This document provides an overview of key chemistry and biochemistry concepts relevant to microbiology. It discusses the structure and function of atoms, molecules, and chemical bonds including covalent bonds, ionic bonds, hydrogen bonds, and hydrophobic interactions. It also examines biologically important elements and how water interacts with hydrophilic and hydrophobic molecules. Additionally, it covers acid-base balance, properties of water, and the role of buffers in maintaining pH.

1. Chemistry and Biochemistry for
Microbiology
Part 1
By BugLady

2. Structure and Function
Atoms and molecules
Chemical bonds strength
Water: the molecule of life
Carbohydrates: (energy source, storage, and cell structure)
Lipids (energy storage and membrane main component)
Proteins (enzymes, transporters, movement)
Nucleic acids (genetic material, small effector molecules)

3. Check point
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Electron
Nucleus
Neutron
Proton

4. Biologically Important Elements
Present in large amounts
C= every organic molecule, CO2
H= water, H+ , all organic molecules
O=water, e- acceptor, all organic molecules
N= proteins, nucleic acids, some vitamins
P= phospholipids, inorganic P, nucleic acids
S= proteins, some vitamins
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5. Why isn’t calcium mentioned as a
most abundant element in this class?
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6. Biologically Important Elements
Elements required in smaller amounts
Electrolyte balance, protein stabilization,
shells of some microorganisms
o Cations: K+ (potassium), Na+ (sodium), Fe2+/3+
(iron), Ca2+ (calcium), Mg2+ (magnesium)
o Anions: Cl- (chlorine), PO4
-3 (inorganic
phosphate)
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7. Biologically Important Elements
Essential elements required in trace amounts
o Metals: Zinc, Vanadium, Chromium, Molybdenum
o Lithium, Iodine, Selenium etc…
Found in vitamins and enzymes
Essential functions in a cell will be
compromised if trace elements are missing in
a diet.
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8. Chemical Bonds
Chemical bond: Atoms form interactions
involving their outer most electrons
There are 4 main types of chemical bonds.
Covalent bonds
Ionic Bonds
Hydrogen Bonds
Hydrophobic (non polar)
Interactions
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9. Comparison of Chemical Bonds
Strength of Bonds:
Covalent>ionic>hydrogen bonds>non polar
interactions
Non covalent interactions are individually
weak but collectively strong.
Ionic bonds and hydrogen bonds are polar
or hydrophilic
Non polar interactions are hydrophobic
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10. Covalent Bonds
Covalent bonds may have an equal or
unequal attraction for the shared electrons
Nonpolar covalent
o Bonds formed between identical atoms or
between atoms that have similar attraction of
electrons
 H-H or C-H
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National Institute of General Medical
Sciences

11. Covalent Bonds
Polar covalent
o One atom has a greater attraction to the
electrons than the other
o Produces a slight charge within the molecule
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12. Ionic Bonds
Loss and gain of electrons leads to charged atoms (ions)
Ionic bonds are attractions between ions of opposite
charge.
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13. Ionic Bonds
Ionic bonds in solution are weaker than
covalent bonds
o Bonds dissociate in water
 Easily broken at room temperature
 Approximately 100 time weaker than covalent
bonds
Important among weak forces holding
biological molecules together
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14. Hydrogen Bonds
Weak bonds formed from the attraction of
positively charged hydrogen atoms
o Hydrogen atoms in polar molecules are attracted
to negatively charged atoms or molecules
 Most commonly oxygen or nitrogen
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15. Hydrogen Bonds
Hydrogen bonds occur between molecules and
within molecules (water, proteins, DNA)
Covalent bonds occur within molecules
o Hydrogen bonds hold molecules together
o Covalent bonds hold atoms together
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16. Nonpolar (Hydrophobic) Interactions
Nonpolar residues (lipids or hydrophobic regions
of proteins) are exposed at the surface of two
different molecules.
It is energetically more favorable for polar water
molecules to approach each other closely
displacing the "oily" non-polar molecules from
between them.
o Example: plasma membranes.
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17. Properties of Water
Stabilized by
hydrogen bonds in
liquid state: H----O
High heat capacity
High surface tension
Excellent solvent
H+ and OH
participate in
chemical reactions
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National Institute of General Medical
Sciences

18. The Universal Solvent
Ionic compounds dissociate, forming solutes
o Salts, acids, bases
Polar compounds form hydrogen bonds with
water molecules
o Small alcohols, aldehydes, quaternary
ammonium compounds, vitamins
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Figure 2.5

19. Interaction with Water
Hydrophilic- “water-loving”
o Readily dissolve in water
o Ions and molecules that contain polar covalent
bonds
 Sugars, salts, protein hormones, some vitamins, small
organic molecules
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Courtesy USDA

20. Interaction with Water
Hydrophobic- “water-fearing”
o Do not readily dissolve in water
o Hydrocarbons, lipids, non polar solvents, lipid
soluble vitamins, steroids
Amphipathic molecule
o Have both polar or ionized regions
at one or more sites and nonpolar
regions at other sites
 Detergents are amphipatic
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Renee Comet NCi
Micelle by SuperManu CC BY-SA 3.0

21. Acid-Base Balance
pH is a measure of the concentration of H+ in a
solution.
pH =7.0 is neutral (pH of freshly distilled water)
pH>7.0 is basic
pH<7.0 is acidic
pH = log 1/[H+]=log[H+]
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22. Jul-14 22

23. Acid-Base Balance
Increasing [H+] increases acidity. pH decreases.
Increasing [OH] or pH increases alkalinity.
Most organisms grow best between pH 6.5 and 8.5
Optimal pH for most bacterial growth is slightly basic.
Optimal pH for fungal growth is slightly acidic.

24. Why Measure pH?
Acidic pH of skin and stomach are part of our
innate immune system
Proper function of proteins requires a specific pH.
DNA stability is dependent on pH.
Most biochemical assays used for bacteria
identification measure changes in pH.
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Lactose positive and
lactose negative
colonies of Escherichia
coli on Mac Conkey
agar

25. Buffers
Buffered solutions resist change in the pH of a
solution when hydrogen ions (protons) or hydroxide
ions are added or removed.
Buffers are critical to the maintenance of life.
Buffered systems mean that organisms can maintain
a suitable pH environment in their cells and tissues.
One of the most common buffers is the bicarbonate
buffer.
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