2
Chemistry Comes Alive: Part A
Matter
•Anything that has mass and occupies space
•States of matter:
1.Solid—definite shape and volume
2.Liquid—definite volume, changeable shape
3.Gas—changeable shape and volume
Energy
•Capacity to do work or put matter into motion
•Types of energy:
•Kinetic—energy in action
•Potential—stored (inactive) energy
Forms of Energy
•Chemical energy—stored in bonds of chemical substances
•Electrical energy—results from movement of charged particles
•Mechanical energy—directly involved in moving matter
•Radiant or electromagnetic energy—exhibits wavelike properties (i.e., visible light, ultraviolet light, and X-rays)
Energy Form Conversions
•Energy may be converted from one form to another
•Conversion is inefficient because some energy is “lost” as heat
Composition of Matter
•Elements
•Cannot be broken down by ordinary chemical means
•Each has unique properties:
•Physical properties
•Are detectable with our senses, or are measurable
•Chemical properties
•How atoms interact (bond) with one another
Composition of Matter
•Atoms
Unique building blocks for each element
•Atomic symbol: one- or two-letter chemical shorthand for each element
Major Elements of the Human Body
•Oxygen (O)
•Carbon (C)
•Hydrogen (H)
•Nitrogen (N)
Lesser Elements of the Human Body
• About 3.9% of body mass:
•Calcium (Ca), phosphorus (P), potassium (K), sulfur (S), sodium (Na), chlorine (Cl), magnesium (Mg), iodine (I), and iron (Fe)
Trace Elements of the Human Body
•< 0.01% of body mass:
•Part of enzymes, e.g., chromium (Cr), manganese (Mn), and zinc (Zn)
Atomic Structure
•Determined by numbers of subatomic particles
•Nucleus consists of neutrons and protons
Atomic Structure
•Neutrons
•No charge
•Mass = 1 atomic mass unit (amu)
•Protons
•Positive charge
•Mass = 1 amu
Atomic Structure
•Electrons
•Orbit nucleus
•Equal in number to protons in atom
•Negative charge
•1/2000 the mass of a proton (0 amu)
Models of the Atom
•Orbital model: current model used by chemists
•Depicts probable regions of greatest electron density (an electron cloud)
•Useful for predicting chemical behavior of atoms
Models of the Atom
•Planetary model—oversimplified, outdated model
•Incorrectly depicts fixed circular electron paths
•Useful for illustrations (as in the text)
Identifying Elements
•Atoms of different elements contain different numbers of subatomic particles
•Compare hydrogen, helium and lithium (next slide)
Identifying Elements
•Atomic number = number of protons in nucleus
Identifying Elements
•Mass number = mass of the protons and neutrons
•Mass numbers of atoms of an element are not all identical
•Isotopes are structural variations of elements that differ in the number of neutrons they contain
Identifying Elements
•Atomic weight = average of mass numbers of all isotopes
Radioisotopes
•Spontaneous decay (radioactivity)
•Similar chemistry to stable isotopes
•Can be detected with scanners
Radioisotopes
•Valuable tools for biological research and medicine
• Cause damage to living tissue:
• Useful against localized cancers
• Radon from uranium decay causes lung cancer
Molecules and Compounds
• Most atoms combine chemically with other atoms to form molecules and compounds
•Molecule—two or more atoms bonded together (e.g., H2 or C6H12O6)
•Compound—two or more different kinds of atoms bonded together (e.g., C6H12O6)
Mixtures
•Most matter exists as mixtures
•Two or more components physically intermixed
•Three types of mixtures
•Solutions
•Colloids
•Suspensions
Solutions
•Homogeneous mixtures
•Usually transparent, e.g., atmospheric air or seawater
•Solvent
•Present in greatest amount, usually a liquid
•Solute(s)
•Present in smaller amounts
Concentration of Solutions
•Expressed as
•Percent, or parts per 100 parts
•Milligrams per deciliter (mg/dl)
•Molarity, or moles per liter (M)
•1 mole = the atomic weight of an element or molecular weight (sum of atomic weights) of a compound in grams
•1 mole of any substance contains 6.02 1023 molecules (Avogadro’s number)
Colloids and Suspensions
•Colloids (emulsions)
•Heterogeneous translucent mixtures, e.g., cytosol
•Large solute particles that do not settle out
•Undergo sol-gel transformations
•Suspensions:
•Heterogeneous mixtures, e.g., blood
•Large visible solutes tend to settle out
Mixtures vs. Compounds
•Mixtures
•No chemical bonding between components
•Can be separated physically, such as by straining or filtering
•Heterogeneous or homogeneous
•Compounds
•Can be separated only by breaking bonds
•All are homogeneous
Chemical Bonds
•Electrons occupy up to seven electron shells (energy levels) around nucleus
•Octet rule: Except for the first shell which is full with two electrons, atoms interact in a manner to have eight electrons in their outermost energy level (valence shell)
Chemically Inert Elements
•Stable and unreactive
•Outermost energy level fully occupied or contains eight electrons
Chemically Reactive Elements
•Outermost energy level not fully occupied by electrons
•Tend to gain, lose, or share electrons (form bonds) with other atoms to achieve stability
Types of Chemical Bonds
•Ionic
•Covalent
•Hydrogen
Ionic Bonds
•Ions are formed by transfer of valence shell electrons between atoms
•Anions (– charge) have gained one or more electrons
•Cations (+ charge) have lost one or more electrons
•Attraction of opposite charges results in an ionic bond
Formation of an Ionic Bond
•Ionic compounds form crystals instead of individual molecules
•NaCl (sodium chloride)
Covalent Bonds
•Formed by sharing of two or more valence shell electrons
•Allows each atom to fill its valence shell at least part of the time
Covalent Bonds
•Sharing of electrons may be equal or unequal
•Equal sharing produces electrically balanced nonpolar molecules
•CO2
Covalent Bonds
•Unequal sharing by atoms with different electron-attracting abilities produces polar molecules
•H2O
•Atoms with six or seven valence shell electrons are electronegative, e.g., oxygen
•Atoms with one or two valence shell electrons are electropositive, e.g., sodium
Hydrogen Bonds
•Attractive force between electropositive hydrogen of one molecule and an electronegative atom of another molecule
•Common between dipoles such as water
•Also act as intramolecular bonds, holding a large molecule in a three-dimensional shape
Chemical Reactions
•Occur when chemical bonds are formed, rearranged, or broken
•Represented as chemical equations
•Chemical equations contain:
•Molecular formula for each reactant and product
•Relative amounts of reactants and products, which should balance
Examples of Chemical Equations
H + H H2 (hydrogen gas)
4H + C CH4 (methane)
Patterns of Chemical Reactions
• Synthesis (combination) reactions
• Decomposition reactions
• Exchange reactions
Synthesis Reactions
•A + B AB
•Always involve bond formation
• Anabolic
Decomposition Reactions
•AB A + B
•Reverse synthesis reactions
•Involve breaking of bonds
•Catabolic
Exchange Reactions
•AB + C AC + B
•Also called displacement reactions
•Bonds are both made and broken
Oxidation-Reduction (Redox) Reactions
•Decomposition reactions: Reactions in which fuel is broken down for energy
•Also called exchange reactions because electrons are exchanged or shared differently
•Electron donors lose electrons and are oxidized
•Electron acceptors receive electrons and become reduced
Chemical Reactions
•All chemical reactions are either exergonic or endergonic
•Exergonic reactions—release energy
•Catabolic reactions
•Endergonic reactions—products contain more potential energy than did reactants
•Anabolic reactions
Chemical Reactions
•All chemical reactions are theoretically reversible
•A + B AB
•AB A + B
•Chemical equilibrium occurs if neither a forward nor reverse reaction is dominant
• Many biological reactions are essentially irreversible due to
• Energy requirements
• Removal of products
Rate of Chemical Reactions
•Rate of reaction is influenced by:
• temperature rate
• particle size rate
• concentration of reactant rate
•Catalysts: rate without being chemically changed
•Enzymes are biological catalysts