1. Basic Chemistry

2. Matter Matter- is anything that takes up space.
Each kind of matter has specific properties, or characteristics, that distinguish it from every other kind of matter.

3. Physical Properties of Matter
Physical properties- are characteristics that can be determined without changing the basic makeup of the substance.

4. Chemical Properties of Matter
Chemical properties- describe how a substance acts when it combines with other substances to form an entirely different substance, with new properties.

5. Atoms Atoms- are the basic building blocks of matter.
Atomic structure of an atom consists of a nucleus and electron cloud.

6. Atomic Structure
Nucleus- central core of the atom, consists of two types of subatomic particles: protons (p+) carry a positive electrical charge of +1, and electrically neutral neutrons (no). The nucleus thus has an overall positive charge.

7. Atomic Structure
Electron cloud- the region or space around the nucleus that is occupied by rapidly moving particles that are negatively charged, called electrons (e-).

8. Ions Ion- atom that has gained or lost one or more electrons.
Anion- atoms that have gained electron(s) and are now negative ions (Cl-, F-)
Cation- atoms that have lost electron(s) and are now positive ions (H+, Na+)

9. Ionic Bonds- formed when one or more electrons are transferred from one atom to another
Forms ions- positively or negatively charged atoms resulting from the gain or loss of electrons

10. Sodium
How many valence electrons does sodium have?

11. Chlorine
How many electrons in Chlorine’s valence shell?

12. If Na has one electron to give, and Cl needs another electron, both would be more stable if Na gave Cl its electron.
What would happen to the charges of both atoms?
Na is losing a negative charge (11+, 10-) so net charge of +1
Cl is gaining a negative electron, so it has a charge of –1
We write these charges as Na+ and Cl-
These are IONS.
What happens between positive and negative charges?
THEY ATTRACT– Forming an IONIC BOND

13. Ionic Bonding

14. Elements
Elements- substance composed of one type of atom, that can not be changed into a simpler substance by chemical means.
The number of protons in the nucleus determines the type of atom/element, and the number of protons in a given element is always constant.

15. Elements There are 92 naturally occurring elements.
Example: Carbon, you can heat it etc., but it will NOT turn into anything other than carbon.

16. Elements About 25 elements are essential to life:
96% of all matter is made of C, O, H, and N
The remaining 4% is made up mostly of P, S, Ca, and K.
Trace elements—those required by organisms in minute quantities(less than 0.01%)
Ex: Fe, I, B, Cr, Co, Mn, Zn

17. Elements (cont.)
Symbols- a simple, standard, abbreviated way of referring to elements. Usually one letter, sometimes two letters, but the second letter is not capitalized.
Ex. H = hydrogen, Na = sodium

18. Elements (cont.)
Isotopes- are atoms of the same element with different numbers of neutrons. Isotopes do have the same chemical properties.

19. Elements (cont.)
Radioactive Isotopes- isotopes with unstable nuclei that break down at a constant rate over time. As a result, they give off radiation which can be harmful. But they can also be used as “labels” or “tracers.”

20. Chemical Formulas
Chemical Formula- a group of symbols that show what type and how many atoms are present in a compound.

21. Chemical Formulas
Subscripts tell how many of each atom is present. No subscript it is understood to represent one atom.
Ex. H2O = 2 Hydrogen atoms and 1 Oxygen atom
Coefficients, a number in front of the entire formula, represents how many molecules you have.
Ex. 2H2O = 2 molecules of water
To find out the number of atoms in more than one molecule, multiply each subscript by the number in front of the formula.
Ex. 2H2O = 4 Hydrogen atoms and 2 Oxygen atoms

22. Compounds
Compounds- are two or more elements that are chemically combined. Each compound has its own special properties, which differ from the properties of the individual elements within that compound.
Ex. Chlorine (poisonous gas) combined with sodium = sodium chloride (table salt).

23. Molecules
Molecules- The smallest particle of a compound or element that can have stable, independent existence.

24. Molecules
A molecule of a compound contains two or more different atoms.
When atoms combine to form molecules a molecule of an element may consist of one, two or more atoms of that element.
Their outer energy levels become more stable.
One way of filling their energy levels is by sharing electrons. The electrons then move in the region between and around the nucleus.

25. Chemical Bonds
Chemical bonds- forces holding atoms in any molecule or compound together.
Chemical bonds result from the interaction of electrons in the outer energy levels of atoms.

26. Ionic Bonds
Ionic Bonds- electrical attraction between positively and negatively charged ions (atom that has gained or lost one or more electrons). The resulting charge of an Ionic Compound will be zero or neutral.

27. Ionic Bonds
Example: Sodium and Chlorine combine to make salt. Sodium atom transfers its single outer electron to the chlorine atom. Transfer gives each atom a stable outer energy level of eight electrons.

28. Covalent Bonds
Covalent Bonds- when atoms combine to form a molecule by sharing electrons.
Single covalent bond- atoms share 2 electrons
Double covalent bond- atoms share 4 electrons

29. Covalent Bonds

30. Covalent Bonds
Example: Methane (CH4 ) The Carbon atom has 4 electrons in its outer energy level, it need 4 more for a stable outer level. Each of the 4 Hydrogen atoms shares its single electron with the carbon atom, completing the carbon’s outer level. At the same time, the carbon atom shares one of its electrons with two of the hydrogen atoms. Altogether 6 electrons are shared in a methane molecule- one contributed by each hydrogen atom and two contributed by the carbon atom.

31. Hydrogen Bonds
Hydrogen Bonds: slight attractions that develop between the oppositely charged regions of nearby molecules. Although these forces are not as strong as ionic or covalent bonds, they can hold molecules together, especially when the molecules are large. Individual these bonds are weak and short lived, but collectively they are very strong.

32. Van der Waals Interactions
Van der Waals Interactions are weak attractions between molecules or parts of molecules that are brought about by localized charge fluctuations.
INTERmolecular bond- in between molecules

33. Chemical Equations
Chemical Equations are mathematical ways to represent chemical reactions.
Ex. C + 4H  CH4
Reactants Products
If 4H’s are shown in the products, they must be found in the reactants.

34. Chemical Reactions
Chemical reactions are due to the making and breaking of chemical bonds leading to a change in the composition of matter.
Ex: formation of H2O
Reactants—starting materials
Products—results
Matter is conserved in a chemical reaction…Law of Conservation of Matter.

35. Mixtures
Mixture- is the molecules of different substances mingling together (physically) with out chemically combining. Each substance retains all its chemical properties and its physical properties, which makes it possible to separate them physically.
Mixtures can contain solids, liquids and gases.

36. Solutions
Solutions- are mixtures that are the same throughout, but have variable compositions, depending on how much of one substance is dissolved in the other.

37. Solutions Two Parts of a Solution:
Solvent- the substance that can dissolve other substances. Water is often called the “Universal Solvent.”
Solute- the substance that dissolves in the solvent.
Examples: Salt Water: Water is the solvent and Salt is the solute. Kool Aid: Water is the solvent and Sugar is the solute.

38. Solutions
Concentration- the amount of solute dissolved in a given amount of solvent.
Molarity- the moles of solute per liter of solution. A 1.0 molar solution of sucrose will contain 1 mole (6.02 x 1023 molecules) of sugar in each liter of solution.

39. Suspensions
Mixtures of water and non-dissolved materials that are in pieces so small that the movement of water molecules does not allow them to settle out but keeps them “suspended.”
Blood is an example of a suspension.

40. Water and The Fitness of the Environment
Overview: The Molecule That Supports All of Life
Water is the biological medium here on Earth
All living organisms require water more than any other substance

41. Water, not Earth
Three-quarters of the Earth’s surface is submerged in water
The abundance of water is the main reason the Earth is habitable
Figure 3.1

42. Water is Polar The water molecule is a polar molecule:
Oxygen is significantly more electronegative than hydrogen.
Although oxygen and hydrogen share electrons, oxygen does not share fairly…it keeps the electrons more of the time.
Water can form four hydrogen bonds with neighboring molecules.

43. Polarity of Water
The polarity of water molecules allows them to form hydrogen bonds with each other and contributes to the various properties water exhibits. These are not as strong as ionic or covalent.
Hydrogen bonds + H
 –
Figure 3.2

44. Hydrogen Bonding

45. Water and Life
Emergent properties of water due to its polarity contribute to Earth’s fitness for life

46. Cohesion Water molecules exhibit cohesion
Cohesion is the bonding of a high percentage of the molecules to neighboring molecules. It is due to hydrogen bonding.
Water is extremely cohesive
Explains why drops of water “beads up” on a smooth surface
Allows water striders to
walk on surface of water

47. Water conducting cells
Cohesion helps pull water up through the microscopic vessels of plants
Water conducting cells
100 µm
Figure 3.3

48. Adhesion Attraction between molecules of different substances
Explains the dip in water level when reading a graduated cylinder

49. Surface tension is a measure of how hard it is to break the surface of a liquid. It is related to cohesion.
Figure 3.4

50. Moderation of Temperature
Water moderates air temperature by absorbing heat from air that is warmer and releasing the stored heat to air that is cooler

51. Evaporative Cooling
Evaporation is the transformation of a substance from a liquid to a gas
Hydrogen bonds are broken and heat is absorbed
Many animals rid their bodies of excess heat by allowing water is evaporate in the form of sweat, or by panting.

52. Insulation of Bodies of Water by Floating Ice
Solid water, or ice is less dense than liquid water. This is why ice floats in liquid water.
The hydrogen bonds in ice are more “ordered” than in liquid water, making ice less dense
Liquid water
Hydrogen bonds constantly break and re-form
Ice
Hydrogen bonds are stable
Hydrogen bond
Figure 3.5

53. Ice and Life
Since ice floats in water life can exist under the frozen surfaces of lakes and polar seas.
Ice formed in the winter in temperate zones does not sink and can be melted by the sun in the spring.
Water is most dense at 38oF or 4oC.

54. The Solvent of Life Some Examples: Salt in water Sugar in water
Water is a versatile solvent due to its polarity
It can form aqueous solutions with polar molecules
Some Examples:
Salt in water
Sugar in water
Carbon Dioxide in soda

55. The different regions of the polar water molecule can interact with ionic compounds called solutes and dissolve them
Negative oxygen regions of polar water molecules are attracted to sodium cations (Na+). +
Cl – –
Na+
Positive hydrogen regions of water molecules cling to chloride anions (Cl–).
Cl–
Figure 3.6

56. Water can also interact with polar molecules such as proteins
This oxygen is attracted to a slight positive charge on the lysozyme molecule.
This oxygen is attracted to a slight negative charge on the lysozyme molecule.
(a) Lysozyme molecule in a nonaqueous environment
(b) Lysozyme molecule (purple) in an aqueous environment such as tears or saliva
(c) Ionic and polar regions on the protein’s Surface attract water molecules. + –
Figure 3.7

57. Hydrophilic and Hydrophobic Substances
A hydrophilic substance is polar. It has an affinity for water. Hydrophilic means it “loves” water.
A hydrophobic substance is non-polar. It does not have an affinity for water. Hydrophobic means it “hates” water.

58. Solute Concentration in Aqueous Solutions
Amount of solute (dissolved particles) in a given amount of solvent
Which beaker has a higher concentration of solute?
Since most biochemical reactions occur in water it is important to learn to calculate the concentration of solutes in an aqueous solution
A mole represents an exact number of molecules of a substance in a given mass
Molarity is the number of moles of solute per liter of solution

59. Acids and Bases
Dissociation of water molecules leads to acidic and basic conditions that affect living organisms.

60. Dissociation of Water
Water can dissociate into hydronium ions and hydroxide ions
Changes in the concentration of these ions can have a great affect on living organisms H
Hydronium
ion (H3O+)
Hydroxide
ion (OH–) + –
Figure on p. 53 of water dissociating

61. Acids and Bases
An acid is any substance that increases the hydrogen ion concentration of a solution
A base is any substance that reduces the hydrogen ion concentration of a solution

62. The pH Scale
Increasingly Acidic
[H+] > [OH–]
Increasingly Basic
[H+] < [OH–]
Neutral
[H+] = [OH–]
Oven cleaner 1 2 3 4 5 6 7 8 9 10 11 12 13 14
pH Scale
Battery acid
Digestive (stomach)
juice, lemon juice
Vinegar, beer, wine,
cola
Tomato juice
Black coffee Rainwater
Urine
Pure water
Human blood
Seawater
Milk of magnesia
Household ammonia
Household bleach
Figure 3.8
The pH of a solution is determined by the relative concentration of hydrogen ions. It is low in an acid and high in a base.
The pH scale and pH values of various aqueous solutions
Ranges from 0 – 14
pH of human stomach is 2
pH of human blood is 7.4

63. Buffers The internal pH of most living cells must remain close to pH 7
Buffers are substances that minimize changes in the concentrations of hydrogen and hydroxide ions in a solution. They consist of an acid-base pair that reversibly combines with hydrogen ions.