1. AP Bio Chapter 2 – Chemistry Review
Notes/homework
Campbell Biology in Focus

2. Matter consists of elements and compounds
Organisms are composed of matter.
Matter is anything that takes up space and has mass.
Matter is made up of elements.
An element cannot be broken down to other substances by chemical reactions.
A compound is a substance consisting of two or more elements in a fixed ratio.
A compound has emergent properties, which are different from those of its elements.

3. Elements of life
Essential elements – needed by an organism to live and reproduce.
Trace elements are required in only small quantities.
Some naturally occurring elements are toxic.
CHNOPS – the primary elements in biochemistry
Carbon
Hydrogen
Nitrogen
Oxygen
Phosphorus
Sulfur

4. An element’s properties depend on the structure of its atoms
An atom is the smallest unit of matter that still retains the properties of that element.
Atoms are composed of subatomic particles
Neutrons – in nucleus; no electrical charge (n0)
Protons – also in nucleus; positive electrical charge (p+)
Electrons – surrounding nucleus; negative electrical charge (e-)
Different elements have different numbers of subatomic particles
Atomic number = the number of protons (this will equal the number of electrons in a neutral atom)
Mass number = the number of protons plus neutrons in the nucleus
Periodic table – arrangement of the elements based on their atomic number

5. Isotopes
All atoms of an element have the same number of protons but may differ in the number of neutrons.
Isotopes are two atoms of an element that differ in the number of neutrons.
Radioactive isotopes decay spontaneously, giving off particles and energy.
Biological applications of radioactive isotopes:
Dating fossils
Tracing atoms through metabolic processes
Diagnosing medical disorders

6. Energy levels of electrons
Energy is the capacity to cause change.
Potential energy is the energy that matter has because of its location or structure.
The electrons of an atom differ in their amounts of potential energy.
An electron’s state of potential energy is called its energy level, or electron shell.
The energy level of each shell increases with distance from the nucleus.
Electrons can move to higher or lower shells by absorbing or releasing energy.

7. Third shell (highest energy level in this model) Energy
Figure 2.5b
Third shell (highest
energy level in this
model)
Energy
absorbed
Second shell (higher
energy level)
First shell (lowest
energy level)
Energy
lost
Figure 2.5b Energy levels of an atom’s electrons (part 2: shell model)
Atomic
nucleus
(b) 7

8. Electron distribution and chemical properties
The chemical behavior of an atom is determined by the distribution of electrons in electron shells.
The periodic table shows the electron distribution for each element.
The chemical behavior of an atom depends mostly on the number of electrons in its outermost shell, or valence shell.
Valence electrons are those that occupy the valence shell.
Atoms with completed valence shells are unreactive or inert.

9. He Figure 2.6 Hydrogen 1H 2 4.00 Atomic number Helium 2He Atomic mass
First
shell
Element symbol
Electron
distribution
diagram
Lithium
3Li
Beryllium
4Be
Boron 5B
Carbon 6C
Nitrogen 7N
Oxygen 8O
Fluorine 9F
Neon
10Ne
Second
shell
Sodium
11Na
Magnesium
12Mg
Aluminum
13Al
Silicon
14Si
Phosphorus
15P
Sulfur
16S
Chlorine
17Cl
Argon
18Ar
Figure 2.6 Electron distribution diagrams for the first 18 elements in the periodic table
Third
shell 9

10. Chemical bonding
Covalent bonds = the sharing of a pair of valence electrons
The shared electrons count as part of each atom’s valence shell
Two or more atoms held by a covalent bond make up a molecule
The notation used to represent atoms and bonding is called a structural formula:
For example: H—H
This can be abbreviated with a molecular formula: H2
In a structural formula, a single line represents the sharing of one pair (2e-) of electrons. A double line shows the sharing of two pairs (4e-) of electrons.
H—H versus O—O

11. Name and Molecular Formula Electron Distribution Diagram Structural
Figure 2.8
Name and
Molecular
Formula
Electron
Distribution
Diagram
Structural
Formula
Space-
Filling
Model
(a) Hydrogen (H2)
(b) Oxygen (O2)
(c) Water (H2O)
Figure 2.8 Covalent bonding in four molecules
(d) Methane (CH4) 11

12. Covalent bonds can share electrons unevenly
Electronegativity – an atom’s attraction to the electrons in a covalent bond
The more electronegative an atom, the more strongly it pulls shared electrons toward itself.
In a nonpolar covalent bond the atoms share the electrons equally
In a polar covalent bond, one atom is more electronegative and the atoms do not share the electrons equally
Unequal sharing of electrons causes a partial positive and a partial positive charge on the molecule

13. Figure 2.9 − O H H  
Figure 2.9 Polar covalent bonds in a water molecule
H2O 13

14. Chemical bonding…
Ionic bonds – one atom loses e- and the other gains it
Both atoms have complete valence shells
Cation = positively charged ion (it has lost e-)
Anion = negatively charged ion (it has gained e-)
Compounds formed by ionic bonds are called ionic compounds or salts – a metal and a nonmetal
Salts are often found in nature as crystals

15. Weak chemical bonds
Van der Waals interactions – result from ‘hot spots’ of positive or negative charge
Are individually weak and occur only when atoms and molecules are very close together
Collectively the interactions can be strong – example is a gecko’s toe hairs and a wall surface

16. Figure 2.UN01
Figure 2.UN01 In-text figure, Van der Waals interactions, p. 27 16

17. Weak chemical bonds…
Many large biological molecules are held in their functional form by weak bonds
Hydrogen bonds form between polar covalent molecules – they are the attraction between the partial positive on one molecule and the partial negative on the other
Hydrogen bonds between water molecules account for many of the properties of water that are important for living things:
Cohesion and adhesion
Ability to moderate temperature
Expansion upon freezing
Versatility as a solvent

18. −  Water (H2O)  Hydrogen bond − Ammonia (NH3)   
Figure 2.12 − 
Water (H2O) 
Hydrogen bond −
Ammonia (NH3)
Figure 2.12 A hydrogen bond    18

19. Cohesion of water molecules
Water molecules are linked by multiple hydrogen bonds = cohesion
This contributes to the transport of water and nutrients against gravity in plants
Adhesion = the clinging of water molecules to another substance
Surface tension = how hard it is to break the surface of a liquid

20. Two types of water-conducting cells
Figure 2.17
Adhesion
Two types of
water-conducting
cells
Cohesion
Direction
of water
movement
Figure 2.17 Water transport in plants
300 m 20

21. Figure 2.18
Figure 2.18 Walking on water 21

22. Moderation of temperature by water
Water can absorb or release a large amount of heat energy with only a slight change in its own temperature
This is related to water’s high specific heat – the amount of heat that must be absorbed or lost for 1g of that substance to change its temperature by 10 C
The specific heat of water is 1 cal/g/0C
Water resists changing its temperature because of its high specific heat
This can be traced to hydrogen bonding: heat is absorbed when hydrogen bonds break; heat is released when hydrogen bonds form
The high specific heat of water keeps temperature fluctuations within limits that permit life

23. Ice floats
Ice floats in liquid water because the hydrogen bonds are more ‘ordered’, making the ice less dense
Water is the most dense at 40C
If ice sank, all bodies of water would eventually freeze solid from the bottom up, making life impossible on Earth

24. Form follows function
A molecule’s shape is usually very important to its function
Molecular shape determines how biological molecules recognize and respond to each other
Biological molecules recognize and interact with each other based on molecular shape
Molecules with similar shapes can have similar biological effects

25. Hydrogen bond Liquid water: Hydrogen bonds break and re-form Ice:
Figure 2.20
Hydrogen bond
Liquid water:
Hydrogen bonds
break and re-form
Figure 2.20 Ice: crystalline structure and floating barrier
Ice:
Hydrogen bonds
are stable 25

26. The ‘universal’ solvent
Solution = homogeneous liquid mixture
Solvent = the dissolving agent
Solute = the substance that is dissolved
Water is a versatile solvent due to its polarity – it will dissolve other polar molecules as well as ionic compounds
Hydrophilic (‘water loving’) substances have an affinity for water and will dissolve
Hydrophobic (‘water fearing’) substances do not have an affinity for water and will not dissolve
“Like dissolves like”

27. Chemical reactions make and break bonds
Chemical reactions are the making and breaking of chemical bonds
The starting molecules are the reactants
The final molecules are the products
All chemical reactions are reversible – chemical equilibrium is reached when the forward and reverse reaction rates are equal
Reactions can release energy: exothermic or absorb energy: endothermic

28. Acids and Bases
Sometimes a hydrogen ion (H+) is transferred from one water molecule to another, leaving behind a hydroxide ion (OH-)
Solutes called acids increase the H+ concentration in water, while bases reduce the concentration of H+
The pH of a solution is defined by the negative logarithm of H+ concentration
For a neutral solution [H+] = 10-7
Acidic solutions have pH less than 7
Basic solutions have pH greater than 7
Most biological solutions have pH values in the range of 6-8
“OH, It’s a base!”

29. Gastric juice, lemon juice
Figure 2.23
pH Scale 1
Battery acid 2
Gastric juice, lemon juice 3
Vinegar, wine,
cola
Increasingly Acidic
[H]  [OH−] 4
Tomato juice
Beer
Acidic
solution 5
Black coffee
Rainwater 6
Urine
Saliva
Pure water
Human blood, tears
Neutral
[H]  [OH−] 7 8
Seawater
Inside of small intestine
Neutral
solution 9
Figure 2.23 The pH scale and pH values of some aqueous solutions 10
Increasingly Basic
[H]  [OH−]
Milk of magnesia 11
Household ammonia 12
Basic
solution
Household
bleach
Oven cleaner 13 14 29

30. Buffers
The internal pH of most living cells must remain close to neutral (pH 7)
Buffers minimize changes in the concentration of H+ and OH- in a solution
Carbonic acid is a buffer that contributes to pH stability in human blood

31. Homework: Define the following terms
Element
Covalent bond
Surface tension
Compound
Molecule
Specific heat
Atom
Electronegativity
Hydrophilic
Nucleus
Polar covalent bond
Hydrophobic
Proton
Ion
Acid
Neutron
Ionic bond
Base
Electron
Hydrogen bond pH
Atomic number
van der Waals forces
Buffer
Isotope
Reactant
Valence shell
Product
Chemical bond
Cohesion

32. Answer the following questions:
p. 20 #1-2
p. 24 #1-3
p. 28 #1-3
p. 29 #1-2
p. 37 (concept check) #1-5
p. 37 (scientific skills exercise) #1-4