1. Chemistry Notes Chapter 2

2. Hydrogen Bonding animation:

3. Energy and chemical reactions:

4. Atoms and bonding:

5. I. Nature of Matter What is Matter?
Matter= anything that has mass and takes up space (has volume).
What is NOT matter?
Light
Sound
Electricity

6. Give some examples of what is matter and what is Not matter.
Think About it:
What is matter?
Give some examples of what is matter and what is Not matter.

7. C. Composition of Matter
The ATOM is the basic “building block” (or unit) of all matter.
What was the basic Unit of Life?
Atoms are the smallest units of matter.

8. II. The Atom A. Atom Anatomy Protons = + Neutrons = 0 (neutral)
Electrons = -
Nucleus = Protons and Neutrons
Electron Cloud (shells, orbital)= Electrons surrounding nucleus
Valence Electrons = electrons on outermost valence of atom.
Only protons and neutrons have substantial mass (= 1 AMU) So they make up an atoms weight or mass.Electrons have negligible mass

9. C B. Atomic # vs. Mass # 6 Carbon 12 Atomic Number = # of Protons
Mass Number = # of Protons and Neutrons
# of Protons 6 C
Carbon 12
Symbol
# of Protons AND Neutrons

10. Electron
Nucleus
Valence Electron
Electron
Proton
CARBON
Neutron

11. III. Elements Elements consist of entirely one type of atom.
The # of protons determines the type of element. (Look at your table)
Organized on the Periodic Table of Elements in order of increasing Atomic #. WHY??

12. What particles are found in the nucleus of an Atom?
Neutrons and Protons!!

13. What subatomic particles are responsible for the bonding between atoms?
Valence Electrons!!

15. IV. Common Biological Elements
You are responsible for knowing the following elements throughout the year:
Carbon ***
Hydrogen
Nitrogen
Oxygen
Phosphorus
Sulfur
* CHON = 96% of living matter
CHNOPS!!!

16. What are the 6 most common biological elements?
Think about it:
What are the 6 most common biological elements?

17. What is the basic unit of all Matter?
The Atom

18. V. Molecules
Molecule = the smallest unit any substance can be divided into (without losing its properties).
Example: 1 Water molecule!
H2O
2 Hydrogen atoms + 1 Oxygen atom

19. VI. Compounds
Compounds = a combination of two or more different elements bonded together.
Example: Glucose!
C6H12O6
C= Black
H = White
O = Red

20. What’s the Difference?
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.

21. Lets practice together!!

22. How many total atoms are found in one molecule of C12H22O­11?
b. 28
c. 35
d. 56
e. 45

23. VII. Isotopes
Isotope = An atom of an element with a different number of neutrons.
# of protons remain the same
Mass # changes as a result
Example:
H-1 vs. H-2

24. Radioactive isotopes:
Occur when the nucleus of an atom begin to break down, releasing energy.
Used to find the age of extremely old organic matter.
Carbon 14 Dating

25. An atom of an element with a different number of neutrons is a ________?
a. Element
b. Carbon
c. Compound
d. Isotope

26. Which of the following is a radioactive isotope of carbon?
c. C-14
d. C-15
e. C-16

27. Bonding and Chemical Reactions Notes

28. I. Bonds Bonds= the force of attraction that holds atoms together.
Occurs between valence electrons that are oppositely charged (+ and -).

29. A. Types of Bonds There are two major types of bonds: Ionic Bond
Covalent Bond
Hydrogen Bond

30. 1. Ionic Bonds
Ions: An atom or group of atoms that has an electrical charge (+ or -).
Ion = an atom that has gained or lost an electron.
Occur between a metal and a non-metal
Positive Ions:
Lose an electron
More p+ than e-
Negative Ions:
Gain an electron
Less p+ than e-

31. Ionic Bonds form from the attraction between oppositely charged atoms.
Example: salt
Na + Cl-
Electrons are transferred.
How do you determine charge and number of electrons in the ion?

32. 2. Covalent Bonds Covalent Bonds form when two atoms SHARE electrons.
Occurs between 2 non-metals
The electrons shared are the outermost electrons—the VALENCE electrons.
There are two types of covalent bonds:
Polar
Non-Polar

33. Polar Covalent Bonds = Covalent bonds where the electrons are SHARED unequally.
Example = H20
Non-Polar Covalent Bonds = Covalent bonds where the electrons are shared EQUALLY.
Example = CH4 (Methane Gas!)

34. Which of the following bonds occurs when there is a transfer of electrons?
a. polar covalent
b. non-polar covalent
c. ionic
d. bipolar

35. Which of the following bonds occurs when there is an equal sharing of electrons?
a. polar covalent c. ionic
b. non-polar covalent d. bipolar

36. 3. Hydrogen Bonds
Hydrogen Bonds= are bonds that form between Hydrogen atoms and Oxygen/ Nitrogen.
These bonds are weaker than covalent bonds.

37. Which of the following is bonded due to ionic bonds?
a. H2O
b. NaCl; metal and nonmetal
c. CH4; metal and metal
d. all of them
e. none of above
b. NaCl; metal and nonmetal

38. II. Chemical Reactions
Definition: A chemical reaction occurs when substances undergo chemical changes to form NEW substances.
Break and form new bonds

39. A. Why Atoms Move
Every chemical reaction involves the rearrangement of elements as they react with one another.

40. sunlight
6CO2 + 6H2O C6H12O6 + 6O2
(Reactants) (Products)
Reactants—the raw materials on the LEFT side of the arrow.
Products—the RESULT of the chemical reaction. Found on the RIGHT side of the arrow.

41. sunlight
6CO2 + 6H2O C6H12O6 + 6O2
(Reactants) (Products)
Photosynthesis : 6 molecules of carbon dioxide react with 6 molecules of water to make 1 molecule of sugar (glucose) and 6 molecules of oxygen.

42. B. Types of Energy (Why Bonds Form)
Catabolic Reaction= breaking bonds and energy is released.
Anabolic Reaction= when atoms join back together to form a new bonds and energy is stored in the bond.

43. Exothermic Reactions exo = Out thermic = Heat
Therefore, exothermic reactions release heat energy as they occur.

44. Endothermic Reactions
Endo = In
Therefore, endothermic reactions require the absorption of heat energy in order to occur.

45. If a reaction occurs and heat is released, is it an exothermic reaction or endothermic reaction?

46. Solution Chemistry

47. II. Solutions
Definition = the even distribution of substances dissolved in H2O.
Example:
H2O and NaCl

48. A. Solvents= the substance that breaks apart other substances.
H2O is the “universal” solvent.
Solutes= the substance that is broken apart.
- Kool Aid!
C. Concentration = how much solute is in the solution.

49. What is the substance that breaks apart other substances?
Solute
b. Solvent
c. H-Bond
d. Suspension

50. III. Acids Acids contain hydrogen ions: H+
Acids donate hydrogen ions (H+) to water (H2O) to form H3O+
H3O+ is known as a hydronium ion

51. A. Properties of Acids Acids taste sour Acids conduct electricity.
Acids turn litmus paper RED.

52. Concentrated acids are dangerous!
Acids can burn your skin and eyes…

53. IV. Bases Properties of Bases Bases contain hydroxide ions: OH-
Bases taste bitter.
Bases feel slippery
Bases turn litmus paper BLUE.

54. V. pH A. pH is the measure of hydronium [H3O+] ions in a solution.
The pH scale is a much more accurate measure of how acidic or basic a solution is.

55. VI. pH Scale The pH scale ranges from 0-14 0-6 is Acidic (more H+)
8-14 is Basic (more OH-)
0 = VERY acidic [H3O+]
7 = neutral [H2O] (equal H+ and OH-)
14 = VERY basic [OH-]
Acidic
Neutral 7
Basic 14

56. Low pH is acid
Lots of H3O+
Little OH-
High pH is base
Little H3O+
Lots of OH-

57. Which of the following has the pH that is most acidic?
3.5
b. 6.3
c. 2.8
d. 1.5
e. 11.0

58. Which of the following has the pH that is most basic?
a b c d e. 12.5

59. Why is pure water neutral? Please explain in words.