1. The Chemical Context of Life Chapter 2 Biology – Campbell Reece

2. Chemical Elements and Compounds  Matter consists of chemical elements in pure form and in combinations called compounds What is matter? What is matter?  Element – a substance that cannot be broken down to other substances by chemical reactions Ex. Oxygen (O), Nitrogen (N), Sodium (Na) Ex. Oxygen (O), Nitrogen (N), Sodium (Na)  Compound – a substance consisting of two or more different elements combined in a fixed ratio Ex. NaCl, H 2 O Ex. NaCl, H 2 O

3. Figure 2.2 The emergent properties of a compound + →

4. Table 2.1 Naturally Occurring Elements in the Human Body  25 of the 92 natural elements are known to be essential for life  C, H, O, N make up 96% of living matter  P, S, Ca, K and a few others make up the other 4%  Trace elements – those required in minute quantities

5. Figure 2.4 Goiter

6. Atoms  Atom – smallest unit of matter that retains the properties of an element Neutrons – neutral charge (n 0 ) Neutrons – neutral charge (n 0 ) Protons – positive charge (p+) Protons – positive charge (p+) Neutrons and protons are found in the nucleusNeutrons and protons are found in the nucleus Electrons – negative charge (e-) Electrons – negative charge (e-) Move in a cloud around the nucleusMove in a cloud around the nucleus

7. Atomic Number & Atomic Mass  Neutrons and protons have a mass of 1 dalton (same as atomic mass unit)  Atomic number – number of protons 2 He 2 He  Mass number – sum of the protons and neutrons 4 He 4 He  Atomic mass – mass of an atom (same as the mass number)

8. Isotopes  All atoms of an element have the same number of protons  Isotopes – have a different number of neutrons (and therefore weigh more) Carbon-12, Carbon-13, Carbon-14 Carbon-12, Carbon-13, Carbon-14  Radioactive isotopes – one in which the nucleus decays spontaneously, giving off particles and energy Used in dating fossils, as tracers in medicine Used in dating fossils, as tracers in medicine

9. Figure 2.7 A PET scan, a medical use for radioactive isotopes

10. Energy Levels of Electrons  An atom’s electrons vary in the amount of energy they possess  The different stages of energy that electrons have are called energy levels, or electron shells Electrons in the shell closest to the nucleus have the lowest energy

11. Electron Configuration  The chemical behavior of an atom is determined by its electron configuration  Valence electrons – the electrons in the outermost shell (valence shell)  An element is stable when its valence shell is full

12. Electron Orbitals  Orbital – the three-dimensional space where an electron is found 90% of the time Each orbital can hold 2 e- Each orbital can hold 2 e-  1s orbital is spherical in shape (1) 1 st shell holds 2 e- 1 st shell holds 2 e-  2s orbital is also spherical (1)  2p orbitals are dumbbell-shaped (3) 2 nd shell holds 8 e- 2 nd shell holds 8 e-

13. Figure 2.11 Electron orbitals

14. Covalent Bonds  The sharing of a pair of valence electrons by two atoms Single covalent bond – one pair is shared Single covalent bond – one pair is shared Double covalent bond – two pairs are shared Double covalent bond – two pairs are shared  Molecule – two or more atoms held together by covalent bonds Structural formula – H-H Structural formula – H-H Lewis dot structure – H:H Lewis dot structure – H:H Molecular formula – H 2 Molecular formula – H 2  Compound – a combination of two or more different elements

16. Nonpolar & Polar Covalent Bonds  The attraction of an atom for the electrons of a covalent bond is called its electronegativity  Nonpolar covalent bond – the electrons are shared equally Ex. O 2 Ex. O 2  Polar covalent bond – the electrons are not shared equally Ex. H 2 O Ex. H 2 O

17. Figure 2.13 Polar covalent bonds in a water molecule

18. Ionic Bonds  Sometimes two atoms are so unequal in their electronegativity that the more electronegative atoms takes the electron forming an ion  Cation – a positive ion (Na + )  Anion – a negative ion (Cl - )  The attraction between the opposite charges forms an ionic bond  Compounds formed by ionic bonds are NOT molecules

19. Other Types of Bonds  Hydrogen bonds – form when a hydrogen atom covalently bonded to one electronegative atom is also attracted to another electronegative atom (weak)  Van der Waals Interactions – attractions formed by the ever- changing “hot spots” of positive and negative charges (as the electrons move around)

20. Figure 2.17 Molecular shapes due to hybrid orbitals Molecular shape is crucial in Biology because it determines how most biological molecules recognize and respond to one another.

21. Chemical Reactions  The making and breaking of chemical bonds, leading to changes in the composition of matter  Some reactions go to completion, but most are reversible 3H 2 + N 2 ↔ 2NH 3 3H 2 + N 2 ↔ 2NH 3 Chemical equilibrium – the relative concentrations of products and reactants stop changing Chemical equilibrium – the relative concentrations of products and reactants stop changing

22. Unnumbered Figure (Page 38) Chemical reaction between hydrogen and oxygen