# Chapter 12 (States of Matter)

12.1 Gases
Using the kinetic – molecular theory to explain the behavior of gases; describe how mass affects the rates of diffusion and effusion; define the different ways to measure gas pressure; calculate the partial pressure of a gas.

kinetic – molecular theory
describes the behavior of matter in terms of particles in motion

Assumption #1 of the k-m theory:
the size of gas particles is negligible compared to the empty space that surrounds each particle, therefore gas particle do not affect each other (there is no attraction or repulsion between individual gas particles)

Assumption #2 of the k-m theory:
gas particles are in constant, random motion; when gas particles collide the collision is “elastic”, meaning that no energy is lost after the collision – the total kinetic energy before and after the collision is the same.

Assumption #3 of the k-m theory:
the kinetic energy (KE) of a gas particle is related to its mass and the square of its velocity: KE = 1/2 mv²
velocity is the direct result of a ges’s temperature.

Why do gases have such low density (compared to liquids and solids)?
The Empty Space between gas particles explains why gases have such low density.
Example: chlorine gas ,Cl₂, molecules are more than 1/3 the mass of gold atoms (71g/mol vs 197g/mol). But a piece of gold that takes up the same space as a sample of chlorine gas will have around 6500 times more mass.This is due to the large amount of empty space around each Cl₂ molecule.

What allows gases to be compressed or expanded so easily?
The Empty Space between gas particles explains why gases can be compressed (increasing its density) or expanded (decreasing its density)

Diffusion
the movement of a gas into an area that already contains another gas. The gas will move from an area where the gas is in a high concentration to where the gas is in a low concentration.

Examples of diffusion:
(1) the scent of cooking food spreads from the kitchen to the rest of the house; (2) the smell of a fart or turd quickly spreads from the bathroom to other parts of the house if the bathroom door is not closed or a candle is not lit right away; (3) a woman sprays perfume on when she goes to a nightclub, and the pleasant scent spreads to all the nearby males within seconds.

Effusion
a gas in an enclosed space escapes through a tiny opening. The rate of effusion (escape) is inversely related to the molar mass of the gas – the less dense a gas is, the more quickly it will esape through the opening.

Examples of Effusion:
(1) Helium balloons tend to deflate more rapidly than balloons with air in them because helium is a much less dense gas than air. (2) over time car or bicycle tires will go flat due to the escape of small amounts of air through tiny openings or leaks in the tire.

Graham’s Law of Effusion:
Rate of Effusion∞ 1 / √molar mass

(Rate of Effusion of Gas “A”) ÷ (Rate of Effusion of Gas “B”) =
√ molar mass of “B” ÷ molar mass of “A”

Gas Pressure:
The amount of force applied by a gas over a given amount of area.

Pressure (in Pascals) =
Force (in Newtons) ÷ Area (meters²)
1 Pascal = 1N/m²

Atmospheric (Air) Pressure
Measured by a devise called a BAROMETER.

Units of Pressure
1000 Pascals = 1 kilopascal (1kPa)
1 Atmosphere = 101.3 kPa = 760 mm of Hg = 760 torrs = 14.7 pounds per square inch (P.S.I.) = 1.01 bars

Dalton’s Law of Partial Pressures
the total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture.

Total Pressure =
P₁+ P₂+ P₃+ P₄= Total Pressure

12.2 Forces of Attraction
Intramolecular vs Intermolecular

Intramolecular Forces: attraction within a molecule or formula unit (strongest of the two types)
Ionic Bonding (metals + nonmetals): strongest of all intramolecular forces.

Metallic Bonding (metallic cations + delocalized electrons): similar to ionic but only involves a single type of metal atom.

Covalent Bonding (nonmetals + nonmetals): weakest of intramolecular forces

Intermolecular Forces:
attraction between two different molecules (weaker than intramolecular)

Dispersion or London Forces (temporary dipole):
weak attractive forces that exist between all particles as a result of temporary shifts in the density of electron clouds when atoms or molecules come into close proximity to one another. Dispersion force is stronger as the size of the particle increases.

Dipole – Dipole Forces:
the attraction between oppositely-charged regions of a polar molecule. The partially positive pole of one polar molecule is attracted (and oriented) towards the partially negative pole of another polar molecule.

Hydrogen Bond:
special type of dipole – dipole attraction. Only occurs in molecules containing hydrogen and either fluorine, oxygen or nitrogen.

The special case of H₂O: why water is more dense as a liquid than as a solid.
H2O is one of only a select few molecules that is more dense in the liquid state than in the solid state. This is because of the Hydrogen bond that exists between the hydrogen atom of one water molecule and the oxygen atom of neighboring water molecules. When liquid H2O freezes, the Hydrogen bond between the hydrogen and the oxygen allows each molecule to remain far enough apart to effectively cause solid water to actually expand from its liquid state (causing a lowering of its density). Which is why ice floats in water!

The special case of H₂O: why water is more dense than substances that have a similar molar mass
Water, methane, and ammonia have very similar molar masses. But methane and ammonia are both gases at room temperature, while water is a liquid at room temperature. This is because water has an extra attractive force holding its molecules together that methane and ammonia do NOT have – the Hydrogen bond that exists between the hydrogen atom of one water molecule and the oxygen atom of neighboring water molecules.

12.3 Liquids & Solids
Contrast the arrangement of particles in liquids and solids by using the kinetic – molecular theory

Properties of Liquids:
(1) Volume is relatively fixed (can’t be compressed very much even at high pressure)
(2) Position of individual particles are NOT fixed (limited motion) but have much less freedom of movement than a gas.
(3) At S.T.P. (standard temperature of 25°C and pressure of 1 atm.) liquids are much denser than gases due to stronger intermolecular forces.
(4) Both gases and liquids are FLUIDS, because they can flow and will diffuse (spread from areas of high concentration to areas of low concentration).But liquids are much LESS fluid and diffuse much less freely than gases.

Special properties of liquids:
(1)Viscocity
(2) Surface Tension

Viscocity – measure of the resistance to flow
Viscocity increases with:
a. the strength of the intermolecular force (honey and glycerol have greater intermolecular forces than water, therefore they are more VISCOUS than water).
b. a decrease in temperature (cold maple syrup flows less easily that warm maple syrup because the lower temperature makes a liquid more VISCOUS).
c. increased size of a molecule.
d. an increase in the complexity of a molecule’s shape.

Surface Tension – a measure of the inward pull on particles at the surface of a liquid by the particles in the interior of the liquid
Surface tension allows small objects (including insects and spiders) to be able to walk or float on the surface of water.

Cohesion – the force of attraction between identical molecules
Adhesion – the force of attraction between molecules that are different.

Miniscus – the concave-curved surface of a liquid when placed in a narrow glass cylinder
a menuscus is produced when the adhesive force between the graduated cylinder’s glass walls and the water molecules is greater than the cohesive force between the water molecules with each other.

Capillary Action – is the ability of a liquid to flow in narrow spaces without the assistance of, and in opposition to external forces like gravity.
Occurs because of inter-molecular attractive forces between the liquid and solid surrounding surfaces; the combination of surface tension (which is caused by cohesion within the liquid) and adhesive forces between the liquid and container act to lift the liquid.

Examples of capillary action:
Examples:
1. water is drawn from the roots of a tree to its leaves through capillary action.
2.Paper towels absorb liquid through capillary action, allowing a fluid to be transferred from a surface to the towel.
3. The small pores of a sponge act as small capillaries, causing it to absorb a comparatively large amount of fluid.
4. Some modern sport and exercise fabrics use capillary action to “wick” sweat away from the skin.
5. Capillary action allows trees to move water from its roots to its leaves.
6. Capillary action allows animals to pump blood throughout its body.

Crystalline Solids –
a solid whose atoms, ions, or molecules are arranged in an orderly, geometric structure.

5 Types of Crystalline Solids:
1. atomic – made of atoms; soft/very soft, very low melting points, poor conductivity (Group 18 elements)
2. Molecular – molecules; fairly soft w/ low to moderately high m.p., poor conductivity (I₂, H₂O, NH₃, CO₂)
3. Covalent Network – atoms connected by covalent bonds; very hard w/ very high m.p., poor conductivity (diamond and quartz)
4. Ionic – ions; hard, brittle, high m.p., poor conductivity (NaCl, KBr, CaCO₃)
5. Metallic – atoms surrounded by mobile valence electrons; soft to hard, low to very high m.p., malleable and ductile, excellent conductivity (all metallic elements)

Crystal Lattice –
a 3 – dimensional geometric arrangement within a crystalline solid.

Types of Crystal Lattices:
a. simple cubic
b. body-centered
c. face-centered

Unit Cells –
smallest arrangement of atoms in a crystal lattice; 7 categories based on shape

7 Types of Unit Cells (Bravais Lattices)
1. Cubic 2. Tetragonal 3. Orthorhombic 4. Triclinic 5. Hexagonal 6. Rhonbohedral 7. Monoclinic

Amorphous Solids – solids without a crystalline structure
Examples include volcanic glass (obsidian), man-made glass (SiO₂), rubber, and plastic