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Basic Chemistry

Basic Chemistry. Balance of Life. A human body is a highly organized, ever changing collection of chemicals. These chemicals must maintain a certain concentration or BALANCE in order for the human organism to survive. Balance of Life. Example of this balance is IRON (Fe)

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Basic Chemistry

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  1. Basic Chemistry

  2. Balance of Life • A human body is a highly organized, ever changing collection of chemicals. • These chemicals must maintain a certain concentration or BALANCE in order for the human organism to survive

  3. Balance of Life • Example of this balance is IRON (Fe) • Blood has 4 iron atoms in each blood cell to help carry oxygen (O2) • If to little iron (Fe) in blood than cells can’t carry O2 • Appears pale, fatigue, short of breath, rapid pulse

  4. Balance of Life Conversely – Too much Iron (Fe) leads to Hemochromatosis – to much iron in blood. • Iron gets stored in organs: heart, liver, pancreas destroying them over time. Disease Symptoms - joint pain, fatigue, weight loss, increased skin pigmentation. Maintaining a Chemical Balance is important for human function

  5. Introduction to Chemistry • Chemistry is the branch of science that considers the composition (what things are made up of) of matter and how this composition changes. • Understanding chemistry is essential for understanding how the body funtions

  6. Matter and Energy • Matter—anything that occupies space and has mass (weight) • Includes – all solids, liquids and gases • Energy—the ability to do work • Chemical • Electrical • Mechanical • Radiant • Potential

  7. Composition of Matter • Elements—fundamental units of matter • Living organisms require 20 different elements • 96% of the body is made from four elements • Carbon (C) • Oxygen (O) • Hydrogen (H) • Nitrogen (N) Element's are composed of Atoms – the smallest unit of an element

  8. Atoms • Atoms of a single element may be similar to each other but vary between elements. • Vary in: • Size • Weight • How the attract to each other - Chemical Bond (some atoms can combine with atoms like themselves but others can not)

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

  10. Atom Structure Since the protons in the nucleus are positively charged the and the electrons are negatively charged – when the same number of electrons match the number of protons – make a neutral charged atom

  11. Atomic Structure of Smallest Atoms Figure 2.2

  12. Identifying Elements • Atomic number—equal to the number of protons that the atom contains • Atomic Weight—sum of the protons and neutrons • IE: • Hydrogen has one proton = Atomic Weight =1 • Lithium has 4 protons and 3 neutrons atomic weight = 7

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

  14. Isotopes and Atomic Weight • Atomic weight • The number of protons plus number of neutrons • Atomic weight reflects natural isotope variation

  15. Radioactivity • Radioisotope – unstable nucleus • Heavy isotope • Tends to be unstable • Decomposes to more stable isotope • Radioactivity—process of spontaneous atomic decay

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

  17. Chemical Reactions • Atoms are united by chemical bonds • Atoms dissociate from other atoms when chemical bonds are broken

  18. 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

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

  20. 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 • Shell 3 can hold a maximum of 18 electrons

  21. 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

  22. Inert Elements – valance shell is complete Figure 2.5a

  23. 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 Figure 2.5b

  24. 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

  25. 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) Figure 2.6

  26. Ionic Bonds Cl Na Sodium atom (Na)(11p+; 12n0; 11e–) Chlorine atom (Cl)(17p+; 18n0; 17e–) Figure 2.6, step 1

  27. Ionic Bonds Cl Na Sodium atom (Na)(11p+; 12n0; 11e–) Chlorine atom (Cl)(17p+; 18n0; 17e–) Figure 2.6, step 2

  28. 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) Figure 2.6, step 3

  29. 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 PLAY Covalent Bonds

  30. Examples of Covalent Bonds Figure 2.7a

  31. Examples of Covalent Bonds Figure 2.7b

  32. Examples of Covalent Bonds Figure 2.7c

  33. Polarity • Covalently bonded molecules • Some are non-polar • Electrically neutral as a molecule • Some are polar • Have a positive and negative side Figure 2.8

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

  35. Hydrogen Bonds Figure 2.9

  36. 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 PLAY Disaccharides

  37. Synthesis and Decomposition Reactions Figure 2.10a

  38. Synthesis and Decomposition Reactions Figure 2.10b

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

  40. Patterns of Chemical Reactions Figure 2.10c

  41. Basic Chemistry

  42. Molecules and Compounds • Molecule – two or more same type of atoms bond together • Gases like hydrogen, oxygen, nitrogen common • Compound – Atoms of different molecules bond together • ie: two atoms of hydrogen bond with atom of oxygen = H2O • A molecule of a compound ALWAYS contains same kind and number of atoms – otherwise it changes – H2O2 = hydrogen peroxide

  43. Molecular Formula Molecular Formula represents the numbers and types of atoms in a molecule. • Displays the element symbol & # of atoms for each element. • ie: Water H2O = each water molecule has 2 atoms of hydrogen and one atom of oxygen • Sugar (glucose) – C6-H12-06 = 6 carbon atoms, 12 hydrogen atoms and 6 oxygen atoms

  44. Chemical Reactions #1 - Synthesis Chemical reactions either FORM or BREAK bonds between atoms, ions or molecules. Synthesis - When two or more atoms (reactants) bond to form a new more complex structure. • ie: H2 O2 synthesis H2O and gives off a single O molecule • A + B AB • Synthesis is important in growth and repair

  45. Chemical Reactions # 2 - Decomposition Decomposition – when bonds within a molecule break into simpler atoms • ie: AB A + B • Important in the breakdown of food for energy

  46. Chemical Reaction # 3 Exchange Reaction • Exchange Reaction – parts of two different types of molecules change position. • ie: AB + CD AD + CB • Example of this type is when acid reacts with a base Reversible Reactions – some reactions can go either way • (opposite arrows)

  47. Independent - Class work • Complete packet questions 1 – 12.

  48. Biochemistry: Essentials for Life • Organic compounds • Contain carbon and hydrogen • Most are covalently bonded • Example: C6H12O6 (glucose) • Inorganic compounds • Lack carbon • Tend to be simpler compounds • Example: H2O (water)

  49. Important Inorganic Compounds • Water • Most abundant inorganic compound in living organism makes up 2/3 of the weight. • Vital properties • High heat capacity • Polarity/solvent properties – many substances dissolve in water – when smaller, chemicals have greater chance or reacting with each other • Chemical reactivity • Cushioning

  50. Important Inorganic Compounds • Salts • Easily dissociate into ions in the presence of water • Vital to many body functions • Include electrolytes which conduct electrical currents • Smaller ions are important to body function ie: • Transport substances in and out of cell • Muscle contraction • Nerve impulse conduction

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