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Chapter 4 Chemical Foundations for Cells

Chapter 4 Chemical Foundations for Cells. College Prep Biology Mr. Martino. Introduction. Certain chemicals enable organisms to function properly and are very useful Life is composed of matter : anything that has mass and occupies space

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Chapter 4 Chemical Foundations for Cells

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  1. Chapter 4Chemical Foundations for Cells College Prep Biology Mr. Martino

  2. Introduction • Certain chemicals enable organisms to function properly and are very useful • Life is composed of matter: anything that has mass and occupies space • Matter is composed of elements: substances that cannot be further broken down to other substances • Approx. 92 elements are naturally occurring

  3. Introduction – con’t • About 25 elements are essential to life • C, H, O, N make up about 96.3% of humans • Remaining 4% are considered trace elements: essential to life, but only in minute quantities (< 0.01)

  4. 4.1 Atoms • Each element consists of one kind of atom: (invisible) the smallest unit of matter that still retains the properties of an element • More than 100 subatomic particles – but only 3 important to us • Proton (+, nucleus: p+) • Neutron (neutral, nucleus: no) • Electron (-, orbits nucleus: e-)

  5. Atoms – con’t • Atoms differ in number of subatomic particles • Atomic number: unique number of protons possessed by each element • Neutral atoms have equal numbers of protons (+) and electrons (-) • Atomic mass:sum of the number of protons and neutrons in the nucleus • Protons and neutrons are almost identical in mass • Electrons are much smaller – approx. 1/2000 the mass of a proton • Atomic weight: same as atomic mass, but a whole number

  6. 4.3 Bonding of Atoms • Electron arrangement mainly determines how an atom behaves • Distance from nucleus determines E • The farther away, the E • Electron shell: (orbits) energy levels around the nucleus of an atom where electrons are found • 1st: 2 e- • 2nd: 8 e- • 3rd: 8 e-

  7. Atomic Electron Configurations • Bohr Model – shows all electrons arranges around a central nucleus

  8. Atomic Electron Configurations con’t.. • Lewis Structure – only shows the valence electrons of an atom of molecule • Valence electrons – outermost electrons that determine an atom’s reactive properties

  9. Electron Arrangement – con’t • When outer shells are full – the atoms are stable (inert) • If not, they react readily • Reactions enable them to fill their outer energy levels and become stable • All atoms “desire” to become stable • Chemical bonds: when atoms either share or transfer electrons in order to become stable • Compound: substance containing two or more different elements in a fixed ratio (H2O) • Mixture: two or more elements “intermingle” in varying proportions – no chemical reaction

  10. Covalent Bonds • Covalent bond: strong chemical bond in which two or more atoms share one or more pairs of outer-shell electrons • 2 or more atoms held together by covalent bonds form a molecule • Ex.: H2, O2, H2O • Very strong bond • Found in most organic compounds - sugars, fats and proteins

  11. Number of single covalent bonds an atom can form is equal to number of additional electrons needed H 1 O 2 N 3 C 4 Double bond: forms when 2 pairs of electrons are shared O2 O2 and H2 are molecules, but not compounds CH4 is a molecule AND a compound Covalent Bonds – con’t

  12. 4.4 Ionic Bonds • Ion: an atom or molecule with an electrical charge resulting from the gain or loss of one or more electrons • Ions of opposite charges attract each other and form an ionic bond • ex: Na+ Cl-

  13. Chemical Reactions • Chemical reaction: process that rearranges matter determining the way it behaves and interacts • 1st Law of Thermodynamics: amount of matter and energy in the universe is constant; they can be transferred and transformed - not destroyed • In other words, equations must be balanced! • 2 H2 + O2 2 H2O reactants = products

  14. Chemical Reactions – con’t • E is needed to make & break chemical bonds • Activation energy: the E required to start a chemical reaction • 2 main types of chemical reactions: • 1. Exergonic: more E is released than used and stored • Ex. Burning wood

  15. Chemical Reactions – con’t • 2. Endergonic: more E is used and stored than is released • Ex. Most cellular reactions • Metabolism: sum of all the reactions that occur within an organism

  16. Have many uses since readily absorbed by cells can be monitored on film or by Geiger counter as tracers: substances with a radioisotope attached 14CO2 to illustrate photosynthesis and other cellular processes PET Scanner – shows organs and cancer Used to: diagnose and treat AIDS, Alzheimer’s, arthritis, cancer, as well as other diseases of the brain, heart, lungs and bones Sterilize medical products Tissue grafts Medical research 4.2 Radioisotopes

  17. Isotopes • Isotopes: variant forms of elements that have a different number of neutrons • Number of protons and electrons remain same • Radioisotopes: nuclei decay spontaneously - giving off particles and energy (E) • 14C is radioactive

  18. Radioisotopes – con’t • Can also pose serious risks • Uncontrolled exposure damages DNA • Also can cause cancer • Naturally occurring radon gas is very dangerous • Found around uranium-bearing rocks

  19. Isotopes – con’t

  20. Isotopes – con’t • 1 out of every 3 people admitted to hospital has a nuclear medical procedure for diagnosis or treatment using radioisotopes • Nuclear medicine is $7-$10 billion/year industry • Isotopes sales are $100 million/year

  21. 4.4 Properties of Water • Water comprises 70 – 95% of living organisms • Water is electronegative: attraction between shared electrons in a bond • Greater the electronegativity, the more strongly an atom pulls electrons towards itself • H2, O2, CH4 exert equal pull and are considered to be nonpolar

  22. Water – con’t • Polar: when an molecule is negative at one end and positive at the other because an element is pulling the electrons more • Oxygen is one of the most electronegative elements • Ex. H2O

  23. Water – con’t • Polarity of water attracts other polar molecules • Hydrophilic substances: (water-loving): other polar molecules that are attracted to water • Ex. Sugars • Water’s polarity repels nonpolar molecules • Hydrophobic substances: (water-fearing) • Ex: oils

  24. Hydrogen Bonding • Water’s polarity leads to hydrogen bonding • H-bonds last only a few trillionths of a second • Each H2O molecule can H-bond with as many as 4 others • Common in all 3 states of matter • Allows for life on Earth

  25. Water is Cohesive • Cohesion: tendency for molecules to stick together • Due to H-bonds • Allows for capillary action • Surface tension: measure of how difficult it is to stretch or break the surface of a liquid • Permits organisms to walk on water • Unusually high in water

  26. Ice, Ice Baby No, not him….this ! • Hydrogen bonds of ice are stable • Each molecule is bonded to four others – forming a crystal • Causes the molecules to spread apart • Ice is less dense than water • Helps to maintain life

  27. Water Moderates Temperature • Temperature: measure of molecular motion • High heat capacity: takes a lot of E to temperature • Helps organisms to maintain constant temp. • Helps environment to moderate temp. • High heat of vaporization: takes a lot of E to cool and evaporate • Heat is released when more H-bonds form, slowing the cooling process • Evaporation also cools – water left behind body temperature

  28. The Universal Solvent • Solution: homogeneous, liquid mixture of two or more substances • Solute: substance that is dissolved • Solvent: dissolving agent • Water easily dissolves ions and polar molecules – so, it is referred to as the universal solvent

  29. Universal Solvent – con’t • Concentration: amount of solute in a solvent • Saturation:solution cannot hold any more solvent • When water is the solvent, results in - aqueous solution • Polarity allows for water’s tremendous solvent abilities –universal solvent

  30. 4.4 pH • Sometimes the H2O molecules separate into ions…H+ and OH- • Proper balance of ions is critical in organisms • Acid: compound that releases H+ in solution • Ex. HCl • Base: compound that removes H+ from solution and releases OH- • Ex. NaOH • Salts: release ions other than H+ and OH- in solution • Strong acid + strong base = salt • HCl + NaOH NaCl + H2O

  31. pH – con’t • pH Scale: describes how acidic or basic a solution is (potential hydrogen) • Each point is a 10x • 7 is neutral • pH inside most living cells is neutral – a change can be deadly

  32. pH – con’t • Buffer: resists change in pH by giving off or taking in H+ • Metabolic rxns are sensitive to very slight shifts in pH • Uncontrolled pH shifts in blood (pH 7.3 - 7.5) can cause a coma.

  33. pH Indicators • Blue Litmus Paper • Turns red in an acid • Red Litmus Paper • Turns Blue in a base • pH Paper • Changes color according to the color chart

  34. 4.5 Acid Rain • Acid precipitation:rain or snow with pH below 5.6 • Occurs mostly from sulfur and nitrogen oxides due to burning fossil fuels • Over past 20 years, several American, Asian and European lakes have “died” as a result • Eastern U.S. has experienced rain with pH as low as 2 – 3 • pH 1.7 fog has been recorded in Los Angeles

  35. Now, it is time to go study!!!!

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