Intensive Chemistry Day 2: Properties of Matter

1. Intensive ChemistryDay 2: Properties of Matter Katy Johanesen Ph.D. Candidate, USC Department of Earth Sciences

2. This talk has two basic parts: Properties of Matter Physical Reactions

3. Properties of matter • All matter has mass, volume, density, and certain physical and chemical properties • Dictated by the atomic or molecular structure • Matter can be malleable, ductile, sectile, magnetic, conduct heat or electricity, react with water or oxygen… • Color, phase at room temperature, and hardness are also properties

4. Properties of compounds • A compound does not have the same properties as the elements that make it up. • Table salt (NaCl) is edible and dissolves in water • Sodium (Na) and Chlorine (Cl) are both poisonous • Na metal reacts strongly with water (our bodies are full of water) • Cl2 was used as a poison gas in WWI • http://www.youtube.com/watch?v=9bAhCHedVB4

5. Density • Density = mass/volume • Atomic mass of elements in the compound (mass) • Phase/crystal structure of the compound (volume) • CAUTION: weight ≠ mass! • Weight is a force • weight = mass times acceleration due to gravity • Two minerals… use density and the periodic table to determine which is which. • Formulas: PbS and FeS2

6. Phases of matter • Phase = a volume of space where all physical and chemical properties are uniform • Solid, liquid, gas, plasma • Supercritical fluids and degenerate gases are other strange phases of matter that do not naturally exist on Earth’s surface • Form of matter depends on space between molecules and energy of the system • Energy = heat

7. Common phases of matter • Solid = molecules are close together and ordered in a crystal; energy of the system is the lowest • Liquid = molecules are close together but are able to move and switch places; energy in the system is higher than in solid • Gas = molecules are widely separated, move around freely, move at high speeds; energy in the system is the highest of all three • Ideal gas law: PV=NkT • http://www.phy.ntnu.edu.tw/ntnujava/index.php?topic=25 Definitions modified from http://www.enchantedlearning.com/physics/Phasesofmatter.shtml

8. Not so common phases of matter • Plasma = gas composed of positively charged ions + electrons; energy is so high matter becomes ionized; makes up 99.9% mass of the solar system • Ex: the sun, stars, fluorescent lamps, arc welding, lightning, plasma TV • Supercritical fluid = when a liquid or a gas is compressed and heated enough, their densities become equal; no distinction between liquid and gas beyond this pressure and temperature • Found: deep within Earth, other planets • Degenerate gas = gas compressed so hard that atoms are touching; behaves like a solid • Ex: white dwarf stars, neutron stars, metallic hydrogen (Jupiter, Saturn) Definitions modified from http://www.enchantedlearning.com/physics/Phasesofmatter.shtml; background from dailymail.co.uk

9. Solids: crystal structure • Atoms or molecules organize into regular, repeating structures • NaCl is ionically bonded • CaCO3 contains ionic and covalent bonds (CO32- is a covalently bonded ionic compound) • Graphite and Diamond are both covalently bonded, but have different properties due to the crystal structure • Color • Hardness • Crystal shape

10. Solids: crystal structure • Metals and Alloys: + ions bonded by delocalized valence electrons • Ductile: can be drawn into a wire (Cu, Au, Ag) • Malleable: can be hammered into a sheet (Au, Pb, Al) • Sectile: can be cut with a knife (Na, Au, Paper) • Electric conductivity- electrons are free to move through the material

11. Physical Reactions • Reversible • Change in state or phase • Change in shape or crystal structure • Mixture of two or more substances which can be separated again • What are some physical reactions you can think of?

12. Phase transitions Phase Diagram • Matter can change between phases with enough change in energy or pressure • Each element (and compound) can have a different boiling point (ptable.com) Figures modified from Wikipedia

13. Ping-pong ball model of matter • Balls represent atoms or molecules • Shaking the container represents energy (heat) • What would happen if the same number of molecules occupied a larger container?

14. Solutions • Solutions are homogeneous mixtures composed of one phase • Solutes are dissolved in solvents • Solutes and solvents can be any forms of matter • Gas in gas (air, O2 in N2) • Liquid in liquid (vodka) • Solid in solid (steel) • Gas in liquid (soda) • Solid in liquid (salt water) • Liquid in gas (water vapor) • etc. Source: http://commons.wikimedia.org/wiki/File:SaltInWaterSolutionLiquid.jpg

15. Solutions • Solutions are physical, not chemical reactions; can be separated to their constituent parts • Can boil (or freeze!) seawater to remove water from salt • Can let soda go flat to remove carbon dioxide from soda water • Can boil amalgam to remove mercury from gold (used in CA gold rush!) Source: http://commons.wikimedia.org/wiki/File:SaltInWaterSolutionLiquid.jpg

16. Solutions: Acids and bases • Acids and bases are ionic compounds that dissolve and break apart in water to form excess H+ or OH- ions • The strength is based on the concentration of ions • Acids are in lemons, vinegar, coffee, milk… • Bases are in peppers, seawater, bleach, soap… • Acids and bases can be solids, liquids, or gases • most acids and bases we know are solutions in water

17. Brønsted-Lowry theory of Acids and bases • Acids are H+ donors; bases are H+ acceptors • Recall that: H2O = H+ + OH- • Another way of representing this is: • H2O + H2O = H3O+ + OH- • Example: Add HCl (gas, acid) to water: • HCl + H2O = H3O+ + Cl- • In this case, HCl is an acid because it donates the H+ to the water molecule, creating H3O+

18. Brønsted-Lowry theory of Acids and bases • Example: Add NH3 (gas, base) to water: • NH3 + H2O = NH4+ + OH- • In this case, NH3 is a base because it accepts the H+ from the water (leaving extra OH- in solution)

19. pH • One way to quantify the strength of acids and bases is the pH scale • pH = potential of Hydrogen • is a scale from 0-14 • Pure water (neutral) is pH = 7 • Acids are < 7 • Bases are >7

20. pH indicators • Indicators can tell us the pH of a solution by changing colors when added • Litmus paper turns red in acid, blue in base • Phenolphthalein solution turns purple when in a base, colorless in an acid/neutral • pH paper strips can accurately measure pH from 0-14

21. pH • Water breaks up to form a very tiny percentage of H+ and OH- ions; most stay as bonded H2O • H2O = [H+] + [OH-] • Scale is a mathematical representation of the concentration of H+ ions

22. Introduction to Scientific Notation • Scientific notation is used to represent very large or very small numbers as exponents • Often chemistry uses 10 as a base number • 10,000,000 = 107 or 10×10×10×10×10×10×10 • 0.0000001 = 10-7 or 10/10/10/10/10/10/10

23. pH • Normally when pure water breaks into ions, it forms a ratio of 10-7 H+ and 10-7 OH- ions for every water molecule • Based on the last slide, that means for every ten million liters of water, there is one gram of H+ (and OH-)

24. pH • pH is a exponential scale based on the strength of H+ ions • pH 0 (acid) = [H+] = 100 or 1 • 1 gram of H+ per every 1 L of water • pH 7 (neutral) = [H+] = 10-7 • 1 gram of H+ per every 10,000,000 L of water • pH 14 (base) = [H+] = 10-14 • 1 gram of H+ per every 100,000,000,000,000 L of water • The higher the pH, the less H+ in solution

25. Acids and Bases are useful • Stomach acids break down our food (digestion) • Geologists use acids to digest rocks in the lab • HF digests silica • HCl digests iron • Bases are great for cleaning (lye, bleach, ammonia, soaps)