1. Basic Chemistry

2. Matter Anything that takes up space and has mass.
Can be found in three states:
Solid= Definite shape and volume
Liquid= No definite shape, definite volume
Gas= no definite shape, no definite volume

3. Energy
Energy is not a state of matter, but is measured based on what it does to matter.
Energy- the capacity to do work or put matter into motion.
Kinetic energy- energy in action.
Potential energy- stored energy.

4. Forms of Energy
Chemical Energy- stored in the bonds of chemical substances. (i.e. in the phosphate bonds of ATP)

5. Electrical energy
Electrical energy- movement of charged particles. (i.e. ions that travel across a cell membrane)

6. Mechanical energy
Energy directly involved in moving matter. (i.e. pedaling a bicycle)

7. Radiant/electromagnetic energy
Energy that travels in waves. This includes UV rays (i.e. solar rays), X rays, radio waves.

9. Atoms and Elements
Elements: unique substances that cannot be broken down into smaller substances. (i.e. phosphorus, carbon)
Each element is composed of smaller pieces called atoms.

10. Structures in an element
Nucleus: contains neutrons and protons
Neutrons: particles with no charge (1 amu)
Protons: particles with a positive charge (1 amu)
Electrons: particles with a negative charge (0 amu)
Electron cloud: space outside of the nucleus where electrons float.

11. Identifying Elements

12. Radioisotopes
Heavy isotopes are unstable. Because of this instability, they begin to decay.
Alpha decay= loss of 2p and 2n
Beta decay= loss of electron like particles.
Gamma decay=loss of electromagnetic radiation

13. Elements and Compounds
When multiple elements come together to form one new unit, a compound is formed. (i.e. H2O)

14. Mixtures and solutions
Mixture: composed of two or more elements or compounds. These elements and compounds are physically mixed, so the amount of each compound may vary depending on the sample.
Solutions: Homogenous: meaning there are the same amount of each element throughout the mixture.

15. Molarity
Molarity represents the concentration of a compound in a solution.
To calculate the molarity of a substance, add the atomic mass of each element in the compound. That total number represents the number of grams needed to obtain one mole of that substance.

16. Colloids A.k.a. emulsions
Heterogeneous, often appear as a translucent, milky color. They can undergo a sol-gel transformation (changing reversibly from solid to liquid).

17. Suspensions
Heterogeneous, large visible solutes that tend to settle out.
A great example of this is blood.

18. Distinguishing mixtures and compounds
Often, no chemical bonds occur in mixtures
Most mixtures can be broken down by physical means
Compounds are homogenenous

19. Electrons and their importance
When elements “bond” what they are doing is sharing electrons to some extent.
Electrons float in different levels, called shells, which are representative of the energy each electron contains.
The highest shell is called the valence shell. The electrons here contain the most energy, and are the ones that are shared.

21. Types of chemical bonds
Ionic
Covalent

22. Ionic Bonds
Formed between a cation, a positively charged atom, and an anion, a negatively charged particle.
When two atoms with a high electronegativity difference interact, the more electronegative (usually on the far right of the periodic table) will strip 1 or more electrons from the valence shell of another atom (usually on the far left of the periodic table).

23. Continued…
Elements like to fill their valence shell with a maximum of eight electrons, so when two elements like sodium and chlorine interact, chlorine will take an electron from sodium.
Sodium having only one electron in its valence shell will now have a filled shell because it’s new valence shell will be the complete shell below.

24. Continued…
Since chlorine has one more electron now, it becomes negatively charged.
Since sodium has one more proton now, it becomes positively charged.
These two opposite charges attract each other, forming the molecule NaCl.

26. Covalent Bonds
When two elements have a similar electronegativity difference, it is possible for each element to “borrow” the electrons of the other through covalent bonds.

27. Continued…
To fill its outer shell, two oxygen atoms will share two of each of their electrons.
These electrons tend to spend roughly the same amount of time with each atom.

28. Polar and nonpolar molecules
Polar: slight charges formed by an electronegativity difference in the elements that cause electrons to stay closer to one of the covalent atoms, meaning one side will more often have a negative charge.
Nonpolar: the shape of the compound causes the electrons to spend equal time with each element.

29. Polar and nonpolar molecules
Polar Nonpolar

30. Hydrogen Bonding
The ability of a hydrogen atom to form a temporary bond with electronegative atoms. DO NOT CONFUSE with hydrogen atoms bonding in a covalent bond.
Hydrogen bonds are only temporary.

31. Why are hydrogen bonds so important?
Inside our DNA, there are four nitrogenous bases. They are adenine, thymine, cytosine, and guanine. These four bases are paired adenine with thymine and cytosine with guanine. These bases do not actually lock together. They are held together with hydrogen bonds that form and break over and over again.

32. Chemical Reactions
When chemical bonds are formed, rearranged, or broken, a chemical reaction occurs.
The symbolic form for a chemical equation is written as a chemical equation.

33. Chemical equations Chemicals reacting= reactants
Chemicals formed= products
Reactants  products
A + B  AB

34. Types of chemical reactions
Synthesis reaction- formation of new bonds.
A + B  AB
A and B were once separate atoms, but have now come together by forming a bond, creating molecule AB.

35. Decomposition Reactions
Decomposition reaction- breaking of previous bonds.
AB  A + B
AB was once a single molecule. After a reaction occurs (like an enzyme splitting a larger molecule) A and B are now separate atoms.

36. Exchange Reactions
Exchange reaction are a combination of synthesis and decomposition reaction.
Single Exchange
AB + C  AC +B
Double exchange
AB + CD  AD + CB
Old bonds between molecules are broken, and the new atoms exchange partners, forming new molecules.

37. Redox reactions Deals with the giving and receiving of electrons.
Oxidation- Loss of electrons (gains + charge)
Reduction- Gain of electron (gains a – charge)

38. Energy flow in a reaction
Exergonic reactions- Energy is released. Usually when bonds are broken.
Endergonic reactions- energy is absorbed. These reactions are when bonds are formed.

39. Chemical equilibrium
If any two compounds can be combined, then they can also be taken apart.
When the rate of compounds being combined and rate of decomposition are equal, then the chemicals have reached equilibrium.

40. Factors that influence chemical reactions
Temperature
Concentration
Particle size
Catalysts

41. Biochemistry
The study of the chemical composition and reactions of living matter.
There are two types of biochemical compounds:
Organic
Inorganic

42. Inorganic compounds
Compounds that do not contain carbon are inorganic compounds.
Water is the most important of all inorganic compounds.

43. Why is water such an important compound?
High specific heat
High vaporization point
Polar solvent properties (can dissolve almost any compounds)
Reactivity (through hydrolysis)
Cushioning

44. Salts
Salts in solutions dissociate from each other and form their respective ions.
Those ions carrying a charge allow for the flow of electricity to occur. Without salts, we would not be able to transmit electrical nerve impulses throughout the body.

45. Electrolytes
Salt is an electrolyte because when it dissociates in a solution, it gains the ability to conduct electricity.
Acids and bases are also electrolytes

46. Acids
Sour taste
Acids- react with metals. Releases hydrogen ions in detectable amounts.
A hydrogen ion is just a proton, so another name for an acid would be proton donor.
Most every acid will start with a “H” in its chemical formula.
i.e. HCl, H2SO4

47. Bases
Bitter taste
They can take up hydrogen in detectable amounts, meaning they are proton acceptors.
Most, but not all bases, have a hydroxide group (-OH-) attached to it.
i.e. NaOH, NH3

48. pH pH is a measure of the acidity or alkalinity of a solution.
Low pH= highly acidic (more hydrogen ions)
High pH= highly basic (less hydrogen ions)

49. The pH Scale

50. Buffers Designed to prevent sharp changes in pH.
Without them, if we ingest anything that has a pH difference from our own blood, we may accidently alter our pH too far and cause ourselves serious bodily harm, if not death.

51. Organic compounds
Contain Carbon in their structures. Usually many carbons are linked together to form a carbon backbone.

52. Carbohydrates Includes sugar and starch
They are classified based on their size and solubility.
can be:
Monosaccharide- one sugar
Disaccharide- two sugars
Polysaccharide- many sugars

53. Lipids A.k.a. fats Insoluble in water.
Allows for a long term way to store energy.
Includes:
Triglycerides
phospholipids
steroids
Eicosaniods

54. Proteins
Any of a large group of nitrogenous organic compounds that are essential constituents of living cells; consist of polymers of amino acids; essential in the diet of animals for growth and for repair of tissues.
Includes:
Amino acids
Globular proteins
Fibrous proteins