1. Chapter 2 Atoms and Molecules: The Chemical Basis of Life

2. Inorganic vs. Organic  Inorganic compounds - simple substances that do not contain carbon Ex. water  Organic compounds – carbon- containing compounds that are large and complex Ex. Glucose

3. Organic compounds all contain: 1. Oxygen 2. Nitrogen 3. Carbon 4. Zinc

4. Elements  Elements – substances that cannot be broken down into simpler substances  4 elements responsible for more than 96% of the mass of most organisms: Oxygen Carbon Hydrogen Nitrogen  See table 2-1 in book p.24

5. Other elements that make up living organisms  Calcium  Phosphorus  Potassium  Magnesium  Sodium  Iron  Sulfur

6. What is the most abundant element in the human body? 1. Oxygen 2. Carbon 3. Hydrogen 4. Nitrogen

7. Chemistry Quick Review  Atom – the smallest portion of an element that retains its chemical properties  Subatomic particles: Electron – negative charge Proton – positive charge Neutron – uncharged  # of electrons = # of protons  Nucleus – protons and neutrons  Electrons – move rapidly through space around nucleus

8. Atomic Number  Each kind of element has a fixed number of protons in the atomic nucleus  Written as a subscript to the left of the chemical symbol Example: 8 O Oxygen nucleus contains 8 protons Determines the atom’s identity and defines the element

9. Which element has an atomic number of 7? 1. Hydrogen 2. Carbon 3. Nitrogen 4. Helium

10. The Periodic Table  Chart in which elements are arranged in order by atomic number  Can be used to determine electron configurations Bohr model – shows the electrons arranged in a series of concentric circles around the nucleus

11. Bohr Model What element is this?

12. Atomic Mass  Mass of protons + neutrons  Mass of electron = 1/1800 the mass of a proton or neutron  Atomic mass number is a superscript to the left of the symbol Example: 16 O

13. Isotopes  Atom with different number of neutrons (different masses)  Most elements mixture of isotopes Ex. Carbon-12, Carbon-14  Mass of element is average of the masses of its isotopes Atomic mass of Carbon = 12.011

14. Isotopes  Radioisotopes – unstable isotopes Tend to break down (decay) to a more stable isotope Emit radiation when they decay Ex. Carbon-14 decays to Nitrogen

15. Electrons move in orbitals  Orbitals are more like “electron clouds”  The farther away from the nucleus, the more energy the electrons have  Valence electrons – the most energetic electrons Occupy valence (outer) shell

16. Chemical Reactions  Valence electrons participate in chemical reactions  When valence shell is full, it is stable  When valence shell is not full, atoms tend to lose, gain, or share electrons

17. To be full, the first electron shell has how many electrons? 1. 1 2. 2 3. 4 4. 8 5. 18

18. Compounds and Molecules  Atoms combine to form compounds and molecules  Compounds - 2 or more different elements combined in a fixed ratio Ex. NaCl (table salt)  Molecules - 2 or more atoms combine chemically Ex. O 2, DNA

19. Are all molecules compounds? 1. Yes 2. No

20. Molecule or Compound? O 2 1. Molecule 2. Compound 3. Both

21. Chemical Formulas  Represents chemical composition  Simplest formula – most simple ratio Ex. NH 2  Molecular formula – actual numbers of each type of atom per molecule Ex. N 2 H 4  Structural formula – shows arrangement of atoms Ex. Water H – O – H

22. Chemical Equations  Reactants – participate in reaction  Products – formed by the reaction  Example – cellular respiration C 6 H 12 O 6 + 6O 2 -> 6CO 2 + 6H 2 O + Energy

23. Chemical Bonds  Valence electrons dictate # of bonds  2 types of chemical bonds: Covalent – atoms share electrons Ionic - attraction between positive cations and negative anions  Transfer electrons

24. Covalent Bonds  Ex. H 2 gas Each atom has 1 electron 2 electrons fill valence shell Both atoms attract the electrons (share) Valence shell is full w/ 2 electrons

25. Types of Covalent Bonds  Single covalent bond – 1 pair of electrons is shared  Double covalent bond – 2 pairs of electrons shared  Triple covalent bond – 3 pairs of electrons shared

26. Covalent Bonds  Electronegativity - measure of atom’s attraction for shared electrons in chemical bonds  Oxygen, Nitrogen, Fluorine, Chlorine very electronegative  Can be polar or nonpolar  Similar electronegativities = nonpolar bonds  Different electronegativities = polar bonds Electrons are pulled closer to the nucleus of the atom with the higher electronegativity

27. Polar Molecules  Molecule with one or more polar covalent bonds  One end with a partial positive charge and other end a partial negative charge  Ex. Water (p.31)

28. A water molecule is polar because 1. The electrons orbit the H atoms more closely 2. The electrons orbit the O atom more closely 3. The electrons orbit all atoms equally

29. Ionic Bonds  Ionic compound – consists of anions and cations bonded together  Ex. NaCl (p.31 & 32) Na – 1 valence electron Cl – 7 valence electrons Cl takes electron from Na to complete valence shell

30. Hydrogen Bonds  Weak attractions  Important in determining the 3-D structure of large molecules DNA Proteins

31. Why are hydrogen bonds essential to the function of DNA? 1. They keep the 2 strands tightly bonded together 2. They allow the 2 strands to separate for replication 3. They are strong bonds

32. Redox Reactions  Reaction that involves electron transfer  Cellular Respiration and Photosynthesis  Oxidation – atom/ion loses electron  Reduction – atom/ion gains electron

33. Water  70% of total body weight  Reactant/product in many chemical reactions  Solvent for most biological reactions Hydrophilic – react with water Hydrophobic – not disrupted/dissolved by water

34. Which of the following substances is hydrophobic? 1. Salt 2. Sugar 3. Oil

35. Properties of Water  Cohesive – water molecules stick to each other  Adhesive – water molecules stick to other substances  Capillary action – cohesion and adhesion working together Water will move against gravity in a narrow tube In plants, water moves from soil to roots

36. Properties of Water  Surface tension – water molecules crowd together at the surface strong layer  High specific heat Maintains a stable temperature  High heat of vaporization Much heat required to change to water vapor

37. Acids and Bases  Acids – proton donors Acid -> H + + Anion Acidic solutions have higher hydrogen ion concentration Turn blue litmus paper red Sour taste HCl – inorganic acid Acetic Acid – from vinegar, Lactic Acid – from sour milk (organic acids)

38. Acids and Bases  Bases – proton acceptors Base -> OH - + Cation Basic solutions have lower hydrogen ion concentration Turn red litmus paper blue Feel slippery to the touch Ex. Sodium Hydroxide, Ammonia – inorganic Purine and Pyrimidine – organic

39. pH scale  Logarithmic expression of the hydrogen ion concentration of solution  7 = neutral  Below 7 = Acid  Above 7 = Base