1. Chapter 2-1 and 2-2 Basic Chemistry

2. 2-1 The Nature of Matter What is everything made of? Anything that has mass and takes of space is called MATTER. The smallest unit that an element can be broken down to and still have the properties of that element is called an ATOM.

3. Elements A substance that cannot be changed into a simpler substance Compose all matter natural (#1-92) & synthetic (>#93) on Periodic Table (appendix G) essential (C, H, O & N) trace (Cu, Mg, Fe, etc.)

5. The Atom Basic unit of matter Made up of sub-atomic particles called –Protons (+) –Neutrons (no charge) –Electrons (-) Atoms like to be neutral- no charge –Equal number of protons and electrons

6. Structure of an Atom NUCLEUS Protons (+) Neutrons (o) Held together by the strong force ELECTRON CLOUD Surrounds the nucleus Electrons (-) - Constantly moving within orbital- attracted to the nucleus by the “weak force”

7. The Electron Orbitals 1 st orbital can only hold 2 electrons (too close to nucleus - not much space) 2 nd orbital can hold up to 8 3 rd orbital can hold up to 8

8. Atomic Number # of protons (and also # of electrons) Chemical symbol Name of Element Atomic Mass The weight Of carbon atom or average weight of all isotopes 6 C Carbon 12.011 atomic mass = protons + neutrons

9. What is the difference between atoms? # of protons, neutrons and electrons (usually all same number) Atoms are defined by their number of protons. If proton number changes the atom changes to new corresponding number (ie. nuclear fission)

10. Isotopes Atoms of same element w/different number of neutrons Examples – carbon; atomic mass usually 12 (C 12 ); other forms are C 13 & C 14

11. Isotopes Radioactive isotopes – give off radiation when nuclei break down; ex. - uranium Some used as diagnostic tool (iodine – for examining thyroid gland; cobalt – treatment for cancer

12. IONS Atoms that have lost or gained electrons –cation – positively charged ion – lost electrons –anion – negatively charged ion – gained electrons

13. Learning Checkpoint What are the three subatomic particles and their charges? What is the only actual difference between gold and mercury? What is the atomic mass of lead? What is the mass number of nitrogen? What is an isotope? What is an ion?

14. Organization of Matter Atoms usually do not occur alone, but exist with other atoms as: –Elements (all of the same atoms) –Molecules or Compounds Same or different atoms bonded

15. Why do atoms bond? An atom wants to have a complete outer shell (called valence level) of electrons. To do this, it can… Share electrons with another atom Give away its electron(s) in this level Receive electrons from another atom *Remember: an atom is when its outer orbital is filled

16. BONDING Atoms need to bond together to make molecules or compounds –“Molecule”is often used to refer to an individual grouping. “singular” –“Compound”- Composed of atoms of different elements –Molecules and compounds are written out in a chemical formula: example- C 6 H 12 O 6

17. Molecules vs. Compounds

18. TYPES OF BONDS 1.IONIC - One atom (very unstable) gives 1, 2 or 3 electrons away to another atom. The atom that loses electrons becomes positively charged. The atom that gains the electrons becomes negatively charged. The opposite charges cause the atoms to “bond” together (opposites attract). 2.COVALENT- atoms share a pair of electrons (sometimes share 2 (double bond) or 3 (triple bond) pairs) STRONGEST OF THE BONDS! 3.HYDROGEN- (will be discussed in detail in next section)

19. Example of a Covalent Bond

20. Na (sodium) is very unstable It only has one e- in its outer orbital. Cl’s (chlorine) outer orbital is almost filled Na gives its lonely e- to Cl. Na become Na+ Cl becomes Cl- Their opposite charges cause them to be attracted to one another- This is an ionic bond. Example of Ionic Bonding-NaCl

21. Ionic Bonds CaCl→

22. ←LiCl

23. Learning Checkpoint What is an ion? Why is it important that atoms bond? What causes atoms to bond? Explain the difference between an ionic bond and a covalent bond.

24. IV. Chemical Reactions A.When they occur: 1. Bonds are formed. 2. Bonds are broken. 3. Substances change/are altered into different substances. B.Metabolism 1. Refers to all the chemical reactions that take place within an organism.

25. Examples of a Chemical Equation C 6 H 12 O 6 + 6O 2 6CO 2 + 6H 2 O + Energy Reactants - Substances that undergo the reaction. Products - Substances which are formed from a reaction. Subscripts - The # of atoms of each element in a molecule. Example: C 6 H 12 O 6 Coefficient - The # before each chemical formula. (The # of molecules of that substance.) Reactants Products Yields

26. Coefficient Example 12 H 2 0 12 Molecules of Water 2 Atoms of H 1 Atom of O H 2 0, H 2 0, H 2 0, H 2 0 How many atoms of hydrogen in 12 molecules of water? 24

27. 2-2 Water Importance of Water: Provides a place for chemical reactions Provides a means of transport Makes up a large portion of living organisms Because of the arrangement of atoms, water has a + end and a – end like a magnet

28. Water Density Ice is less dense than liquid water When water freezes air is trapped within the frozen ice making the cube larger and less dense Benefits: –Fish and plant life can survive in liquid layers of water under ice

29. Polarity Water is polar Although the compound is neutral overall there is a shift of charge within the compound The much larger atom, Oxygen, pulls more on the shared e- This end of the compound becomes slightly more negative. Hydrogen ends become slightly positive

30. Water Hydrogen bond – a weak bond that forms when H 2 O molecules attract each other (+H to –O)

31. Hydrogen Bonding Due to polarity, water compounds attract to one another Slightly negative oxygen attracts slightly positive hydrogen from another compound This attraction among water is COHESION. Water is also attracted to other materials. This is ADHESION.

33. Cohesion An attraction b/t molecules of the same substance Because of hydrogen bonds water molecules are attracted to each other Examples – water molecules on the surface of a body of water have a strong attraction called “surface tension”

34. Examples of Cohesion Insect walking on water → ←water beading on smooth surface

35. Adhesion An attraction b/t molecules of different substances Occurs b/t water & other polar molecules like glass Occurs when water is drawn up in small tubes (ex. - from the roots to the stems of plants) Called capillary action

36. Water The greatest solvent on Earth! Water’s polarity allows it to break ionic bonds of other compounds…creating free ions.

37. Mixtures Mixture – composed of 2 or more elements that are physically mixed but not chemically combined Ex. – salt & pepper, sugar & cinnamon, different colored M&Ms Two types: 1. solution 2. suspension

38. Solution A mixture in which the components are evenly distributed throughout Components: a. solute – the substance that is dissolved b. solvent – the substance in which the solute dissolves Example – salt in water

39. Suspension Created when materials do not dissolve in water but separate into pieces & do not settle out of solution; ex. – blood (water + dissolved cmpds.), sand in water

40. Matter All matter is composed of atoms and groups of atoms bonded together, called molecules. –Substances that are made from one type of atom only are called pure substances. –Substances that are made from more than one type of atom bonded together are called compounds. –Compounds that are combined physically, but not chemically, are called mixtures.

41. Elements, Compounds, Mixtures Sodium is an element. Chlorine is an element. When sodium and chlorine bond they make the compound sodium chloride, commonly known as table salt.  Compounds have different properties than the elements that make them up.  Table salt has different properties than sodium, an explosive metal, and chlorine, a poisonous gas.

42. Element, Compound or Mixture? Mixture of elements.

43. Element, Compound or Mixture? Pure compound like Salt.

44. Element, Compound or Mixture? Mixture of elements and compounds.

45. Element, Compound or Mixture? Pure element like gold.

46. Element, Compound or Mixture? Mixture of compounds.

47. Element, Compound or Mixture? This is a molecule like O 2 or N 2.

48. Element, Compound or Mixture? This is a compound like water.

49. Element, Compound or Mixture? This is an element like Argon.

50. Learning Checkpoint Why does ice float on a lake? Explain the polarity of water – how are the charges distributed? What is the difference between adhesion and cohesion? Explain the difference between a solution and a suspension.

51. Acids and Bases Water can react to form individual ions: H 2 O H+ + OH- In pure water this occurs naturally but the amount of H+ is always = to the amount of OH- so water remains neutral

52. pH scale: “the power of Hydrogen” Some solutions made with water become acidic or basic. This is determined by the amount of H+ (hydrogen ions) in the solution pH = - log [H+] Because it’s logarithmic, each pH unit represents a tenfold difference in concentration of H+ ions. This means something with a pH of 4 is 10 times more acidic than something with a pH of 5.

53. Acids pH range from 0 – 6.99 Any compound that forms H+ ions in solution H+ ions > OH- ions The closer to 0 the more acidic the solution Examples: stomach acid, lemon juice

54. Bases (Alkaline) pH ranges from 7.01 to 14 Any compound that forms 0H- ions in solution OH- ions > H+ ions The closer to 14 the more basic the solution Examples: lye, bleach, oven cleaner

55. ACID:Any compound that forms H+ ions in solution BASE: Any compound that forms 0H- ions in solution

56. pH and Living Things pH values in living cells are usually kept between 6.5 and 7.5 –Optimal pH for chemical reactions to take place in the body –Any switch in pH could cause serious/fatal problems

57. Buffers Weak acids or bases that can react with strong acids or bases Used to regulate pH and prevent sharp sudden changes in pH There are natural buffers in your blood that keep the pH at 6.5 to7.5

58. Learning Checkpoint What makes a solution acidic or basic? How is acidity measured? A solution with pH 8.5 is considered…. A solution with a pH of 3.4 is considered… What is a buffer and why is it important?