Unit 3

1. Unit 3 • The Chemistry of Life • Unit Objective: • To identify the carbon-based organic molecules of life and understand the roles they play in life processes. Only copy down underlined material. Everything else is there for support.

2. What’s Important About the Nutrition Label? • What does the label tell you? • Why is this important?

3. What Are Organisms Made Of? • Every organism is different depending upon it’s role or needs but there are basic molecules that are universal.

4. The Basic Compounds of Life. • Regardless of the amount, most organisms use these elements to form the macromolecules of life. • Fats • Proteins • Sugars • Nucleic Acids. • These are the fundamental molecules of life that you will be required to know. • These carbon-based molecules are known as Organic Compounds. • Another molecule you will learn about in this mini-unit is water.

5. Essential Questions Objectives • How does hydrogen, carbon, and oxygen combine to form molecules that participate in living systems? • What do these molecules do to support life? • To review/identify what makes up matter. • To know the elements that form most organic molecules. • To review/understand why atoms form bonds. • To understand how and why these bonds are essential for life by forming the carbon-based organic molecules of life. • This information will help you soon be able to explain how these molecules participate in important processes in living things.

6. Vocabulary • Atoms • Element • Valence electrons • Compound • Molecule • Ion

7. Before You Learn About the Organic Macromolecules… • You must understand that all bigger molecules are built of smaller pieces. • Everything of atoms… • Sometimes arranged into basic molecules. • These smaller pieces are called… • These will sometimes be called… • OR • What is a subunit? • Something that is combined with other subunits to build something bigger. BUILDING BLOCKS SUBUNITS MONOMERS

8. Atoms • Every living and nonliving thing is made of matter. • Matter is anything that has mass and takes up space. • All matter is made of very small particles called atoms. • An atomis the smallest unit of matter that cannot be broken down by chemical means. • These are the most basic subunit of matter & life.

9. Atoms • The atom is composed of three main types of smaller particles. • Protons: positively charged particles. • Neutrons: particles with no charge. • Electrons: negatively charged particles.

10. Atoms • The particles are in two specific areas: • The nucleus. • The electron cloud. • Protons and neutrons are in the nucleus. • Electrons are in various energy levels contained in the electron cloud around the nucleus.

11. B What are the parts of an atom? C A D

12. COLUMN How Is an Atom Built? • Find carbon on the periodic table. • This is the information for the element CARBON • Notice its row and column. • Every box represents a different element. • An element is any quantity of a substance that is 100% the same type of atom. • Being the same type means they have the same # of protons. • Ex: Diamonds are always the element carbon, regardless of the size, because every atom has 6 protons. ROW

13. CHNOPS: What are the Most Abundant Elements in Biology? • The six highlighted below are the most abundant elements in the bodies of most biological organisms. • Mark these in your Periodic Table.

14. Reading the Periodic Table • Atomic Number • Symbol • Name • Atomic Mass

15. Atomic Numbers, Mass, Electrons… • The boxes give a lot of information. • How do you know how many protons, electrons, and neutrons an element has? • It’s all based upon the atomic number found in the periodic table. • Atomic # = # protons • # electrons = # protons • # neutrons = atomic mass (whole) - # protons.

16. IONS: What Happens When an Atom Gains or loses an Electron. • Ions are atoms that gain or lose electrons. • This results in unequal numbers of p+ & e-. • These particles are now considered charged, or ions. • Why? Some elements steal electrons.

17. + - Na Cl Formation of Ions: Count the p+ & e- before & after the exchange… Anion: A negatively charged ion. Cation: A positively charged ion.

18. Special Circumstance: Isotopes Elements • Recall, elements are all the same type of atom because every atom has the same number of protons. • For example, every atom for any amount of the element carbon has six protons. • There are several types of carbon though. • Atoms of an element can have different numbers of neutrons. • Isotopes are atoms of elements that have a different number of neutrons. • Isotopes create challenges and benefits because bonds form differently and they are usually radioactive to various degrees.

19. Isotopes • Because Isotope atoms have the same number of protons, and also electrons, isotopes have the same chemical properties. • The extra neutrons makes the isotope radioactive.

20. The Numbers Really Matter • The numbers represent an atom’s proton, neutron, and electron number when it is un-bounded and electrically neutral. • Atoms are rarely like this, stable, in nature. • Mainly because the electrons are always flying around. electron

21. The Energy Levels • Atoms are usually represented as a nucleus surrounded by rings. • The rings are energy levels. • The row (going from left to right) tells you how many rings the atom has. • Hydrogen is in the 1st row = one ring • Nitrogen is in the 2nd row = two rings • Shown to the right. • These are the atom’s energy levels.

22. The Outer Ring… The Valence Shell • The outer ring holds the valence shell electrons. • For most groups of atoms, you can determine the # of valence shell electrons from the column it’s in. • How many does this element have? • What element is it? • Knowing the valence shell is important because in the valence shell bonding happens.

23. Reading the Periodic Table • The Column tells you how many electrons in the outer shell (valence electrons). • The Row tells you how many rings (Energy levels)

24. Concept Check • On your handouts, complete the blank atom for the element carbon. • Draw the: • Correct # of electrons • Correct # of electron in the valence shell • Correct # of protons • Correct # of neutrons • Put them in the correct places.

26. Try It Again… • Try Phosphorus • Draw the: • Correct # of electrons • Correct # of electron in the valence shell • Correct # of protons • Correct # of neutrons • Put them in the correct places.

27. Chemical Bonds • Remember, electrons are orbiting the nucleus in the region called the electron cloud... In different energy levels. • The outer most edges of this cloud is called the valence shell. • There is a strict rule for how many electrons are in this valence shell & each group is slightly different. • It’s based upon the group the element is in. • Group 1A (hydrogen, etc.) has one valence electron. • Group 2A (beryllium, etc.) has two.

28. Chemical Bonds • Remember, electrons in the outermost level, or shell, are called valence electrons. • When more than one atom combines, a force called a chemical bond holds them together in the valence shell. • There are three types of bonds that you need to know. • Covalent bonds • Ionic bonds • Hydrogen bonds. • The RULE OF BONDING = Atoms tend to combine with each other such that eight electrons will be in the valence shell. • The reason is stability!

29. Valence electrons:

30. Chemical Bonds, continued • Every other element will bond with other elements to get to 8 valence electrons. • Chemical bonds form between groups of atoms because atoms become stablewhen they have eight electrons in the valence shell. • When atoms of different elements combine, a compound forms. • A compoundis a substance made of the bonded atoms of two or more elements.

31. Common Compounds • Ammonia (NH3) • Water (H2O) • Methane (CH4) • Glucose (C6H12O6) • Salt (NaCl)

32. Ionic Compounds Ionic Bonding • Atoms can sometimes achieve a stable valence level by losing or gaining electrons. • When this happens, the charge of the atom changes slightly and an ion is formed. • An ion is an atom or group of atoms that has an electric charge because it has gained or lost electrons. • Opposite charges attract. • The attractive force between oppositely charged ions is an ionic bond.

33. How Do You Know How Many Electrons Get Exchanged? • It is all based upon the number of valence electrons for the element in its basic form. • The metals on the left usually give electrons because they are closer to 8 if they lose a few. • The non-metals on the right usually gain electrons because they will achieve 8 if they gain just a few.

34. What happens when the electron leaves sodium and goes to chlorine? • The electrons no longer equal the protons and they form into oppositely charged ions, like mini-magnets. These opposite charges attract each other!

35. Find These Elements on you Periodic Table Gains e- Carbon Gainsor Loses e- Loses e- How close are these elements to achieving 8 in their valence shell? Is it faster to gain a few or lose a few? The numbers of e- gained/lost are variable & correspond to how many they need. What about carbon?

36. Try Drawing Example B1: Sodium & Chlorine Na Cl

37. + - Na Cl Formation of Ions: Count the p+ & e- before & after the exchange…

38. How Do You Show the Bonds? • Ionic Bohr Model Or… Lewis Structure

39. Try Drawing Example B2: Magnesium & Iodine Mg I

40. Since each chlorine received one electron what should each charge be? I I Since Magnesium donated 2 electrons what should its charge be? +2 -1 -1 When Magnesium loses its outer electrons it exposes its next lower energy level, which happens to have 8 electrons.

41. Chemical Bonds, continued Covalent Bonding • One way that atoms bond is by sharing valence electrons to form a covalent bond. • A molecule is a group of atoms held together by covalent bonds. • A water molecule, H2O, forms when an oxygen atom forms covalent bonds with two hydrogen atoms.

42. Poor Oxygen… • Ah, I’m sad because my valence shell isn’t filled. • If there was only some way to get more. • Wait. Hey Buddy. If we shared some electrons, then we could both have 8…kinda. • Now we both can have eight! • Sometimes… and that’s enough to form a covalent bond.

43. How Do You Show the Bonds? • Covalent Bohr Model Or… Lewis Structure

44. Illustrate Ionic & Covalent Bonding • 2 minutes… • With Lewis Structures, complete C2.

45. Polarity • Polar Molecules result when the resulting molecule has partial charges on opposite ends because of electrons are not shared equally. • In some covalent molecules, the electrons are shared equally between the atoms in the molecule. • In some covalent bonds, the shared electrons are attracted more strongly to one atom than to the other. • It’s due to electronegativity (something you don’t need to know right now). • As a result, one end of the molecule has a partial negative charge, while the opposite end has a partial positive charge.

46. Same Atoms sharing electrons equally Make it Non-Polar - - - - - + - - + - - - - -

47. Not Sharing Electrons Equally Results in Polar Molecules Partially – The electrons spend more time on this side of water. - - - - - Partially + The electrons spend less time on this side of water - - -

48. Not Sharing Electrons Equally can sometimes be 2 atoms, sometimes more. - - Partially + - - - Partially - - - -

49. Hydrogen Bonding • A hydrogen bondis a bond that forms between the positive hydrogen atom of one molecule and the negative pole of another molecule. • Represented as dashed lines. • We will see these again…

50. Polarity, continued Hydrogen Bonds • When bonded to an oxygen, nitrogen, or fluorine atom, a hydrogen atom has a partial positive charge nearly as great as a proton’s charge. • It attracts the negative pole of other nearby molecules. • This attraction is stronger than attractions between other molecules, but not as strong as covalent bonds. • However, hydrogen bonding plays an important role in many of the molecules that make up living things.