1. Chemistry Comes Alive: Part A2
Chemistry Comes Alive: Part A

2. Anything that has mass and occupies space States of matter:
Solid—definite shape and volume
Liquid—definite volume, changeable shape
Gas—changeable shape and volume

3. Capacity to do work or put matter into motion Types of energy:
Kinetic—energy in action
Potential—stored (inactive) energy
PLAY
Animation: Energy Concepts

4. Forms of Energy
Chemical energy —stored in bonds of chemical substances
Electrical energy —results from movement of charged particles
Mechanical energy —directly involved in moving matter
Radiant or electromagnetic energy —exhibits wavelike properties (i.e., visible light, ultraviolet light, and X-rays)

6. Energy Form Conversions
Energy may be converted from one form to another
Conversion is inefficient because some energy is “lost” as heat

7. Composition of Matter Elements
Cannot be broken down by ordinary chemical means
Each has unique properties:
Physical properties
Are detectable with our senses, or are measurable
Chemical properties
How atoms interact (bond) with one another

8. Atomic symbol: one- or two-letter chemical shorthand for each element
Composition of Matter
Atoms
Unique building blocks for each element
Atomic symbol: one- or two-letter chemical shorthand for each element

9. Major Elements of the Human Body
Oxygen (O)
Carbon (C)
Hydrogen (H)
Nitrogen (N)
About 96% of body mass

10. Lesser Elements of the Human Body
About 3.9% of body mass:
Calcium (Ca), phosphorus (P), potassium (K), sulfur (S), sodium (Na), chlorine (Cl), magnesium (Mg), iodine (I), and iron (Fe)

11. Trace Elements of the Human Body
< 0.01% of body mass:
Part of enzymes, e.g., chromium (Cr), manganese (Mn), and zinc (Zn)

12. Atomic Structure – Review Write this down if needed!
Determined by numbers of subatomic particles
Nucleus consists of neutrons and protons

13. video

14. r

15. Atomic Structure Neutrons Protons No charge
Mass = 1 atomic mass unit (amu)
Protons
Positive charge
Mass = 1 amu

16. Atomic Structure Electrons Orbit nucleus
Equal in number to protons in atom
Negative charge
1/2000 the mass of a proton (0 amu)

17. Nucleus Nucleus Helium atom Helium atom 2 protons (p+) 2 neutrons (n0)
2 electrons (e–)
2 protons (p+)
2 neutrons (n0)
2 electrons (e–)
(a) Planetary model
(b) Orbital model
Proton
Neutron
Electron
Electron
cloud
Figure 2.1

18. Video

19. Identifying Elements – do not write this
Atoms of different elements contain different numbers of subatomic particles
Compare hydrogen, helium and lithium (next slide)

20. Hydrogen (H) (1p+; 0n0; 1e–) Helium (He) (2p+; 2n0; 2e–) Lithium (Li)
Proton
Neutron
Electron
Hydrogen (H)
(1p+; 0n0; 1e–)
Helium (He)
(2p+; 2n0; 2e–)
Lithium (Li)
(3p+; 4n0; 3e–)
Figure 2.2

21. Identifying Elements
Atomic number = number of protons in nucleus

22. Mass number = mass of the protons and neutrons
Identifying Elements
Mass number = mass of the protons and neutrons
Mass numbers of atoms of an element are not all identical
Isotopes are structural variations of elements that differ in the number of neutrons they contain

23. Identifying Elements
Atomic weight = average of mass numbers of all isotopes

24. Valuable tools for biological research and medicine
Radioisotopes
Valuable tools for biological research and medicine
Cause damage to living tissue:
Useful against localized cancers
Radon from uranium decay causes lung cancer

25. Molecules and Compounds
Most atoms combine chemically with other atoms to form molecules and compounds
Molecule—two or more atoms bonded together (e.g., H2)
Compound—two or more different kinds of atoms bonded together (e.g., C6H12O6)

26. Most matter exists as mixtures
Two or more components physically intermixed
Three types of mixtures
Solutions
Colloids
Suspensions

27. Usually transparent, e.g., atmospheric air or seawater
Solutions
Homogeneous mixtures
Usually transparent, e.g., atmospheric air or seawater
Solvent
Present in greatest amount, usually a liquid
Solute(s)
Present in smaller amounts

28. Colloids and Suspensions
Colloids (emulsions)
Heterogeneous translucent mixtures, e.g., cytosol
Large solute particles that do not settle out
Undergo sol-gel transformations
Suspensions:
Heterogeneous mixtures, e.g., blood
Large visible solutes tend to settle out

29. Solution Colloid Suspension Solute particles are very
tiny, do not settle out or
scatter light.
Solute particles are larger
than in a solution and scatter
light; do not settle out.
Solute particles are very
large, settle out, and may
scatter light.
Solute
particles
Solute
particles
Solute
particles
Example
Mineral water
Example
Gelatin
Example
Blood
Figure 2.4

30. Mixtures vs. Compounds Mixtures Compounds
No chemical bonding between components
Can be separated physically, such as by straining or filtering
Heterogeneous or homogeneous
Compounds
Can be separated only by breaking bonds
All are homogeneous

31. Chemical Bonds – review only
Electrons occupy up to seven electron shells (energy levels) around nucleus
Octet rule: Except for the first shell which is full with two electrons, atoms interact in a manner to have eight electrons in their outermost energy level (valence shell)

32. Chemically Reactive Elements – review only
Outermost energy level not fully occupied by electrons
Tend to gain, lose, or share electrons (form bonds) with other atoms to achieve stability

33. (b) Chemically reactive elements
Outermost energy level (valence shell) incomplete 4e 1e 2e
Hydrogen (H)
(1p+; 0n0; 1e–)
Carbon (C)
(6p+; 6n0; 6e–) 1e 6e 8e 2e 2e
Oxygen (O)
(8p+; 8n0; 8e–)
Sodium (Na)
(11p+; 12n0; 11e–)
Figure 2.5b

34. Types of Chemical Bonds
Ionic
Covalent
Hydrogen

35. Ions are formed by transfer of valence shell electrons between atoms
Ionic Bonds
Ions are formed by transfer of valence shell electrons between atoms
Anions (– charge) have gained one or more electrons
Cations (+ charge) have lost one or more electrons
Attraction of opposite charges results in an ionic bond

36. + – Figure 2.6a-b Sodium atom (Na) (11p+; 12n0; 11e–)
Chlorine atom (Cl)
(17p+; 18n0; 17e–)
Sodium ion (Na+)
Chloride ion (Cl–)
Sodium chloride (NaCl)
(a) Sodium gains stability by losing one electron, and
chlorine becomes stable by gaining one electron.
(b) After electron transfer, the oppositely
charged ions formed attract each other.
Figure 2.6a-b

37. Formation of an Ionic Bond
Ionic compounds form crystals instead of individual molecules
NaCl (sodium chloride)

38. (c) Large numbers of Na+ and Cl– ions
CI–
Na+
(c) Large numbers of Na+ and Cl– ions
associate to form salt (NaCl) crystals.
Figure 2.6c

39. Covalent Bonds
Formed by sharing of two or more valence shell electrons
Allows each atom to fill its valence shell at least part of the time

40. + Reacting atoms Resulting molecules or Structural formula shows
single
bonds.
Hydrogen
atoms
Carbon
atom
Molecule of
methane gas (CH4)
(a) Formation of four single covalent bonds:
carbon shares four electron pairs with four
hydrogen atoms.
Figure 2.7a

41. + Reacting atoms Resulting molecules or Structural formula shows
double
bond.
Oxygen
atom
Oxygen
atom
Molecule of
oxygen gas (O2)
(b) Formation of a double covalent bond: Two
oxygen atoms share two electron pairs.
Figure 2.7b

42. + Reacting atoms Resulting molecules or Structural formula shows
triple
bond.
Nitrogen
atom
Nitrogen
atom
Molecule of
nitrogen gas (N2)
(c) Formation of a triple covalent bond: Two
nitrogen atoms share three electron pairs.
Figure 2.7c

43. Sharing of electrons may be equal or unequal
Covalent Bonds
Sharing of electrons may be equal or unequal
Equal sharing produces electrically balanced nonpolar molecules
CO2

44. Figure 2.8a

45. Covalent Bonds
Unequal sharing by atoms with different electron-attracting abilities produces polar molecules
H2O
Atoms with six or seven valence shell electrons are electronegative, e.g., oxygen
Atoms with one or two valence shell electrons are electropositive, e.g., sodium

46. Figure 2.8b

47. Figure 2.9

48. Hydrogen Bonds
Attractive force between electropositive hydrogen of one molecule and an electronegative atom of another molecule
Common between dipoles such as water
Also act as intramolecular bonds, holding a large molecule in a three-dimensional shape
PLAY
Animation: Hydrogen Bonds

49. Hydrogen bond (indicated by dotted line)+ –
Hydrogen bond
(indicated by
dotted line) + + – – – + + + –
(a) The slightly positive ends (+) of the water molecules become aligned with the slightly negative ends (–) of other water molecules.
Figure 2.10a

50. (b) A water strider can walk on a pond because of the high surface tension of water, a result of the combined strength of its hydrogen bonds.
Figure 2.10b

51. Occur when chemical bonds are formed, rearranged, or broken
Chemical Reactions
Occur when chemical bonds are formed, rearranged, or broken
Represented as chemical equations
Chemical equations contain:
Molecular formula for each reactant and product
Relative amounts of reactants and products, which should balance

52. Examples of Chemical Equations
H + H  H2 (hydrogen gas)
4H + C  CH4 (methane)
(reactants)
(product)

53. Patterns of Chemical Reactions
Synthesis (combination) reactions
A + B  AB
Decomposition reactions
AB  A + B
Exchange reactions
AB + C  AC + B

54. Synthesis Reactions A + B  AB Always involve bond formation
Anabolic ( creation of a product, the building up of organs and tissues, requires energy! ATP)
Anabolism is powered by Catabolism!

55. (a) Synthesis reactions
Smaller particles are bonded
together to form larger,
more complex molecules.
Example
Amino acids are joined together to
form a protein molecule.
Amino acid
molecules
Protein
molecule
Figure 2.11a

56. Decomposition Reactions
AB  A + B
Reverse synthesis reactions
Involve breaking of bonds
Catabolic (set of metabolic pathways that break down molecules and release energy)

57. (b) Decomposition reactions
Bonds are broken in larger
molecules, resulting in smaller,
less complex molecules.
Example
Glycogen is broken down to release
glucose units.
Glycogen
Glucose
molecules
Figure 2.11b

58. Exchange Reactions AB + C  AC + B Also called displacement reactions
Bonds are both made and broken

59. + + (c) Exchange reactions Bonds are both made and broken
(also called displacement reactions).
Example
ATP transfers its terminal phosphate
group to glucose to form glucose-phosphate. +
Glucose
Adenosine triphosphate (ATP) +
Glucose
phosphate
Adenosine diphosphate (ADP)
Figure 2.11c

60. Oxidation-Reduction (Redox) Reactions
Decomposition reactions: Reactions in which fuel is broken down for energy
Also called exchange reactions because electrons are exchanged or shared differently
Electron donors lose electrons and are oxidized
Electron acceptors receive electrons and become reduced

61. All chemical reactions are either exergonic or endergonic
Exergonic reactions —release energy
Catabolic reactions
Endergonic reactions —products contain more potential energy than did reactants
Anabolic reactions

62. Chemical Reactions – review only
All chemical reactions are theoretically reversible
A + B  AB
AB  A + B
Chemical equilibrium occurs if neither a forward nor reverse reaction is dominant
Many biological reactions are essentially irreversible due to
Energy requirements
Removal of products

63. Rate of Chemical Reactions – review only
Rate of reaction is influenced by:
 temperature   rate
 particle size   rate
 concentration of reactant   rate
Catalysts:  rate without being chemically changed
Enzymes are biological catalysts