Chemistry Comes Alive (Chapter 2)
Matter
Mass is constant, weight varies w/ gravity
Anything that has mass and takes up space
States of matter
Solid – has definite shape and volume
Liquid – has definite volume, changeable shape
Gas – has changeable shape and volume
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Energy
The capacity to do work (put matter into motion)
Types of energy
Kinetic – energy in action
Potential – energy of position; stored (inactive)
energy
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Energy Concepts
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Forms of Energy
Chemical – stored in the bonds of chemical substances
(e.g. ATP
ADP)
P+
Electrical – results from the movement of charged particles
(e.g. flow of electrons down a wire or ions moving
along/across cell membrane
Mechanical – directly involved in moving matter
Radiant or electromagnetic – energy traveling in waves
(i.e., visible light, ultraviolet light, and X-rays)
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Energy Form Conversions
Energy is easily converted from one form to
another
During conversion, some energy is “lost” as heat
Thus, energy conversion is inefficient due to “loss”
of unusable energy
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Composition of Matter
Elements – unique substances that cannot be
broken down by ordinary chemical means (e.g.
oxygen, carbon, gold, etc… There are 112
elements, 92 existing in nature, that we know of)
Atoms – more-or-less identical building blocks for
each element
Empirical (Can be measured) e.g. H+ is 0.1 nm in
diameter.
Atomic symbol – one- or two-letter chemical
shorthand for each element
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Properties of Elements
Each element has unique physical and chemical
properties
Physical properties – those detected with our
senses
Chemical properties – pertain to the way atoms
interact with one another
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Major Elements of the Human Body
Oxygen (O)
Carbon (C)
Hydrogen (H)
Nitrogen (N)
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Lesser and Trace Elements of the Human Body
Lesser elements make up 3.9% of the body and
include:
Calcium (Ca), phosphorus (P), potassium (K),
sulfur (S), sodium (Na), chlorine (Cl), magnesium
(Mg), iodine (I), and iron (Fe)
Trace elements make up less than 0.01% of the
body
They are required in minute amounts, and are
found as part of enzymes
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Atomic Structure
The nucleus consists of neutrons and protons (very
dense due to concentrated mass)
Neutrons – have no charge and a mass of one atomic
mass unit (amu)
Protons – have a positive charge and a mass of
1 amu
Electrons are found orbiting the nucleus
Electrons – have a negative charge and 1/2000 the
mass of a proton (0 amu)
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Models of the Atom
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
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Size and space in an atom
The number of protons and electrons in an atom is
the same (overall neutral charge)
E.g. if an atom were the size of a football field, the
nucleus would be the size of ref’s whistle and the
electrons would be the size of a fly sitting on the
goal post.
Thus, there is a lot of empty space in an atom!
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Models of the Atom
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Figure 2.1
Identification of Elements
Atomic number – equal to the number of protons (e.g. 1H;
also the number of electrons)
Mass number – equal to the # of Protons + # of Neutrons
(e.g.42He)
Isotopes possess the same # of protons & electrons but
differ in the number of neutrons (e.g. 1H, 2H, 3H possessing
0, 1, and 2 neutrons respectively)
Atomic weight – average of the mass numbers of all
isotopes of an element. Relatively equal to the mass
number of its most abundant isotope in nature.
What distinguishes the elements is the number of protons,
neutrons, and electrons
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Identification of Elements
Isotope – atoms with same number of protons but a
different number of neutrons
Radioisotopes – Heavy isotopes are unstable and
decompose into stable forms (radioactivity)
E.g. α (alpha) particles (packets of 2 protons, 2
neutrons, lowest penetrating power)
β (beta) particles (electron-like negative
particles
γ (gamma) rays (electromagnetic energy, greatest
penetrating power)
All three can be ejected from the nucleus during decay
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Identification of Elements: Atomic Structure
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Figure 2.2
Identification of Elements: Isotopes of Hydrogen
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Figure 2.3
Molecules and Compounds
Molecule – two or more atoms held together by
chemical bonds
Molecule of the element-2 or more same atoms
Compound – two or more different kinds of atoms
chemically bonded together
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Mixtures and Solutions
Mixtures – two or more components physically
intermixed (not chemically bonded)
Solutions – homogeneous mixtures of components
be it gases, liquids, or solids
Solvent – substance present in greatest amount
(dissolving medium)
Solute – substance(s) present in lesser amounts
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Concentration of Solutions
Percent of the solute in the total solution. (parts per 100
parts)
Molarity, or moles per liter (M)
A mole of an element or compound is equal to its atomic
or molecular weight (sum of atomic weights) in grams
E.g. a 1 M solution is that # weighed out in grams (gram
molecular weight) and made up in a final volume of 1 liter.
One mole of any substance contains exactly the same # of
solute particles, Avogadro’s Number (6.02 x 1023) except
for ions.
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Colloids and Suspensions
Colloids (emulsions) – heterogeneous mixtures
whose solutes do not settle out and thus appears
milky or cloudy. E.g. Jell-O
“ Colloid Cosby”
Suspensions – heterogeneous mixtures with visible
solutes that tend to settle out
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Mixtures Compared with Compounds
Mixtures
No chemical bonding takes place in mixtures
The properties of the atoms/molecules do not change
Most mixtures can be separated by physical means
Mixtures can be heterogeneous or homogeneous
Compounds
Chemical bonding occurs between the components
Compounds cannot be separated by physical means, only
by chemical means (e.g. breaking of bonds)
All compounds are homogeneous
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Chemical Bonds
Up to 7 electron shells form the electron cloud
(valence shells)
Each electron shell represents a different energy
level
Electrons furthest from the nucleus have the
greatest potential energy & are most likely to
interact chemically w/ other atoms (form bonds).
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Chemical Bonds
Valence shell – outermost energy level containing
chemically active electrons
Each electron shell can hold a specific # of
electrons e.g.
valence shell 1, 2
electrons
valence shell 2, 8 electrons
valence shell 3, 18 electons
etc…
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Chemical Bonds
Valence shells are filled consecutively
Bonding electrons are found in the outermost shell
When valence shell is filled to capacity or contains
8 electrons, the atom is stable and chemically inert,
e.g. Noble gases
Octet rule – except for the first shell which is full
with two electrons, atoms interact in a manner to
have eight electrons in their valence shell
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Chemically Inert Elements
Inert elements have their outermost energy level
fully occupied by electrons
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Figure 2.4a
Chemically Reactive Elements
Reactive
elements do
not have their
outermost
energy level
fully occupied
by electrons
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Figure 2.4b
Types of Chemical Bonds
Ionic
Covalent
Hydrogen
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Ionic Bonds
Atoms are electrically neutral
Ions form when the balance between + & - charge
is lost due to the transfer of electrons
Thus, Ions are charged atoms resulting from the
gain or loss of electrons
Anions have gained one or more electrons (net –
charge)
Cations have lost one or more electrons (net +
charge)
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Formation of an Ionic Bond
Ionic bonds are chemical bonds that form between
atoms involving the transfer of one or more
electrons from one atom to another
Ionic compounds form crystals instead of
individual molecules
Example: NaCl (sodium chloride)
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Formation of an Ionic Bond
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Figure 2.5a
Formation of an Ionic Bond
Na-2-8-1
7-8-2-Cl
NaCl achieves stability by chemically binding Na and
Cl ions
Most ionic compounds are salts like NaCl
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Figure 2.5b
Covalent Bonds
Covalent bonds are chemical bonds formed
between atoms involving the sharing of one or
more electrons from one atom to another in order
to fill the outer valence shell at least part of the
time.
Shared electrons occupy a single common orbital
common to both atoms
e.g. H-1 1-H
Hydrogen gas
H-2-H
Electron sharing produces molecules
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Covalent Bonds
Types of covalent bonds:
Single bond
H-H
Double bond
O=O
Triple bond
N≡N
Single, double, and triple refer to the number of
electrons shared
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Single Covalent Bonds
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Figure 2.7a
Double Covalent Bonds
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Figure 2.7b
Triple Covalent Bonds
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Figure 2.7c
Polar and Nonpolar Molecules
Electrons shared equally between atoms produce
nonpolar molecules
Unequal sharing of electrons produces polar
molecules
Atoms with six or seven valence shell electrons are
electronegative
Atoms with one or two valence shell electrons are
electropositive
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Comparison of Ionic, Polar Covalent, and
Nonpolar Covalent Bonds
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Figure 2.9
Hydrogen Bonds
Polar molecules, like water, orient themselves
toward other dipoles and charged particles
Hydrogen bonds form when a hydrogen atom
already covalently linked to a electronegative atom
is attracted by another electronegative atom.
E.g. Water molecule & surface tension
Helix formation in DNA
Secondary structural component in proteins
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Hydrogen Bonds
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Hydrogen Bonds
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Figure 2.10a
Chemical Reactions
Occur when chemical bonds are formed,
rearranged, or broken
Written in symbolic form using chemical equations
Chemical equations contain:
Number and type of reacting substances, and
products produced
Relative amounts of reactants and products
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Examples of Chemical Reactions
(5 Unjoined atoms)
(4 Chemical bonds)
Thus, 4 moles of hydrogen + 1 moles carbon yield
4 moles of methane
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Patterns of Chemical Reactions
Combination reactions:
Synthesis reactions which always involve bond
formation
A + B → AB Anabolic: Bond Formation
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Patterns of Chemical Reactions
Decomposition reactions: Molecules are broken
down into smaller molecules
AB → A + B
Catabolic: Bonds broken
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Patterns of Chemical Reactions
Exchange reactions: Bonds are both made and
broken
AB + C → AC + B Anabolic & Catabolic
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Oxidation-Reduction (Redox) Reactions
Decomposition rxns. Food
ATP conversion
Reactants losing electrons are electron donors and
are oxidized
Lose Electron Oxidized (LEO)
Reactants taking up electrons are electron
acceptors and become reduced
Gain Electron Reduced (GER)
Leo the Lion
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GER!
Energy Flow in Chemical Reactions
Exergonic reactions – reactions that release energy
e.g. catabolic & oxidative rxns
Endergonic reactions – reactions that absorb
energy
e.g. anabolic rxns.
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Reversibility in Chemical Reactions
All chemical reactions are theoretically reversible
A + B → AB
AB → A + B
If neither a forward nor reverse reaction is
dominant, chemical equilibrium is reached
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Factors Influencing Rate of Chemical Reactions
Must overcome the repulsion of electrons
Temperature – chemical reactions proceed quicker
at higher temperatures
Particle size – the smaller the particle the faster the
chemical reaction
Concentration – higher reacting particle
concentrations produce faster reactions
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Factors Influencing Rate of Chemical Reactions
Catalysts – increase the rate of a reaction without
being chemically changed
Enzymes – biological catalysts
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