1. Lesson 2 Nature of Molecules/Properties of Water
September 1, 2015

2. Nature of an Atom
Matter is defined as any substance that has mass and takes up space
Mass vs. Weight
All matter is comprised of various chemical elements
An atom is the building block of an element.
There is one atom for each corresponding element
Each element is made of certain atoms that give it its unique characteristics
The element, gold, is made of only gold atoms

3. Composition of an Atom
Electron- particle that travels around the nucleus of an atom. Contains a negative charge
Proton- particle located inside the nucleus of an atom. Contains a positive charge
Neutron-particle located inside the nucleus of an atom. Contains no charge
Mass measured in daltons or atomic mass units (amu)

4. Structure of an Atom Cloud of negative charge (2 electrons) Electrons
Nucleus - - + + + +
Figure 2.4 Simplified models of a helium (He) atom
(a)
(b)
Helium (He)

5. Characteristics of an Atom
The protons and neutrons of an atom are held together by a force called the nuclear binding energy
Electrons orbit the nucleus due to the electrostatic forces b/w its negative charge and protons’ positive charge
When an atom has the same number of electrons as protons it is electrically neutral (no charge)

6. Characteristics of an Atom
Atoms are listed by their atomic number which is the number of protons in the nucleus.
Unique to each element
The sum of the masses of protons and neutrons in an atom is called the atomic mass
Both neutrons and protons have approximately the same weight which is measured in daltons
1.7x10-24 grams = 1 dalton
Neutrons and protons weigh roughly 2,000x the weight of the electron
Weight of an electron does NOT contribute to Atomic Mass

7. Atomic number/Atomic mass Designation

8. Ions
Electrons circle the nucleus of an atom in regions termed orbitals
Strong forces can dislodge electrons from its orbital or cause the atom to gain electrons
Atoms that contain more/less number of electrons than protons are called ions. Ions have a net charge (either + or -)
Cation—atom that has lost an electron (+ charge)
Anion—atom that has gained an electron (- charge)
Examples
Na+ = Cation
Cl- = Anion

9. Isotopes
Atoms of a single element that possess different number of neutrons are isotopes
The number of protons remain the same
Most elements in nature exist as a mixture of isotopes
Some isotopes are very unstable (radioactive isotopes)
Nucleus breaks apart giving off energy (radioactive decay)

10. Carbon Isotopes

11. Electrons determine how atoms interact
B/c electrons are attracted to the protons of a nucleus, it requires energy to keep the e(-) in the orbital of that atom
The greater the distance the electron is from the nucleus, the more energy is required
This makes the electron LESS stable (More Reactive)
Electrons within the atom have discrete energy levels and are denoted by rings
Not to be confused with orbitals!!!!

12. Electron Configurations
In atoms, electrons’ energy levels are denoted by electron shells
Each electron shell has a different energy level
The further the shell is from the nucleus, the higher the energy level
The outermost shell (valence shell), because of its energy level, can readily accept or lose electrons
The arrangement of an atom’s electrons in the electron shells is referred to as the electron configuration

13. Electron Shells
Only for denoting the energy of the electron and not the location of the electron!!!!

14. *Each shell can hold a specific number of electrons*
*Outermost shell is called the Valence Shell*

15. Can you Draw The Electron Configuration????
Lithium
Atomic Number = 3
Oxygen
Atomic Number = 8
Carbon
Atomic Number = 6
Aluminum
Atomic Number = 13

16. Section 2: Elements Found in Living Systems
92 elements occur naturally, each with a different number of protons and a different arrangement of electrons
The most common elements in living organisms are Oxygen(65.0%), Carbon(18.5%), Hydrogen(9.5%), Nitrogen(3.3%), Phosphorus(1.0%), Sulfur(0.3%)
When the elements are arranged in respect to their atomic numbers, they exhibit a certain pattern of chemical properties that repeats in groups of eight
This is referred to as periodicity

17. Element Periodicity
The eight-element periodicity is based on the interaction of the electrons in the outermost energy level of the different elements (valence electrons)
For most atoms, the outermost energy level can contain no more than eight electrons

18. Element Periodicity
Elements possessing a full outer shell are inert (non-reactive)
Helium, Neon, Argon
Elements possessing 7 electrons in its outer shell are highly reactive (readily accepts e-)
Fluorine, Chlorine, Bromine
Elements possessing 1 electron in its outer shell is also are highly reactive (readily loses e-)
This concept of an atoms tendency to completely fill their outer shell is known as the octet rule

19. Molecular Bonding
Atoms obtain their full complement of electrons in their valence shells by combining with other atoms to form molecules
Molecules that contain atoms of more than one element are called compounds
Molecules and compounds are held together through the attractive forces of their valence electrons. This force is called a chemical bond
Covalent Bond
Ionic Bond
Hydrogen Bond
Van der Waals Interactions

20. Covalent Bond
Bond where the valence electrons are shared between two atoms
Contains no net charge

21. Name and Molecular Formula
Electron
Distribution
Diagram
Lewis Dot
Structure and
Structural
formula
Space-
Filling
Model
(a) Hydrogen (H2) H H
(b) Oxygen (O2) O O
(c) Water (H2O) O H H
Figure 2.10 Covalent bonding in four molecules
(d) Methane (CH4) H H C H H

22. Polar vs. Non-polar Covalent Bonds
Atoms differ in their affinity for electrons. This is called electronegativity
Electronegativity increases left to right across a row of the periodic table
Electronegativity decreases down the column of the periodic table

24. Polar vs. Non-polar Covalent Bonds
Electrons are more likely move closer to the atoms with a greater electronegativity
This creates a region of partial negative charge near the more electronegative atom
Negative charge is relatively small
Molecules with similar atoms have no charge b/c atoms have same electronegativity (O2)
Molecules containing atoms of different electronegativities exhibit a charge (H2O)

25. Ionic Bond Atoms have gained or lost their valence shell electrons
Electronegativity between two atoms are so unequal that electrons are stripped
The attraction b/w ions with opposite charges holds them together (ionic bonds)
Na+ and Cl-
The combination of these to atoms are termed ionic compounds or salts.

26. Figure 2.12-2 Na Cl Na Cl Na Sodium atom Cl Chlorine atom+ - Na Cl Na Cl
Na Sodium atom Cl
Chlorine atom
Na+ Sodium ion (a cation)
Cl- Chloride ion (an anion)
Figure Electron transfer and ionic bonding (step 2)
Sodium chloride (NaCl)

27. Figure 2.13
Na+
Figure 2.13 A sodium chloride (NaCl) crystal
Cl-

28. Hydrogen Bond
Bond where a hydrogen atom that is covalently bonded to oxygen or nitrogen is attracted another oxygen or nitrogen atom
Results from atoms having different affinities for electrons. (electronegativity)
Weakest of the three bonds

29. Van der Waals Interactions
If electrons are distributed asymmetrically in molecules or atoms, they may accumulate by chance in one part of a molecule
Van der Waals Interactions are attractions between molecules that are close together as a result of these charges

30. Section 3: Properties of Water
One of the most important inorganic compounds is water.
Water, on average, makes up 65-75% of every living cell.
Water is the medium for most chemical rxns in the cell.
Because of water’s polarity (direction of charge) it acts as a great solvent (dissolving medium)
Stabilizes charges of substrates/intermediates
Water is both cohesive (attracted to itself) and adhesive (attracted to other polar molecules)

31. How water stabilizes ions
Sodium chloride crystal

32. Properties of Water
Water has a high specific heat. It resists changes in temperatures
Since most of our cells are composed of water, this helps us maintain a relatively constant internal temperature (homeostasis)
Water has a high heat of vaporization, which is defined as the amount of energy required to change 1 gram of a substance from a liquid to a gas
Many organisms dispose of excess body heat by evaporative cooling (sweating)

33. Section 4: Acids and Bases
Acid– is a substance that dissociates into one or more hydrogen ions (H+)
Also known as a proton donor
Examples are HCl and HNO3
Base—is a substance that dissociates into one or more hydroxide ions (OH-)
Proton acceptor
Examples are NaOH and KOH
The more H+ ions in a solution the more acidic it is and the more OH- ions in a solution the more basic it becomes

34. Measurements of Acidity
A solutions acidity is measure by a logarithmic pH scale.
pH scale range between 0-14.
0 is most acidic. 14 is most basic
An increase of 1 on the scale represents a ten fold increase in H+ ions
A solution at pH 5 contains 10X more H+ ions than a solution that is pH 6
Buffers are substances that minimizes the changes of H+ and OH-

35. Buffering Within Blood

36. Increasingly Acidic [H+] > [OH−] Increasingly Basic [H+] < [OH−]
Figure 3.10
pH Scale 1 2 3 4 5 6 7 8 9 10 11 12 13 14
Battery acid
Gastric juice, lemon juice H+ H+ H+
Vinegar, wine, cola H+
OH–
OH– H+ H+
Increasingly Acidic [H+] > [OH−] H+ H+
Acidic solution
Tomato juice Beer
Black coffee Rainwater
Urine
Saliva
OH–
OH–
Neutral [H+] = [OH−]
Pure water
Human blood, tears H+ H+
OH–
OH–
OH– H+ H+ H+
Seawater Inside of small intestine
Neutral solution
Figure 3.10 The pH scale and pH values of some aqueous solutions
Increasingly Basic [H+] < [OH−]
Milk of magnesia
OH–
OH–
OH– H+
OH–
Household ammonia
OH–
OH– H+
OH–
Basic solution
Household bleach
Oven cleaner

37. Critical Thinking
People take antacids wherever they exhibit any sign of acid reflux
The main ingredient in most antacids is Mg(OH)2. Therefore the chemical reaction that takes place in our digestive tract is:
Mg(OH)2 + 2HCl
MgCl2 + H2O
Which is the acid? Which is the base?