A clear understanding of chemistry is essential for the study of physiology.

This is because organ functions depends on cellular functions, which occur as a result of chemical reactions. Watson & Crick first proposed the double helix structure of DNA

Biochemistry

= Chemistry of living things

Matter

= Anything that has mass and takes up space (Solids, liquids, gasses)

Element

= Fundamental substance of matter (e.g. Carbon, Hydrogen, Oxygen)

Compound

= Two or more

different

elements chemically bonded together (e.g. H 2 O = water, C 6 H 12 O 6 = glucose)

Molecule

= two or more atoms chemically joined together. Molecules may be compounds (H 2 O = water molecule), or Molecules may be of the same element (H 2 = hydrogen molecule)

Our body consists of 11 bulk elements and 7 trace elements.

Bulk elements

make up 99.9% of our body: Hydrogen (H) Nitrogen (N) Sodium (Na) Chlorine (Cl) Oxygen (O) Sulfur (S) Potassium (K) Phosphorus (P) Carbon (C) Magnesium (Mg) Calcium (Ca)

Trace elements

make up less than 0.1% of our body: Cobalt (Co) Iron (Fe) Fluorine (F) Zinc (Zn) Iodine (I) Manganese (Mn) Copper (Cu) Learn each bulk element and trace element along with their atomic symbols shown in parentheses

All elements are arranged onto a Periodic table

Atoms

Atoms are the smallest particles of an element that still have the properties of that element.

Atoms are composted of 3 subatomic particles:

Proton

– carries a single positive charge

Neutron

– carries no electrical charge

Electron

– carries a single negative charge An atom contains a central nucleus composed of protons and neutrons. Electrons orbit the nucleus.

Subatomic Particles

Electrical Charge:

Proton: +1 charge. Electron: -1 charge.

Neutron: 0 charge

Atomic Mass:

Proton: 1 dalton Neutron: 1 dalton Electron: 0 Most atoms contain equal number of protons and electrons, so an atom contains no overall net charge and is neutral.

Subatomic Particles

Atomic Number:

The number of protons in one atom. Atomic number identifies an element.

Example. The atomic number of oxygen is 8. Oxygen,

and only oxygen

has 8 protons.

Atomic Weight:

The sum of protons and neutrons in one atom. Remember, the weight of electrons is negligible.

Isotopes

Isotopes are atoms with the same atomic number, but different atomic weights. Isotopes occur because the number of neutrons of an element varies between atoms.

Two isotopes of oxygen

: Oxygen 16 (O 16 ) protons: 8 electrons: 8 neutrons: 8 Atomic Number: Atomic Weight: 8 16 Oxygen 17 (O 17 ) protons: 8 electrons: 8 neutrons: 9 8 17 *The atomic weight of an element is an average of the isotopes present.

Understand the notations on a periodic table.

End of Section 1, Chapter 2

Section 2 of Chapter 2 Bonding of Atoms

Properties of electrons

Electron Shells: Electrons encircle the nucleus in discrete orbits, called electron shells. Each shell can contain only a fixed number of electrons.

1st shell holds 2 electrons

2 nd

shell holds 8 electrons

3 rd

shell holds 8 electrons

Octet rule:

Except for the 1 st shell, each electron shell holds up to 8 electrons * Lower shells are filled first.

Helium

Atomic number = 2 Atomic weight = 4 (2 electrons fill the 1st electron shell)

Carbon

Atomic number = 6 Atomic weight = 12 (The first 2 electrons fill the inner shell, and the remaining 4 electrons are placed the 2 nd electron shell).

Ions Ions are atoms that readily gain or loose electrons

Cation:

an ion that looses electrons • Cations are positively charged ions

Anion

: an ion that gains electrons •Anions are negatively charged ions

Example of a cation

Sodium (Na)

atomic number = 11 atomic weight = 23

Na+ = Sodium cation

Only 1 lone electron sits in the outer shell. This electron is unpaired and is easily lost, forming the sodium cation.

Example of an anion

Chlorine (Cl)

atomic number = 17 atomic weight = 35

Cl = Chloride anion

7 electrons fill the outer shell of chlorine, leaving room for 1 more electron. Chlorine readily accepts one electron, creating the chloride anion.

Ionic Bond

Ionic bonds are formed when the oppositely charged particles attract. Figure 2.4 (a) An ionic bond forms when on atom gains and another atom looses electrons, and then (b) oppositely charged ions attract.

Covalent Bonds Covalent bonds are formed when atoms share electrons.

Example: A hydrogen molecule (H 2 ) is formed when two hydrogen atoms share their single electron.

H + H H 2

Covalent Bonds of water Water consist of oxygen covalently bonded to two hydrogen atoms.

Structural Formula: depicts the covalent bonds of a molecule as lines.

Molecular Formula: is a shorthand notation for representing molecules.

Oxygen joined to two hydrogen atoms by single bonds Two oxygen atoms joined by a double bond.

A Carbon atom joined to hydrogen by a single bond and to nitrogen by a triple bond.

Nonpolar covalent bonds Nonpolar covalent bonds occur when the atoms share the electrons equally, so the molecule has no overall charge. Two hydrogen atoms share their electrons equally. Thus, the hydrogen molecule has no overall charge and is nonpolar.

Polar covalent bonds Polar bonds have an unequal distribution of electrons. One portion of the atom has a higher affinity for electrons than the rest of the molecule (electronegative).

Slightly negative end Slightly positive end

Water is a polar molecule because the oxygen atom (with 8 protons) tends to pull the electrons away from hydrogen. The oxygen end has a slight negative charge, while the hydrogen end has a slight positive charge.

Hydrogen bonds Occur when the slightly positive (hydrogen) end of a polar molecule weakly attracts to the slightly negative end of another molecule.

• Hydrogen Bonds: Form weak bonds at room temperature, but are strong enough to form ice • Stabilize large proteins, DNA, and RNA

End of Section 2, Chapter 2

section 3 of chapter 2 chemical reactions

Chemical Reactions

activation energy: energy required to start a reaction A

catalyst

reduces the amount of energy needed to initiate a reaction.

Catalysts increase the rate of reactions, but are not consumed by the reaction- reusable

Acids, Bases, and Salts

Electrolytes – are substances that dissociate in water to release ions.

Example: NaCl → Na + + Cl -

Acids - electrolytes that dissociate to release protons (H + ) in water Example: HCl → H

+

+ Cl Bases- electrolytes that absorb H+ from water, or electrolytes that dissociate to release hydroxide ions (OH ) in water Examples: NaOH → Na + + OH

-

Salt – electrolyte formed by the reaction between an acid and base Example: Acid + Base → Salt + water H Cl + Na OH

→ NaCl +

H 2 O

acid and base concentrations

pH pH measures the concentration of hydrogen ions [H+] in a solution.

As pH decreases, [H+] increases = solution is more acidic pH 0 acidic property increasing 7 neutral alkaline property increasing 14

Small changes in pH reflect large changes in [H + ] change of 1 pH = 10 fold change in [H + ] change of 2 pH = 100 fold change in [H+] change of 3 pH = 1000 fold change in [H+]

Blood Average blood pH = 7.35 - 7.45

Acidosis

= blood pH less than 7.3

Symptoms include fatigue, disorientation, and difficulty breathing.

Alkalosis

= blood pH greater than 7.5

Symptoms include agitation and dizziness Blood contains several buffers Buffer = resists changes to pH

Chemical components of cells

Organic Vs. Inorganic Molecules

Organic molecules Compounds with carbon May form macromolecules Includes proteins, carbohydrates, lipids, nucleic acids Inorganic molecules Compounds that lack Carbon (exception is CO2) Usually dissociate in water

Inorganic Substances Water (H2O) 2/3 of weight in a person Transports gasses, nutrients, wastes, hormones, ect.

Oxygen (O2) Used in cellular respiration Carbon Dioxide (CO2) Waste of metabolic reactions Inorganic Salts Na + , Cl , K + , Ca 2+ , HCO 3 , PO4 2-

End of Section 3, Chapter 2

Section 4, Chapter 2 Organic Molecules

Organic Molecules Molecules that contain carbon Organic Synthesis Small molecules (monomers) join together to form larger molecules (polymers) Monomer portion of a polymer

Covalent Bonds formed by Carbon C 12.01

6 Atomic Number of Carbon = 6 2 electrons in 1 st shell 4 electrons in 2 nd shell Note there are 4 empty spaces in the 2 nd shell available for covalent bonds.

Examples of covalent bonds formed by carbon Carbon can form 4 covalent bonds Carbon can also form double or even triple bonds Carbon to Carbon bonds can form long chains hydrocarbon

Polymers and Monomers Large organic molecules, called polymers consist of repeating subunits, called monomers.

Example: Starch is a polysaccharide composed of many glucose molecules (monosaccharides) joined together.

major organic macromolecules of the cell Monomer Polymer Monosaccharide (simple sugars) Disaccharides (double sugars) Polysaccharides (complex carbohydrates) Amino Acids Proteins Fatty Acids + Glycerol Nucleotides Fats* * Not truly a polymer Nucleic Acids

Carbohydrates Simple carbohydrates = sugars Monosaccharides Disaccharides Complex Carbohydrates Also called Polysaccharides Composed of several simple carbohydrates

monosaccharides Twice as many Hydrogen as Oxygen atoms Example: Glucose (C 6 H 12 O 6 )

disaccharides 2 monosaccharides bonded together Examples of disaccharides

polysaccharide

Built of simple carbohydrates

examples of polysaccharides

Starch – easily digested Cellulose- Plant polysaccharide, indigestible by humans Glycogen – storage form of energy, synthesized by liver Glycogen

Glycerol Molecule OH (in red) represents sites of fatty acid attachments

Unsaturated fat

Proteins

Proteins have many functions: Proteins provide structural material.

They are a source of energy. Some act as chemical messengers (hormones, neurotransmitters).

Many proteins are receptors.

Most enzymes are proteins.

Proteins: enzymes

Enzymes catalyze reactions (increases rate), but are not consumed by the reaction (reusable).

Synthesis reaction involving an enzyme

Proteins: amino acids

All amino acids consists of:

An amino group (-NH 2 ) A Carboxyl Group (-COOH) A single Carbon atom An “R” group (R = rest of the molecule) 1 of 20 possible “R” groups = determines amino acid

Peptide bond (red) joins two amino acids.

4 Levels of Protein Structure

A protein’s shape, or conformation, determines its function. Therefore, it’s important to understand a protein’s shape at 4 levels.

4 Levels of protein structure

Red dots indicate hydrogen bonding

4 Levels of protein structure

4 Levels of protein structure

4 Levels of protein structure

Protein Structure

Conformation Complex 3 dimensional fold of a protein Conformation determines a protein’s function Denature Treatment that alters the shape of a protein to make it nonfunctional Heat, pH changes, radiation, certain chemicals may denature proteins

Nucleic acids: overview Nucleic Acids Includes DNA and RNA Genetic information Consists of monomers, called nucleotides

RNA Contains the sugar ribose (

ribo

nucleic acid) RNA is a single-stranded nucleic acid.

Transcribes DNA for protein synthesis RNA also may act as an enzyme DNA DNA contains a sugar, called deoxyribose (

deoxyribo

nucleic acid) Double-stranded helix Encodes genetic information for protein synthesis.

Nucleotides Nucleotides are the monomers of Nucleic Acids 3 Components of a Nucleotide 5 Carbon Sugar (S) Nitrogenous Base (B) Phosphate Group (P) RNA Sugar = ribose DNA Sugar = deoxyribose

H bonds Antiparallel