Basic Chemistry and Properties of Water

1. Basic Chemistry and Properties of Water • Matter: anything that occupy space and has mass • All solids, liquids, and gases • All living and nonliving things are composed of atoms

2. Atoms • Atoms: The basic unit of matter • Atoms contain: • Sub-atomic Particles: • Made of protons, neutrons, electrons

3. Atoms Continued: • Protons: have a positive charge, found in the atomic nucleus • Neutrons: are neutral and found in atomic nucleus • Electrons: have a negative charge and orbit the nucleus in shells • Atoms form Elements: • Elements are pure substances consisting entirely of one type of atom • example H2 = hydrogen, O2 = oxygen

4. All the elements are listed in the periodic table Atomic number Element symbol Mass number Figure 2.2

5. Atomic Number = number of protons, which also equals the number of electrons. • Atomic Mass (atomic weight) or mass number = # of protons + # of neutrons.

6. Elements and Isotopes • Elements: are pure substances consisting of entirely one type of atom. • Example: H = Hydrogen, Na = Sodium • Isotope: atoms of an element with the same number of protons but a different number of neutrons. • Example: isotopes of carbon, nonradioactive Carbon-12 (C12) has 6 electrons, 6 protons, and 6 neutrons. • Carbon 13 (C13) nonradioactive has 6 electons, 6 protons and 7 neutrons. • Radioactive Carbon-14 (C14) has 6 electrons, 6 protons, and 8 neutrons.

7. Carbon-12 and Carbon - 13

8. Electrons and Electrons Shells • Orbit the atomic nucleus at fixed distances • Negatively charged • Orbits called shells • In order to become stable, they fill their shells • First shell is full with 2 electrons • The last or valence shell holds 8 electrons

9. How atoms interact to form molecules • Atoms can interact with one another when the outermost shell of an atom is not at full capacity • Example: Carbon has 4 electrons in outmost shell but can hold eight. This will allow carbon to share up to 4 electrons.

10. Atoms of the four elements most abundant in life First electron shell (can hold 2 electrons) Outermost electron shell (can hold 8 electrons) Electron Hydrogen (H) Atomic number = 1 Carbon (C) Atomic number = 6 Nitrogen (N) Atomic number = 7 Oxygen (O) Atomic number = 8 Figure 2.7

11. Atoms Interact • Atoms interact with one another according to two basic principles. 1. Atoms do not interact if the outermost electron shell is completely full or empty. Example: Noble gases are stable because full outer shells. 2. Atoms do react with other atoms when the outer shell is only partially full. Example: Carbon

12. Chemical Bonds • Atoms in compounds are held together by chemical bonds. • Types of chemical bonds: Ionic bonds and covalent bonds.

13. Chemical Bonds • Valence Shell Electrons: Atoms want to be stable like the noble gases, so they want full shells. The first shell fills with 2 and the rest with 8 electrons. • Ionic Bonds: formed when one or more electrons are transferred from one atom to another. They are held together by electrical attractions between positive and negative charges. Example: Formation of sodium chloride (NaCl)

14. Ionic Bonds: Sodium Chloride Sodium gives up an election to chlorine and thus becomes a sodium ion with a positive charge. Chlorine accepts an electron becomes the chloride ion and has a negative charge. (form sodium chloride held together with ionic bond – ionic bonds are weak)

15. Covalent Bonds • Covalent Bonds: A chemical bond between atoms that share electrons • When atoms share 2 electrons it is a single bond • Sharing 4 electrons is a double bond.

17. Covalent Bonds and Ionic Comparison Figure 2.9

18. The Structure of Water • Studied in isolation, the water molecule is deceptively simple • Its two hydrogen atoms are joined to one oxygen atom by single covalent bonds H H O Unnumbered Figure 2.2

19. But the electrons of the covalent bonds are not shared equally between oxygen and hydrogen. This is called a Polar-Covalent Bond. • This unequal sharing makes water a polar molecule () () () () Figure 2.11a

20. The polarity of water results in weak electrical attractions between neighboring water molecules () Hydrogen bond () () () • These interactions are called hydrogen bonds () () () () (b) Figure 2.11b

21. Hydrogen Bonds Water molecules carry slight charges Electrons favor O over H Hydrogen bonds form between O of one and H of another + – O H + H – O H + HydrogenBonds + H +

22. Water’s Life-Supporting Properties • The polarity of water molecules and the hydrogen bonding that results explain most of water’s life-supporting properties • Water’s cohesive nature • Water’s ability to moderate temperature • Versatility of water as a solvent

23. The Cohesion of Water • Water molecules stick together as a result of hydrogen bonding • This is called Cohesion • Cohesion is vital for water transport in plants Microscopic tubes Figure 2.12

24. Surface Tension is the measure of how difficult it is to stretch or break the surface of a liquid • Hydrogen bonds give water an unusually high surface tension Figure 2.13

25. Cohesion and Surface Tensionamong Water Molecules

26. Adhesion and Capillary Action Adhesion:Attraction between water molecules and different surfaces. Capillary Action: Ability of a liquid to flow against gravity and rise in a narrow thin tube.

27. Water as the Solvent of Life • Because water is polar and has both positive and negative poles, it works well as a solvent • When salt is put into water, the positive charged hydrogen ends of water are attracted to the negatively charged chloride ions (Cl-) • Negatively charged oxygen ends of water surround the positively charged sodium ions (Na+) • Sodium and chloride ions are surrounded, crystals separate Salt crystal Ion in solution Figure 2.16

28. NaCl Dissolved by Water NaCl is ionicly bonded Water is polar and charged O sticks to Na+ H sticks to Cl-

29. Acids, Bases, and pH • Acids have greater amount of H+ ions then pure water and have a pH < 7 • Pure water has a ph of 7, neutral • Base (Alkaline) • Has less H+ ions and more OH (hydroxide ions)then pure water and has a pH > 7 • Acid

31. Oven cleaner Household bleach Household ammonia Basic solution Milk of magnesia Seawater Human blood Pure water Urine Neutral solution Tomato juice Grapefruit juice Lemon juice; gastric juice Acidic solution pH scale Figure 2.17

33. Buffers are substances that resist pH change. Help maintain a solution at a relatively constant pH. They are weak acids or bases that can react with a strong acids or base to prevent sudden pH changes. • They accept H+ ions when they are in excess • They donate H+ ions when they are depleted • Buffering is not foolproof • Example: acid precipitation Figure 2.18