Transcript 2 Chemistry Comes Alive Part A Sixth Edition

Chemistry Comes Alive
Part A
Human Anatomy & Physiology, Sixth Edition
Elaine N. Marieb
2
Matter
 Has mass & takes up space
 States of
 Solid – definite shape & volume
 Liquid –definite volume, changeable shape
 Gas – changeable shape & volume
Energy
 Capacity to do work
 Types of energy
 Kinetic – motion/action
 Potential – position; stored energy
Forms of Energy




Chemical – atomic bonds
Electrical – movement of charged particles
Mechanical – moving matter
Radiant – energy traveling in waves
 Conversion between forms
Composition of Matter
 Subatomic particles
 Electrons (e-)
 Neutrons
 Protons
 Atoms
 Unique arrangements of subatomic particles
 Can’t break down chemically
 Elements
 Matter of a single type of atom
 Atomic symbols
Properties of Elements
 Periodic table
 Elements grouped by properties
 Unique physical & chemical properties
 Physical properties –detected by senses
 Chemical properties –way atoms interact
Elemental Composition of the Human Body
 Major Constituents
 O ~ 65%
 C ~ 18.5%
 H ~ 9.5%
 N ~ 3.2 %
 Remaining ~ 3.8%
 Ca, P, K, S, Na, Cl, Mg, I, Fe
 Trace elements
 Zn, Mn, Cu
Atomic Structure
 Nucleus
 Neutrons
 no charge
 mass = 1 atomic mass unit (amu)
 Protons
 +1 charge
 mass = 1 amu
 Electron orbitals
 Electrons
 orbit nucleus in energy levels (orbitals)
 -1 charge
 mass = 0.0005 amu
Atomic Models
Planetary Model
Electrons move around the nucleus in
fixed, circular orbits
Orbital Model
Regions around the
nucleus in which
electrons are most likely
to be found
Characteristics of Atoms & Elements
 Atomic number –
 # protons
 Atomic mass –
 # protons + # neutrons
 Isotope –
 neutron # can vary
 different # of neutrons
 Atomic weight – average mass of all isotopes
 11C, 12C, 13C, 14C or 235U, 238U
 Radioisotopes – atoms that undergo spontaneous decay
called radioactivity
 131I, 99mTc, 60Co, 14C, or 235U
Subatomic Configuration of Elements
Figure 2.2
Isotopes of Elements
Figure 2.3
Chemically Inert Elements
 Inert elements have full outer e- orbitals
 Full = 8 e- (except for He)
Figure 2.4a
Chemically Reactive Elements
 Reactive elements
have unfilled outer
orbitals
Figure 2.4b
Molecules & Compounds
 Molecule – ≥ 2 atoms bonded
 Compound – ≥ 2 different kinds of atoms bonded
 ALL Compounds Are Molecules
 BUT
 All Molecules NOT Compounds
Chemical Bonds
 Chemical bonds formed by e- in outer orbitals
(valence shell)
 The Octet rule
 Atoms interact to have 8 e-s in valence shells
 (except H only 2 e-s)
 lose, gain or share
 Types of Bonds
 Ionic – loss or gain e-s
 Covalent – sharing e-s
Ionic Bonds
 Ions
 charged atoms
 Anions - charge = gained e Cations + charge = lost e Example: NaCl (sodium chloride)
Formation of an Ionic Bond
 Ionic compounds form a crystal structures
Covalent Bonds
 Sharing e-s fill outer orbitals
Single bond
each atom donates 1 e-
Double & Triple Covalent Bonds
 Atoms share 2 or 3 e-s
Polar & Nonpolar Bonds
 Polar bonds
 Unequal e- sharing
 Atoms w/ 6 -7 valence shell e-s = electronegative
 Atoms w/ 1-2 valence shell e-s = electropositive
 Electronegative & electropositive atoms form ionic
bonds
 Nonpolar bonds
 Equal sharing of e-s
 Atoms with 3-5 valence electrons form covalent
bonds
Comparison of Ionic, Polar Covalent, &
Nonpolar Covalent Bonds
Hydrogen Bonds
 Very important type of bond for life functions
 Allows reversible interactions between molecules
 Due to unequal sharing of H’s electron with N or O
 Responsible for properties of H2O
 Very important bonds between large macromolecules
(ie proteins & nucleic acids)
Hydrogen Bonds in H2O
Figure 2.9
Properties of Water
 High heat capacity – absorbs & releases large amounts
of heat before changing temperature
 High heat of vaporization – changing from a liquid to
a gas requires large amounts of heat
 Polar solvent properties – dissolves ionic substances,
forms hydration layers around large charged
molecules, & serves as the body’s major transport
medium
 Reactivity – is an important part of hydrolysis &
dehydration synthesis reactions
Mixtures & Solutions
 Mixtures – two or more components physically
intermixed but not chemically bonded
 Solutions – homogeneous mixtures of components
 Solvent – substance present in greatest amount
 Solute – substance(s) present in smaller amounts
 Colloids - heterogeneous mixtures whose solutes do
not settle out
 Suspensions - heterogeneous mixtures with visible
solutes that settle out
Concept of Concentration
 Concentration refers to the amount of a substance in
a given volume
 A critical concept to master
Units of Concentration
 Percent – parts per hundred
 PPM – parts per million
 PPB – parts per billion
 Molarity – moles per liter
 Molality – moles particles per kg solvent
 Mass/volume
 g/ml (gram per milliliter)
 mg/ml (milligram per milliliter )
 g/ml (microgram per milliliter)
 Mass/Mass
 mg/kg body mass
 g/kg body mass
Molecular Weight, Moles & Molarity
 Molecular weight
 The mass of a mole of atoms or molecules
 Has units of g/mole
 = atomic weight of an atom in grams
 1 mole of C = 12 g OR C is 12g/mole
 Molecule’s molecular weight = sum of the atomic
weights of its atoms
 1 mole of H2O = 1g + 1g + 16g = 18 g
 MW of H2O = 18g/mole
Molecular Weight, Moles & Molarity
 What’s a mole??
 The number of molecules in the gram molecular
weight of that molecule
 Always 6.02 x 1023
 This magical number is Avogadro’s Number
 MW of C = 12
 there are 12g C/mole C
 there are 6.02 x 1023 C atoms in 12g of C
 MW of H2O = 18
 there are 18g H2O/mole H2O
 there are 6.02 x 1023 H2O molecules in 18g of
H2 O
Molarity – The Standard of Chemical Concentration
Terms
 Molarity = moles of solute per liter (L) of solvent
 Abbreviated with M
 A 5M NaCl solution contains 5 moles of NaCl
molecules per liter of solution
 So 1 L of a 5 M NaCl solution contains 3.01 x 1024
molecules of NaCl
Calculating Molarity
 Chemical formula for glucose is C6H12O6
 MW = 180 g/mole
 What is concentration of glucose in a can of coke?
 42g of glucose/355ml H2O
 How many moles of glucose?

42g / 180g/mole = 0.233 moles
 How many liters of coke?

355 ml / 1000ml/L = 0.355 L
 How many moles/liter?

0.233 moles/0.355 L = 0.656 M
Examples using Molarity
 Ion concentrations in cells & body fluids
 [Na] = 0.15M outside cells & 0.015M inside cells
 [K] = 0.005M outside cells & 0.15M inside cells
 Molar is often converted to millimolar (mM)
 M = moles/L
 mM = millimoles/L
 0.15M = 150mM
 Simply multiply M by 1000 to convert M to mM
Mass/Volume & Percent Concentration Terms
 Many molecules in the body are measured in mass/volume
 Normal glucose = 100mg/dL
 or 100mg/100ml
 or 1mg/ml
or 1g/L
 or 0.001g/ml
 or 0.1g/100ml
 Cholesterol should be below 200mg/dL
 or 0.2g/dL
 or 0.2g/100ml
 Many substances are described in percentages
 Percentage is g/100g or g/100ml
 Blood [glucose] would = 0.1%
 Blood [cholesterol] would = 0.2%
Chemical Reactions
 Forming or breaking chemical bonds
 Chemical equations show:
 Reactants & products & their relative amounts
Patterns of Chemical Reactions
 Combination reactions: Synthesis reactions which
always involve bond formation
A + B  AB
 Decomposition reactions: Molecules are broken
down into smaller molecules
AB  A + B
 Exchange reactions: Bonds are both made & broken
AB + C  AC + B
Oxidation-Reduction (Redox) Reactions
C6H12O6 + 6O2
glucose
6CO2 +
6H2O
carbon
dioxide
 Reactants losing electrons are electron donors & are
oxidized
 Reactants taking up electrons are electron acceptors
& become reduced
Energy Flow in Chemical Reactions
 Exergonic reactions – reactions that release energy
 Endergonic reactions – reactions whose products
contain more potential energy than did its reactants
Rates & Reversibility of Chemical Reactions
A+B
AB
 Chemical reactions proceed with measurable rates
 All reactions are theoretically reversible
 Equilibrium (dynamic)
 Forward & reverse reactions proceed at same rate
Factors Influencing Rate of Chemical Reactions
 Temperature
 higher temperatures increase rates
 Concentration
 higher [reactant] increases rates
 Catalysts
 Molecules that increase reaction rates
 Enzymes
 Biological catalysts with high specificities