Transcript Slide 1

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. H2O = water, C6H12O6 = glucose)
Molecule = two or more atoms chemically joined together.
Molecules may be compounds (H2O = water molecule), or
Molecules may be of the same element (H2= 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)
Zinc (Zn)
Manganese (Mn)
Iron (Fe)
Iodine (I)
Copper (Cu)
Fluorine (F)
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 (O16)
Oxygen 17 (O17)
protons: 8
electrons: 8
protons: 8
electrons: 8
neutrons: 8
neutrons: 9
Atomic Number:
8
8
Atomic Weight:
16
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
2nd shell holds 8 electrons
3rd shell holds 8 electrons
Octet rule: Except for the 1st 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 2nd
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 (H2) is formed when
two hydrogen atoms share their single electron.
H+H
H2
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
Acid + Base → Salt + water
Example:
HCl + NaOH → NaCl + H2O
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
acidic property increasing
pH
0
alkaline property increasing
7
neutral
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+, Ca2+, HCO3-, PO42-
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 1st shell
4 electrons in 2nd shell
Note there are 4 empty spaces in the
2nd 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
Fats*
*Not truly a polymer
Nucleotides
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 (C6H12O6)
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 (-NH2)
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 (ribonucleic 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
(deoxyribonucleic 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