GPS S8P1. Students will examine the scientific view of the nature of matter. b. Describe the difference between pure substances (elements and compounds) and mixtures. c. Describe the movement of particles in solids, liquids, gases, and plasmas states. d. Distinguish between physical and chemical properties of matter as physical (i.e., density, melting point, boiling point) or chemical (i.e., reactivity, combustibility).

Matter Anything that has mass and takes up space. Makes up everything in the universe.

Kind of Matter Elements –A substance that can’t be broken down into any other substances by chemical or physical means. –Elements are the building blocks of matter. –Each element is made up of tiny particles called atoms.

Compounds –A substance made of two or more elements chemically combined in a specific ration. Examples: water, chalk, carbon dioxide

Mixtures –A material made up of two or more substances – elements, compounds, or both – that can be easily separated by physical means.

Physical Properties Can be observed or measured without changing the matter’s identity.

Examples 1.Thermal conductivity – rate at which a substance transfers heat 2.State – solid, liquid, gas 3.Density – mass per unit volume 4.Solubility – ability of a substance to dissolve in another substance.

5.Ductility – the ability of a substance to be pulled into a wire. Example: copper 6.Malleability – the ability of a substance to e rolled or pounded into thin sheets.

Chemical Properties Matter’s ability to change into new matter that has different properties. 1.Flammability – ability of a substance to burn 2.Reactivity – two or more substance combine and form a new substance

Changes in Matter 1. Physical Change: a change in size, shape, or state of matter (three states of matter are solids, liquids, and gases). May involve energy changes but the kind of substance – the identity of the element or compound – does not change.

Examples: water, Boiling of water (liquid water, ice and steam are just the liquid, solid and gas forms of H 2 O) Freezing of water to form ice Chewing of food Sharpening of a pencil Crystallization of sugar from a sugar solution Melting of gold

2.Chemical change: A change in one substance to another. Example: antacid tablet in a glass of water and the smell in the air after a thunderstorm In some chemical changes, a rapid release of energy – detected as heat, light, and sound – is a clue that changes are occurring.

Clues such as heat, cooling, or the formation of bubbles or solids in a liquid are helpful indicators that a reaction is taking place. However, the only real proof is that a new substance is produced.

A chemical change can be expressed as a chemical equation. The same elements (and same number of atoms of each element) will be present on each side of the equation.

Measuring Matter Mass –The measurement of how much matter it contains. –SI unit for mass is kilogram

Volume –The amount of space that matter occupies. –For rectangular objects Volume = length x width x height Example: 3 cm x 3 cm x 12 cm = 108 cm3 - for objects with irregular shapes, put the object in a graduated cylinder containing water and measure the change in the volume of the water.

Density - the measurement of how much mass is contained in a given volume. - Density = mass/volume

States of Matter A. The Kinetic Theory 1. All matter is composed of small particles (atoms, molecules, or ions). 2. They are in constant, random motion. 3. They constantly collide with each other and with the walls of their container.

B. Phase Properties Particle Properties

Phase Proximity EnergyMotionVolumeShape SolidCloseLittle Vibrational Definite LiquidClose Moderate Rotational DefiniteNot Definite GasesFar apartA lot Transitional Not definite Not Definite Particle Properties

C. Other States 1.Solids with particles in repeating geometric patterns are crystals. Those with particles arranged randomly are amorphous.

2. Plasma a. Hot, ionized gas particles. b. Electrically charged. c. Most common state in universe.

D. Thermal E x p a n s i o n 1. Particles in any state expand when heated (generally). 2. Examples of solids: a. Expansion joints b. Power lines c. Thermostats

4. Mercury and alcohol are liquids that expand in thermometers 5. Air expands when heated (becoming less dense) 6. Water reaches maximum density at about 4 C.

Ice particles are farther apart than liquid water (so it floats).

Changes in State (phase changes) 1. Melting - solid to liquid a. Particles get more kinetic energy and begin rotating around each other. b. There isn’t enough energy to break the inter-particular attractions, so the particles remain close (liquid). c. The energy required to melt a solid is called the heat of fusion.

2. Freezing - liquid to solid a. Particles lose kinetic energy and slow down. b. Attractive forces between particles become stronger than the particles’ motion, so the particles begin merely vibrating in place. c. The amount of heat the particles must lose to turn into a solid is called the heat of fusion.

3. Vaporization - liquid to gas a. Types: 1) Boiling - rapid; gas bubbles are produced throughout. 2) Evaporation - slow; occurs at the surface. b. Liquid particles gain enough kinetic energy to overcome forces between the particles and they begin translational motion; this energy is called the heat of vaporization.

Evaporation is a cooling process. a. Particles in a liquid gain kinetic energy. b. They leave as gas particles (taking the energy away with them). c. This leaves less energy in the liquid, therefore cooling down what is left

4. Condensation - gas to liquid a. Particles lose kinetic energy, slow down, and come closer together. b. Inter-particular forces become strong enough to make particles merely rotate around each other. c. The energy they lose to turn into a liquid is the heat of vaporization.

5. Sublimation - solid to gas or gas to solid a. Dry ice - carbon dioxide b. Iodine c. Frost During phase changes there is no change of temperature.

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