Biochemistry Chapter 2

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THE NATURE OF MATTER

Section 2-1

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Biochemistry

• Study of chemicals of living things

• Investigates the structure and reactions of organic compounds

• Organic compounds – compounds that contain carbon:

• Carbohydrates, lipids, nucleic acids, proteins

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Atoms

• Composed of sub-atomic particles • Protons and neutrons

- make-up the nucleus • Electrons – negatively

charged and orbit the nucleus • Valence Electrons

• Atomic number = number of protons

• Atomic mass = number of p’s + n’s

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Atoms • Atoms normally have as many

electrons as protons

• Opposite charges balance leaving atom neutral

• Electrons are attracted to the positive nucleus

• ·Revolve around nucleus in orbitals

• ·Can be pushed into higher orbitals with energy

• ·Release that energy when they fall back to lower orbital

• ·Different energy levels referred to as electron shells

• electrons - outermost electron called Valence Electrons

• ·Inner-most energy level can only hold 2 electrons

• ·Every other energy level can hold as many as 8

• ·Every atom wants 8 electrons in outermost shell - Octet Rule

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Elements and Isotopes

• Elements • Pure substance that

consists entirely of one type of atom

• More than 100 known elements organized in Periodic table according to the elements atomic number (#of P’s)

• Isotopes • Atom of an element

with a different number of neutrons

• Won’t change chemical properties of the element

• Radioactive Isotopes • Subatomic particles

from decay of nucleus • Used in nuclear

medicine • Cancer treatment --

Radiation therapy • Labels or “tracers” to

follow movement of substances within an organism

• Half-life – calculates age of fossils

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Periodic Table of Elements

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An Element on the Periodic Table

• On the periodic table, the atomic mass listed is “weighted”, meaning the abundance of each isotope is taken into consideration to get an “average” mass for that element. • EXAMPLE: Carbon

appears as 12.011 because Carbon-12 is extremely common and the other two isotopes are pretty rare.

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Atomic Number

Element Name

Element Symbol Atomic Mass (average)

Isotopes of Carbon

Nonradioactive carbon-12 Nonradioactive carbon-13 Radioactive carbon-14

6 electrons 6 protons 6 neutrons

6 electrons 6 protons 8 neutrons

6 electrons 6 protons 7 neutrons

Figure 2-2 Isotopes of Carbon

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Carbon 14 Dating

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Chemical Compounds

• Compounds – Substances formed by combination of two or more elements

• Chemical formula – shows the number and types of elements in the compound • H2O, C6H12O6

• Chemical Bonds – • Ionic Bonds – Oppositely charged ions are attracted to each other

and form an ionic bond • Covalent Bonds – Electrons are shared between atoms to form

molecules

• Single covalent bond – atoms share two electrons

• Double covalent bond – atoms share four electrons

• Triple covalent bond (more rare) – atoms share six electrons

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Chemical Compounds • substance formed by chemical

combination of 2 or more elements in definite proportions

• ·Chemical formula scientific notation indicates proportions of elements in compound

• ·water - H2O

• ·salt - NaCl

• ·physical & chemical properties of a compound very different than elements alone

• ·ex. Na(s), Cl(g) H(g), O(g)

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Chemical Compounds

• Most elements are combined with others (in nature) – very few exist as “pure” elements (but carbon does)

• Physical and chemical properties of compounds are very different than the elements they are made of • EXAMPLE – H2O – hydrogen and oxygen are both a gas at room

temperature but combine to form liquid water

• EXAMPLE – NaCl – Sodium is a very soft metal (explosive in water) and Chloride is a poisonous, greenish gas (used in WWI) but combine to form table salt that dissolves easily and harmlessly in water.

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Covalent Bonds

• Electrons are shared between atoms to form molecules

• The structural formula of a molecule indicates a shared pair of electrons by a line between the two atoms e.g. single covalent bond (H–H), or double covalent bond (O=O)

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Hydrogen gas

Oxygen gas

Methane

Electron Model Chemical

Formula

Structural

Formula

Chemical Compounds

• Van der Waals Forces

• Intermolecular forces of attraction

• Created by polar molecules • Molecules that have

positive and negative regions caused by uneven sharing of electrons

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Ionic Bonding

Sodium atom (Na) Chlorine atom (Cl) Sodium ion (Na+) Chloride ion (Cl-)

Transfer of electron

Protons +11

Electrons -11

Charge 0

Protons +17

Electrons -17

Charge 0

Protons +11

Electrons -10

Charge +1

Protons +17

Electrons -18

Charge -1

Figure 2-3 Ionic Bonding

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Ionic Bonding

Sodium atom (Na) Chlorine atom (Cl) Sodium ion (Na+) Chloride ion (Cl-)

Transfer of electron

Protons +11

Electrons -11

Charge 0

Protons +17

Electrons -17

Charge 0

Protons +11

Electrons -10

Charge +1

Protons +17

Electrons -18

Charge -1

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Van der Waals Forces • attraction between

oppositely charged regions of nearby molecules

• ·due to polarity of molecules

• ·able to hold large molecules together

• ·Ex. Gecko's toes -- covered by half mill. tiny hairlike projections each divided into hundreds of tiny fibers

• ·structure allows foot to come into contact with surfaces on molecular level and molecular attractions create adhesive forces

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PROPERTIES OF WATER

Chapter 2-2

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Water & Hydrogen Bonds

• H2O - Polar molecule –

• Strong Solvent – polarity makes it able to dissolve (break apart) compounds – strongest solvent • Solute – the substance that’s dissolved

• Hydrogen bond – An attraction between polar molecules

• Responsible for waters special properties: • Cohesion – attraction between molecules of the

same substance • Adhesion – attraction between molecules of

different substances – water’s surface curves up sides of glass

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Water

• ·Oxygen pulls on

electrons stronger than

Hydrogen can

• ·Creates (-) charge on O

end of molecule and (+)

charge on H end

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Solutions & Suspensions

• Solutions – evenly distributed mixture of two or more substances

• Suspensions -- mixture of water and non-dissolved material • Some materials do not dissolve but separate into

pieces and stay in solution • Blood – mostly water (plasma) contains dissolved

compounds • When salt

is dissolved in water, the water is the solvent

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Acids, Bases, & pH

• Acids • An acid is any compound

that forms hydrogen ions (H+) in solution

• Acids have a pH less than 7

• Strong acids have a pH from 1 to 3

• EXAMPLE – Stomach acid (hydrochloric acid – HCl) has a pH of about 1.5

• Bases • A base is any compound

that forms hydroxide ions (OH⁻) in solution

• Often called alkaline solutions

• Bases have a pH greater than 7

• Strong bases have a pH from 11 to 14

• EXAMPLE – Bleach (sodium hypochlorite – NaOCl) has a pH of about 11

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Acids, Bases, & pH

• The pH scale • Shows the concentration of

hydrogen ions (H+) in solution • Ranges from 0 to 14 • At pH 7, the concentration of

hydrogen ions (H+) and hydroxide ions (OH-) are equal. Pure water has a pH of 7.

• Solutions above pH 7 are basic and have more hydroxide ions (OH-)

• Solutions below pH 7 are acidic and have more hydrogen ions (H+)

• Each step on the pH scale has a factor of 10 • pH 5 is 10 times more acidic than pH

6 • pH 4 is 100 times more acidic than

pH 6 (10 times 10)

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Acids, Bases, & pH

• Buffers • pH of fluids in living things must generally be kept between 6.5

and 7.5

• Controlling pH is important for maintaining homeostasis

• Buffers are weak acids or bases that react with strong acids or bases to prevent sharp or sudden changes in pH

• EXAMPLE – bicarbonate (secreted by the pancreas), electrolytes, proteins, blood

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CARBON COMPOUNDS

Chapter 2-3

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The Chemistry of Carbon

• Organic Chemistry is the study of all compounds that contain bonds involving carbon atoms

• Carbon can make up to four strong covalent bonds with many different elements

• Carbon can also bond with other carbon atoms to make long chains • Carbon bonds to one another can be single, double or even triple

bonds

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Macromolecules

• Macromolecules are “giant” molecules made up of many smaller molecules

• Formed by polymerization (joining of smaller compounds together) • The smaller units are called monomers

• Monomers join together to make polymers

• Macromolecules are grouped to make their study easier • Carbohydrates

• Lipids

• Nucleic acids

• Proteins 28

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Summary of Macromolecules

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Carbohydrates Proteins Nucleic Acids Lipids

·Polymers of sugars which form starches ·single sugars such as glucose, galactose and fructose are monsacharides ·Main source of energy for all living things

·Polymers of

amino acids

·many are

enzymes which

control rates of

chem. rx's

·build tissues of

body

·Polymers of

nucleotides

·Molecules of

heredity

·DNA & RNA

·polymers of

fatty acids and

glycerol

·form biological

membranes &

waterproof

coverings

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Carbohydrates

• Compounds made up of carbon, hydrogen and oxygen

• Usually in a ratio of 1 : 2 : 1

• They are the main source of energy for living things

• Plants and some animals use carbohydrates for structural purposes

• The breakdown of sugars (like glucose) supplies immediate energy for all cell activities

• Extra sugar is stored as complex carbohydrates known as starches

• Simple sugar molecules are also called monosaccharides

• Glucose, galactose (in milk), fructose (in fruit)

• The macromolecules formed from many monosaccharides are called polysaccharides

• Animals often store excess sugar in a polysaccharide called glycogen (usually in the liver and muscles)

• Plants also use the polysaccharide called cellulose for structural purposes

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Lipids

• Generally not soluble in water • Made from carbon and hydrogen atoms • Most common types are fats, oils and waxes • Used to store energy • Some lipids are part of biological membranes and waterproof

coverings • Steroids are also lipids and can serve as chemical messengers • Many are formed when a glycerol molecule combines with a fatty

acid • A fatty acid is made from carbon, hydrogen and oxygen atoms • If all the carbon atoms are joined with a single covalent bond, then it is saturated (it is

saturated with hydrogen atoms) • Tend to be solid at room temperature

• If there is at least one carbon-carbon double bond in a fatty acid, then it is unsaturated • Tend to be liquid at room temperature

• If there is at more than one carbon-carbon double bond in a fatty acid, then it is polyunsaturated

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Nucleic Acid

• Macromolecules containing hydrogen, oxygen, nitrogen, carbon and phosphorus

• Polymers assembled from monomers called nucleotides • Nucleotides are made of three parts

• A 5-carbon sugar

• A phosphate group

• A nitrogen base

• Store and transmit hereditary, or genetic, information

• Two kinds of nucleic acids • Ribonucleic acid

• Contains the sugar ribose

• Deoxyribonucleic acid • Contains the sugar deoxyribose

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Proteins

• Macromolecules that contain nitrogen, carbon, hydrogen, and oxygen • They are polymers of molecules called amino acids

• Amino acids are compounds with an amino group on one end ( -NH2) and a carboxyl group on the other (-COOH)

• A third part, called the R-group is what makes the amino acids different • Some are acidic, some basic • Some are polar, some non-polar

• Proteins have different roles • Some control the rate of reactions and regulate cell processes • Some are used to form bones and muscles • Some transport substances into or out of cells • Some help to fight disease

• Proteins have up to four levels of organization • Sequence of amino acids in a protein chain • Amino acids can be twisted or folded • The chain itself is folded • If it has more than one chain, it has a specific arrangement

• NOTE – van der Waals forces and hydrogen bonds help maintain a protein’s shape

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CHEMICAL REACTIONS AND ENZYMES

Chapter 2-4

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Chemical Reactions

• Everything that happens with life – its growth, its interaction with the environment, its reproduction and even its movement – is based on chemical reactions

• A chemical reaction is a process that changes one set of chemicals into another set of chemicals • Some occur slowly (combining iron and oxygen to form an iron oxide called

rust) • Some occur quickly (hydrogen gas being ignited in the presence of oxygen –

explosive – the Hindenburg)

• The elements or compounds that enter into a chemical reaction are called reactants

• The elements or compounds that are produced are called products

• Chemical reactions always involve breaking bonds in reactants and forming new bonds in products • EXAMPLE – Ridding the body of Carbon Dioxide

• From Cells to lungs - CO2 + H2O H2CO3 (carbonic acid – very soluble) • Once in lungs – H2CO3 CO2 + H2O

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Chemical Reactions

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Synthesis, Decomposition, Single Replacement, Double Replacement

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Sample Chemical Reaction

• Acetic Acid: CH3COOH -> CH3COO- + H+

• Sodium Bicarbonate: NaHCO3 -> Na+ + HCO3 –

• H++ HCO3- -> H2CO3

Carbonic Acid (highly unstable so…)

• H2CO3 -> H2O + CO2 • **This is actually TWO reactions -

a double replacement and a decomposition reaction!

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Energy In Reactions

• Energy Changes • Energy is either released or absorbed when chemical bonds are

broken or formed

• Reactions that release energy can occur spontaneously

• Reactions that absorb energy will not occur without a source of energy

• 2H2 + O2 2H2O – this releases heat (and sometimes light and sound)

• 2H2O 2H2 + O2 – requires so much energy to get started that an electric current is usually used

• So what does this mean for living things?

• Since all living things carry out reactions, all living things need a source of energy

• Plants get from sunlight

• Animals get it from plants or animals that they consume

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Energy in Reactions (part 2)

• Activation Energy • Many chemical reactions that release energy do not occur

spontaneously

• The energy that is needed to start a reaction is called activation energy

• Typically energy absorbing reactions need more energy to get started than energy releasing reactions C

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Enzymes

• Some reactions are too slow to be of use to life or require too much activation energy to make them practical for living tissue

• A catalyst is a substance that speeds up the rate of a chemical reaction • They lower the amount of activation energy required

• Enzymes are proteins that act as a catalyst for biological reactions (those reactions that take place in cells) • EXAMPLE – Ridding the body of Carbon Dioxide

• From Cells to lungs - CO2 + H2O H2CO3 (carbonic acid – very soluble)

• BUT reaction is so slow that cells may produce more CO2 than body can handle

• So it is sped up with an enzyme (called carbonic anhydrase) which increases speed by a factor of 10 million

• Enzymes generally only work for one type of chemical reaction and typically get their name from the reaction they perform

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Enzyme Action

• The Enzyme-Substrate Complex • Enzymes provide a specific site where reactants are brought together

• The reactants of an enzyme catalyzed reaction are called substrates

• The substrates bind to an area of the enzyme called the active site

• The active site only holds specific substrates and fit them like a lock and key

• When the reaction is over, the products are released and the enzyme can speed up another reaction (of the same type)

• Regulation of Enzyme Activity • Enzymes can be affected by any variable that influence a chemical reaction

• Most work best at a particular pH level or a particular temperature

• Most cells contain proteins that turn key enzymes “on” and “off”

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