Matter
• The “stuff” that makes up the universe
– anything that takes up space
• States of matter
– Solid
• has definite shape and volume
– Liquid
• has definite volume, changeable shape
– Gas
• has changeable shape and volume
Composition of Matter
• Elements
– unique substances that cannot be broken down by
ordinary chemical means
• carbon, oxygen, helium, uranium, gold, iron…
• Only 24 elements have a role in our body
– 98.5% of body weight consists of
• O, C, H, N, Ca, P
• Atoms
– building blocks for each element
The Chemical Elements
• Periodic table
– atomic symbols of elements arranged by atomic
number
• Atomic number of each element
– number of protons in its nucleus
• Only 24 elements have a role in our body
– 98.5% of body weight consists of
• O, C, H, N, Ca, P
Atomic Structure (Bohr or Planetary Model)
• Nucleus, the center of atom which contains:
– protons
• positive charge, mass (weight) of 1 atomic mass
unit (amu)
• determines atomic number
– neutrons
• neutral (no) charge, mass of 1 amu
• Electron shells that surround the nucleus and contain:
– electrons
• negative charge, mass of 0 amu
• valence electrons are in the outermost shell
–furthest from the nucleus
–interact with other atoms
• All atoms have:
– an equal number of protons and electrons
• atoms are neutral (have no net charge)
– an atomic mass = total mass of protons + neutrons
Planetary Model of an Atom
Electron Shells
• The electron shell closest to the nucleus can hold up
to 2 electrons
– additional electrons are located in shells outside of
the first shell
• All other electron shells outside of the first shell can
hold up to 8 electrons
• octet rule
–an electron shell is full when there are:
• 2 electrons in the first shell
• 8 electrons in 2nd, 3rd, 4th,… shell
• If the valence shell of an atom is not completely FULL,
then that atom is UNSTABLE
– valance electrons of unstable atoms interact with
valance electrons of other unstable atoms to create
chemical bonds
• allows both atoms to become stable
Atomic Structure
• Nucleus, the center of atom which contains:
– protons
• positive charge, mass (weight) of 1 atomic mass
unit (amu), determines atomic number
– neutrons
• neutral (no) charge, mass of 1 amu
– atomic mass = total # of protons + neutrons
• Electron shells
– electrons: negative charge
– # of electrons = # of protons, atoms have neutral
charge
– electrons further from nucleus have higher energy
– valence electrons are in the outermost shell
• interact with other atoms
• determine chemical behavior
Planetary Model of an Atom
Electron Shells
• The electron shell closest to the nucleus can hold up
to 2 electrons
– additional electrons are located in shells out side of
the first shell
• All other electron shells outside of the first shell can
hold up to 8 electrons
• Valance electrons of an atom interact with valance
electrons of other atoms in order to obey the octet rule
– octet rule: atoms react to obtain a full valence shell
• 2 electrons in the first shell or 8 electrons in 2nd,
3rd, 4th,… shell
• an atom with a full valence shell is STABLE
Planetary Models of Atoms
p+ represents protons, no represents neutrons
Chemical Bonds
• The reaction between 2 atoms results in the formation
of a chemical bond
• Bonds are formed between 2 atoms using the
electrons in the valence shell of each atom
• An atom is stable when the valence shell is completely
full (satisfying the “octet rule”)
• 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 (thus becoming
stable)
Chemically Inert (Non-reactive) Elements
• Inert elements have their valence shell fully occupied
by electrons
Chemically Reactive Elements
• Reactive elements
do not have their
outermost energy
level fully occupied
by electrons
Types of Chemical Bonds
• Ionic
• Covalent
• Hydrogen
Molecules and Compounds
• Molecules
– two or more atoms of same element covalently
bonded
• Compounds
– two or more atoms of different elements covalently
bonded
• Structural formula
– shows arrangement of atoms
• Molecular formula
– itemizes each element present and its quantity
Molecular and Structural Formula
Formation of an Ionic Bond
• Ionic bonds form between atoms after the transfer of
one or more electrons from one atom to another atom
in a process called ionization
– results in the creation of 2 ions
• Ions
–atoms that have an unequal numbers of
protons and electrons
–also known as electrolytes
Anions and Cations
• Cation
– atom lost one or more electrons
– more protons than electrons = net positive charge
• Anion
– atom gained one or more electrons
– more electrons than protons = net negative charge
Anions and Cations
• The anion and the cation are held together by the
ATTRACTION between a positively charged
substance and a negatively charged substance
– 2 substances that have the same charge REPELL
one another
• Atoms bound by ionic bonds form crystals (salts)
– NaCl (sodium chloride)
• Ionic bonds cannot exist in water
– the bond breaks (dissociates)
• the components of the molecule exist in their
ionic form (anion and cation) in the body
–NaCl + H2O→ Na+ and Cl- +H2O
Covalent Bonds
• Formed by sharing valence electrons between 2
atoms
• Types of covalent bonds include:
– polar and nonpolar covalent bonds
• Are stable (Do NOT dissociate) in water
• Nearly all biologically important molecules are made
from atoms bonded together with covalent bonds
Polar Covalent Bonds
• Polar (electrically charged) bonds occur between an
electronegative atom (O or N) and an atom that is
neither O nor N
– the nucleus of an electronegative atom has a
stronger “pull” on the shared electrons, pulling the
shared electron(s) closer to it
– this causes the electronegative atom to become
partially negative while the other atom in the
covalent bond becomes partially positive
• similar to a battery
electrons shared
unequally
Nonpolar Covalent Bonds
• Nonpolar bonds occur between two atoms neither of
which are O or N
– the shared electrons are pulled equally between the
nuclei of the 2 atoms
– both atoms remain neutral
• the nonpolar covalent bond is neutral
electrons shared
equally
Ions
• Ions – atoms that carry a charge (unequal
numbers of protons and electrons)
• Ionization - transfer
of electrons from one
atom to another
(provides stability of
the atom)
Anions and Cations
• Anion - atom gained electron, net negative charge
• Cation - atom lost an electron, net positive charge
Formation of an Ionic Bond
• Ionic bonds form between atoms by the transfer of one
or more electrons
• One atom donates one or more electrons, becoming a
cation (+), the other atom accepts the donated
electron(s), becoming an anion (-)
• The anion and the cation are attracted to one another
and are held together by an ionic bond
• Example: NaCl (sodium chloride)
• In water (such as in your body) the ionic bond
dissociates (breaks) and the components of the
molecule exist in their ionic form (anion and cation)
• NaCl → Na+ and Cl-
Ionic Bonds
• Ionic bonds are weak and dissociate (break) in water
• These compounds tend to form crystals (salts)
Formation of an Ionic Bond
Formation of an Ionic Bond
Formation of an Ionic Bond
Sodium Chloride Crystal
Sodium Chloride Crystals
Covalent Bonds
• Formed by sharing valence electrons between 2
atoms
• Different types of covalent bonds
– single covalent bond
– double covalent bond
– nonpolar covalent bond
– polar covalent bond
• DO NOT dissociate in water
– because all molecules that are formed by ionic
bonds dissociate in water, only molecules that are
formed by covalent bonds exist (and function) in the
body
Single Covalent Bond
• One pair of electrons are shared
Single Covalent Bond
Double Covalent Bond
• Two pairs of electrons are shared
Sharing of electrons
• The sharing of electrons in a covalent bond can be
either “equal” or “unequal”
• When the shared electrons between 2 atoms is such
that the electrons are located exactly between the
nuclei of 2 atoms, the sharing is said to be equal
– In other words, the shared electrons are not closer
to the nucleus of either atom
• However, when the shared electrons between 2 atoms
is such that they are closer to one of the 2 atoms, the
sharing is said to be unequal
– In other words, the shared electrons are closer to
the nucleus of one atom and further away from the
nucleus of the other atom
Nonpolar and Polar Covalent Bonds
• Electrons shared equally between atoms produce
nonpolar bonds because the negative charge of the
electron is spaced evenly between the 2 atoms
• Unequal sharing of electrons produces polar bonds
– Polar bonds occur between an electronegative
atom (O or N) and an atom that is not O or N
– Electronegative atoms have a stronger “pull” on the
shared electrons, pulling the electrons closer to it
– This causes the electronegative atom to become
partially negative while the other atom in the
covalent bond becomes partially positively
Nonpolar and Polar Covalent Bonds
electrons shared
equally
electrons shared
unequally
Functional Groups
• small groups of covalently
bonded atoms arranged
in a very specific manner
• parts of large compounds
(carbohydrates, proteins,
fats and nucleotides)
• determine the chemical
properties of large
compounds (polar vs
nonpolar, acid vs base)
• react with functional
groups on other
compounds
Acid/Base Concentration (pH)
• pH is the measurement on a scale ranging from 0 to
14 of H+ concentration in a solution
– H+ is the ionized form of a hydrogen atom
• the only electron has been removed, leaving a
single proton
• H+ = hydrogen ion = proton
• pH = -log[H+]
– [H+] = molar concentration of H+ in a solution
– the greater the [H+] the lower the pH, the lower the
[H+] the higher the pH
• Acidic solutions have higher [H+]
– a lower pH
• Alkaline (basic) solutions have lower [H+]
– a higher pH
pH Scale
• Acidic: pH 0 – 6.99
• Basic: pH 7.01 – 14
• Neutral: pH 7.00
Acids and Bases
• Acids are molecules that are capable of increasing the
number of H+ in a solution
– called proton donors
– decrease the pH of a solution
• Bases are molecules that are capable of decreasing
the number of H+ in a solution are
– called proton acceptors
– increase the pH of a solution
Buffers
• Substances that are capable of resisting large
changes in the pH of a solution
– allow pH to remain relatively constant
• buffers in the body allow the body pH to remain
at 7.4 (slightly basic)
Water as a Solvent
• water molecules overpower the ionic bond above
between Na+Cl- by forming hydration spheres
• note orientation of water molecules: negative pole
faces Na+, positive pole faces Cl-
Liquid Mixtures
Substances that are
physically blended but
not chemically
combined
Water is the most
abundant compound in
biological mixtures
• Solutions
• Colloids
• Suspensions
Solutions
• Consists of:
– Solute (less abundant
substance in mixture)
which is dissolved in
the solvent
– Solvent (more
abundant substance
in mixture)
• always water in the
body
• Transparent
• e.g. copper sulfate
solution
Concentration of Solutions
Measurement of the amount of solute(s) in a solution
• Weight per Volume
– weight of solute in a given volume of solution
• 8.5 grams of dextrose in 1 liter of water = 8.5 g/L
• Percentages
– weight or volume of solute in solution
• 2 grams of dextrose + 8 grams of H2O is a
solution containing a total of 10 grams
• 2 of the 10 grams is dextrose, therefore this
solution is a 20% dextrose solution
Concentration of Solutions
• Molarity
– number of moles of solute/liter in solution
– based on molecular weight (MW)
• the addition of the atomic masses of a molecule
• H2O = H (1 amu) x 2 + O (16 amu) = 18 amu MW
• for a known MW, weigh out that many grams, this
gives you its gram molecular weight or 1 mole
• 1 mole always contains the same number of
molecules (6.02 × 1023 = Avogadro’s number)
• MW of glucose is 180, so one mole of glucose is
180g, a one molar solution of glucose contains
180g/L
Percentage vs. Molar Concentrations
• Percentage
– # of molecules
unequal
– weight of solute
equal
• Molar
– # of molecules equal
– weight of solute
unequal
Concentration of Solutions
• Osmolarity
– total molarity of all solute particles in a solution
– used when measuring concentrations of solutions
that contain multiple solutes (such as a dissociated
salt)
• 2 moles of NaCl + 3 moles of glucose in 1L of
water = 7 osmolar solution
–in water, 2 moles of NaCl = 2 moles of Na+
and 2 moles of Cl- (4 moles together)
Suspensions
• The solute is not
dissolved in the solvent,
but rather is suspended
within the solvent
• The solute settles to the
bottom of the container,
separating from the
solvent
• Cloudy or opaque
• Separate on standing
• e.g. blood
Colloids
• The solute is not
dissolved in the solvent,
but rather is suspended
within the solvent
• The solute does not
settle to the bottom of
the container,
separating from the
solvent
• Cloudy
• e.g. milk