Chapter 2
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Chapter 2. Chemical Foundations for Cells. Chapter Outline. Review of elements and atomic structure Radioactive elements and health/medicine Chemical bonding Ionic Covalent: nonpolar and polar Hydrogen “bonding” Properties of water Acids, bases, and buffers Chemical change.

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Chapter 2

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Chapter 2

Chemical Foundations for Cells


Chapter Outline

  • Review of elements and atomic structure

    • Radioactive elements and health/medicine

  • Chemical bonding

    • Ionic

    • Covalent: nonpolar and polar

    • Hydrogen “bonding”

  • Properties of water

  • Acids, bases, and buffers

  • Chemical change


Elements (2.1, 2.3)

  • Living organisms are composed of matter

    • Matter is composed of elements

      • Element - substance that cannot be broken down into other substances by chemical means

    • Elements are made up of atoms.

    • Atoms join together to make compounds.

      Atoms

      Compounds


Elements (2.1)

  • 92 naturally occurring elements

    • Life requires ~25 of these

    • ~96% of human body is made up of:

      • Carbon (C)

      • Hydrogen (H)

      • Oxygen (O)

      • Nitrogen (N)


Compounds (2.1)

  • Atoms of one element can join with atoms of other elements to form compounds.

    • A given compound is always made of the same elements combined in the same ways.

      • NaCl – table salt

      • H2O - water

      • C6H12O6 - glucose


Compounds of Life

  • Only living organisms have the ability to make the compounds of life:

    • Carbohydrates: C, H, O

    • Lipids: C, H, O

    • Proteins: C, H, O, N, S

    • Nucleic acids: C, H, O, N, P


Atoms (2.3)

  • An atom is the smallest unit of an element

  • Atoms are composed of 3 subatomic particles:

    • Protons

    • Neutrons

    • Electrons


Subatomic Particles


Subatomic Particles and the Elements

  • Each element has a unique number of protons.

    • Atomic number - number of protons in an atom

      • Elements are arranged by atomic number on the periodic table.

    • Atoms are neutral, therefore # p = # e


Isotopes

  • Number of neutrons is NOT on the periodic table for most elements….

    • Isotopes - atoms of a given element that differ in the number of neutrons in the nucleus

      • Mass number – sum of the protons and neutrons in an atoms’ nucleus

      • The periodic table shows the average of the mass numbers for the isotopes of an element.


Describing Isotopes

Mass number 12C

  • Isotopes of carbon

    • 12C carbon-12__ neutrons

    • 13Ccarbon-13__ neutrons

    • 14Ccarbon-14__ neutrons

      • All contain ____ protons and electrons.

        Carbon on the

        Periodic table


Isotopes and Radioactivity (RA)

  • RA isotope has an unstable nucleus

    • Nucleus emits energy and particles in an effort to become more stable

    • May change the number of protons in the nucleus and become a different element.


Radioactive Isotopes

  • Possible to target the energy and detect the radioactivity.

  • RA isotopes are used:

    • in research to track/follow molecules

    • in medicine to treat cancer and diagnose disease

      • Radiation therapy – treatment of localized cancer

      • PET - diagnosis


Radioactive Isotopes

  • Overexposure to RA isotopes is HARMFUL.

    • Energy emitted damages cells.

      • radiation therapy takes advantage of this, goal is to damage and kill cancer cells

    • Exposure to RA can also cause mutations that lead to cancers

      • Eg – exposure to RA element radon is the 2nd leading cause of lung cancer


Diagnosis - PET Scans

  • A radioactive tracer is put into the body.

    • Often RA glucose

  • The RA glucose goes to the parts of the body that use glucose for energy.

    • Cancers use glucose differently from normal tissue

  • As the radiotracer is broken down positrons are made. This energy appears as a 3-dimensional image on a computer monitor.


Electron Arrangement (2.5)

  • When compounds form, the electrons of the bonding atoms interact in attempt to obtain a more stable state.

  • Some electron arrangements are more stable than others…….see board


Chemical Bonding (2.6-2.7)

  • Chemical bonding – atoms gain, lose, or share electron(s) to obtain a stable number of electrons

    • Can be ionic bond or covalent bond


Chemical Bonding - Ionic

  • Ionic Bond – strong attractive force between oppositely charged ions

    • Atoms form ions by losing or gaining enough electron(s) to obtain a stable # of electrons in their outer shell


electron transfer

SODIUM

ATOM

11 p+

11 e-

CHLORINE

ATOM

17 p+

17 e-

SODIUM

ION

11 p+

10 e-

CHLORINE

ION

17 p+

18 e-


Ionic Bonding


Chemical Bonding - Covalent

  • Covalent Bond – bonded atoms share pair(s) of electrons and form molecules.

    • Occurs between nonmetals such as: C, O, H, N, P, S

    • Covalent bonding occurs in

      • H2

      • O2

      • H2O


Two Classes of Covalent Bonds

  • Nonpolar Covalent Bond – bonded atoms share electrons equally

    • Occurs between like atoms or between atoms with a similar ability to attract shared electrons

  • Polar Covalent Bond – unequal sharing of electrons by the bonded atoms

    • Occurs between atoms with very different ability to attract shared electrons


Two hydrogen atoms,

each with one proton,

share two electrons in a single nonpolar

covalent bond.

molecular hydrogen (H2)

H—H

Fig. 2-8b(1), p.25


water (H2O)

H—O—H

Two oxygen atoms share four electrons in a nonpolar double

covalent bond.

molecular oxygen (O2)

O=O


Types of Covalent Bonds

  • Nonpolar covalent – bonded atoms share the electrons equally

    • Examples of nonpolar bonds:

      • H2


  • Atoms with different electronegativity values form polar covalent bonds.

    • Electronegativity (EN) – measure of an atom’s ability to attract shared electrons in a covalent bond

    • Oxygen and nitrogen have fairly large EN values – often d -

    • Carbon and hydrogen have low EN values – often d +


  • Polar Covalent – unequal pull on shared electrons by the bonded atoms

    • Results in partial charges on the bonded atoms

d -

O

H

H

d +

d +


Common Polar Covalent Bonds

O-H N-H

C-O C=O

Label the polarity in each bond.


Forces between Molecules

  • Molecules are weakly attracted to each other by intermolecular (IM) forces,

  • The most important IM force in biology is the hydrogen “bond” (2.8)

    • Attractive force between d + H and d – O, N or F


  • Hydrogen “bond” is a weak attractive force between a d + hydrogen and a d-O, N, or F in a second polar bond

Water is a polar molecule

capable of hydrogen bonding.


Properties of Water

  • Water is cohesive and has high surface tension.

    • Cohesion – ability of molecules to stick together

    • Surface tension - ability to resist rupturing when under tension


Properties of Water

  • Water resists changes in temperature.

    • When heat is applied to an aqueous solution much of the heat (energy) is used to break hydrogen bonds, not to increase the movement of the molecules.


Properties of Water

  • Solid water (ice) is less dense than liquid water

    • Ice floats

      • Therefore, ice forms on the top of lakes and insulates the liquid water below.


  • Water is a good solvent for ionic compounds and small polar molecules.

    Water H bonds to polar

    molecules like ethanol


Water as solvent

  • Water pulls ions apart and hydrates them


Related Terms

  • Hydrophilic

    • Water loving

    • Capable of hydrogen bonding to water (polar)

  • Hydrophobic

    • Water “fearing”

    • Cannot hydrogen bond to water (nonpolar)


  • Acids, Base, and Buffers (2.14)

    • Many ions are dissolved in the fluids in/outside of cells – called electrolytes

      • Na+, Ca+2, K+

      • H+

    • Level of each ion is critical

      • Our focus is on H+ (hydrogen ions)


    Acids, Base, and Buffers

    • Acid: Substance that produces H+ when dissolved in water……….

      • Examples:

        • Hydrochloric acid – stomach acid

        • Lactic acid – made when cells run out of oxygen

        • Amino acids – building blocks of proteins


    Acids, Base, and Buffers

    • Base: substance that accepts H+1 (hydrogen ions) in water

      • Examples:

        • Sodium hydroxide - NaOH

        • Most nitrogen containing compounds

          • Ammonia – NH3

          • Urea – in urine

          • Amino acids – building blocks for proteins


    Acids, Base, and Buffers

    • Classify substances as acid, base or neutral by their pH

      • Acids: pH < 7

      • Base: pH > 7

      • Neutral: pH = 7

        • Pure water has a pH of 7

        • See page 28


    Acids, Base, and Buffers

    • How the pH scale works

      • The lower the pH the more acidic

      • The higher the pH the more basic (alkaline)

      • A difference of 1 pH unit is a 10-fold difference in acidity or alkalinity


    Why is pH important?

    • Most cells require a pH near 7.

    • Above or below this pH for too long and they die.

      • Proteins function only at specific pHs.

        • In lab you will determine the optimal pH for a protein that is needed to breakdown hydrogen peroxide in cells


    Acids, Base, and Buffers

    • Buffers: solution that resists changes in pH even when acid or base is added

      • Buffers can both produce H+ and neutralize H+

      • Buffers are key to maintaining pH homeostasis

        • Most body solutions are buffered


    Why is pH important?

    • Blood has a pH of 7.3 – 7.4

      • If the pH is above or below this range for more than a couple of days death occurs.

        • The blood buffer system helps keep blood pH in a range that supports life.


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