1. Lecture 2 Atoms and Molecules

2.  Every atom has the same basic structure  Core nucleus of protons and neutrons  Orbiting cloud of electrons Atoms  Matter is any substance in the universe that has mass and occupies space  All matter is composed of extremely small particles called atoms

3. Atoms  Atomic number  Number of protons  Atomic mass  Number of protons and neutrons  Element  A substance that cannot be broken down by ordinary chemical means

4. Isotopes  Isotopes are atoms with the same number of protons but different numbers of neutrons Different atomic mass Same atomic number 99% of all carbon

5. Radioactive isotope dating Radioactive Decay  The nucleus of an unstable isotope breaks down into particles with lower atomic numbers  Radioactive isotopes are used in  1. Medicine  Tracers are taken up and used by the body  Emissions are detected using special lab equipment  2. Dating fossils  The rate of decay of a radioactive element is constant  The amount of decay can be used to date fossils

6. Energy  The capacity to do work (put matter into motion)  Types of energy  Kinetic – energy in action  Potential – energy of position; stored (inactive) energy Energy Concepts PLAY

7.  Electrons have energy due to their relative orbital position (potential energy) Electrons  Electron shells, or energy levels, surround the nucleus of an atom  Valence shell – outermost energy level containing chemically active electrons  Bonds are formed using the electrons in the outermost energy level

8. The Octet Rule  Inert elements have their outermost energy level fully occupied by electrons  Reactive elements do not have their outermost energy level fully occupied by 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

9. Molecules  A molecule is a group of atoms held together by energy  The holding force is called a chemical bond  There are three kinds of chemical bonds 1.Ionic bonds 2.Covalent bonds 3.Hydrogen bonds

10. Ionic Bonds  Ionic bonds form between atoms by the transfer of one or more electrons  Ionic compounds form crystals instead of individual molecules  Example: NaCl (sodium chloride)  Two key properties 1.Strong: But not as strong as covalent bonds 2.Not directional: They are not formed between particular ions in the compound

11. Ionic Bonds

12. Covalent Bonds Water molecules contain two covalent bonds  Covalent bonds are formed by the sharing of two or more electrons  Electron sharing produces molecules  Two key properties 1.Strong: The strength increases with the number of shared electrons 2.Very directional: They are formed between two specific atoms

13. Comparison of Bonds  Electrons shared equally between atoms produce nonpolar molecules  Electrons shared unequally produces polar molecules  Atoms with six or seven valence shell electrons are electronegative  Atoms with one or two valence shell electrons are electropositive

14. Hydrogen Bonds Hydrogen bonding in water molecules  Formed by the attraction of opposite partial electric charges between two polar molecules  Too weak to bind atoms together  Common in dipoles such as water  Responsible for surface tension in water  Important as intramolecular bonds, giving the molecule a three-dimensional shape

15. Chemical Reactions  Occur when chemical bonds are formed, rearranged, or broken  Are 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

16. 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 and broken AB + C  AC + B  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

17. Oxidation-Reduction (Redox) Reactions  Reactants losing electrons are electron donors and are oxidized  Reactants taking up electrons are electron acceptors and become reduced  Generally, the atom that is reduced contains the most energy

18. Energy Flow in Chemical Reactions  Exergonic reactions – reactions that release energy  Endergonic reactions – reactions whose products contain more potential energy than did its reactants

19. 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

20. Hydrogen Bonds Give Water Unique Properties  Water molecules are polar molecules  They can thus form hydrogen bonds with each other and with other polar molecules  Each hydrogen bond is very weak  However, the cumulative effect of enormous numbers can make them quite strong  Hydrogen bonding is responsible for many of the physical properties of water  Heat Storage  A large input of thermal energy is required to disrupt the organization of liquid water  This minimizes temperature changes  Ice Formation  At low temperatures, hydrogen bonds don’t break  Water forms a regular crystal structure that floats  High Heat of Vaporization  At high temperatures, hydrogen bonds do break  Water is changed into vapor Water Transport PLAY

21. Hydrogen Bonds Give Water Unique Properties  Cohesion  Attraction of water molecules to other water molecules  Example: Surface tension  Adhesion  Attraction of water molecules to other polar molecules  Example: Capillary action Water strider

22. Hydrogen Bonds Give Water Unique Properties  High Polarity  Polar molecules are termed hydrophilic  Water-loving  All polar molecules that dissolve in water are termed soluble  Nonpolar molecules are termed hydrophobic  Water-fearing  These do not form hydrogen bonds and are therefore not water soluble

23. Water Ionizes  Covalent bonds within a water molecule sometimes break spontaneously H2OH2O + OH – hydroxide ion H+H+ hydrogen ion  This process of spontaneous ion formation is called ionization  It is not common because of the strength of covalent bonds

24.  A convenient way to express the hydrogen ion concentration of a solution pH = log [H + ] _  The pH scale is logarithmic  A difference of one unit represents a ten-fold change in H + concentration  Acid: Dissociates in water to increase H + concentration  Base: Combines with H + when dissolved in water

25.  Hydrogen ion reservoirs that take up or release H + as needed  The key buffer in blood is an acid-base pair Buffers