2. Chemistry Overview Chemistry = The study of matter and how it changes Matter = anything that has mass and occupies space MatterNot Matter

3. Categorizing Matter Stuff MatterEnergy MixturePure substance Heterogeneous mixture Homogeneous mixture compoundElement

4. 2-1 What is Matter? Matter - made of atoms – the smallest particles that have the properties of an element. -The properties and characteristics of matter all depend upon what atoms the matter is made of.

5. Classifying Matter Pure Substances – has a constant composition 1. Elements – made of only one type of atom - cannot be separated by ordinary chemical or physical means - ex: iron (Fe), copper (Cu), aluminum (Al)

6. 2. Compound – made of two or more elements in set proportions - ex: water (H 2 0), salt (NaCl)

7. Note: A molecule is formed when two or more atoms join together chemically. A compound is a molecule that contains at least two different elements. All compounds are molecules but not all molecules are compounds. Oxygen gas (O 2 ) is a molecule, but not a compound.

8. Chemical Formulas - C 16 H 10 N 2 O 2 is the formula for a molecule of indigo, a dye - It has 16 carbon atoms, 10 hydrogen atoms, 2 nitrogen atoms, and 2 oxygen atoms - If a number appears before the formula, 6 C 6 H 12 O 6 it means there are 6 molecules of glucose

9. Mixtures – blend of pure substances - 2 or more combined substances in no definite amounts - each substance keeps its original properties

10. 1. Heterogeneous mixture - “hetero” means different… mixture is different - ex: trail mix, salad

11. 2. Homogenous mixture - “homo” means the same… mixture is the same throughout - ex: salt water, tea

12. Liquid Mixtures Miscible – gasoline is made of more than 100 compounds dissolved Immiscible – oil and water do not mix together well Note: gases can mix with liquids – carbonated drinks - homogeneous

13. States of Matter There are three basic components or parts to each state of matter that can vary: 1. Temperature 2. Closeness of particles 3. Energy of particles (E) These components are closely related. Think about how the kinetic theory explains the relationships.

14. “Most matter found naturally on Earth is either a solid, a liquid, or a gas, but matter also exists in other states” (p46).

15. Plasma – free electrons and ions of element; electrically charged gas Temperature – higher than gases Particle closeness – far apart; little or no order Energy – High energy Examples – fluorescent light bulb, neon signs, stars, lightning, plasma TV screens

16. Gases Temperature Above the boiling point of the substance Steam burns are worse than water burns because boiling water = 100°C where steam can be 125°C

17. Particle Closeness With available space, gas particles spread apart as more Energy is added Distance between particles allows for a lot of movement

18. Energy Level Higher E level than liquid or solid Requires a lot of E to evaporate  heat

19. Other Gases take the shape of their containers. Gases exert pressure on closed containers Molecules of gas moving fast in balloon constantly hitting each other and sides of balloon.

20. Liquids Temperature Between freezing and boiling points Water is liquid from 0°C to 100°C

21. Particle Closeness Particles are close together, but not as attracted to each other as in solid form Particles can move independently of each other Energy Level Energy is absorbed or released to form solid or gas

22. Other Liquids vary in rate at which they spread This property is called viscosity If liquid molecules have a strong attraction to each other  slower spread like molasses Both liquids and gases can spread, so are called fluids Some substances do not have a liquid form at normal temperature and pressure

23. Solids Temperature Lower temp to achieve freezing Freezing point – melting point

24. Particle Closeness Particles are close allowing almost no movement… particles can vibrate in place Organized arrangement because molecules bond Solids maintain their own shape

25. Energy Level Lowest E level  there is still energy present There is never NO energy in a system

26. Energy’s Role in Phase Changes Energy is the ability to change or move matter Examples: electricity, batteries (stored), candles, food, etc

27. Phase Changes

28. Energy must be added to cause melting or evaporating. The molecules gain E and then move faster. Think of the kinetic theory. Energy is transferred from substance to surroundings when condensing and freezing.

29. Sublimation – change from solid directly to gas Dry ice = CO2 solid to smoky gas Ice in freezer will sublime and ice gets smaller

30. A change in the state of a substance does not change its composition.

31. Law of Conservation of Mass Mass cannot be created or destroyed. Mass before change = Mass after change

32. 1. Ice melting Change of state Same mass Remember Δ volume ≠ Δ mass

33. 2. Combining 2 liquids no mass is lost 20 g + 20 g = 40 g

34. 3. Burning a match Mass A = Mass B Match = burnt mass + gas

35. 4. Chemical reaction XXX + YYY  ZZZ Reactants  Products 3 atoms + 10 atoms  13 atoms The amount of atoms entering into a reaction will be the amount that come out of the reaction.

36. Law of Conservation of Energy Energy cannot be created or destroyed. Energy before change = Energy after change Energy is stored in many ways = food, gasoline, batteries

37. 1. Heating water  provides water molecules with enough energy to evaporate

38. 2. Energy from gasoline and you  car drives a certain distance