States of Matter: Pure Substances and Mixtures

Module 1 - States of Matter

Key Concepts

• Matter is anything with mass and takes up space (volume)
• The universe contains four major states of matter

Classifications of Matter

• Matter is classified into pure substances and mixtures
• Pure substances have particles of only one kind
• Mixtures contain more than one kind of particle, with each component retaining its properties
• Elements are the building blocks of matter that contain only one type of atom
• Compounds are pure substances with two or more different atoms joined chemically by chemical bonds
• Mixtures can be homogenous
  • Homogeneous mixtures have evenly distributed particles that cannot be easily separated (e.g., sugar dissolved in water)
• Mixtures can be heterogeneous
  • Heterogeneous mixtures have unevenly distributed particles which can be distinguished (e.g., sand and water)

Physical Properties

• Physical properties can be observed without changing the substance's identity
• Examples of physical properties include color, luster, hardness, density, and state at room temperature
• Testing these properties does not change the substance's composition

Chemical Properties

• Chemical properties describe how a substance reacts with others
• These properties are observed when a substance's identity changes
• Flammability, rusting, and reaction with acid are examples
• Testing chemical properties alters the substance's identity

Physical Changes

• Physical changes alter the shape or state of a substance without changing its identity
• Tearing paper, sawing wood, heating an iron nail or water changing state are examples
• Composition of the substance does not vary through these changes

Chemical Changes

• Chemical changes occur when new substances with new properties are formed
• A change in color, fizzing/bubbling, heat release (exothermic), heat absorption (endothermic), or precipitate formation indicates a chemical change
  • Exothermic changes release heat
  • Endothermic changes absorb heat

Four States of Matter

• The four states of matter are solid, liquid, gas, and plasma

Properties of Solid, Liquid, and Gas

• Solids have definite shape and volume, typically are very dense, not easily compressed or diffused.
• Liquids take the shape of the container, have definite volume, usually less dense than solids, not easily compressed but are easily diffused.
• Gases take the shape and volume of the container, are usually much less dense than solids and liquids, easily compressed and diffused.

Diffusion

• Diffusion is the ability of substances to move through other substances
• Gases diffuse most easily, followed by liquids, then solids

Plasma

• Plasma is a gaseous mixture of positive ions and electrons, existing at temperatures over 100 million degrees Celsius
• Plasma is the most abundant form of matter in the universe, but the least abundant on Earth
• Examples of plasmas include aurora borealis and stars

Kinetic Molecular Theory

• Chemists use the Kinetic Molecular Theory to explain how matter behaves, especially gases

Main Points of Kinetic Molecular Theory

• Particles in matter are very small and have spaces between them
• Gases have much larger spaces between particles compared to the particles themselves
• Particles are in constant motion, moving in straight lines and changing direction only upon collision
• Forces of attraction between particles exist, but are negligible in gases
• Speed of particles increases with increasing temperature, decreasing with decreasing temperature

Elastic Collisions

• Kinetic Molecular Theory states that collisions between particles and their container are elastic (no loss of energy)

Average Kinetic Energy

• Experiments by Clerk Maxwell and Ludwig Boltzmann showed that at a given temperature, particles do not have the same energy amount

Kinetic Energy Distribution Curve

• At any temperature, there is a distribution of energies for the particles
• Temperature indicates the average kinetic energy of the particles, and also increases with temperature
• Each particle has varying kinetic energy and moves at a different speed

Gases (according to Kinetic Molecular Theory)

• Gases are easily compressed, indicating particles are far apart and loosely packed, with densities being low
• Gas particles occupy less than one-tenth of one percent of a gas container's total volume, resulting in +99.9% empty space
• Gases fill any volume, meaning attraction forces between particles are low, and that particles move freely in rapid straight-line motion

Solids (according to Kinetic Molecular Theory)

• Solids cannot be compressed easily due to the particles being close together, resulting in high densities
• Solids have definite shape and volume, leading to the conclusion that particles are tightly held together and only vibrate in place
• Attraction forces between particles in a solid are very high due to intermolecular forces (imfs)
• Solids include crystalline and amorphous
• Crystalline solids (or crystals) have an orderly, geometric, three-dimensional arrangement of molecules, ions, or atoms
• A covalent network solid has atoms covalently bonded in a crystal
• Ionic solids have a regular arrangement of positive and negative ions in a crystal
• Molecular solids have molecules as the lattice points
• Amorphous materials have an irregular arrangement of particles and do not have a definite melting point

Liquids (according to Kinetic Molecular Theory)

• Liquids are not easily compressed because particles are close and densely packed
• Liquids do not have a definite shape
• Attraction forces must be strong enough to hold particles but not strong enough to maintain a constant shape
• Liquid particles are able to slide past each other

Liquid crystals

• Liquid crystals are substances that lose the fixed position of solid particles in only one or two dimensions when heated or exposed to an electric current

Freezing and Melting

• Freezing (or solidification) changes liquid to solid and melting (or fusion) changessolid to liquid

Phase Changes

• Melting (or fusion) occurs when a solid absorbs heat and becomes a liquid, where energy is needed to be able to move the solid particles apart
• Vaporization is when a liquid turns into a gas, so energy is needed for the particles to overcome the imfs
• Evaporation is gas conversion, happening on the surface of the liquid
• Boiling occurs throughout liquid when vaporization occurs
• Condensation (or liquefaction) is when a gas turns to a liquid after the particles get closer
• Freezing (or solidification) when a liquid turns to a solid where energy is released
• Sublimation is a solid changing directly to a gas (e.g., dry ice) and deposition is when a gas changes to a solid

Freezing Point and Melting Point Definitions

• Freezing and melting point of a substance is the temperature at which a substance freezes, or melts/fuses
• Freezing: as energy is lost, particles slow down, intermolecular forces increase, they then arrange in repeating patterns in a crystal
• Melting: intermolecular forces are slowly overcome with the substance slowly becoming molecules as more heat is absorbed

Heating Curves

• The temperature remains constant during melting until all solid has disappeared
• The constant temperature on the graph of heating is called a plateau and indicates the melting point
• Energy overcomes intermolecular forces instead of increasing temperature

Cooling Curves

• Cooling curve is the time versus temperature graph showing cooling of substance
• Plateau occurs as the liquid cools, the particles move closer, and organize so that they can create crystal
• Energy is then released as particles organize so that the temperature remains constant during crystallization as temperatures then start to to fall

Intermolecular Forces in Freezing and Melting

• Greater intermolecular forces must be overcome for a substance to melt
• Greater temperatures and more energy indicates the amount a substance melts or freezes

Important Notes about Melting

• A normal melting point is temperature of a substance melting or freezing at standard pressure
• Mass increases show increase melting point comparing similar covalent as a result in energy to change states than the compounds
• Ionic compounds typically have higher melting points when comparing molecular
• lons held together in a crystal lattice by electrostatic forces
• The melting point of sodium chloride is 801°C

Special things to know about water

• Stronger attraction called hydrogen-bonding holds the water molecules together
• Higher melting point than carbon monoxide, ammonia with similar masses
• Water molecules arrange in six-sided crystals, being less dense than the liquid form
• Melting is endothermic, while freezing is exothermic

Applications

• Freezing is an exothermic process that releases heat when freezing
• Farmers spray crops with water for heat release to protect plants from freezing during the coldest part of the night

Evaporation

• Evaporation is defined as changing a liquid to a gas on the surface (at a temperature below boiling point)
• Particles will have enough energy in the surface to overcome between liquid particles

Cooling

• Particles evaporate with most energy and have lower energy
• Kinetic energy and temperature are connected, lowering leads to lowering temperature
• This is named evaporative cooling

Applications

• Ever wonder why sweating cools your body or you feel cold climbing out of the shower
• Water absorbs heat and evaporates
• Standing is cooler with breeze during a hot day as the water molecules absorb heat and evaporates and air prevents being saturated with water vapor
• Early natives prevent foods from spoiling by putting into clay pots soak in water to have heat absorbed and have the food left cool

Vaporizing

• Vaporizing occurs from liquid to gas
• Condensing is change going to liquid

Key notes for boiling and condensing

• Boiling is liquid to a gas and boiling point is the temperature which the substance boils
• Boiling happens until vapour bubbles form going to the top and being released
• Temperature remains constant through the state and water that boils remain at 100C
• Condensation is the change of state with gas becoming liquid with the particles becoming close

Gas and cooling

• As gas particles cool, the particles start to have movement to together
• Intermolcular forces start to hold and particles move to liquid phase to require energy

Boiling point

• Heat supplied to stove overcomes imf rather than increase
• When water particles vaporize, the gas increases

Boiling points and relation to other aspects

• Temperature of water will not change as long as pressure remains constant which then shows constant temperatiure on the graph which shows boiling point of liquid
• Strongeter the force of attravtion between the particles that increases need for energy
• Molecular mass increases with an increased boiling point
• Ionic compounds have higher boiling points due to high electrostatic forces between the positively and negatively charged oils

Boiling point of water

• Has a special imf of attraction and water molecules that are pulled togehter are callled hydrogen bonding
• Water hydrogen compunds with group v11 elements can have a boiling poitn near -80c

Atmospheric pressure

• Atmospheric Pressure is defined as the force exerted on a surface by the air above it as gravity pulls it to Earth.
• Atmospheric Pressure = Force / Area, with standard units of pressure being Pascals (Pa)
• The measurement is done via Barometer and Manometer devices

Vapour pressure vs temperature

• As temperature increases, amount of vapour increases
• Rate of evaporation increases which then increases vapor particles as then the pressure increases

Vapor pressure and boiling point

• When vaporiziation occurs, the liquid temo remains the same
• Boiling occurrs a substance vapor pressure is equal to the atmospheric pressure