General Chemistry 2

Solid

– is formed when the temperature of a liquid is low and the pressure is sufficiently high causing the particles to come very close to one another.
– they are rigid.
– their particles hardly diffuse.

Nature of Solids
1. Crystalline Solids
2. Amorphous Solids

Crystalline Solids

– atoms, ions, or molecules are arranged in well-defined arrangement
– having flat surface and sharp edges
– example: gems, salts, sugar and ice.

* Ionic Crystalline Solids
* Molecular Crystalline Solids
* Covalent Crystalline Solids
* Metallic Crystalline Solids

Types of Crystalline Solids (Based on bonds)

Ionic Crystalline Solids

– composed of (+) and (-) ions

Molecular Crystalline Solids

– composed of atoms and molecules

Covalent Crystalline Solids

– atoms connected in a network of covalent molecules

Diamond

– all carbons connected by covalent bonds; one of the strongest materials.

Metallic Crystalline Solids

– composed of atoms and molecules

Unit Cell

– the smallest portion of the crystal which shows the complete pattern of its particles.

Crystal Lattice

– the repetition of unit cells in all directions.

Graphite

– sheets of covalent networks of carbon weakly bonded to another sheet (used in pencils).

Amorphous Solids

– from the Greek word for “without form”

Amorphous Solids

– solid particles which do not have orderly structures. Opposite of crystalline.

Amorphous Solids

– they have poorly defined shapes.

Amorphous Solids

– are rigid, but they lack repeated periodicity or long-range order in their structure.

Amorphous Solids

– examples include thin film lubricants, metallic glasses, polymers, and gels.

– they lack regular arrangement of atoms.

Amorphous Solids

– examples are charcoal, powder and glass.

Amorphous Solids

– transparent fusion of inorganic materials (mainly SiO2) that have been cooled without crystallizing.
– behaves more as a liquid than a solid.

Glass

is a familiar and important amorphous solid.

It can be noted that as temperature of crystalline solid is increased (energy is produced), the particles vibrate back and forth about its lattice point. The crystal becomes less ordered. The heat added increases the kinetic motion of the particles. Until the crystalline structure is completely destroyed by the vibrations of the particles, melting is achieved.

How does melting happens?

Metallic Crystalline Solids

– soft to hard, low to high melting point, malleable, ductile and good thermal and electrical conduction

Covalent Crystalline Solids

– held together by covalent bonds

Molecular Crystalline Solids

– composed of atoms and molecules

Ionic Crystalline Solids

– composed of (+) and (-) ions

Metallic Crystalline Solids

– all metallic elements: Cu, Na, Zn, Fe, and Al.

Metallic Crystalline Solids

– metals to metals

Covalent Crystalline Solids

– very hard, very high melting point and often poor thermal and electrical conductivity.
– examples: Plastics, Allotropes of carbon, silicon carbide, diamond, graphite, Quartz, Sand, Glass (SiO2).

Molecular Crystalline Solids

– held together by: H-Bond, Dipole-dipole, and London Dispersion forces (intermolecular)

Molecular Crystalline Solids

– soft, low to moderate melting point and poor thermal and electrical conductivity
– examples: (diatomic gases) CH4, C12H22O11, CO2, H2O (Water), and Br2. Ice, Dry Ice, Sugar Crystals.

Ionic Crystalline Solids

– held by electrostatic attractions (ionic bonding)
– they are hard, brittle and poor
– electrical and thermal conduction
– example: NaCl

Phase

– a homogenous part of a system that is separated from the rest of the system by a well-defined boundary.

Phase change

– transition from one phase to another
– caused by the removal or addition of energy.
– energy involved is usually in the form of heat.

• Melting (Fusion) – solid to liquid
• Sublimation – sold to gas
• Vaporization – liquid to gas

As the temperature goes down: (gain energy/heat)

• Freezing – liquid to solid
• Deposition – gas to solid
• Condensation – gas to liquid

As the temperature goes up: (lose energy/heat)

Vaporization or Evaporation

Liquid - Vapor Phase Transition

Vaporization or Evaporation

- the process in which liquid particles with sufficient energy escape from liquid surface and enter the gas phase.

Boiling point

– the temperature at which the vapor pressure of liquid equals the atmospheric pressure. The boiling temperature varies with the external pressure.

Molar heat of vaporization (∆Hvap)

– it is the amount of heat required to vaporize one mole of a substance at its boiling point. It is usually expressed in kJ/mol.

Molar heat of vaporization (∆Hvap)

is dependent on the strength of intermolecular forces. High IMF = High Molar heat of vaporization

Condensation

Vapor - Liquid Phase Transition

Condensation

– opposite of vaporization; process when a molecule in the gas phase hits the liquid's surface, it can transfer some of its energy to the other liquid particles and remain in liquid phase.

True

The energy transferred out of the system upon condensation is equal to the energy transferred into the system upon vaporization.

negative

The gained energy from vaporization is going to go back, which means the condensation has _______ values.

True

The enthalpy change required to completely vaporize 1 mol of liquid water once it has reached the boiling point of 100ºC at 1 bar (the molar vaporization) and the enthalpy change when 1 mol of water vapor at 100ºC condenses back to liquid water at 100ºC (the molar enthalpy of condensation) have the same values but opposite signs: ΔHºvap = – ΔHº cond

• Cooling the sample.
• Applying pressure on it.

The condensation (liquefaction) process can be achieved either by:

Critical temperature (Tc)

– the temperature above which a gas cannot be liquefied by applying pressure.

Critical pressure (Pc)

– the pressure that must be applied to liquefy a gas at Tc.

Supercritical fluid

– the fluid that exists above Tc and Pc.

Freezing or Solidification

Liquid - Solid Phase Transition

Freezing or Solidification

– liquid becomes a solid if heat energy is removed from it. When the temperature is lowered, the motion of the particles in the liquid slows down. The attractive forces between particles can now overcome the motion of those particles with low kinetic energy. As a result, these particles are held together in fixed position in a lattice structure.

Freezing

– transformation of liquid to solid.

Melting (fusion)

Solid – Liquid Phase Transition

Melting (fusion)

– opposite of freezing. (Solid to liquid)

Melting

– is the process when the temperature of the hot solid increased, the vibrations of the particles become more violent. Finally, at a certain temperature, the vibrations of the particles become violent that the particles in the solid have acquired sufficient kinetic energy to overcome the attractive forces between the particles. The particles then break away from one another and move away from their original positions, changing phase from solid to liquid. Melting (fusion) happens.

Melting point of solid (or freezing point of liquid)

– temperature at which the solid and liquid phases coexist in equilibrium.

Dynamic equilibrium

in which the forward and reverse processes are occurring at the same rate.

Molar heat of fusion (∆Hfus)

– it is the amount of heat required to melt one mole of a solid. It is usually expressed in kJ/mol. ΔHºfus = 6.02 kJ/mol

Sublimation

Solid - Vapor Phase Transition

Sublimation

a process in which the substance goes directly from the solid to the gaseous states.

Deposition

Vapor - Solid Phase Transition

Deposition

: a process in which the substance goes directly from the gas to solid state. Reverse of sublimation. Values are negative.

Deposition

When the gas particles lose energy, and changes directly into solid.

Molar heat of sublimation (∆Hsub)

- it is the amount of energy required to sublime one mole of solid usually in kJ/mol.
∆Hsub = ∆Hfus + ∆Hvap

Heating Curve

- the changes in the state of a matter can be represented by a heating curve, which is a plot of temperature vs. time for a process when an energy is added in a constant rate to a specific amount of a substance. Plotting to see the transition.

Stoichiometry

– it is the quantitative relation between the number of moles (and therefore mass) of various products and reactants in a chemical reaction.

Avogadro’s Number

- in the SI system, the mole (mol) is defined as the amount of substance containing the same numbers of particles as their atoms in exactly 12 g of carbon – 12 isotopes.

Amedeo Avogadro

he originated the Avogadro's Number

6.02214076 × 10^23 or 6.02 × 10^23

one mole of a substance is equivalent to the Avogadro’s number of particles:

Molar Mass

- is the mass in grams of one mole of a substance (g/mol)

Mole (n = Mol)

- amount of substance.

X × (1 mol of X)/(MM of X) × (6.02 ×10²^3atoms or molecules)/(1 mol of X)

X × (1 mol of X)/(MM of X) × (1 mol of X)/(6.02 ×10²^3atoms or molecules)

Avogadro's Number Formula

Molar mass

We can calculate the mole & mass of a substance using ______

Mixtures

- composed of two or more substances, either elements or compounds combined physically in variable proportion.
- components retain their properties and can be separated by physical means.

1. Homogeneous Mixture
2. Heterogeneous Mixture

Types of Mixtures

Homogeneous Mixture

– uniform distribution of particles (air, steel, wine, rain)

Solutions

- homogeneous mixture of molecules, ions, or atoms of two or more different substances in a single phase.

a. Solvent
b. Solute

Solutions is composed of:

Solvent

– dissolving medium

Solute

– substance to be dissolved (ex. sugar/salt in water, air, brass, soda)

Solvation

– is the process of surrounding solute particles with solvent particles to form a solution. Hindi nadidisolve.

Hydration

– process where solute is dissolved in water.

• Particle Size
• Stirring/ Agitation
• Temperature (Heat)

Factors that Affect the Rate of Formation of a Solution

Particle Size

– small particles dissolved rapidly than larger ones as a result of greater surface that is exposed to solvent.

Stirring/ Agitation

– stirring or shaking increases rate of dissolving.

Temperature (Heat)

– increase in temperature increases the kinetic energy of the molecules resulting to an increase in rate of dissolving.

Saturated Solutions

- solutions can be described as saturated, unsaturated or supersaturated.

Saturated

– a solution that contains the max. amount of dissolved solute at a given temperature.

Unsaturated

– a solution that contains less solute than a saturated solution at a given temperature.

Supersaturated

– a solution that contains more solute then the usual maximum amount.

Solubility

- refers to the amount of solute that dissolves in a given amount of solvent to form a saturated solution at a specified temperature.

Insoluble

a substance that does not dissolve in a solvent is said to be ____________

Miscible

- when a liquid dissolve in another liquid it is said to be miscible.

Immiscible

is when two liquids don’t dissolve and form layers.

• Nature of Solute & Solvent
• Temperature
• Pressure

Factors Affecting Solubility

Heterogeneous Mixture

– non-uniform distribution of particles (ice water, cereal in milk, soil, oil and vinegar)

Nature of Solute & Solvent

“Likes dissolves like”
Solutes tend to be more soluble in solvents of their kind.
Polar in polar, nonpolar in nonpolar.

Temperature

Solubility of a solid in liquid increases with temperature.
There are some solid whose solubility remains constant.
Gas is less soluble in liquid if temperature is increased.

Pressure

Change in pressure have little effect on dissolving solids.
Increase in pressure will increase the solubility of gases in liquids.