Transcript Bio_130_files/Chemistry Review

Chemistry Review
Atoms and Elements
• All types of matter (solids, liquids and gases) are composed of atoms.
• A substance that is composed of only one type of atom is called an
element.
– Elements are the simplest form of matter with unique chemical
properties. They are charted on the periodic table based on some of
their chemical characteristics.
• There are 24 major elements that have various roles in the body.
– These include structural, enzymatic, and homeostatic balance.
• Compounds, like water, are formed by combining the atoms of different
elements together.
• Atoms may create various types of chemical bonds. 3 types of bonds
include:
– Ionic
– Covalent
– Hydrogen
Major Elements of the Human Body
• Oxygen (O): Required for energy production during
cellular respiration
• Carbon (C): Organic back bone for fats,
carbohydrates, amino acids and nucleic acids.
• Hydrogen (H): Vital for energy( ATP) production
• Nitrogen (N): Most abundant in atmosphere,
Characteristic element of protein
• Phosphorus (P): found in
– DNA: Blue print for life
– RNA : Vital for protein production
– ATP : Cellular energy
The Atom
• Atoms – identical building blocks for each element
• Atomic symbol – one- or two-letter chemical short
hand for each element
• Atomic number : # of protons in nucleus
– periodic table
• elements arranged by atomic number
• Atomic weight – equal to the mass of the protons
and neutrons
• Isotope – atoms with same number of protons but a
different number of neutrons
Atomic Structure
• The nucleus consists of neutrons and protons
– 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
– valence electrons are located in the outermost shell
• interact with other atoms to form bonds
– Electrons – have a negative charge and 1/2000 the mass of
a proton (0 amu)
Chemical Bonds
• Electron shells, or energy levels, surround the
nucleus of an atom.
• Bonds are formed using the electrons in the
outermost energy level
– Valence shell – outermost energy level containing
chemically active electrons
• 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.
Planetary Models of Elements
p+ represents protons, no represents neutrons
Chemically Reactive Elements
• Reactive elements do not have their outermost energy level
fully occupied by electrons therefore are able to interact with
other elements
Chemically Inert Elements
• Inert elements have their outermost energy level fully occupied by
electrons therefore don’t interact with other elements.
Types of Chemical Bonds
• Ionic:
• Covalent
• Hydrogen
Formation of an Ionic Bond
• Ionic bonds form between atoms by the transfer of one
or more electrons
• Ionic compounds form crystals instead of individual
molecules
• Example: Na+Cl-(sodium chloride)
Formation of an Ionic Bond
• A valance electron from Na is transferred to Cl
• Cl now has 18e and 17p resulting in a – charge
• Na has 10e and 11P resulting in a + charge.
Ion Formation
• Ions are charged atoms resulting from the gain or loss of
electrons.
– Anions have gained one or more electrons therefore
are negatively charged(-)
– Cations have lost one or more electrons giving them a
positive charge(+)
• Typically occur between elements on opposite sides of
the periodic table.
Electronegativity and Bond Formation
• Elements on opposite ends of the periodic tables have a greater
electronegative gradient. Ionic bonds result.
• Elements that are closer to each other have smaller
electronegative gradient thus form covalent bonds.
Covalent bonds
• Covalent bonds are formed by the sharing of two
or more electrons.
– Covalent bonds are classified as Polar or Nonpolar.
• When two atoms with similar electronegativities
they share their valance electrons.
– Nonpolar( neutral charge) bond results.
• CO2, O2, N2
• If there is a larger electronegative gradient
between the atoms.
– a polar covalent bond (charged compound) results.
• H 2O
Nonpolar Bonds
• Electrons shared equally between atoms produce nonpolar
bonds.
• The negative charged electrons are spaced evenly between
the 2 atoms resulting in a neutral charge.
Covalent Bonds
Double Covalent Bonds
Polar Covalent Bonds
• Uneven sharing of electrons produces polar bonds
• One atom has a greater electronegativity.
– This atom will have stronger pull on the shared
electrons
– The shared electrons spend more time closer to
the nucleus of electronegative atom.
– The addition of the shared electrons makes the
electronegative atom partially negative charged,
while the atom with a lower electronegativity
becomes partially positively charged
– Polar bonds occur between an electronegative
atom (mostly O or N)
• ex. H2O
Inorganic compounds
• Do not contain carbon
• Water, salts, and many acids and bases
• Minerals such as magnesium and calcium.
• Salts : NA+CL–
– contain cations other than H+ and anions other than
OH–
• Are electrolytes; they conduct electrical currents
and function in various metabolic reactions.
– Electrical activity of the nervous system
– Vital for bone formation
Acid-Base Concentration (pH)
• pH scale ranges from 0 to 14.
• Acidic solutions have higher [H+] and a lower pH.
– Considered proton donors
– pH less than 7
• Alkaline (basic) solutions have lower [H+] and a
higher pH
– Considered proton acceptors
– pH greater than 7
• Neutral solutions have equal H+ and OH–
concentrations
– pH = 7
pH Scale
• Acids release H+ and are therefore proton donors
HCl  H+ + Cl –
• Bases release OH– and are proton acceptors
NaOH  Na+ + OH–
Buffers
• The body has many mechanisms devoted
to resist abrupt and large swings in the pH
of body fluids.
• These systems allow pH to remain
relatively constant .
– Approximately 7.4 (slightly basic)
– Maintaining a stable pH is critical for
creating an environment necessary for
metabolic reactions.
Chemical Reactions
• Chemical reactions in the body act by forming,
breaking or rearranging bonds.
• Chemical equations contain:
– Relative amounts of reactants (starting
chemicals) and products (finishing chemicals)
– Number and type of reacting substances, and
products produced
Synthesis and Decomposition Reactions
Oxidation-Reduction (Redox) Reactions
• Reactants losing electrons are electron donors
and are oxidized
• Reactants taking up electrons are electron
acceptors and become reduced
• Na + Cl → Na+ + Cl– Na is oxidized and Cl has been reduced
• LEO THE LION SAYS GER
Forms of Energy
• Chemical – stored in the bonds of chemical substances
– Energy from food
• Electrical – results from the movement of charged
particles
– Household Appliances run on it
• Mechanical – directly involved in moving matter
– Machines such as cranes or bull dozers
• Radiant or electromagnetic – energy traveling in waves
– visible light, ultraviolet light, and X rays
First Law of Thermodynamics
– Energy cannot be created or destroyed, but only
change form.
• During each conversion, some of the energy
dissipates into the environment as heat.
–Heat is defined as the measure of the
random motion of molecules.
– the second law states that "energy systems have a tendency
to increase their entropy"
Energy
• The capacity to do work (put matter into
motion)
• Types of energy
– Kinetic – energy in action. Ball rolling down a hill.
– Potential – energy of position; stored (inactive)
energy. Ball sitting on top of hill.
Fig. 8.2 (TEArt)
Potential energy
Energy - the capacity to do work
–kinetic - energy of motion
–potential - stored energy
•
Kinetic energy
Factors Influencing Rate of Chemical
Reactions
• 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
• Catalysts – increase the rate of a reaction
without being chemically changed
• Enzymes – biological catalysts