1. Chemistry Review

2. Atoms and Elements
All types of matter (solids, liquids and gases) are composed of atoms.
A substance that is composed of only one type of atom is called an element.
Elements are the simplest form of matter with unique chemical properties. They are charted on the periodic table based on some of their chemical characteristics.
There are 24 major elements that have various roles in the body.
These include structural, enzymatic, and homeostatic balance.
Compounds, like water, are formed by combining the atoms of different elements together.
Atoms may create various types of chemical bonds. 3 types of bonds include:
Ionic
Covalent
Hydrogen

5. Major Elements of the Human Body
Oxygen (O): Required for energy production during cellular respiration
Carbon (C): Organic back bone for fats, carbohydrates, amino acids and nucleic acids.
Hydrogen (H): Vital for energy( ATP) production
Nitrogen (N): Most abundant in atmosphere, Characteristic element of protein
Phosphorus (P): found in
DNA: Blue print for life
RNA : Vital for protein production
ATP : Cellular energy

6. The Atom Atoms – identical building blocks for each element
Atomic symbol – one- or two-letter chemical short hand for each element
Atomic number : # of protons in nucleus
periodic table
elements arranged by atomic number
Atomic weight – equal to the mass of the protons and neutrons
Isotope – atoms with same number of protons but a different number of neutrons

7. 6

8. Atomic Structure The nucleus consists of neutrons and protons
Neutrons – have no charge and a mass of one atomic mass unit (amu)
Protons – have a positive charge and a mass of 1 amu
Electrons are found orbiting the nucleus
valence electrons are located in the outermost shell
interact with other atoms to form bonds
Electrons – have a negative charge and 1/2000 the mass of a proton (0 amu)

9. Chemical Bonds
Electron shells, or energy levels, surround the nucleus of an atom.
Bonds are formed using the electrons in the outermost energy level
Valence shell – outermost energy level containing chemically active electrons
Octet rule – except for the first shell which is full with two electrons, atoms interact in a manner to have eight electrons in their valence shell.

10. Planetary Models of Elements
p+ represents protons, no represents neutrons

11. Chemically Reactive Elements
Reactive elements do not have their outermost energy level fully occupied by electrons therefore are able to interact with other elements

12. Chemically Inert Elements
Inert elements have their outermost energy level fully occupied by electrons therefore don’t interact with other elements.

13. Types of Chemical Bonds
Ionic:
Covalent
Hydrogen

14. Formation of an Ionic Bond
Ionic bonds form between atoms by the transfer of one or more electrons
Ionic compounds form crystals instead of individual molecules
Example: Na+Cl-(sodium chloride)

15. Formation of an Ionic Bond
A valance electron from Na is transferred to Cl
Cl now has 18e and 17p resulting in a – charge
Na has 10e and 11P resulting in a + charge.

16. Ion Formation
Ions are charged atoms resulting from the gain or loss of electrons.
Anions have gained one or more electrons therefore are negatively charged(-)
Cations have lost one or more electrons giving them a positive charge(+)
Typically occur between elements on opposite sides of the periodic table.

17. Electronegativity and Bond Formation
Elements on opposite ends of the periodic tables have a greater electronegative gradient. Ionic bonds result.
Elements that are closer to each other have smaller electronegative gradient thus form covalent bonds.

18. Covalent bonds
Covalent bonds are formed by the sharing of two or more electrons.
Covalent bonds are classified as Polar or Nonpolar.
When two atoms with similar electronegativities they share their valance electrons.
Nonpolar( neutral charge) bond results.
CO2, O2, N2
If there is a larger electronegative gradient between the atoms.
a polar covalent bond (charged compound) results.
H2O

19. Nonpolar Bonds
Electrons shared equally between atoms produce nonpolar bonds.
The negative charged electrons are spaced evenly between the 2 atoms resulting in a neutral charge.

20. Covalent Bonds

21. Double Covalent Bonds

22. Polar Covalent Bonds Uneven sharing of electrons produces polar bonds
One atom has a greater electronegativity.
This atom will have stronger pull on the shared electrons
The shared electrons spend more time closer to the nucleus of electronegative atom.
The addition of the shared electrons makes the electronegative atom partially negative charged, while the atom with a lower electronegativity becomes partially positively charged
Polar bonds occur between an electronegative atom (mostly O or N)
ex. H2O

23. Inorganic compounds Do not contain carbon Salts : NA+CL–
Water, salts, and many acids and bases
Minerals such as magnesium and calcium.
Salts : NA+CL–
contain cations other than H+ and anions other than OH–
Are electrolytes; they conduct electrical currents and function in various metabolic reactions.
Electrical activity of the nervous system
Vital for bone formation

24. Acid-Base Concentration (pH)
pH scale ranges from 0 to 14.
Acidic solutions have higher [H+] and a lower pH.
Considered proton donors
pH less than 7
Alkaline (basic) solutions have lower [H+] and a higher pH
Considered proton acceptors
pH greater than 7
Neutral solutions have equal H+ and OH– concentrations
pH = 7

25. pH Scale Acids release H+ and are therefore proton donors
HCl  H+ + Cl –
Bases release OH– and are proton acceptors
NaOH  Na+ + OH–

26. Buffers
The body has many mechanisms devoted to resist abrupt and large swings in the pH of body fluids.
These systems allow pH to remain relatively constant .
Approximately 7.4 (slightly basic)
Maintaining a stable pH is critical for creating an environment necessary for metabolic reactions.

27. Chemical Reactions
Chemical reactions in the body act by forming, breaking or rearranging bonds.
Chemical equations contain:
Relative amounts of reactants (starting chemicals) and products (finishing chemicals)
Number and type of reacting substances, and products produced

28. Synthesis and Decomposition Reactions

29. Oxidation-Reduction (Redox) Reactions
Reactants losing electrons are electron donors and are oxidized
Reactants taking up electrons are electron acceptors and become reduced
Na + Cl → Na+ + Cl-
Na is oxidized and Cl has been reduced
LEO THE LION SAYS GER

30. Forms of Energy Chemical – stored in the bonds of chemical substances
Energy from food
Electrical – results from the movement of charged particles
Household Appliances run on it
Mechanical – directly involved in moving matter
Machines such as cranes or bull dozers
Radiant or electromagnetic – energy traveling in waves
visible light, ultraviolet light, and X rays

31. First Law of Thermodynamics
Energy cannot be created or destroyed, but only change form.
During each conversion, some of the energy dissipates into the environment as heat.
Heat is defined as the measure of the random motion of molecules.
the second law states that "energy systems have a tendency to increase their entropy"

32. Energy The capacity to do work (put matter into motion)
Types of energy
Kinetic – energy in action. Ball rolling down a hill.
Potential – energy of position; stored (inactive) energy. Ball sitting on top of hill.

33. Fig. 8.2 (TEArt) Energy - the capacity to do work
Potential energy
Kinetic energy
Energy - the capacity to do work
kinetic - energy of motion
potential - stored energy

34. Factors Influencing Rate of Chemical Reactions
Temperature – chemical reactions proceed quicker at higher temperatures
Particle size – the smaller the particle the faster the chemical reaction
Concentration – higher reacting particle concentrations produce faster reactions
Catalysts – increase the rate of a reaction without being chemically changed
Enzymes – biological catalysts