Download Chap 2 - mcvts

Your 3 week old patient had a heel stick for blood work
and has a Potassium level of 6.8meq (normal 3.5-5 meq).
The infant is cooing at mom and drinking a bottle. The
nurse should?
1. Ignore it
2. Have MD order more potassium
3. Give them a dose of magnesium sulfate STAT
4. Repeat the test
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Chapter 02
Chemistry of Life
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Lesson 2.1
Levels of Chemical Organization and Chemical Bonding
Define the term and describe the structure of an
atom.
2. Define the terms element, molecule, and compound.
3. Compare and contrast the major types of chemical
bonding.
1.
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3
Elements (cont.)
Matter
 The substances from which the universe is made
Elements
 All of the different types of matter
 Identified by names or chemical
symbols
 Also identified by number
 Described and organized in the
periodic table
11
Sodium
Na
22.99
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What is the symbol and number for carbon?
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Elements (cont.)
Living matter contains
26 of 92 natural
elements.
The body’s chemical composition by weight.
 96% of body weight—
four elements
 4% of body weight—
nine elements
 0.1% of body weight—
13 elements
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Levels of Chemical
Organization
 Atom—smallest unit of matter
 Nucleus—central core of atom




Proton—positively charged particle in nucleus
Neutron—uncharged particle in nucleus
Atomic number—number of protons in nucleus
Atomic mass—number of protons and neutrons
combined
From Sugimoto Y et al: Chemical identification of individual surface atoms by atomic force microscopy,
Nature 446:64-67, 2007.
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7
Atoms

Energy levels—orbital regions surrounding
atomic nucleus that contain electrons
Electron—negatively charged particle
Each region has space for a specific number of
electrons.

The first energy level has room for two electrons.

The second energy level has room for eight
electrons.

•
•

An atom is most stable when its energy levels are filled
with electrons.
Energy level increases the farther away it is from the
nucleus
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8
Models of the Atom
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9
Elements (cont.)
Hydrogen
Carbon
Total number of electrons
1
6
Number of electrons in first energy level
1
2
Number of electrons in second energy level
0
4
Energy Levels (cont.)
 Hydrogen has only one energy level with room for one
more electron.
 Carbon’s first energy level is full.
 Carbon’s second energy level has room for four more
electrons.
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Isotopes and Radioactivity (cont.)
Examples
Proton Number
Neutron Number
Atomic Weight
Carbon-12
6
6
12
Carbon-13
6
7
13
Carbon-14
6
8
14
Isotope
Isotopes
 Forms of an element that have the same atomic number but
different atomic weight

Different atomic weight because of a different number of
neutrons
 May be stable or unstable (radioactive)
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Elements, Molecules, and
Compounds



Element—a pure substance; made up of only one
kind of atom
Molecule—a group of atoms bound together in a
group
Compound—substances whose molecules have
more than one kind of atom
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12
Name
Oxygen
Symbol
O
Function
Part of water; needed to metabolize
nutrients for energy
Carbon
C
Basis of all organic compounds; component
of carbon dioxide, the gaseous byproduct of
metabolism
Hydrogen
H
Part of water, participates in energy
metabolism; determines the acidity of body
fluids
Nitrogen
N
Present in all proteins, ATP (the energystoring compound), and nucleic acids (DNA
and RNA)
Calcium
Ca
Builds bones and teeth; needed for muscle
contraction, nerve impulse conduction, and
blood clotting
Phosphorus
P
Active ingredient in ATP; builds bones and
teeth; component of cell membranes and
nucleic acids
Potassium
K
Active in nerve impulse conduction; muscle
contraction
Sulfur
Sodium
S
Na
Part of many proteins
Active in water balance, nerve impulse
conduction, and muscle contraction
Iron
Fe
Part of hemoglobin, the compound that
carries oxygen in red blood cells
The elements are listed in decreasing order by
weight ©2014
in the by
body.
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Chemical Bonding

Chemical bonds form to make atoms more
stable.


Atoms react with one another in ways that make
their outermost energy level full.
Atoms may share electrons or donate or borrow
them to become stable.
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14
Ionic Bonds

Ions form when an atom gains or loses
electrons in its outer energy level to become
stable


Positive ion—has lost electrons; indicated by
superscript positive sign(s), as in Na+ or Ca++
Negative ion—has gained electrons; indicated by
superscript negative sign(s), as in Cl–
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15
Ionic Bonds
 Ionic bonds form when positive and negative
(oppositely charged) ions attract each other
 Electrolyte—molecule that dissociates (breaks
apart) in water to form individual ions; an ionic
compound
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16
Ionic Bonding
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17
Covalent Bonds



Covalent bonds form when atoms share their
outer energy ions to complete the energy
level and thus become stable.
Covalent bonds do not ordinarily easily
dissociate in water.
Covalent bonding is used to form all of the
major organic compounds found in the body.
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18
Covalent Bonding
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19
Hydrogen Bonds



Hydrogen bonds do not create new
molecules.
Hydrogen bonds weakly bond to neighboring
molecules.
Hydrogen bonds are present in water, DNA,
and proteins.
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20
Hydrogen Bonds
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21
Lesson 2.2
Inorganic and Organic Chemistry
4. Distinguish between organic and inorganic chemical
compounds.
5. Discuss the chemical characteristics of water.
6. Explain the concept of pH.
7. Discuss the structure and function of the following
types of organic molecules: carbohydrate, lipid,
protein, and nucleic acid.
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22
Inorganic Chemistry


Organic molecules contain carbon-carbon
covalent bonds and/or carbon-hydrogen
covalent bonds; inorganic molecules do not.
Organic molecules are generally larger and
more complex than inorganic molecules.
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23
Water
 Water is an inorganic compound essential to
life.
 Water is a solvent (liquid into which solutes
are dissolved), forming aqueous solutions in
the body.
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24
Mixtures
Table 2-2 Mixtures
Type
Definition
Example
Solution
Homogeneous mixture formed when
one substance (solute) dissolves in
another (solvent)
Table salt (NaCl)
dissolved in water;
table sugar (sucrose)
dissolved in water
Suspension
Heterogeneous mixture in which one
substance is dispersed in another but
will settle out unless constantly mixed
Red blood cells in
blood plasma; milk of
magnesia
Colloid
Heterogeneous mixture in which the
suspended particles remain evenly
distributed based on the small size and
opposing charges of the particles
Blood plasma; cytosol
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Mixtures (cont.)
The Importance of Water
 Most abundant compound in body
 Critical in all physiologic processes
 Deficiency (dehydration) threatens health
 Universal solvent
 Stable liquid at ordinary temperatures
 Participates in body’s chemical reactions
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Water
 Water is involved in chemical reactions:
 Dehydration synthesis
 Hydrolysis
 Chemical reactions always involve energy
transfers, as when energy is used to build ATP
molecules
 Chemical equations show how reactants interact
to form products; arrows separate the reactants
from the products
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27
Acids, Bases, and Salts
Acid
 A substance that releases hydrogen ions
HCl  H+ + Cl−
Base
 A substance that releases hydroxide ions and accepts
hydrogen ions
NaOH  Na+ + OH−
Salt
 A substance formed by a reaction between an acid and a
base
HCl + NaOH  NaCl + H2O
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Water-based Chemistry
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29
Acids, Bases, and Salts
 Water molecules dissociate to form equal
amounts of H+ (hydrogen ion) and OH–
(hydroxide ion)
 Acid—substance that shifts the H+/OH– balance
in favor of H+; opposite of base
 Base—substance that shifts the H+/OH– balance
against H+; also known as an alkaline; opposite of
acid
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30
Acids, Bases, and Salts (cont.)
The pH Scale
 Measures the relative concentrations of hydrogen and
hydroxide ions in a solution.
 Scale from 0 (most acidic) to 14 (most basic).
 Each unit represents a 10-fold change.
 Normal body fluid pH range is between 7.35 and 7.45.
 Acidosis: Body fluid pH less than 7.35
 Alkalosis: Body fluid pH greater than 7.45
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pH

pH—Mathematical expression of relative H+
concentration in an aqueous solution


7 is neutral (neither acid nor base)
pH values above 7 are basic; pH values below 7
are acidic
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32
pH
 Neutralization occurs when acids and bases mix
and form salts.
 Buffers form chemical systems that absorb excess
acids or bases and thus maintain a relatively
stable pH. (proteins, Bicarbonate, or
hemeoglobin)
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33
Organic Chemistry
 Carbohydrates—sugars and complex
carbohydrates
 Contain carbon (C), hydrogen (H), and oxygen
(O)
 Monosaccharides—basic unit of carbohydrate
molecules (for example, glucose)
 Disaccharide—double sugar made up of two
monosaccharide units (for example, sucrose,
lactose)
 Polysaccharide—complex carbohydrate made up
of many monosaccharide units (for example,
glycogen; stored by the body)
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34
Carbohydrates
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Lipids—Fats and Oils

Triglycerides


Formed by a glycerol unit and joined to three
fatty acids
Store energy for later use
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36
Phospholipids



Similar to triglyceride structure, but have
phosphorus-containing units—each with a
head and two tails
The head attracts water and the double tail
does not, thus forming stable double layers
(bilayers) in water
Form membranes of cells
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Phospholipids
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38
Cholesterol



Molecules have a steroid structure made up
of multiple rings
Stabilizes the phospholipid tails in cellular
membranes
Also converted into steroid hormones by the
body
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39
Proteins


Very large molecules made up of amino acids
held together in long, folded chains by
peptide bonds
Structural proteins



Form essential structures of the body
Collagen is a fibrous protein that holds many
tissues together
Keratin forms tough, waterproof fibers in the
outer layer of the skin
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40
Protein
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41
Functional Proteins



Participate in chemical processes of the body
Examples include hormones, cell membrane
channels and receptors, and enzymes
Enzymes—chemical catalysts


Help chemical reactions occur
Enzyme action sometimes called lock-and-key
model
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42
Enzyme Action
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43
Nucleic Acids

Made up of nucleotides



A phosphate unit
A sugar (ribose or deoxyribose)
A nitrogen base (adenine, thymine or uracil,
guanine, cytosine)
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44
DNA (Deoxyribonucleic Acid)



Used as the cell’s “master code” for
assembling proteins
Uses deoxyribose as the sugar and A, T (not
U), C, and G as bases
Forms a double helix shape
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45
DNA
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46
RNA (Ribonucleic Acid)


Used as a temporary “working copy” of a
gene (portion of the DNA code)
Uses ribose as the sugar and A, U (not T), C,
and G as bases
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47
ATP (Adenosine Triphosphate)
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48
Key Terms
acid
chemistry
ion
salt
amino acid
colloid
isotope
solute
anion
compound
lipid
solution
aqueous
denaturation
molecule
solvent
atom
electrolyte
nucleotide
steroid
base
electron
neutron
substrate
buffer
element
pH
suspension
carbohydrate
enzyme
protein
valence
catalyst
glucose
proton
cation
glycogen
radioactive
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Word Anatomy (cont.)
Word Part
Meaning
Example
Chemical Bonds
co-
together
Covalent bonds form when atoms share
electrons.
Solutions and Suspensions
aqu/e
water
In an aqueous solution, water is solvent.
heter/o-
different
Heterogeneous solutions are different (not
uniform) throughout.
hom/o-
same
Homogeneous mixtures are the same
throughout.
hydr/o
water
Dehydration is a deficiency of water.
phil
to like
Hydrophilic substances “like” water—they mix
with or dissolve it.
phobia
to fear
Hydrophobic substances “fear” water—they
repel and do not dissolve it.
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Word Anatomy (cont.)
Word Part
Meaning
Example
Organic Compounds
-ase
suffix used in naming
enzymes
A lipase is an enzyme that acts on lipids.
de-
remove
Denaturation of a protein removes its ability to
function (changes its nature).
di-
twice, double
A disaccharide consists of two simple sugars.
glyc/o-
glucose, sweet
Glycogen is a storage form of glucose. It breaks down
to release glucose.
mon/o-
one
In a monosaccharide, “mono-” refers to one.
poly-
many
A polysaccharide consists of many simple sugars.
sacchar/o
sugar
A monosaccharide consists of one simple sugar.
tri-
three
Triglycerides have one fatty acid attached to each of
three carbon atoms.
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