1 MATTER States of Matter Gas - no fixed volume or shape - easily compressed or expanded Examples: air, neon, acetylene Liquid - has fixed volume - shape of liquid is shape of container - not easily compressed Examples: water, gasoline, blood Solid - has fixed volume and shape - not easily compressed Examples: salt, steel, wool Plasma - the fourth state of matter where some electrons are separated from the rest of the matter - a gas of charges created from high voltage. Examples: sun, neon light CLASSIFICATIONS OF MATTER I. Homogeneous – same composition throughout sample A. Pure substances 1. Elements - cannot be decomposed into simpler substances - building blocks of all matter - only 115 elements, listed in periodic table - individual constituents are atoms Examples: oxygen, iron, helium 2. Compounds - made of two or more elements - millions of compounds - individual constituents are molecules or ions - atoms (elements) bind together to form molecules (compounds) Examples: sugar, water, rust B. Solutions - perfect blending of two or more pure substances with each other - no region of matter has only one type of molecule - examples a) air N2 in O2 b) sea water salt in water c) brass Copper (Cu) in Zinc (Zn) d) soda pop CO2 and sugar in water e) cocktail alcohol dissolved in water f) millions of others possible 2 II. Heterogeneous -must be imperfect mixture -properties are different throughout sample -examples a) sand b) wood c) meat d) potato e) butter PROPERTIES OF MATTER Physical properties – identity of substance does not change when measured intensive properties – independent of amount of substance – Examples: color, density, melting point, temperature extensive properties – dependent on amount of substance – Examples: mass, volume, energy Chemical properties – how substance reacts with other substances Consider how wood changes when we mix it with other substances - When we heat it with oxygen (burning), the wood emits smoke (carbon dioxide, water vapor, etc.) and leaves behind charcoal and ash (carbon). - When we heat it with lead, no chemical change occurs. (Lead melts on surface of wood.) - When we put a drop of concentrated sulfuric acid on it, smoke is released and smells nasty (hydrogen sulfide) and wood turns brown. - When we put a drop of water on it, no chemical change occurs. (It only gets wet.) CHANGES IN MATTER Physical change – composition remains the same - ice melting - adjustable tint sunglasses - dissolving sugar in water Gas g ltin me zing e fre de po su s bli ition ma tio n Phase changes - when one phase of matter changes to another - a phase change is a physical change Solid evaporation condensation Liquid 3 Melting point – temperature where solid melts - same as freezing point Tm(H2O) = 0 C Tm(O2) = -223 C Tm(Fe) = 1536 C Boiling point – temperature where liquid boils - same as condensation point Tb(H2O) = 100 C Tb(N2) = -196 C Tb(ethanol) = 78 C Chemical change – composition is altered Examples: - burning of paper - baking cake - cooking egg NOTES ON DENSITY Riddle: Which is heavier, a pound of feathers or a pound of lead? Answer: Neither, they both weigh one pound. Question: Which has more volume, a pound of lead or a pound of feathers? Answer: A pound of feathers Since feathers take more space than lead, the feathers are less dense than lead. Definition of density density mass volume - in chemistry, density often has the units of g/mL or g/cm3. Example: What is the density of motor oil? A 755 mL sample of motor oil weighs 733 g. Density can be considered a conversion factor between mass and volume. d(H2O) = 1.00 g/mL 1.00 g = 1.00 mL 100 . g H2O mL H 2 O or 100 . mL H 2O g H 2O d(Hg) = 13.6 g/mL 13.6 g = 1.00 mL 13.6g Hg 1.00mL Hg or mL Hg 13.6g Hg 4 Example: The density of kerosene is 0.82 g/mL. How much mass does 250 mL contain? Example: How much space does 4.76 kg of lead occupy? The density of lead is 11.3 g/mL. Note: we will have to include a conversion from kg to g in our calculation. ENERGY IN CHEMICAL REACTIONS Energy is the capacity to do work (to move an object through a distance). Forms of Energy Kinetic – energy which causes motion - A moving ball has kinetic energy. - Electrical energy is kinetic energy. (Electrical charge is moving.) - Hot water has kinetic energy. (Water molecules are moving.) Potential – energy which has ability (i. e. potential) to cause motion - Apple on a branch has potential energy. (Gravity is pulling on it. When it drops, it has motion.) - Battery has potential energy. (When circuit is complete, electricity will flow.) TYPES OF ENERGY Light Energy – as in energy from a light bulb Heat Energy – as in energy from a fire Electrical Energy – as in putting your finger in an electrical socket Ouch! Mechanical Energy – as in energy from an engine or water wheel Chemical Energy – as in energy from burning gasoline (chemical energy is converted to heat and light) - Chemical energy is stored inside a substance until a substance undergoes a chemical change. Then it can be released or sometimes absorbed. - Chemical energy is stored in chemical bonds between atoms. (More about bonds later) Energy exchange in chemical reactions Exothermic - energy is released when reaction occurs - most reactions are exothermic Endothermic - energy is absorbed when reaction occurs - example is an instant cold pack. 5 Calculation of heat energy Units of energy Joule – SI unit (J) 1 kg m 1 m Definition: 1J 1 kg s s2 2 2 A 0.250 kg baseball going 60 mi/hr (27 m/s) has 90 J of energy (kinetic). The amount of heat energy needed to start 1 quart of water boiling is about 10,000 J. Calorie – traditional (English) unit (cal) Definition a) original – amount of heat needed to raise 1 gram of water 1 C b) modern – 1 cal = 4.184 J (exactly) Note: 1 dietary calorie (Cal) = 1 kilocalorie (kcal) HEAT CHANGES AND SPECIFIC HEAT Heat energy within a substance is proportional to both mass and temperature difference. - when mass increases, more heat is needed to raise temperature to same level. - when the temperature difference increases (i. e. to make temperature hotter), the amount of heat needed increases Heat Equation q = m c T q – heat m – mass T – temperature difference = Tfinal - Tinitial c – specific heat Specific Heat - shows how much heat is needed to raise temperature of a substance - compare specific heats of Cu and H2O. c(Cu) = 0.3844 J/gC c(H2O) = 4.184 J/gC - specific heats show that copper heats and cools quickly while water heats and cools slowly. 6 Example: How much heat is lost when 3780 g of H2O (1 gallon) is cooled from 61.4 C to 30.0 C? c(H2O) = 4.184 J/gC Negative sign means heat is lost. (More about this later.) Example: 98.3 g of Aluminum at 25 C is given 15,500 J of heat. What is the final temperature of the metal? c(Al) = 0.895 J/gC LAW OF CONSERVATION OF MASS In a chemical reaction, mass is not created or destroyed. We use the law of conservation of mass later to account for all the mass in a chemical reaction. LAW OF CONSERVATION OF ENERGY In a chemical reaction, energy is not created or destroyed. Question: If this is true, then where does energy from burning gasoline come from? Answer: From the chemical energy inside the gasoline