Chemistry Chapter 13 Christen Rowland Harding Charter Prep

Intermolecular Forces Intermolecular forces are the forces that occur among molecules that cause them to aggregate to form a solid or liquid. If not ionic bonding, collectively they are called Van der Waals forces. Van der Waals forces Dipole-dipole attractions – molecules with dipole moments can attract each other by lining so that the positive and negative ends are close to each other About 1% the strength of ionic or covalent bonding Increasing distance between dipoles equals weaker attraction Hydrogen Bonding – exceptionally strong dipole-dipole attraction between hydrogen covalently bound to a very electronegative atom and another highly electronegative atom of another molecule (such as oxygen, nitrogen, or fluorine) London dispersion forces – are forces of attraction that can occur between nonpolar molecules Weakest of the Van der Waals forces

Intermolecular Forces Dipole-dipole actractionsLondon dispersion forces

Changes of State Continued Changes from solid to liquid and liquid to solid are physical changes The energy required to melt on mole of a substance is called the molar heat of fusion 6.02 kJ/mol for water The energy required to change one mole of a liquid to a vapor is called the molar heat of vaporization 40.6 kJ/mol for water Heating Cooling Curve for Water Conducted at 1 ATM of pressure (Standard Pressure)

Heat Capacity and Specific Heat Heat capacity is the amount of heat required to increase the temperature of an object exactly 1ºC Specific heat is the amount of heat it takes to raise the temperature of 1g of a substance 1ºC Units of specific heat are J/(g·ºC) or cal/(g·ºC)

Molar Heat Sample Problem Calculate the energy (in kJ) required to heat 25 g of liquid water from 25 ºC to 100 ºC and change it to steam at 100 ºC. The specific heat capacity of liquid water is 4.18 J/gºC, and the molar heat of vaporization is 40.6 kJ/mol

Changes of State Sublimation occurs when substance changes from a solid to gas without passing through the liquid state Sublimation occurs in solids with vapor pressures that exceed atmospheric pressure at or near room temperature Iodide (pictured left) and Dry Ice (CO 2 ) are examples of substances that sublime

Phase Diagrams The relationships among the solid, liquid, and vapor states of substance can be represented in a single graph called a phase diagram Notice the triple point of water at 0.61 kPa and ºC Triple point describes the only set of conditions at which all three phases exist in equilibrium with each other

Nature of Liquids: Model Particles in both gases and liquids have kinetic energy Unlike gases, particles making up liquids are attracted to each other, keeping the particles close resulting in the definite volume of the liquid Increased pressure does not affect a liquids volume The interplay between the disruptive motions of particles in a liquid and the attraction among the particles determines the physical properties of liquids

Nature of Liquids: Evaporation The conversion of liquid to gas or vapor is vaporization When vaporization occurs at the surface of a liquid that is not boiling, it is called evaporation During evaporation, only those molecules with a certain minimum kinetic energy can escape from the surface of the liquid Heating liquids increases their kinetic energy causing faster evaporation –Evaporation is an endothermic process which causes the liquid to cool

Nature of Liquids: Vapor Pressure Vapor pressure is a measure of the force exerted by a gas above a liquid In a system at a constant vapor pressure, a dynamic equilibrium exists between the vapor and the liquid. The system is in equilibrium because the rate of evaporation of liquid equals the rate of condensation of vapor pressure LiquidGas Evaporation Condensation An increase in temperature of a contained liquid increases the vapor pressure –Vapor pressure is measured using a manometer

Boiling Point When a liquid is heated to a temperature at which particles throughout the liquid have enough kinetic energy to vaporize, the liquid begins to boil The temperature at which the vapor pressure of the liquid is equal to the external pressure on the liquid is the boiling point Because a liquid boils when its vapor pressure is equal to the external pressure, liquids don’t always boil at the same temperature Lower atmospheric pressure = lower boiling point Higher atmospheric pressure = higher boiling point Because liquid can have various boiling points depending on pressure, the normal boiling point is defined as the boiling of a liquid at a pressure of kPa (atmospheric pressure at sea level)

The Nature of Solids Model for Solids Orderly and fixed arrangement particles When heated to melting point, disruptive vibrations of particles are strong enough to overcome the attractions that hold them in fixed positions resulting in the solid transitioning to a liquid Melting point and freezing point are at the same temperature At melting and freezing points, the liquid and solid phases are in equilibrium SolidLiquid Melting Freezing

Solids: Crystal Structure and Unit Cells Most solid substances are crystalline Crystals have particles arranged in an orderly, repeating, three-dimensional pattern called a crystal lattice The shape of a crystal reflects the arrangement of the particles within the solid Type of bonding between particles in crystals determines melting points Ionic solids have the strongest bonds and highest melting points Held together by the strong forces that exist between oppositely charged ions Molecular Solids compounds have relatively lower melting points Formed by molecules held together by Van der Waals forces Atomic Solids vary considerably A diamond is a special kind of atomic solid called a network solid Some substances decompose rather than melt

Solids: Crystal Systems Seven types of crystal systems that differ in terms of the angles between the faces and in the number of edges of equal length on each face Cubic, Tetragonal, Orthorhombic, Monoclinic, Triclinic, Hexagonal, and Rhombohedral The smallest group of particles within a crystal that retains the geometric shape of the crystal is known as a unit cell

Solids: Allotropes Allotropes are two or more different molecular forms of the same element in the same physical state Carbon, oxygen, phosphorus, sulfur, boron and antimony are the only elements known to have allotropic forms

Bonding in Metals Metals are made up of closely packed cations rather than neutral atoms The valence electrons of metal atoms can be modeled as a sea of electrons The valence electrons are mobile and can drift freely from one part of the metal to another Metallic bonds are the forces of attraction between free-floating valence electrons and positively charged metal ions Metallic bonding of metals can explain three properties of metals Electrical conductivity (electrons entering metal are countered by electrons freely flowing out of metal) Ductility (sea of electrons act as ball bearings around the positively charged metal ions) Malleability (sea of electrons act as ball bearings around the positively charged metal ions) When ionic crystal have a force applied to them, the force tends to push ions of like charge into contact causing them to repel each other that results in the crystal shattering.

Crystalline Structure of Metals Metal atoms in crystals are arranged into very compact and orderly patterns Label the following arrangements of atoms with the correct name Chromium Zinc Gold Each atom has 8 neighbors Each atom has 12 neighbors

Crystalline Structure of Metals Metals with body centered cubic lattice patterns include sodium, potassium, iron, chromium, and tungsten Metals with face centered-centered cubic lattice patterns include copper, silver, gold, aluminum, and lead Metals with hexagonal close-packed structures include magnesium, zinc, and cadmium

Alloys A mixture of two or more elements, at least one of which is a metal, is called a(n) alloy. Alloys are important because their properties are often superior to those of their component elements Alloys tend to be harder and more durable than their constituent metals The most common use for nonferrous alloys is coins A very important allow, steel is important for its: Corrosion resistance Ductility Hardness Toughness

Components of some common alloys Sterling silversilver and copper Brasscopper and zinc Surgical steeliron, chromium, nickel, and molybdenum Cast ironiron and carbon

Alloy Types Interstitial alloys have smaller atoms that fit into the spaces between larger atoms. Substitutional alloys have component atoms that are roughly equal in size Interstitial Substitutional Solvent atom Solute atom Solvent atom Solute atom

Non-Crystalline Solids Amorphous solids lack an ordered internal structure