Chapter 2 Classification of Matter

Classification of Matter Concept Map

Pure Substances Composition Constant Element (one kind of atom) Compound (two or more atoms chemically combined)

Mixtures of Matter Mixture- a combination of two or more pure substances in which each pure substance retains its individual chemical properties. Composition is variable - (the number of mixtures that can be created by combining substances is infinite)

Types of Mixtures Homogeneous Mixtures Heterogeneous Mixtures Has a constant composition throughout Also called a “solution” (i.e., salt water) Heterogeneous Mixtures Do not blend smoothly throughout and in which the individual substances remain distinct. Examples: mixture of sand and water, fresh-squeezed orange juice, pizza

Types of Solutions Solid-solid (alloy-steel) Solid-liquid (sugar in water) Liquid-liquid (vinegar) Liquid-gas (water vapor in air) Gas-liquid (carbonated drinks) Gas-gas (air is a mixture of nitrogen, oxygen, argon)

Types of Heterogeneous Mixtures Colloid- contains tiny particles that do not settle (i.e., milk, gelatin) Suspension- contains larger particles that eventually settle out. Particles have to be re-suspended. (i.e., chocolate milk, orange juice)

Techniques for Separating Mixtures Filtration- uses a porous barrier to separate a solid from a liquid Distillation- A mixture is heated until the substance with the lowest boiling point boils to a vapor that can then be condensed into a liquid and collected. Crystallization - separation technique that results in the formation of pure solid particles of a substance from a solution containing the dissolved substance. Chromatography- separates the components of a mixture (called the mobile phase) on the basis of the tendency of each to travel or be drawn across the surface of another material (called the stationary phase). A magnet can be used to separate magnetic particles from others (Ex. separate sand and iron filings using a magnet to extract the iron based on physical property of magnetism)

Techniques for Separating Mixtures Filtration- uses a porous barrier to separate a solid from a liquid Distillation- A mixture is heated until the substance with the lowest boiling point boils to a vapor that can then be condensed into a liquid and collected. Crystallization - separation technique that results in the formation of pure solid particles of a substance from a solution containing the dissolved substance. Chromatography- separates the components of a mixture (called the mobile phase) on the basis of the tendency of each to travel or be drawn across the surface of another material (called the stationary phase). A magnet can be used to separate magnetic particles from others (Ex. separate sand and iron filings using a magnet to extract the iron based on physical property of magnetism)

Properties & Changes in Matter Physical Properties Physical Changes Chemical Properties Chemical Changes Physical and chemical properties depend on temperature and pressure.

Physical Properties Can be observed or measured without changing the substance’s identity. Ex. Density, color, odor, taste, hardness, melting point, boiling point, solubility, state of matter (s,l,g), temperature

Physical Changes Changes that may dramatically alter the appearance yet leave the composition unchanged Examples: split, bend, crush, grind, changes in state of matter (boil, freeze, condense, vaporize, melt), sharpening a pencil, cutting a sheet of paper, breaking a crystal, crumpling a piece of paper

Chemical Properties The ability of a substance to combine with or change into one or more other substances having different properties. Examples: Flammable Supports combustion the ability of iron to rust when exposed to air Inability of iron to react with nitrogen gas at room temp.

Chemical Changes A process that involves one or more substances changing into new substances Commonly referred to as a chemical reaction Examples: Fermentation of grape juice Rusting Exploding Oxidizing Corroding Tarnishing Burning Rotting

Evidence of a Chemical Reaction Change in color, odor, appearance, production of a gas Formation of a solid called a precipitate – a solid that “falls out” of solution) Changes in energy (production of heat, light, or sound)

Law of Conservation of Mass Mass is neither created nor destroyed during a chemical reaction. Massreactants = Massproducts

Chapter 3 States of Matter & Gas Laws

Kinetic Theory of Matter The idea that all matter is made up of tiny particles in constant motion.

Kinetic Theory of Matter – 4 assumptions: Matter is made of atoms and molecules These atoms and molecules are always in motion. The higher the temperature, the faster the particles move. At the same temperature, more massive particles move slower than less massive ones.

States of Matter Solid Liquid Gas Plasma

Solid Has a definite shape and volume Particles are packed very closely together Particles are incompressible Expand when heated

Types of Solids Crystalline solids- Amorphous solids

Crystalline solids have a definite shape or pattern a solid having a distinct shape because its atoms are arranged in orderly, repeating, three dimensional geometric patterns. Examples: sodium chloride (table salt), sucrose (table sugar), diamonds

Amorphous solids a solid having no form or shape; lacks an ordered internal structure sometimes called slow-moving liquids: (i.e. wax, plastics, glass) Glasses are supercooled liquids.(No regular pattern when shattered.)

Liquids Have a constant volume, but take the shape of the container they are in Particles are not rigidly held in place and are less closely packed than a solid Thus, a liquid “flows”

Gases Have no definite shape or volume Flow to conform to the shape of its container and fill the entire volume of the container

Gases (cont’d.) Particles can be very far apart, or close together Gases are easily compressed Examples: neon, methane, air

Temperature A measure of the average kinetic energy of the particles.

Standard Temperature and Pressure Is defined as 25˚C and 1 atmosphere of pressure

Temperature Scales Celsius- most desirable for laboratory work (°C) Kelvin- the absolute scale (also the SI unit of temperature) (K) Fahrenheit- the old English scale (never used in the lab) (°F)

Kelvin SI base unit of temperature Absolute temperature scale Absolute zero = zero degrees Kelvin, is the theoretical point at which all particle motion stops. Absolute zero is -273˚C K = ˚C + 273 Do not use degree (°) symbol, just K

Celsius Water freezes at 0˚C and boils at 100˚C ˚C = K- 273 ˚C = 5/9 (˚F-32)

Fahrenheit Will not be used in the science lab Water freezes at 32˚F and boils at 212˚F ˚F = 9/5˚C + 32

Will change with temperature Properties of Water Will change with temperature

WATER (at standard temperature and pressure--STP) Liquid Density is 1.00 g/cm3 Chemically unreactive

WATER (at 0°C) Solid Density is 0.9167 g/cm³ The crystalline structure of solid water makes it less dense than it is in its liquid form, therefore, it floats. Water is unusual because its solid form is less dense than its liquid form. For mmost matter, this is not true

Water (at 100˚C and up) Is a gas Has a density of 0.0006 g/cm3 Reacts readily with nitrogen gas

Kinetic Energy and Temperature The average kinetic energy of the particles of a substance is proportional to the temperature of the substance.

Kinetic Theory of Gases A gas is composed of particles, usually molecules or atoms Gas particles move rapidly in constant random motion.

Changes in the State of Matter Are a result of changes in temperature (the average kinetic energy of the particles) or changes in pressure.

Changes in the State of Matter Melting-Solid to liquid Freezing-Liquid to solid Vaporization-Liquid to gas Sublimation-Solid to gas Condensation-Gas to liquid Deposition-Gas to solid

Changes in the State of Matter are caused by: Changes in temperature (average kinetic energy of the particles that make up matter) Changes in pressure (i.e., decrease the volume at constant temperature causes an increase in pressure. Higher pressure = more particle collisions with the side of the container, and thus, temperature increases, causing a change in physical state)

Vapor Refers to the gaseous state of a substance that is solid or liquid at room temperature Example: steam

Heat of Vaporization The amount of energy needed to change a material from a liquid to a gas. For water, it is 2260 kJ/kg.

Condensation The change of a substance from a gaseous state to a liquid state due to a loss of heat. Ex: Water vapor condensing on a cold surface.

Sublimation A change of a solid directly to a gas without going through the liquid state. Ex: dry ice forming “smoke”, solid iodine

Heat of Fusion The amount of energy needed to change a material from the solid state to the liquid state. For water, the heat of fusion us 334 kJ/kg.

Evaporation or Vaporization The gradual change of a substance from a liquid to a gaseous state at temperatures below the boiling point.

GAS LAWS

Biography of Robert Boyle Boyle’s Law P1V1 = P2V2 When temperature is held constant, volume and pressure are inversely proportional. Biography of Robert Boyle (1627-1691) Image URL-- http://www.kingscollections.org/exhibitions/specialcollections/to-scrutinize-nature/boyle-and-hooke/boyles-air-pump

Boyle’s Law

Gay-Lussac’s Law The pressure of a gas increases as the temperature increases, if the volume of the gas does not change. The pressure decreases as temperature decreases. Biography of Gay-Lussac

Gay-Lussac’s Law

Charles’ Law V1 = V2 T1 T2 When pressure is held constant, volume and temperature and are directly proportional. Charles' Biography Image Credit: http://upload.wikimedia.org/wikipedia/commons/thumb/9/98/Jacques_Alexandre_C%C3%A9sar_Charles.jpg/200px-Jacques_Alexandre_C%C3%A9sar_Charles.jpg

Charles’ Law