Collin County Community College BIOL 2401

Anatomy &Physiology I Chapter 2

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Matter §

The “stuff” of the universe

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Anything that has mass and takes up space

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States of matter §

Solid – has definite shape and volume

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Liquid – has definite volume, changeable shape

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Gas – has changeable shape and volume

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Energy §

The capacity to do work (put matter into motion)

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Types of energy

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Kinetic – energy in action

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Potential – energy of position; stored (inactive) energy

Forms of Energy §

Chemical – stored in the bonds of chemical substances

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Electrical – results from the movement of charged particles

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Mechanical – directly involved in moving matter

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Radiant or electromagnetic – energy traveling in waves (i.e., visible light, ultraviolet light, and X-rays)

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Energy  Form  Conversions §

§

Energy is easily converted from one form to another During conversion, some energy is “lost” as heat

Energy state 1

Energy State 2

Heat

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Why  should  we  care  ,  why  does  it  Matter  ? § §

Because matter has energy Because life depends on the inter-conversion of one energy form into another §

Solar energy into plant material

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Plant matter into animal matter

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Animal matter into animal matter

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Good chemicals into bad chemicals ( such as free radicals)

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Composition   of  Matter § §

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Matter is made up of elements Elements are unique substances that cannot be broken down by ordinary chemical means. Elements are made up from Atoms Each element has unique physical and chemical properties Each Atom is represented by a symbol – one- or two-letter chemical shorthand for each element.

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Major  Elements   of  the  Human  Body §

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Major elements of that make up 96.1 % of the human body are §

Oxygen (O),

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Carbon (C),

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Hydrogen (H),

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Nitrogen (N)

Lesser elements make up 3.9% of the body and include: §

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Calcium (Ca), phosphorus (P), potassium (K), sulfur (S), sodium (Na), chlorine (Cl), magnesium (Mg), iodine (I), and iron (Fe)

Trace elements make up less than 0.01% of the body §

They are required in minute amounts, and are found as part of enzymes

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Atomic  Structure §

The nucleus of an atom consists of neutrons and protons §

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Neutrons – have no charge and a mass of 1 atomic mass unit (amu) Protons – have a positive charge and a mass of 1 amu

Electrons are found orbiting the nucleus §

Electrons – have a negative charge and 1/2000 the mass of a proton (~ 0 amu)

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Models   of  the  Atom §

Planetary Model : §

§

electrons move around the nucleus in fixed, circular orbits

Orbital Model : §

regions around the nucleus in which electrons are most likely to be found

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Figure  2.1

Identification   of  Elements §

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Each element has a specific Atomic number : equal to the number of protons ( in a neutral atom, number of protons = number of electrons) Each element has a specific Mass number : equal to the mass of the protons and neutrons Isotope : are similar atoms ( they thus have the same number of protons, same atomic number) with a different number of neutrons (different mass) Radioisotopes : atoms that undergo spontaneous decay called radioactivity Atomic weight of an atom : average of the mass numbers of all isotopes

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Identification   of  Elements §

Protons and neutrons reside in the nucleus

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Electrons “buzz” around the nucleus in electron shells

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The first electron shell holds a maximum of 2 electrons

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The next electrons shells hold a maximum of 8 electrons

Each electron shell has a different energy level. The further away from the nucleus, the more energy in those electrons to keep them away from the nucleus. The outer shell is called the Valence shell Remember that the number of electrons is always equal to the number of protons in a neutral atom

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Identification  o f  Elements:  Atomic  Structure

Identification  o f  Elements:  Isotopes  o f  Hydrogen

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Atom  notation Atomic number

5

Atomic weight

10.811

B

Symbol

Atomic weight is not exactly 11 because in nature boron exists as a mixture of different isotopes. The fact that it is close to 11, indicates that the major naturally occuring boron isotope atom has an atomic weight of 11. Number of neutrons = Atomic weight minus atomic number

protons neutrons electrons Copyright  ©  2 006  Pearson  Education,  Inc.,  p ublishing  a s  Benjamin  C ummings

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Molecules   and  Compounds § §

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Molecule – two or more atoms held together by chemical bonds Compound – two or more different kinds of atoms chemically bonded together Molecules and compounds are described by their chemical formula

Examples §

Molecule of Oxygen = O2

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Molecule of Nitrogen = N2

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Molecule of carbon dioxide = CO2

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Molecule of glucose = C6H12O6

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Mixtures   and  Solutions §

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Mixtures – two or more components physically intermixed (not chemically bonded) Solutions – homogeneous mixtures of components §

Solvent – substance present in greatest amount

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Solute – substance(s) present in smaller amounts

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Concentration   of  Solutions §

Solutions can be expressed in different ways:

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Percent, or parts per 100 parts. §

Example : a 5 % NaCl solution refers to 5 grams of NaCl in a total of 100 ml of water

Molarity, or moles per liter (M)

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A mole of an element or compound is equal to its atomic or molecular weight (sum of atomic weights) in grams

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MW of NaCl ( atomic weight of Na=23, a.w. of Cl = 35.5) = 58.5

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1 mole of NaCl = 58.5 grams

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1 Molar solution = 58.5 grams per Liter ( or 5.85 g / 100 ml )

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Calculate   atomic  weight  of  a  compound

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Explain what O18 is ! ?

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What is the MW of Lactic acid ? (Chemical Formula = C3H6O3)

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How do you make a 2 molar solution of Lactic Acid ?

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What Molar concentration is 5 g of Lactic acid in 100 ml water ?

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Answers §

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Just like C14 is the isotope of carbon, with 2 extra neutrons, O18 is is the isotope of oxygen with 2 extra neutrons (thus atomic mass 18 instead of 16) MW of lactic acid C3H6O3 §

(3 x C = 3 x 12 = 36) + (6 x H = 6 x 1= 6) + (3 x O = 3 x 16= 48)

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= 36 + 6 + 48 = 90

1 Molar of Lactic acid is 90 grams per Liter §

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2 Molar = 2 x 90 = 180 gr / Liter (or 18 gr/ 100 ml or 1.8 gr/ 10 ml)

5 gr in 100 ml = 50 gr/1000 ml §

We know that 90 gr/ 1000 ml = 1 Molar; thus 1 gr/1000 ml = (1/90) Molar

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Hence 50 gr/1000 ml = 50 x 1 gr /1000 ml = 50 x (1/90) Molar

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Avogadro’s  number §

§

One mole always contains 6.0 x 10 23 numbers of molecules ( = Avogrado’s number) Thus a 1 Molar Glucose or a 1 Molar Lactic Acid solution have the same number of molecules of glucose or Lactic Acid.

Question : how many molecules of fructose are there in a 5 milliMolar ( 5 mM) solution ? Answer = 0.005 M x 6.0 x 1023 = 5 x 6 x 1020

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Mixtures   Compared  with  Compounds §

No chemical bonding takes place in mixtures

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Most mixtures can be separated by physical means

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Mixtures can be heterogeneous or homogeneous

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Compounds cannot be separated by physical means All compounds are homogeneous

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Chemical   Bonds §

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The nucleus of an atom is surrounded by energy levels ( also called electron shells) in which the orbiting electrons reside Valence shell : is the outermost energy level containing the chemically active electrons Chemical Bonds between atoms are formed using the electrons in the outermost energy level. Octet rule – except for the first shell which is full with two electrons, atoms interact in a manner to have eight electrons in their valence shell. With 8 electrons in the outer shell, the valence shell is full.

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Chemically   Reactive   Elements Reactive elements do not have their outermost energy level fully occupied by electrons

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Examples are :

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Hydrogen ( missing one electron for a complete valence shell) Oxygen ( missing two electrons)

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Carbon (missing 4 )

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Sodium ( missing 7 !)

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Chemically   Inert  Elements §

Inert elements have their outermost energy level fully occupied by electrons

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Figure  2.4a

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Types   of  Chemical   Bonds §

Ionic

§

Covalent

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Hydrogen

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Ionic  Bonds Ions are charged atoms resulting from the gain or loss of one or two electrons in order to obtain a full outer valence shell.

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Anions have gained one or more electrons and are negatively charged

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Cations have lost one or more electrons and are thus positively charged

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Formation   of  an  Ionic  Bond §

Ionic bonds form between atoms by the transfer of one or more electrons

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Ionic compounds form crystals instead of individual molecules

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Example: NaCl (sodium chloride)

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Formation   of  an  Ionic  Bond

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Sodium atom has one electron in the outer shell. It can have a full outer shell if it gets rid of that lonely electron. Chloride atom has 7 electrons in the outer shell. It becomes “happy’ if it gains one more. Ionic bonds form between these atoms by the transfer of one electron from sodium to chloride.

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Salts  are  Components   with  Ionic  bonds Salts form crystals in dry state, but dissolve easily in a watery environment and fall apart into their respective ions. Many elements from the 1st and 2nd column form salts with the 7th column of the periodic table. The physiological important ones are known as electrolytes. Examples §

NaCl

Na+

+

Cl-

§

KCl

K+

+

Cl-

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CaCl2

Ca++

+

2Cl-

§

MgCl2

Mg++ +

2Cl-

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Covalent   Bonds §

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Covalent bonds are formed by the sharing of two or more electrons in order to obtain a full outer electron shell Electron sharing produces molecules. The sharing of two electrons between two atoms creates a single covalent bond, indicated by a single line In the example below, each hydrogen needs one more electron to have a full outer shell and carbon requires 4 to fill its outer shell. By sharing and creating 4 singe covalent bonds, all requirements are met !

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Double  Covalent   Bonds 2 pairs of electrons are shared, indicated by two lines

Triple  Covalent   Bonds 3 pairs of electrons are shared, indicated by three lines

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Covalent   Bonds  by  Major  elements §

Hydrogen : §

can only share one electron and thus only forms one covalent bond

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Also can lose one electron and become a “free floating” positively charge ion.

(the basis of pH) §

Carbon §

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Oxygen § §

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Always forms 4 covalent bonds; never becomes an ion Always forms two covalent bonds Sometimes releases the Hydrogen attached to it and becomes a negatively charged oxygen atom within a molecule

Nitrogen § §

Forms mostly 3 covalent bonds Sometimes it forms an extra covalent bond with hydrogen ; in that case it becomes positively charged

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Polar  and  Nonpolar  Molecules §

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Electrons shared equally between atoms produce non-polar molecules Unequal sharing of electrons produces polar molecules; they act like small magnets. Atoms with six or seven valence shell electrons are electronegative and pull the electrons closer to them. Typical such atom is oxygen. Atoms with one or two valence shell electrons are electropositive. Typical such atom is hydrogen.

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Comparison   of  Ionic,   Polar  Covalent,   and   Nonpolar  Covalent   Bonds

Water is a typical dipolar molecule ! It is therefore the ultimate solvent for polar molecules. Life as we know it depends on this chemical quality of water ! Copyright  ©  2 006  Pearson  Education,  Inc.,  p ublishing  a s  Benjamin  C ummings

C.  Hydrogen  Bonds §

Too weak to bind atoms together

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Common in dipoles such as water

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However, many such bonds can form a strong lattice network. Important as intra-molecular bonds, giving larger molecules a three-dimensional shape (such as DNA and proteins).

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Responsible for surface tension in water

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Chemical   Reactions §

Occur when chemical bonds are formed, rearranged, or broken

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Written in symbolic form using chemical equations

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Chemical equations contain: §

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Number and type of reacting substances, and products produced Relative amounts of reactants and products

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Patterns  of  Chemical  Reactions §

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Combination reactions: Synthesis reactions which always involve bond formation A + B → AB Decomposition reactions: Molecules are broken down into smaller molecules AB → A + B

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Exchange reactions: Bonds are both made and broken AB + C → AC + B

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Patterns  of  Chemical  Reactions §

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When confronted with chemical reactions, always keep in mind that elements are neither created nor destroyed All elements need to be represented in equal amount on both sides of the reaction. Example: C6 H12 O6 + O2 → CO2 + H2 O What’s wrong here ? Analyze the elements…

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Oxidation-­Reduction   (Redox)  Reactions § §

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Reactants losing electrons are electron donors and are oxidized Reactants taking up electrons are electron acceptors and become reduced In biological reactions , losing or gaining electrons usually involves the losing or gaining of a proton atom. NADH + H+

NAD+ + 2e- + 2H+

1/2 O2 + 2e- + 2 H+

H2O

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Oxidation-­Reduction   (Redox)  Reactions §

Many oxidation - reduction reactions are coupled in biochemical reactions Ethanol

Acetaldehyde + 2e- + 2H+

NAD+ + 2e- + 2H+

NADH + H+

Ethanol + NAD+

Acetaldehyde + NADH + H+

• In the example above, NAD+ gains protons and thus becomes reduced in the process. • This means that ethanol becomes oxidized, since the protons that NAD+ gained must have come from ethanol.

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Energy  Flow  in  Chemical   Reactions §

Exergonic reactions – reactions that release energy

AB + C → AC + B + Energy §

Endergonic reactions – reactions whose products contain more potential energy than did its reactants; in other words, energy has to be invested in such reactions.

AC + B + Energy → AB + C §

§

All chemical reactions are theoretically reversible unless the exchange of energy is too large to overcome If neither a forward nor reverse reaction is dominant ( small exchange of energy), chemical equilibrium is reached. In such cases, the change in reactants or products will drive the reaction to right or left direction

A + B §

AB

Adding A or B will drive the reaction to the right until a new equilibrium is reached. A sudden increase in AB will drive the reaction to the left…..

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Metabolism §

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In all living organisms, exergonic reactions drive endergonic reactions. The sum of all reactions in your cells = metabolism The breakdown of substances (= catabolism) releases energy and this energy is used to drive other energy requiring reactions ( such as the synthesis of building materials for your cells = anabolism). The making of covalent bonds requires energy. The larger a molecule, the more covalent bonds, the higher the energy content of that molecule. Breaking down large molecules thus releases energy ( food = large molecules with many covalent bonds)

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Metabolism §

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The larger a molecule, the more covalent bonds, the higher the energy content of that molecule. Breaking down large molecules thus releases energy FOOD = large molecules with many covalent bonds By breaking down the covalent bonds in our food, we obtain the energy, which is transformed into ATP The oxidation of glucose, results in the breaking of 5 covalent bonds between the carbons, yielding 6 single carbons in the form of CO2

C 6 H12 O6 + 6 O2

6 CO2 + 6 H2 O + (many) ATP’s

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Metabolism § §

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And why do we need this energy ? To make new molecules and structural proteins that are the essence of cellular life ( for repair, growth, basic functions,…) The making of covalent bonds and larger molecules requires energy. Catabolic pathways generate the energy (and the building blocks) to keep the anabolic pathways going. No food chemicals….no energy…no repair, no maintenance breakdown of homeostasis

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Enzymes  are  essential   in  Metabolism In general chemistry, chemical reactions are influenced by the following §

Temperature – chemical reactions proceed quicker at higher temperatures

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Particle size – the smaller the particle the faster the chemical reaction

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Concentration – higher reacting particle concentrations produce faster reactions

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Catalysts – increase the rate of a reaction without being chemically changed

Living systems require rapid responses and the metabolic machinery is dramatically accelerated by the availability of proteins that work as catalysts ENZYMES !

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Enzymes  are  essential   in  Metabolism Enzymes are very specific in that they only catalyze one specific reaction. It is thus not surprising to see that we have thousands of enzymes in our body, each with a specific role and function. Many diseases can be traced back to the abnormal function of an enzyme. Copyright  ©  2 006  Pearson  Education,  Inc.,  p ublishing  a s  Benjamin  C ummings

Enzymes  are  essential   in  Metabolism Inborn   errors  of  metabolism  comprise  a  large  class  of  genetic  diseases   involving disorders of metabolism. The majority are due to defects of single genes coding for a single enzymes. The absence or malfunctioning of the enzymes results in accumulation of substances that can be toxic, or failure to make essential products for normal metabolic operation. Examples  :   •Phenylketonuria (PKU): non functional phenylalanine hydroxylase • Phenylalanine   accumulates   and   hinders  development   of  the   brain  becomes   converted   to  phenylpyruvate • Untreated,   it  results  in  mental  retardation,   seizures  and   even  death

•Glucose-6-phosphate dehydrogenase deficiency • The G6PD / NADPH pathway is the only source of reduced glutathione in red blood cells • The role of red cells as oxygen carriers puts them at substantial risk of damage from oxidizing free radicals (glutathione protects against free radicals) • People with G6PD deficiency are therefore at risk of increased damage to RBC (resulting in various degrees of hemolytic anemia).

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