
1 g = 1 ml = 1cm 3
Standard Values for Water
Melting / freezing ……………..0
o C Boiling / condensing ……………100 o C Molar heat of vaporization ……...9.7 kcal [Heat required to convert 1 mole of water (18 g) into a gas]
From the definition of boiling, we can see that as the pressure the boiling point .
Putting this to practical application………..
Vacuum Evaporator Industrial device that reduces the b.p. by reducing the pressure on the liquid.
• • • Evaporated milk Eagle Brand milk Soups
Structure and Properties of Water
Bent Molecule O + H H+ 104.5
o •Oxygen and hydrogen form a
polar covalent
bond [dipole]
In a container…
Equal forces in all directions
On the surface…
Perpendicular force downward on the surface The perpendicular force causes the phenomena of
Surface tension
Pulls surface molecules together and make it less penetrable than expected.
Capillary rise
Rise of a liquid in a tube of small diameter. The attraction between the water and the glass relieves the perpendicular force.
For many substances the 1) structure and the 2) molecular weight can be used to predict the general changes in state.
Generally @ room temperature…
Liquids
Polar covalent bonds
Solids Gases
ionic bonds
m.p.
Pure covalent bonds
NaCl [58.5] 801 o C C 6 BrCH 3 288 o Toluene [487] C C N 2 2 H 4 -169 o C -210 o C
Some substances do not change their state within the ranges expected
Contain H in a
polar covalent bond
H bonded to atoms like N, O, or F These atoms are highly electronegative
Electronegative atoms pull shared electrons closer to them than to the other atom in the bond.
+ Hydrogen
Shared e- Pair Nitrogen The uneven sharing of the electron pair causes slight charges (
dipoles
) to form around the bond.
The Hydrogen Bond
A
weak attractive force
[not a bond] between the hydrogen of one molecule and an electronegative atom in another molecule.
Hydrogen
Electronegative atom
Hydrogen “bond”
Hydrogen bonds are sensitive to temperature. The higher the temperature the
fewer
the hydrogen bonds.
The effect of hydrogen bonding is cumulative. [Like pieces of paper/ threads/hair] Hydrogen bonding causes water to be a liquid @ room temperature. [It would otherwise be a gas due to its low molecular weight.]
0 o C O H bonds between
the water molecules are
rigid.
As the temperature increases….
together.
4 o C
The This flexibility allows the
O H
bonds become
Maximum density.
flexible.
molecules to move closer
At Higher Temperature Molecules begin to move apart. Liquid expands Hydrogen bonding explains why water expands when it freezes.
Chemical Behavior of Water
Extremely stable molecule Does not begin to decompose until temperature is > 2700 o C.
Reacts with metals to form hydrogen gas.
H 2 O + X
H 2 + XOH
Hydrogen gas + Metal hydroxide @ rm temp
H 2 O + X
H 2 + XOH
Hydrogen gas + Metal oxide @ temp > 170 o C [Exception: K can react with cold water to form a metal oxide.]
Basic anhydride
a metal oxide that forms a basic solution when combined with water BASIC solution [due to OH-] Slippery Bitter taste Red litmus Blue Reacts with metal oxides to form metal hydroxides.
H 2 O + XO
XOH XOH
X + + OH -
The hydroxide separates - forms the metal in solution [ X+ and OH- result].
Acid Anhydride
a nonmetal oxide that forms an acid solution when combined with water ACID solution [due to H 3 O + ] Sour taste Blue litmus Red Reacts with nonmetal oxides to form acidic solutions.
H 2 O + CO 2
H 2 CO 3
[carbonic acid] The hydrogen separates form the nonmetal in solution [ H + and HCO 3 —result].
H+ reacts with the water to form the
hydronium ion.
[
H 3 O + ]
Capturing and Releasing Water
Na 2 CO 3 10H 2 O Na 2 CO 3 H 2 O + 9H 2 O Loosely attached Water of hydration
You go to the grocery store and buy a box of washing soda that is marked
14 oz.
You weigh the box when you get home. The balance reads
10 oz.
Assuming that the company is honest and that no mistakes were made in packaging, how can you explain this discrepancy?
Efflorescence
Water of hydration that is lost to the atmosphere.
The higher the vapor pressure of the substance the greater the efflorescence of the hydrate.
Deliquescence
Absorption of water molecules from the atmosphere.
The vapor pressure of the substance is lower than the vapor pressure of the water vapor in the air.
Forms a solution with the water in the air until the vapor pressure of each is equal.
Hydroscopic
Takes water form atmosphere and traps it in pores and imperfections in its surface.
Ex: hair, wool, tobacco, potato chips