1 Chapter 2A The Chemical Context of Life Aim: Why is chemistry so important in the living world? Do Now: Describe a chemical reaction that occurs in the body? Homework: Study for regents level exam on Friday.

2 Matter Takes up space and has mass Exists as elements (pure form) and in chemical combinations called compounds

3 Elements Can ’ t be broken down into simpler substances by chemical reaction Composed of atoms: smallest unit of element Essential elements in living things include carbon C, hydrogen H, oxygen O, and nitrogen N making up 96% of an organism

4 Other Elements A few other elements Make up the remaining 4% of living matter Table 2.1

5 Deficiencies (a) Nitrogen deficiency (b) Iodine deficiency (Goiter) If there is a deficiency of an essential element, disease results Figure 2.3

6 Trace Elements Trace elements Are required by an organism in only minute quantities Ex: Fe, Zn

7 Compounds SodiumChloride Sodium Chloride + Are substances consisting of two or more elements combined in a fixed ratio Have characteristics different from those of their elements Figure 2.2

8 Properties of an Element… Depend on the structure of the atoms An atom is the smallest unit of matter that retains the properties of an element

9 Subatomic Particles Neutrons, which have no electrical charge Protons (+), which are positively charged Electrons (-), which are negatively charged

10 Subatomic Particle Location Protons and neutrons –Are found in the atomic nucleus Electrons –Surround the nucleus in a “ cloud ”

11 Simplified models of an Atom Nucleus (a) (b) In this even more simplified model, the electrons are shown as two small blue spheres on a circle around the nucleus. Cloud of negative charge (2 electrons) Electrons This model represents the electrons as a cloud of negative charge, as if we had taken many snapshots of the 2 electrons over time, with each dot representing an electron ‘ s position at one point in time. Figure 2.4

12 Atomic Number & Atomic Mass Atoms of the various elements differ in their number of subatomic particles The number of protons in the nucleus = atomic number The number of protons + neutrons = atomic mass Neutral atoms have equal numbers of protons & electrons (+ and – charges)

13 Atomic Number Is unique to each element and is used to arrange atoms on the Periodic table Carbon = 12 Oxygen = 16 Hydrogen = 1 Nitrogen = 17

14 Isotopes Different forms of the same element Have the same number of protons, but different number of neutrons May be radioactive spontaneously giving off particles and energy May be used to date fossils or as medical tracers

15 Other uses –Can be used in medicine to find and help treat tumors Cancerous throat tissue Figure 2.6

16 Energy Levels of Electrons An atom ’ s electrons Vary in the amount of energy they possess Electrons further from the nucleus have more energy Electron ’ s can absorb energy and become “ excited ” Excited electrons gain energy and move to higher energy levels or lose energy and move to lower levels

17 Energy –Is defined as the capacity to cause change Potential energy - Stored energy Kinetic Energy - Is the energy of motion

18 Why do some elements react? Valence electrons –Are those in the outermost, or valence shell –Determine the chemical behavior of an atom

19 Aim: How can we identify the different types of chemical bonding that exists? Do Now: Are chemical bonds important for life? Why? Homework: Study for regents level exam Friday.

20 Covalent Bonds Figure 2.10 Sharing of a pair of valence electrons Examples: H 2 Hydrogen atoms (2 H) Hydrogen molecule (H 2 ) In each hydrogen atom, the single electron is held in its orbital by its attraction to the proton in the nucleus. 1 When two hydrogen atoms approach each other, the electron of each atom is also attracted to the proton in the other nucleus. 2 The two electrons become shared in a covalent bond, forming an H 2 molecule. 3

21 Covalent Bonding A molecule –Consists of two or more atoms held together by covalent bonds A single bond –Is the sharing of one pair of valence electrons A double bond –Is the sharing of two pairs of valence electrons

22 Multiple Covalent Bonds (a) (b) Name (molecular formula) Electron- shell diagram Structural formula Space- filling model Hydrogen (H 2 ). Two hydrogen atoms can form a single bond. Oxygen (O 2 ). Two oxygen atoms share two pairs of electrons to form a double bond. HH O O Figure 2.11 A, B

23 Compounds & Covalent Bonds Name (molecular formula) Electron- shell diagram Structural formula Space- filling model (c) Methane (CH 4 ). Four hydrogen atoms can satisfy the valence of one carbon atom, forming methane. Water (H 2 O). Two hydrogen atoms and one oxygen atom are joined by covalent bonds to produce a molecule of water. (d) H O H HH H H C Figure 2.11 C, D

24 Covalent Bonding Electronegativity –Is the attraction of a particular kind of atom for the electrons in a covalent bond The more electronegative an atom –The more strongly it pulls shared electrons toward itself

25 Covalent Bonding In a nonpolar covalent bond –The atoms have similar electronegativities –Share the electron equally

26 Figure 2.12 This results in a partial negative charge on the oxygen and a partial positive charge on the hydrogens. H2OH2O –– O H H ++ ++ Because oxygen (O) is more electronegative than hydrogen (H), shared electrons are pulled more toward oxygen. In a polar covalent bond –The atoms have differing electronegativities –Share the electrons unequally Covalent Bonding

27 Ionic Bonds In some cases, atoms strip electrons away from their bonding partners Electron transfer between two atoms creates ions Ions –Are atoms with more or fewer electrons than usual –Are charged atoms

28 Ions An anion –Is negatively charged ions A cation –Is positively charged

29 Ionic Bonding Cl – Chloride ion (an anion) – The lone valence electron of a sodium atom is transferred to join the 7 valence electrons of a chlorine atom. 1 Each resulting ion has a completed valence shell. An ionic bond can form between the oppositely charged ions. 2 Na Cl + Na Sodium atom (an uncharged atom) Cl Chlorine atom (an uncharged atom) Na + Sodium on (a cation) Sodium chloride (NaCl) Figure 2.13 An ionic bond –Is an attraction between anions and cations

30 Ionic Substances Na + Cl – Figure 2.14 Ionic compounds –Are often called salts, which may form crystals

31 Weak Chemical Bonds Several types of weak chemical bonds are important in living systems: –Reinforce the shapes of large molecules –Help molecules adhere to each other

32 Hydrogen Bonds  – –  + +  + + Water (H 2 O) Ammonia (NH 3 ) O H H  + +  – – N H H H A hydrogen bond results from the attraction between the partial positive charge on the hydrogen atom of water and the partial negative charge on the nitrogen atom of ammonia. ++ ++ Figure 2.15 A hydrogen bond –Forms when a hydrogen atom covalently bonded to one electronegative atom is also attracted to another electronegative atom

33 Molecular Shape and Function Structure determines Function! The precise shape of a molecule –Is usually very important to its function in the living cell –Determines how biological molecules recognize and respond to one another with specificity

34 Morphine Carbon Hydrogen Nitrogen Sulfur Oxygen Natural endorphin (a) Structures of endorphin and morphine. The boxed portion of the endorphin molecule (left) binds to receptor molecules on target cells in the brain. The boxed portion of the morphine molecule is a close match. (b) Binding to endorphin receptors. Endorphin receptors on the surface of a brain cell recognize and can bind to both endorphin and morphine. Natural endorphin Endorphin receptors Morphine Brain cell Figure 2.17

35 Chemical Reactions Chemical reactions make and break chemical bonds A Chemical reaction –Is the making and breaking of chemical bonds –Leads to changes in the composition of matter

36 Chemical Reactions Reactants Reaction Products 2 H 2 O2O2 2H2O2H2O + + Chemical reactions –Convert reactants to products

37 Chemical Reactions Photosynthesis –Is an example of a chemical reaction Figure 2.18

38 Chemical Reactions Chemical equilibrium –Is reached when the forward and reverse reaction rates are equal