Chapter 2 part 1 : Basic Chemistry

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 • Isotopes • 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 • 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 + – Cl Na Cl Na Sodium atom (Na)(11p+; 12n0; 11e–) Chlorine atom (Cl)(17p+; 18n0; 17e–) Sodium ion (Na+) Chloride ion (Cl–) Sodium chloride (NaCl)

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

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

23. Ionic Bonds + – Cl Na Cl Na 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 • 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