Chemistry
Anatomy and Physiology | Tutorial Notes
Chemistry
Learning objectives
After study of today’s learning, the student will:
1. Describe the relationship among matter, atoms, and compounds
2. Describe how atomic structure determines how atoms interact
3. Explain how molecular and structural formulas symbolize the composition of compounds.
4. Describe three types of chemical reactions
5. Describe the differences among acids, bases, and salts
6. Explain the pH scale
7. Explain the function of buffers
tutorial outline
I. Definitions
A. chemistry – investigates the composition and interactions of matter
B. matter – anything that has mass (weight) and occupies space.
Forms of matter include solid, liquid, and gas
C. Element – fundamental substance of matter
D. Compound – combination of two or more different elements (eg. H2O, CO2)
E. Atom – smallest functional particle of an element
II. Bulk Elements – makeup over 99.9% of the body’s mass (see table 2.2)
A. Oxygen – O
B. Carbon – C
C. Hydrogen – H
D. Nitrogen – N
E. Calcium – Ca
F. Phosphorus – P
G. Potassium – K
H. Sulfur – S
I. Chlorine – Cl
J. Sodium – Na
K. Magnesium – Mg
III. Trace Elements – less than 0.1% (see table 2.2)
A. Cobalt – Co
B. Copper – Cu
C. Fluorine – F
D. Iodine – I
E. Iron – Fe
F. Manganese – Mn
G. Zinc – Zn
IV. Atomic Structure (see figure 2.1)
A. Subatomic Particles
1. Proton - Charge of one proton = +1
Relative mass of one proton = 1
2. Neutron - Charge of one neutron = 0 (uncharged)
Relative mass of one neutron = 1
3. Electron - Charge of one electron = -1
Relative mass of one electron = 0
Nucleus – contains protons and neutrons
Electrons orbit the nucleus in distinct electron shells
B. Atomic Charge – number of protons equals the number of electrons, so the charges
cancel and the atom has no net charge (It is electrically neutral)
C. Atomic Number – number of protons in an atom of a particular element.
Each element has a unique atomic number
1. Hydrogen has 1 proton, therefore its atomic number = 1
2. Carbon has 6 protons, therefore its atomic number = 6
3. The atomic number of Oxygen = 8. How many protons does it have?
D. Atomic Weight (Mass) – number of protons + number of neutrons in an atom of an element.
1. The atomic weight of Carbon with 6 protons + 6 neutrons + 6 electrons = 12
*Remember the weight of electrons is so small we don’t calculate it in the atomic weight!
V. Isotopes
A. Isotopes are atoms of an element with different number of neutrons. They have the same atomic number (number of protons) but different atomic weights.
B. Example of isotopes
1. Carbon (C) has three isotopes
a. Carbon-12 has 6 protons and 6 neutrons. About 99% of Carbon is Carbon-12
b. Carbon-13 has 6 protons and 7 neutrons. About 1% of Carbon is Carbon-13
c. Carbon-14 has 6 protons and 8 neutrons. Less than 1% of Carbon is Carbon-14
C. Radioactive Decay
1. The nucleus of an unstable isotope may react to form a more isotope, releasing energy
in the process.
2. Decay of Carbon-14
a. Carbon-14 has 6 protons and 8 neutrons
b. A neutron from C-14 may decay into a proton and electron
c. The additional proton changes the atomic number from 6 into 7, thus
converting Carbon into Nitrogen
VI. Molecules and Compounds
A. Molecule – particle of two or more atoms chemically joined by covalent bond.
1. Compound – molecule of two or more different elements
a. Water molecule H2O (molecular formula)
b. Carbon Dioxide molecule CO2
2. Molecule of an element – molecule of identical elements
a. Hydrogen molecule H2
b. Nitrogen molecule N2
c. Oxygen molecule O2
VII. Properties of Electrons
A. Electron Shells
1. 1st shell holds 2 electrons (closest to the nucleus)
2. 2nd shell holds 8 electrons
3. 3rd shell holds 8 electrons
B. Electron rules
1. Octet rule – atoms react in a way to fill the outermost shell completely. Because the 2nd and 3rd shell hold 8 electrons, we think of this as the octet rule.
2. Electron pairing – electrons tend to react to form electron pairs. That is, lone electrons tend to react to pair with other lone electrons.
VIII. Ions
A. Ion is an atom that gains or loses electrons and becomes electrically charged.
B. Cation
1. ions that tend to loose electrons are called cations
2. cations are positively charged
3. Example: Sodium tends to release 1 electron to make the sodium cation (Na+)
C. Anion
1. ions that gain electrons are called anions
2. anions are negatively charged
3. Example: Chlorine tends to accept 1 electron to make the chloride anion (Cl-)
IX. Chemical Bonds
A. Ionic Bonds (see figure 2.4)
1. Ions with opposite charges attract each other forming ionic bonds.
2. Ionic bonds are formed between Cations (+) and Anions (-)
3. Example: Na+ and Cl- unite to form NaCl (sodium chloride)
4. Ionic bonds tend to form well-organized arrays (crystals) because ions attract opposite
charges from all directions.
B. Covalent Bonds (see figure 2.5)
1. Covalent bonds are formed when atoms share electrons, rather than gaining or losing
them.
2. Covalent bonds may be single, double, or triple Structural formula
a. Single bond: atoms share 1 pair of electrons H—H
b. Double bond: atoms share 2 pairs of electrons O=C=O
c. Triple bond: atoms share 3 pairs of electrons N≡N
3. Polarity
a. Non-polar bonds
§ Atoms share electrons equally, so the molecules are uncharged
§ non-polar bonds are mostly water insoluble
§ hydrocarbons (carbon and hydrogen) are an important group of
non-polar molecules
b. Polar bonds
§ Unequal sharing of electrons
§ One atom has a higher affinity for the electrons than the other atoms
§ polar bonds are common between oxygen and hydrogen (or nitrogen and hydrogen)
§ The oxygen end tends to be negatively charged, while the hydrogen end tends to be positively charged.
C. Hydrogen Bonds (see figure 2.8)
1. Attraction between a slightly positive hydrogen end of one molecule and a negative
oxygen or nitrogen end of a different molecule.
2. Hydrogen bonds are relatively weak, but can form ice at 0°C
3. Hydrogen bonds maintain the structure of proteins and DNA
IX. Chemical Reactions
Reactants (starting material) → Products (end molecules)
A. Synthesis = molecules join together A + B → AB
B. Decomposition = molecules break apart AB → A + B
C. Exchange = bonds are broken and new molecules are formed AB + CD → AC + BD
D. Reversible = products may be reactants and reactants may be products A + B ↔ AB
X. Activation Energy
A. Activation energy is the amount of energy required to initiate a chemical reaction
B. Catalyst
1. A catalyst increases the rate of a reaction without being consumed by the reaction
2. Catalysts work by lowering the activation energy
XI. Acids, Bases, Salts
A. Electrolyte (see figure 2.9)
1. Electrolytes dissociate in water releasing ions
2. Example: sodium chloride in water: NaCl → Na+ + Cl-
3. The polar nature of water forms hydration shells, with negative oxygen ends surrounding the Na+ and the positive hydrogen ends surrounding the Cl-
B. Acids (see table 2.4)
1. An acid is an electrolyte that releases H+ (hydrogen ions/protons) in water.
2. Example: Hydrochloric Acid: HCl → H+ + Cl-
C. Alkaline (Base) (see table 2.4)
1. An alkaline (base) absorbs H+ from water
2. Many bases release OH- (hydroxide ions), the OH- combines with H+ forming H2O
3. Example: Sodium Hydroxide: NaOH → Na+ + OH- … then … OH- + H+ → H2O
D. Salt (see table 2.4)
1. Acids and Bases may react to form salt and water
2. Example: HCl + NaOH → NaCl + H2O
XII. Acid and Base Concentrations (see table 2.5)
A. pH scale is used to measure concentration of H+ in a solutions
B. pH scale measures decimal places instead of measuring concentration (grams/L)
C. Each change in 1 pH represents a 10-fold change in [H+]
D. Examples:
pH [H+] grams/Liter
0 1 (10-0)
1 0.1 (10-1)
2 0.01 (10-2)
3 0.001 (10-3)
4 0.0001 (10-4)
5 0.00001 (10-5)
6 0.000001 (10-6)
7 0.0000001 (10-7)
8 0.00000001 (10-8)
Etc. etc. etc.
E. In other words, decrease 1 pH = 10 fold increase [H+]
decrease 2 pH = 100 fold increase [H+]
decrease 3 pH = 1000 fold increase [H+]
F. Water weakly ionizes into H+ and OH - H2O ↔ H+ + OH-
G. Water has pH = 7.0 (neutral)
H. Acid has pH < 7.0 (acidic)
I. Alkaline has pH > 7.0 (basic)
J. pH 0__←increasingly acidic__ ___7_neutral_ increasingly basic →_14
K. pH of Blood
1. Blood pH 7.35 – 7.45
2. Alkalosis
a. Alkalosis is pH > 7.5
b. caused by excess vomiting, rapid breathing at high altitudes
c. Symptoms include dizziness and agitation
3. Acidosis
a. Acidosis is pH < 7.3
b. caused by severe vomiting, diabetes, lung and kidney disease
c. Symptoms include disorientation and fatigue
4. Buffer
a. A buffer is a substance that resists changes to pH.
b. Blood contains several buffers
XIII. Inorganic Substances – substances that usually lack carbon (see table 2.6)
A. Water (H2O)
1. Important solvent for chemicals in the body
2. Water transports dissolved substances throughout the body
3. Water absorbs and carries heat throughout the body
B. Oxygen (O2) - Organelles within our cells use oxygen to efficiently release energy from food
C. Carbon Dioxide (CO2) - Carbon Dioxide is a waste product from the reactions of cell
respiration
D. Inorganic Salts
sodium (Na+) chloride (Cl-) potassium (K+) calcium (Ca2+) magnesium (Mg2+) phosphate (PO42-) carbonate (CO32-) bicarbonate (HCO3-) sulfate (SO42-)
XIV. Organic Substances
A. Carbon-to-carbon bonding
1. Organic compounds have carbon (and usually hydrogen)
2. Each Carbon can form 4 covalent bonds with other atoms
3. Carbon-to-carbon bonds form hydrocarbon chains and hydrocarbon rings
B. Small molecules (monomers) join together to form larger molecules (polymers)
monomers polymer
C. Examples of Organic Substances include Carbohydrates, Lipids, Proteins, and Fats
XV. Carbohydrates
A. Monosaccharides (simple sugars)
1. monosaccharides (simple sugars) are the building blocks of larger carbohydrates (polymers)
2. examples of monosaccharides
a. glucose (C6H12O6)
b. fructose (C6H12O6)
c. galactose (C6H12O6)
B. Disaccharides
1. sucrose (table sugar)
glucose + fructose
2. lactose (sugar found in milk)
Glucose + galactose
C. Polysaccharides (complex carbohydrates)
1. Starch – easily digested for energy
2. Cellulose – dietary fiber
3. Glycogen – storage form of energy made by our liver
XVI. Lipids
A. Overview
1. Lipids include fats, phospholipids, and steroids
2. Lipids are non-polar (water insoluble)
B. Fats (triglycerides)
1. Building blocks of fats include 1 glycerol molecule + 3 fatty acid molecules
glycerol
fatty acids - may be saturated (all C-C bonds) or unsaturated (C=C)
2. triglyceride (fat)
Glycerol + 3 Fatty Acids
a. unsaturated fat has one or more unsaturated fatty acids
b. saturated fat contains all saturated fatty acids
C. Phospholipids
1. Building blocks include 1 glycerol + 2 fatty acids + 1 phosphate group
2. Polar Head is water soluble Non-polar tail is water insoluble
D. Steroids
1. Steroids are composed of hydrocarbon rings
2. Examples include
a. sex hormones: estrogen, testosterone, progesterone…
b. cholesterol
XVII. Proteins
A. Proteins have many functions
1. Proteins provide structural material
2. Some proteins act as chemical messengers (hormones and neurotransmitters)
3. Many proteins are receptors
4. Proteins provide a source of energy
5. Most enzymes are proteins
Enzyme – biological catalyst
A. Amino Acids
1. Amino Acids are the building blocks of proteins
2.. Amino Acids include: Amino Group (NH2), Carboxyl Group (COOH), Hydrogen, and an
R-group
3. 20 different R groups correspond to 20 possible amino acids
B. Peptides
1. Amino acids are joined together by peptide bonds
2. 2 amino acids = dipeptide, 3 amino acids = tripeptide, many amino acids = polypeptide
Polypeptide
C. Protein Structure
1. Primary Structure – sequence of amino acids
2. Secondary Structure – local shapes within a polypeptide chain, such as alpha helices
and beta sheets
3. Tertiary Structure – three dimensional shape of a polypeptide chain
4. Quaternary Structure – three dimensional shape of entire protein
4 levels of a protein’s structure
D. Protein Conformation
1. Conformation – shape of a protein. The shape of a protein determines its function
2. Denature
a. Denaturing destroys the function of a protein by altering its conformation
b. Chemicals, Heat, UV rays, radiation may denature a protein
XVIII. Nucleic Acids
A. Nucleotides
1. Nucleotides are the building blocks of nucleic acids
2. Nucleotides include a 5-carbon sugar (S), a nitrogenous base (B),
and a phosphate (P) group
3. Nucleotides are joined together by phosphodiester bonds
2 nucleotides = dinucleotide, many nucleotides = polynucleotide
B. Nucleic Acids
1. Deoxyribonucleic Acid (DNA)
a. sugar of DNA = deoxyribose
b. DNA is a double-stranded nucleic acid
c. DNA contains the genetic instructions of each cell
2. Ribonucleic Acid (RNA)
a. sugar of RNA = ribose
b. RNA is a single-stranded nucleic acid
c RNA has many functions: coding, regulation, enzymatic activity
C. Nucleic Acids will be discussed in further details in cell metabolism.
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