CHEMISTRY Concepts and connections Ch. 2

1. CHEMISTRYConcepts and connectionsCh. 2

2. Atoms, Elements and molecules • ALL MATTER IS MADE UP OF ATOMS • ALL MATTER IS ORGANIZED INTO ELEMENTS (in the Periodic Table) • The fundamental building blocks of life are CHEMICALS • ELEMENTS: fundamental form of matter that have mass, occupies space and cannot be decomposed into other substance(elements are made up of atoms that are unique for that substance)

3. Atoms • ATOM: the smallest unit of an element that still retains the properties of that element • Living things are composed of about 25 chemical elements.

4. MOLECULE: • MOLECULE: two or more atoms bonded together in fixed proportions. Example: a water molecule. 2 hydrogen atoms bonded with an oxygen atom H2O • Two molecules : 2H2O • Cells function and do what they do as a result of chemical reactions

5. ELEMENTS FOUND IN ALL LIVING ORGANISMS: • Living organisms contain about 25 chemicals but the four found inALLare :CHON (carbon, hydrogen, oxygen and nitrogen) • These four make up over 96% of our bodies • Elements found in most life: CHONPS • Others important for life are Mg, Na, K, Cl, Ca, I, Fe • Elements found in very small amounts are “trace elements” and function as catalysts and co-enzymes • Ex: iodine (I) • Chemical symbols are universal. C= carbon H= hydrogen O= oxygen • N= nitrogen P=phosphorous S= sulfur Mg= magnesium Na= sodium K= potassium Ca=calcium Fe= iron

6. 0 • Carbon, hydrogen, oxygen, and nitrogen • Make up the bulk of living matter Table 2.1

7. 0 CONNECTION • Trace elements are common additives to food and water • Dietary deficiencies in trace elements • Can cause various physiological conditions Figure 2.2A

8. ATOMS • Atoms are made up of: PROTONS (p+) (positive charge) in the nucleus NEUTRONS (n) (no charge) in the nucleus ELECTRONS (e-) (negative charge) in the energy levels around the nucleus

9. LE 2-8 Helium 2He Hydrogen 1H 2 He 4.00 Atomic number Atomic mass Element symbol First shell Electron-shell diagram Lithium 3Li Beryllium 4Be Boron 5B Carbon 6C Nitrogen 7N Oxygen 8O Fluorine 9F Neon 10Ne Second shell Sodium 11Na Magnesium 12Mg Aluminum 12Al Silicon 14Si Phosphorus 15P Sulfur 16S Chlorine 17Cl Argon 18Ar Third shell

10. Electron Configuration and Chemical Properties • The chemical behavior of an atom is determined by the distribution of electrons in electron shells • The periodic table of the elements shows the electron distribution for each element

11. Atoms are electrically neutral • The number of protons(+) in the nucleus = the number of electrons(-) around it # of protons= # of electrons • Atomic number = the number of protons in the nucleus • Atomic mass= the number of protons + the number of neutrons added together

12. Subatomic Particles • Atoms are composed of subatomic particles • Relevant subatomic particles include: • Neutrons (no electrical charge) • Protons (positive charge) • Electrons (negative charge) • Neutrons and protons form the atomic nucleus • Electrons form a cloud around the nucleus

13. Cloud of negative charge (2 electrons) Electrons LE 2-4 Nucleus

14. Valence electrons are those in the outermost shell, or valence shell • The chemical behavior of an atom is mostly determined by the valence electrons

15. Third energy level (shell) LE 2-7b Energy absorbed Second energy level (shell) First energy level (shell) Energy lost Atomic nucleus

16. ISOTOPES • Isotopes are atoms of a given element that have the samenumber of protons and electrons but have a different number of neutrons. • They have the same atomic number but different mass number. Ex: Carbon (C) there are 6 protons and 6 electrons in carbon but it can have different # of neutrons such as 12C, 13C, 14C

17. 0 • Medical Diagnosis • Radioactive tracers are often used for diagnosis • In combination with sophisticated imaging instruments Figure 2.5A Figure 2.5B

18. LE 2-6 Cancerous throat tissue

19. Uses of Isotopes • In medicine and research as labels and tracers. For Diagnosis: can take pictures with PET scans and other devices. we can detect them with scintillating machines and follow their movements For treatment: we use radium and cobalt to destroy cancerous tumors • To date ancient objects: Since we know the half life of an element ( time it takes for half of it to decay to a stable form) we can calculate how long it has been around.

20. Outermost electron shell (can hold 8 electrons) First electron shell (can hold 2 electrons) Electron Hydrogen (H) Atomic number = 1 Oxygen (O) Atomic number = 8 Nitrogen (N) Atomic number = 7 Carbon (C) Atomic number = 6 0 • Electron arrangement determines the chemical properties of an atom • Electrons in an atom • Are arranged in shells, which may contain different numbers of electrons Figure 2.6

21. 0 • Atoms whose shells are not full • Tend to interact with other atoms and gain, lose, or share electrons • These interactions • Form chemical bonds

22. Chemical bonds in living organisms • The three major chemical bonds used by elements to build compounds are: • Ionic bonds • Covalent bonds • Hydrogen bonds

23. IONS • Atoms become IONS or ionized when they lose or gain electrons so the number of protons is different than the number of electrons. • When an atom gains electrons it has more –(negative ) charges. It becomes a negative ion • When it loses electrons now it has more positive charges. It becomes a positive ion

24. – + Transfer of electron – – Cl Na Cl Na ClChlorine atom NaSodium atom Cl–Chloride ion Na+Sodium ion Sodium chloride (NaCl) 0 • Ionic bonds are attractions between ions of opposite charge • When atoms gain or lose electrons • Charged atoms called ions are created Figure 2.7A

25. 0 • An electrical attraction between ions with opposite charges • Results in an ionic bond

26. Na+ Cl– 0 • Sodium and chloride ions • Bond to form sodium chloride, common table salt Figure 2.7B

27. 0 • Covalent bonds join atoms into molecules through electron sharing • In covalent bonds • Two atoms share one or more pairs of outer shell electrons, forming molecules

28. A molecule consists of two or more atoms held together by covalent bonds • A single covalent bond, or single bond, is the sharing of one pair of valence electrons • A double covalent bond, or double bond, is the sharing of two pairs of valence electrons • Covalent bonds can form between atoms of the same element or atoms of different elements Animation: Covalent Bonds

29. LE 2-11b Name (molecular formula) Electron- shell diagram Structural formula Space- filling model Oxygen (O2)

30. 0 • Molecules can be represented in many ways Table 2.8

31. 0 • Unequal electron sharing creates polar molecules • A molecule is nonpolar • When its covalently bonded atoms share electrons equally

32. (–) (–) O H H (+) (+) 0 • In a polar covalent bond • Electrons are shared unequally between atoms, creating a polar molecule Figure 2.9

33. 0 • Hydrogen bonds are weak bonds important in the chemistry of life • The charged regions on water molecules • Are attracted to the oppositely charged regions on nearby molecules

34. (–) Hydrogen bond (+) H O (–) (+) H (+) (–) (–) (+) 0 • This attraction forms weak bonds • Called hydrogen bonds Figure 2.10

35. WATER Life began in water and land organisms essentially carry their wet environment inside themselves. Water accounts for 75 to 85% of a cell’s weight, and most cells are surrounded by it. About 71% of Earth’s surface is covered by oceans and human bodies are about 66% water. If you weigh 128bl, 85 pounds of you is water. • Water plays a major role in many of life’s processes.

36. 0 WATER’S LIFE-SUPPORTING PROPERTIES • Hydrogen bonds make liquid water cohesive (molecules “stick” to each other) • Due to hydrogen bonding • Cohesiveness makes water molecules move from a plant’s roots to its leaves

37. 0 • Insects can walk on water due to surface tension • Created by cohesive water molecules Figure 2.11

38. 0 • Water’s hydrogen bonds moderate temperature • Water’s ability to store heat • Moderates body temperature and climate

39. Hydrogen bond Liquid water Hydrogen bondsconstantly break and re-form Ice Hydrogen bonds are stable 0 • Ice is less dense than liquidwater • Hydrogen bonds hold molecules in ice • Farther apart than in liquid water Figure 2.13

40. 0 • Ice is therefore less dense than liquid water • Which causes it to float • Floating ice • Prevents lakes and oceans from freezing solid. The water under the ice remains warmer • protects the living things beneath the ice

41. Na+ – Na+ – + Cl– + – – + – + Cl– – + – + – + – – Ion insolution Saltcrystal 0 • Water is the solvent of life • Polar or charged solutes dissolve when water molecules surround them, forming aqueous solutions Figure 2.14

42. WATER PROPERTIES • Water is the best solvent for living things. It can dissolve more substances in greater mounts than any other liquid. Everything that goes In and out of cells is dissolved in H2O This is because of its polarity and hydrogen bonds. • Water is a polar molecule (slightly negative at the oxygen end and slightly positive at the hydrogen end). Polar molecules are hydrophilic (water loving) A hydrophobic substance is one that dislikes water, it is nonpolar .Ex: oils • Because water forms hydrogen bonds with other water molecules ,they are spaced apart when frozen, so ice is lessdense than liquid water and it floats on it. • Ice on the surface insulates the water beneath it, keeping it warmer.

43. Water Properties (continued….) • Water has temperature stabilizing effect. This is because it can absorb a lot of heat before its temperature changes. Heat increases the molecular motion so much that hydrogen bonds are broken and water molecules escape into the air. This is evaporation, cooling the surface. • Water’s cohesion creates surface tension. Cohesion is the capacity to resist rupturing. Water “sticks together” as a result of its hydrogen bonds. • Cohesion is very important for plants to pull water from the soil.

44. 0 • The chemistry of life is sensitive to acidic and basic conditions • A compound that releases H+ ions in solution is an acid • And one that accepts H+ ions in solution is a base • Acidity is measured on the pH scale • From 0 (most acidic) to 14 (most basic) • Neutral pH is 7

45. Acids, Bases and Buffers: The pH scale PH is the measure of the H+ (hydrogen ions) concentration in a solution. The greater the H+ conc. The lower the pH. • The scale goes from 0( acidic) to 7 (neutral) to 14 (basic) • What is an acid? A substance that releases hydrogen ions in solution is an acid. Ex: HCl (hydrochloric acid) • What is a base? Substances that release OH-(hydroxyl ions) in solution are called bases and are alkaline.

46. pH scale 0 1 H+ H+ H+ OH– H+ 2 Lemon juice, gastric juice OH– H+ H+ H+ H+ Increasingly ACIDIC(Higher concentration of H+) 3 Grapefruit juice, soft drink Acidic solution 4 Tomato juice 5 6 Human urine OH– OH– NEUTRAL[H+]=[OH–] OH– 7 Pure waterHuman blood H+ H+ OH– OH– H+ H+ 8 H+ Seawater Neutral solution 9 10 Increasingly BASIC(Lower concentration of H+) Milk of magnesia 11 Household ammonia OH– OH– 12 OH– OH– H+ Household bleach OH– OH– 13 OH– H+ Oven cleaner 14 Basic solution 0 The pH scale Figure 2.15

47. 0 • The pH of most cells • Is kept close to 7 (neutral) by buffers • Buffers are substances that resist pH change

48. What are buffers? • Buffers are substances that the body uses to prevent big shifts in pH. • Either will release H+ ions or combine with H+ ions to keep the balance. Bicarbonate is an example of one of the body’s important buffers

49. 0 CONNECTION • Acid precipitation threatens the environment • Some ecosystems are threatened by acid precipitation • Acid precipitation is formed when air pollutants from burning fossil fuels • Combine with water vapor in the air to form sulfuric and nitric acids • These acids • Can kill trees and damage buildings

50. O2 2 H2 2 H2O 0 CHEMICAL REACTIONS • Chemical reactions change the composition of matter • In a chemical reaction • Reactants interact, atoms rearrange, and products result Figure 2.17A