1. Chapter 2 part 1: Basic Chemistry

2. What is Chemistry?
Chemistry is sometimes called the central science, as its principles are central to understanding all aspects of science, including biology and physiology.
Pharmacology is the science of
drugs, including their composition, uses, and effects on the body.
Drugs are chemical compounds that have specific effects on the body’s mechanisms.

3. Matter and Energy
Matter—anything that occupies space and has mass (weight)
Energy—the ability to do work
Chemical
Electrical
Mechanical
Radiant

4. Composition of Matter Elements—fundamental units of matter
Atom-is the smallest unit of an element that retains its chemical properties.
Every pure element is composed of only one kind of atom. For example carbon is composed of only carbon atoms.

5. Chemical Composition of the Body
CHEMICAL ELEMENTS % BODY COMPOSITION
Carbon (C) Nitrogen (N) 96%
Oxygen (O) Hydrogen (H)
Calcium (Ca) Phosphorus (P) 3%
Potassium (K) Sulfur (S)
Iron (Fe) Chlorine (Cl) Trace quantities
Iodine (I) Sodium (Na)
Magnesium (Mg) Copper (Cu)
Manganese (Mn) Cobalt (Co)
Zinc (Zn) Chromium (Cr)
Fluorine (F) Molybdenum (Mo)
Silicon (Si) Tin (Sn)

6. Atomic Structure Every atom is composed of three different particles.
1) Protons- which are positively charged particles, located inside the nucleus of the atom.
2) Neutrons- which is a particle with no charge. Also located inside the nucleus.
Because there is no negative charges inside the nucleus to counter act the positively charged protons, the overall charge of the nucleus is positive.
3) Electrons- These are particles that have a negative charge and are found outside the nucleus.

7. Atomic Structure of Smallest Atoms

8. Identifying Elements
Atomic number—equal to the number of protons that the atom contains
Atomic mass number—sum of the protons and neutrons
Chemical symbol- A 1,2 or 3 letter abbreviation used to represent different elements.
Atomic #
Chemical Symbol
Atomic mass #

9. Isotopes and Atomic Weight
Have the same number of protons
Vary in number of neutrons

10. Radioactivity Radioisotope
Heavy isotope
Tends to be unstable
Decomposes to more stable isotope
Radioactivity—process of spontaneous atomic decay.
Isotopes have important medical uses.
Radioisotopes are frequently used by radiologists and oncologists to diagnose and treat diseases.

11. Molecules and Compounds
Molecule—two or more like atoms combined chemically
Compound—two or more different atoms combined chemically
Figure 2.4

12. Chemical Bonding Atoms are united by chemical bonds
Atoms dissociate from other atoms when chemical bonds are broken

13. Electrons and Bonding
Electrons occupy energy levels called electron shells
Electrons closest to the nucleus are most strongly attracted
Each shell has distinct properties
The number of electrons has an upper limit
Shells closest to the nucleus fill first

14. Electrons and Bonding
Bonding involves interactions between electrons in the outer shell (valence shell)
Full valence shells do not form bonds

15. Inert Elements
Atoms are stable (inert) when the outermost shell is complete
How to fill the atom’s shells
Shell 1 can hold a maximum of 2 electrons
Shell 2 can hold a maximum of 8 electrons

16. Inert Elements
Atoms will gain, lose, or share electrons to complete their outermost orbitals and reach a stable state
Rule of eights
Atoms are considered stable when their outermost orbital has 8 electrons
The exception to this rule of eights is Shell 1, which can only hold 2 electrons

17. Inert Elements
Figure 2.5a

18. Reactive Elements Valence shells are not full and are unstable
Tend to gain, lose, or share electrons
Allow for bond formation, which produces stable valence

19. Chemical Bonds Ionic bonds Ions
Form when electrons are completely transferred from one atom to another
Ions
Charged particles
Anions are negative (gain e)
Cations are positive (lose e)

20. Ionic Bonds + – Na Cl Na Cl Sodium atom (Na) (11p+; 12n0; 11e–)
Chlorine atom (Cl) (17p+; 18n0; 17e–)
Sodium ion (Na+)
Chloride ion (Cl–)
Sodium chloride (NaCl)

21. Ionic Bonds Na Cl Sodium atom (Na) (11p+; 12n0; 11e–)
Chlorine atom (Cl) (17p+; 18n0; 17e–)

22. Ionic Bonds Na Cl Sodium atom (Na) (11p+; 12n0; 11e–)
Chlorine atom (Cl) (17p+; 18n0; 17e–)

23. Ionic Bonds + – Na Cl Na Cl Sodium atom (Na) (11p+; 12n0; 11e–)
Chlorine atom (Cl) (17p+; 18n0; 17e–)
Sodium ion (Na+)
Chloride ion (Cl–)
Sodium chloride (NaCl)

24. Chemical Bonds Covalent bonds
Atoms become stable through shared electrons
Single covalent bonds share one pair of electrons
Double covalent bonds share two pairs of electrons

25. Examples of Covalent Bonds

26. Examples of Covalent Bonds

27. Polarity Covalently bonded molecules Some are non-polar Some are polar
Electrically neutral as a molecule
Some are polar
Have a positive and negative side

28. Chemical Bonds Hydrogen bonds Weak chemical bonds
Hydrogen is attracted to the negative portion of polar molecule
Provides attraction between molecules

29. Patterns of Chemical Reactions
Synthesis reaction (A + BAB)
Atoms or molecules combine
Energy is absorbed for bond formation
Decomposition reaction (ABA + B)
Molecule is broken down
Chemical energy is released

30. Synthesis and Decomposition Reactions
Figure 2.10a

31. Synthesis and Decomposition Reactions
Figure 2.10b