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Basic Chemistry

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  1. Basic Chemistry

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

  3. Composition of Matter • Elements • Fundamental units of matter • 96% of the body is made from four elements • Carbon (C) • Oxygen (O) • Hydrogen (H) • Nitrogen (N) • Atoms • Building blocks of elements

  4. Atomic Structure • Nucleus • Protons (p+) • Neutrons (n0) • Outside of nucleus • Electrons (e-) Figure 2.1

  5. Atomic Structure • Nucleus • Protons (p+) • Neutrons (n0) • Outside of nucleus • Electrons (e-) Figure 2.1


  6. Identifying Elements • Atomic number • Equal to the number of protons that the atoms contain • Atomic mass number • Sum of the protons and neutrons

  7. The Periodic Table

  8. Isotopes and Atomic Weight • Isotopes • Have the same number of protons • Vary in number of neutrons Figure 2.3

  9. Isotopes and Atomic Weight • Atomic weight • Close to mass number of most abundant isotope • Atomic weight reflects natural isotope variation

  10. Radioactivity • Radioisotope • Heavy isotope • Tends to be unstable • Decomposes to more stable isotope • Radioactivity • Process of spontaneous atomic decay

  11. Molecules and Compounds • Molecule – two or more like atoms combined chemically • Compound – two or more different atoms combined chemically

  12. Chemical Reactions • 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 • Number of electrons has an upper limit • Shells closest to 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 • Have complete valence shells and are stable • Rule of 8s • Shell 1 has 2 electrons • Shell 2 has 10 electrons • 10 = 2 + 8 • Shell 3 has 18 electrons • 18 = 2 + 8 + 8 Figure 2.4a

  16. Reactive Elements • Valence shells are not full and are unstable • Tend to gain, lose, or share electrons • Allows for bond formation, which produces stable valence Figure 2.4b

  17. Chemical Bonds • Ionic Bonds • Form when electrons are completely transferred from one atom to another • Ions • Charged particles • Anions are negative • Cations are positive • Either donate or accept electrons PRESS TO PLAY IONIC BONDS ANIMATION

  18. Chemical Bonds • Covalent Bonds • Atoms become stable through shared electrons • Single covalent bonds share one electron • Double covalent bonds share two electrons Figure 2.6c

  19. Examples of Covalent Bonds PRESS TO PLAY COVALENT BONDS ANIMATION Figure 2.6a–b

  20. Polarity • Covalent bondedmolecules • Some are non-polar • Electrically neutralas a molecule • Some are polar • Have a positiveand negative side Figure 2.7

  21. Electronegativity and Polarity • Electronegativity – the affinity (want) for electrons. • A water molecule is polar because there is an uneven distribution of electrons between the oxygen and hydrogen atoms.

  22. TIMEOUT from WATER

  23. Van Der Waals Forces • Forces of attraction between molecules. • Not nearly as strong as ionic or covalent bonds. • Molecule – smallest unit of most compounds.

  24. A type of Van Der Waals force. Not as strong as Ionic or covalent bonds But the strongest of all other bonds. Hydrogen Bonds

  25. Cohesion The attraction between molecules of the same substance

  26. Adhesion The attraction between molecules of different substances.

  27. Capillary Action Combination of adhesion and cohesion

  28. Why does polarity and hydrogen bonding makes water so special? There are many reasons, but two main reasons are:

  29. #1 ICE Floats • ICE FLOATS • Water is one of the only substances that expands when is freezes. • Almost all other substance constricts when temperature drops. • This occurs because of the polarity of the molecule. • Ice is less dense than water.

  30. Water (Acids and Bases) What makes up water? H (hydrogen) + O(oxygen) What is the chemical formula for water? H2O What is the chemical equation for water? 2H + O H2O

  31. Do you think water can do this? H2O 2H + O Yes, but its more likely to do this: H2O  H+ + OH- In pure water, about 1 water molecule in 550 million will react to form hydrogen ions and hydroxide ions

  32. The pH Scale • measurement system used to indicate the concentration of hydrogen ions (H+) in solution; ranges from 0 to 14

  33. Acids  Acidic solutions contain higher concentrations of H+ ions than pure water and have pH values below 7.

  34. Bases  Basic, or alkaline, solutions contain lower concentrations of H+ ions than pure water and have pH values above 7.

  35. Acid and Base Buffers • Cells in our body like to keep their pH between 6.5 and 7.5 on the pH scale. • A buffer helps to do just that! • Buffer - weak acid or base that can react with strong acids or bases to help prevent sharp, sudden changes in pH

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

  37. 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

  38. Synthesis and Decomposition Reactions Figure 2.9a–b

  39. Patterns of Chemical Reactions • Exchange reaction (ABAC+B) • Involves both synthesis and decomposition reactions • Switch is made between molecule parts and different molecules are made Figure 2.9c