1. Chapter 2 Introductory Chemistry
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Chemical Elements
Fundamental unit in chemistry
112 elements total
Use 1-2 letter symbols for each
Examples: C= carbon, Na = sodium, Cl = chorine.
26 elements present in human body
4 major ones (O, C, H, and N) make up 96%
8 others significant also. See Table 2.1.
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Atoms
Smallest unit of an element that retains characteristics of an element
contains
Nucleus that has protons (+), neutrons (0)
Electrons (–) surrounding nucleus
Total charge is neutral:
Protons # = electron #
Atomic number = number of protons = number of electrons
Mass number = number of protons + number of neutrons
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4. Ions, Molecules and Compounds
When an atom gives up or gains an electron, it becomes an ion
When atoms share electrons, they form a molecule
Two or more different atoms held together with chemical bonds = a compound
Described by the molecular formula
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Molecular Formula
O2 = oxygen
Molecule: has 2 atoms bound together
H2O = water
Compound has 2 different atoms:
H (hydrogen): 2 atoms
O (oxygen): 1 atom
Subscript indicates # of atoms of element
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Molecules
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Chemical Bonding
Attraction between atoms to form attachments
Electrons are grouped into shells
Number of electrons in outer shell determines type of bonding
Types of bonds:
Ionic
Covalent
Hydrogen
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Ionic Bonds
Electron is donated or accepted from another atom  ion
Typically occurs between atoms in which:
One has just 1 or 2 electrons in outer shells
Other has almost full outer shell (6 or 7 electrons)
Electrons are negative (–) so:
If electron is accepted, atom  negative ion: anion
If electron is donated, atom  positive ion: cation
Opposite charges attract  ionic bonding
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Ionic Bonds
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Covalent Bonds
Sharing of electrons in outer shell  covalent bonds
Typically occurs between atoms in which outer shells are about half full.
Example: bonds involving carbon (C) atoms (with 4 electrons in outer shell). These are organic compounds.
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Covalent Bonds – single bond
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Covalent Bonds – double bonds
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Covalent Bonds – triple bonds
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Covalent Bonds – 4 bonds
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Hydrogen Bonds
Form when a hydrogen atom (with a partial positive charge) attracts the partial negative charge of neighboring atoms, such as oxygen or nitrogen.
Contribute strength and stability within large complex molecules such as
DNA
Proteins
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Chemical Reactions
Occur when old bonds break and new bonds form
Types:
Synthesis
Decomposition
Exchange
Reversible
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17. Chemical Reactions: Synthesis
Putting atoms together to form larger molecules
A + B  AB
Example: 2H2 + O2  2 H2O
Synthesis in the body = anabolism
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18. Chemical Reactions: Decomposition
Splitting molecules apart
AB  A + B
Example: CH4  C + 2H2
Decomposition in the body = catabolism
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19. Chemical Reactions: Exchange
Involve both synthesis and decomposition
AB + CD  AD + BC
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20. Chemical Reactions: Reversible
Can go in either direction: synthesis or decomposition or exchange
Examples:
A + B ↔ AB
AB ↔ A + B
AB + CD ↔ AD + BC
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Classes of Chemicals
Inorganic
Structure: lack C-H bonds; structurally simple
Examples
Water, carbon dioxide, bicarbonate, acids, bases, and salts
Organic
Structure:
All contain C-H bonds
Structurally complex (include polymers composed of many units = monomers)
Classes: carbohydrates, lipids, proteins, nucleic acids
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22. Inorganic Compounds: Water
Characteristics of water
Most abundant chemical in human body
Good solvent and lubricant
Takes part in chemical reactions
Absorbs and releases heat slowly; regulates body temperature
Involved in digestion, circulation, and elimination of wastes
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Acids, Bases and Salts
Acid dissolves  H+ (1 or more)
Base dissolves  OH- (1 or more)
Acid + base  salt
Example: HCl + NaOH  NaCl + H2O
acid + base  salt + H2O
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pH Concept
The concentration of H+ or OH– expressed on the pH scale
pH scale: 0–14
pH 7.0: H+ concentration = OH– concentration
pH < 7.0 = more H+ (acid)
The smaller the number, the more H+
pH > 7.0 = more OH– (alkaline)
The larger the number, the more OH–
Alkalosis
Acidosis
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25. Organic Compounds: Carbohydrates
Most common sources of energy for humans
Three major classes: mono-, di-, poly-
Monosaccharide: simple sugar. Common examples:
Glucose (blood sugar) and fructose (fruit sugar)
Disaccharides: two bonded monosaccharides
Larger carbohydrates formed by dehydration synthesis and broken down by hydrolysis
Glucose + fructose ↔ sucrose (table sugar)
Glucose + galactose ↔ lactose (milk sugar)
Glucose + glucose ↔ maltose
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Disaccharide
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Polysaccharides
Monosaccharides (monomers) in long chains
Complex branching structures not usually soluble in water
Examples
Glycogen: carbohydrate stored in animals (liver, muscles)
Starch: carbohydrate stored in plants (potatoes, rice, grains)
Cellulose: plant polymer (indigestible fibers)
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Polysaccharides
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29. Organic Compounds - Lipids
Characteristics
Insoluble in water = hydrophobic
Functions: protect, insulate, provide energy
Classes
Triglycerides
Most plentiful in diet and body
Each composed of 3 fatty acids + 1 glycerol
May be saturated, monounsaturated, or polyunsaturated
Phospholipids: form lipid bilayer in membranes
Steroids based on ring-structure of cholesterol
Fat-soluble vitamins: A, D, E, and K
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Lipids: Triglycerides
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Lipids: Phospholipids
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Lipids: Steroids
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Cholesterol
Used to make steroid hormones
Estrogen, testosterone, cortisone
Help make plasma membranes stiff
Made in liver
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34. Organic Compounds - Proteins
Structure: composed of amino acids (monomers)
20 different amino acids
Amino acid structure: central carbon with
Acid (carboxyl) group (COOH)
Amino group (NH2)
Side chain (varies among the 20 amino acids)
Amino acids joined in long chains
By dehydration synthesis to form peptide bonds  dipeptide  tripeptide  polypeptide
Ultimately, form large, complex structures
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Length of AA chain
Peptide – 2-9 AA’s
Polypeptide – AA’s
Protein – 100-thousands of AA’s
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Protein Structure
Primary (10) – sequence of AA’s
Secondary (20) – twisting of AA’s due to H-bonding
Tertiary (30) – folding of AA chain due to ionic bonds, disulfide bridges, & hydrophobic interactions
Quaternary (40) – interactions between different AA chains
* A protein must be in Quaternary structure to be functional!
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37. Denaturation of Proteins: loss of 3-dimensional conformation (shape)
Extreme pH
Extreme T
Harsh chemicals
High salt concentrations
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Functions of Proteins
Structure – keratin in hair, nails & skin
Transport – hemoglobin
Chemical messengers – hormones, neurotransmitters
Movement – actin & myosin in muscle
Defense – antibodies
Catalysts - enzymes
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Amino Acids
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Enzymes
Proteins that serve as chemical catalysts
Highly specific: one enzyme works on a specific substrate  product
Efficient: one enzyme used over and over
Names
Most end in “-ase”
Many give clues to functions: sucrase, lipase, protease, dehydrogenase
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Enzymes
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Nucleic Acids
DNA or RNA
Huge polymers composed of nucleotides
Each nucleotide (monomer) consists of
Sugar (5-C monosaccharide: ribose or deoxyribose)
Phosphate
Nitrogen-containing (nitrogeneous) base
In DNA: adenine (A), guanine (G), cytosine (C), or thymine (T)
In RNA: adenine (A), guanine (G), cytosine (C), or uracil (U) (which replaces T of DNA)
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DNA Molecule
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Nucleic Acids: DNA
Nucleotides are connected into long chains that are bonded by bases:
C – G, G – C, T – A, or A – T
Two chains form double helix (spiral ladder)
Function: stores genetic information in genes (found in chromosomes) that:
Direct protein synthesis and therefore regulate everyday activities of cells
Carry this genetic information to the next generation of cells
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Nucleic Acid: RNA
Nucleotides are connected into a long, single chain (one side of a ladder)
Function:
Carries out protein synthesis by correctly sequencing amino acids, so helps to regulate everyday activities of cells
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46. ATP – Adenosine TriPhosphate
Function: the main energy-storing molecule in the body
ATP contains 3 phosphates
Carries energy in high-energy chemical bonds between terminal phosphate groups
Energy released from those bonds when they break: ATP  ADP + phosphate + energy
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Structure of ATP and ADP
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End of Chapter 2
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