Matter & Energy

1. Matter & Energy Honors Chemistry

2. I. Science • Science is a body of knowledge collected by scientists over many years & the methods used to obtain the knowledge • Chemistry is the study of the composition, structure and properties of matter & the changes it undergoes • Chemical = any substance that has a definite composition • It is through the analysis of much information on matter that we can solve problems & answer question • What, how much, how it can be changed, & how fast

3. II. States of Matter Solid Gas Liquid

5. B. Properties • Physical property - can be observed without changing the identity of the substance • Intensive is independent of amount • mp, bp, density, conducts electricity/heat, temp • Extensive is dependent of amount • mass, volume, amount of energy, heat • Chemical property – relates to a substance’s ability to undergo changes that transform it into different substances

6. C. Changes • Physical – does not involve a change in the identity of a substance; may change the appearance • Chemical – one or more substances are converted into different substances with different properties • Alters identity of substance. Produces a new substance • The new substance (product) has different properties than the beginning materials (reactants).

7. Signs of a Chemical Change • Color • Gas (change in odor) • Precipitate • Change in temperature (may include light) • Endothermic vs. Exothermic reactions • Note: all chemical and physical changes involve energy

8. What is the 3rd change? • Nuclear Change - changes the composition of the atom’s nucleus • tremendous amount of energy involved • Fission vs. Fusion • Radioactive decay • Where is uranium? • Ground • Refined for nuclear power plants

9. Radioactive Decay

10. Conservation Matter and Energy • Cannot be created or destroyed, only changes form in a chemical or physical change • Burning magnesium • Burn Mg – heavier product, why? • Mg + O2 MgO • Heater – electrical energy to heat energy

11. MATTER Anything that has mass and volume Pure Substances Fixed composition; characteristic chemical & phys properties Mixtures Blend of 2/more kinds of matter, each of which retains its own identity & properties Elements Periodic table; smallest particle to retain all properties - atom Compounds Can be decomposed into 2/more simpler compounds or elements Homogeneous (Solution) Uniform in composition - same proportion of components throughout Heterogeneous Not uniform throughout E. Classification

12. The Periodic Table

13. Metals • Location: to the left of the staircase • At room temp, all are solid except for Hg • Ductile - can be drawn out into thin wires • Malleable - can be hammered into thin sheets • Luster (A.K.A. Shininess) • Good conductors of heat and electricity • High density • High melting points • Ion formation – tend to lose electrons resulting in positive charges

14. Nonmetals • Location: to the right of the staircase • At room temperature, they are solids, liquids, or gases. • Dull – no luster • Insulators of heat and electricity. • Brittle - Neither malleable or ductile • Lower bp and mp than metals. • Ion formation – tend to gain electrons resulting in negative charges

15. Metalloids (Semimetals) • Located between the metals and nonmetals, ALONG the staircase . • Have properties of both metals and nonmetals. • There are 7 metalloids in the periodic table: Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), Tellurium (Te), & Astatine (At).

16. Check for Understanding • List the nonmetals in the 5th period. • Iodine and Xenon • Metalloid(s) in group 5A (15)? • Arsenic and Antimony • Liquid metal? Liquid nonmetal? • Mercury Bromine • Symbol for the ion in group 6A and period 3? • S-2

17. Compounds • 2 or more elements chemically combined through covalent or ionic bonding • Examples: • Na and Cl2 react to form NaCl • C and O2 react to form - CO2 • How many atoms in NH4Cl? • How many H atoms (NH4)2SO4 • How many H atoms in 5 (NH4)2SO4

18. Separation Techniques • Heterogeneous Mixtures • Filtration: Pour liquid through filter paper to collect solid • Centrifuge: separates solid-liquid mixtures • Decanting

19. Separation Techniques • Homogeneous Mixtures • Crystallization: evaporate liquid and solid will crystallize • Chromatography – used to separate pigments of ink on a strip of paper. • Distillation

20. Distillation - separation of a solution based on differences in boiling point

21. Compounds • Decomposition – compound breaks down into two or more simpler compounds or elements

22. Electrolysis - decomposes a compound with electricity

23. Energy Concepts • Thermochemistry: the study of the changes in energy that accompany a chemical reaction and physical changes. • Chemical Reactions involve changes in energy that result from • Bond breaking that requires energy (absorbs) from the surroundings. • Bond making that produces energy (releases) to the surroundings. • Changes in energy result in an energy flow or transfer.

24. Heat vs. Temperature • Heat: (q) is the energy transferred due to changes in temperature. • Temperature (T) is a measure of the average particle motion or the average kinetic energy. • Heat flows spontaneously from a higher to a lower temperature. • Heat vs. Temp Simulation - Eureka

25. Calorimeter Heat is measured in a calorimeter. Changes in temperature are measured in a known quantity of water in an insulated vessel.

26. surroundings Exothermic Reaction (system) surroundings surroundings surroundings Types of Reactions • Exothermic: releases heat into their surroundings. • Heat is a product and temperature of the surroundingsincrease. • This occurs during bond formation.

27. surroundings Endothermic Reaction (system) surroundings surroundings surroundings Types of Reactions • Endothermic: absorbs heat from the surroundings. • Heat acts as a reactant and temperature of the surroundingsdecreases. • This occurs during bond breaking.

28. Exothermic Example: Dissolving calcium chloride in water • Combustion reactions areALWAYSexothermic: C3H8 (g) + 5O2(g) → 3CO2(g) + 4H2O(g) + 2043 kJ • Endothermic Example: 2NH4Cl (s) + Ba(OH)2·8H2O (s) + 63.9 kJ BaCl2 (s) + 2NH3 (g) + 10H2O (l) • Physical states are written – influences the overall energy exchanged. Very specific!

29. Forms of Energy Mechanical, Heat, Chemical, Electrical, Radiant, Sound, Nuclear

30. Changes of State • Energy • Types • Potential energy is the energy of position • As particles move apart, the PE increases • The PE of a gas is greater than the PE of a liquid which in turn is greater than the PE of a solid • During condensation, the PE decreases and energy is released. This is an exothermic change.

31. Changes of State • Kinetic energy is the energy of motion. • Except at 0 K, all particles are in constant motion • Temperature is a measure of the avg KE of the particles in a sample. • When temperature is increased, the KE of the particles increases. • In a liquid, the particles must have a minimum KE (Em) in order to overcome the intermolecular attractions of neighboring particles to escape. • The stronger the intermolecular forces in a liquid, the higher the Em.

32. Heating and Cooling Curvesgraph of temp of a substance • Label the Heating Curve of Water • Evaluate the energy changes that occur during a heating curve. • Hf – heat of fusion: energy needed to melt an amount of a substance at its mp • Hv – heat of vaporization: energy needed to vaporize an amount of a substance at its bp • Hf and Hv Units: J/g or kJ/mol or cal/g • Hf and Hv are physical properties of a substance

33. Heating Curve for Water D E 100 Temperature (ºC) B 0 C A Energy

34. Problems • Calculate the energy (in cal) needed to melt 125.0 g of ice at 0.0°C • How much energy (in kJ) is needed to warm 180.0g of ice at -20.0°C to water at 75.0°C? • If 275.0 g of liquid water at 100.0°C and 475.0 g at 30.0°C of water are mixed in an insulated container, what is the final temperature? 9978 cal 124.1 kJ 55.7°C

35. Physical Properties of Gases: 1. Gases consist of small particles that have mass. These particles are usually molecules, except for the noble gases.

36. Physical Properties of Gases: • Gases have mass. The density is much smaller than solids or liquids, but they have mass. (A full balloon weighs more than an empty one.) • The particles in gases are separated by relatively large distances. Gases can be compressed. It is very easy to reduce the volume of a gas.

37. Unlike liquids, gases completely fill their containers. • The particles in gases are in constant rapid motion (random).

38. Gases can move through each other rapidly - diffusion (ex. food smells and perfume)

40. 7. Gases exert pressure because their particles frequently collide with the walls of their container and each other.

41. Inelastic Collision 8. Collisions of gas particles are elastic. Elastic Collision

42. Gas particles do not slow down when hitting each other or the walls of their container.

44. 9. Gas particles exert no force on one another. Attractive forces are so weak between particles they are assumed to be zero.

45. 10. Temperature of a gas is simply a measure of the average kinetic energy of the gas particles. High temp. = high KE Low temp. = low KE

46. The pressure of a gas depends upon temperature high temp. = more collisions, high pressure low temp. = less collisions, low pressure Low pressure High pressure

47. Boyle’s Law Pressure - Volume Relationship • The pressure & volume of a sample of gas at constant temperature are inversely proportional to each other. Law assumes n is constant. P1V1 = P2V2 Inverse

48. Boyle’s Law