Chapter 3 Matter – Properties & Changes

Introduction: What is matter?

Introduction Matter – anything that has mass and volume  The atom is the building block of all matter

States of Matter What are the 3 main states of matter?

States of Matter Solid  Has definite shape & volume  Particles are tightly packed  Can expand when heated

States of Matter Liquid  Has constant volume but takes the shape of its container  Fluid  Less closely packed particles than solid particles  Can expand when heated

States of Matter Gas  Expands to fill its container & takes the shape of its container  Fluid  Much less closely packed than solid particles  Expands when heated

States of Matter 1) 2) 3)

States of Matter - Plasma Plasma  Does not naturally occur on Earth (except in lightning) Present in the Sun  An ionized gas Sufficient energy is provided to free electrons from atoms or molecules and to allow both ions and electrons to coexist

Phase Changes When a substance freezes, boils, evaporates, or condenses, it undergoes a phase change  The addition or removal of energy is what causes phase changes  This change is also a physical change

Phase Changes

Matter Mixtures Pure Substances Heterogeneous Homogeneous Element Compound

Types of Matter Mixtures 1. Physical combinations of 2 or more substances 2. Components do not chemically react with each other Pure Substances 1. Elements 2. Compounds

Mixtures Mixture – a combination of two or more pure substances in which each pure substance retains its individual chemical properties  2 kinds: Heterogeneous Homogeneous

Heterogeneous Mixtures Mixture that does not blend smoothly throughout – individual substances remain distinct

Homogeneous Mixtures Mixture that has constant composition throughout  Examples: salt water, soda water, brass A.k.a. a solution  Solute in a solvent (salt dissolved in water)

Separating Mixtures Filtration – separates undissolved solids from liquids  Used to separate heterogeneous mixtures

Separating Mixtures Distillation – separates a solution of liquids or a homogeneous mixture

Separating Mixtures Crystallization  Separates dissolved solids from liquids Crystallization Demo & Explanation Crystallization Demo

Separating Mixtures Chromatography – separates components of a solution into its components based on tendency to travel on another surface

Elements Elements – pure substance that cannot by separated into simpler substances by physical or chemical means  Made of only 1 type of atom  All elements are found on the Periodic Table of Elements  Each has a unique name & symbol Ex.) Carbon (C), oxygen (O), platinum (Pt)

Elements All elements have a unique symbol  Each symbol is either 1 or 2 letters One letter = capitalized Two letters = first is capitalized, second is not Symbols are often similar to the name but sometimes they are derived from the Latin name  Ex.) Mercury’s symbol = Hg (from Latin hydrargyrum)

Compounds Compound – combination of 2 or more elements chemically bonded together or combined  Written with a chemical formula Letters indicate the elements in the compound Subscripts indicate the # of each type of atom in the molecule

Compounds The properties of a compound are VERY DIFFERENT from the properties of the individual elements they contain Ex.) Sodium Chloride (NaCl) vs. Sodium & Chlorine Sodium: Sodium in PondSodium in Pond Sodium in 40 Gallon Trashcan

Compounds A molecule is the smallest particle of a compound Compounds have constant composition from one sample to another

Matter Mixtures Pure Substances Heterogeneous Homogeneous Element Compound In Review…

Classify the following as element, compound, heterogeneous mixture, or homogeneous mixture: 1) Bean salad 2) Phosphorous 3) Air 4) Salt water 5) Sugar (C 12 H 22 O 11 ) 6) Copper 7) Alcohol in water 8) Sand (SiO 2 ) 9) Iron filings 10) Rocks in sand 11) Copper sulfate (CuSO 4 ) in oil 12) Carbonated water

Law of Definite Proportions Elements in a compound combine in definite proportions by mass  All samples of a compound have the same proportion by mass of the elements present.

Example – Law of Definite Proportions If 1.0 g of hydrogen reacts completely with 19.0 g of fluorine, what is the percent by mass of hydrogen in the compound that is formed?

Example – Law of Definite Proportions If you have a sample of sucrose (a type of sugar), you will always have 42.2% of carbon, 6.50% of hydrogen, and 51.30% of oxygen, no matter where you find that sample. Therefore, if you are given 50.0 g of sucrose, how many grams of each element will you have?

Law of Multiple Proportions When different compounds are formed by the combination of the same elements, they will combine in small whole number ratios  Ex.) H 2 O & H 2 O 2 ; CuCl & CuCl 2 ; CO & CO 2

Properties of Matter When talking about the properties of matter, there are two types: PHYSICAL properties OR CHEMICAL properties

Physical Properties of Matter Characteristic that can be observed or measured without changing the sample’s composition (or identity)  Ex.) color, boiling point, state of matter, density Physical properties describe appearance or behavior

Physical Properties of Matter Extensive property  Depends on the amount of substance present  Examples: Mass Length Volume

Physical Properties of Matter Intensive property  Independent of the amount of substance present  Examples: Color Density Smell Boiling point

Physical Properties of Matter Intensive or Extensive? 1. Phase 2. Size 3. Odor 4. Flammability 5. Energy 6. Melting point

Chemical Properties of Matter Chemical property – the ability of a substance to combine with or change into one or more other substances  Examples: rusting, flammability, light sensitivity 2 H 2 O 2 (l)  2 H 2 O (l) + O 2 (g)

Observing Properties of Matter When recording physical or chemical changes, note the conditions. Many properties of substances can depend on temperature or pressure!

Physical Changes Physical changes – altering a substance without changing its identity or composition Examples:  Cutting  Breaking  Phase changes (evaporation, condensation, sublimation, etc.)

Chemical Changes Chemical changes - one or more substance changes into a new substance  i.e. a new substance is formed Examples: burning, rotting, rusting 4 Fe + 3 O 2  2 Fe 2 O 3

Evidence of a Chemical Change Properties of new substance formed does not have the same properties as the original substance Signs a chemical change has occurred:  Heat  Cooling  Formation of bubbles (gas produced)  Formation of solids  Color change

Chemical or Physical Change? 1) Butter melting 2) Hot glass cracking when placed in cold water 3) Dissolving sugar in water 4) Burning sugar (like on crème brulee) 5) Wood rotting 6) Burning gasoline in your car 7) Silver tarnishing

Weathering of the Earth – Chemical or Physical Changes? Physical changes include:  The splitting of rocks by freeze thaw cycles does not change the make up of the rock  The cutting through and wearing away of softer rock by the actions of water

Weathering of the Earth – Chemical or Physical Changes? Chemical changes include:  The dissolving of rock by slightly acidic water  Limestone easily dissolves in acidic water When acidic water seeps through limestone and some other minerals, it dissolves some of the material and takes it away.

Conservation of Mass Law of conservation of mass:  Mass is not created or destroyed in a chemical reaction (mass is conserved) Mass products = mass reactants

However, is mass conserved when a log burns in a fire??

By observing a burning log turn to ash, you may be tempted to think that during the chemical change that occurred, some mass was lost. It was not!!! If you gathered all of the gases that formed, soot that was carried away, and water that evaporated out when the log was hot, you would find that there is exactly the same amount of mass after the log was burned as before the log was burned

Conservation of Mass Problems 68.5 grams of H 2 S contains 4.06 g of hydrogen  What mass of sulfur is in 68.5 grams of H 2 S?  What mass of hydrogen is in 200 grams of H 2 S?  What mass of sulfur is in 200 grams of H 2 S?