The Nature of Matter

1. The Nature of Matter Chapter 1 BIOLOGY 391

2. What is everything made of? • MATTER Anything that has mass and takes up space • ATOM  The smallest unit of matter

3. ATOMS • Basic unit of matter • Size: 1,000,000 (million) side by side = 1 cm • Atoms like to be neutral- no charge • Equal number of protons and electrons • There are specific numbers of “sub-atomic” particles that the atom wants • Special cases are called isotopes or ions

4. STRUCTURE OF ATOM • NUCLEUS - protons and neutrons held together by the “strong force” • Protons (+) • Neutrons (o) • ELECTRON CLOUD (orbitals) – electrons surround nucleus • Electrons (-) only contains about 1/1836 the mass of proton or neutron • Constantly moving within orbital- attracted to the nucleus by the “weak force”

5. ATOMS • The smallest particle of an element that can exist and still have the properties of that element. • An atom has many subatomic particles • Atomic mass unit or Dalton • ** The mass of an e- is so small it’s negligible.

6. Electron orbitals 1st orbital can only hold 2 electrons (too close to nucleus - not much space) 2nd orbital can hold up to 8 3rd orbital can hold up to 8

7. How many atoms are there?PERIODIC TABLE

8. Atoms are grouped as Elements • Are listed on the Periodic Table • Dimitri Mendele’ev 1869 • Arranged in order of their atomic #’s • Table is divided into Groups and Periods • The atomic number and atomic mass are given for each element. Example: Atomic # 6 C 12.01 Element Atomic Mass

9. “Groups” Vertical Column Numbered 1-8 All elements of same group have the same # of e-’s in their valence shell & have similar chemical properties Atomic # increases by 1 from left to right Example: H, Li, Na, K, Rb, Cs, Fr All have 1 e- in valence (outer) shell

10. “Periods” • Is a horizontal row of elements • There are 7 periods • Elements in a period have the same # of energy levels Example: H and He are in period 1 They have 1 energy level

11. The Elements are often described as Metallic and Nonmetallic

12. 6 C Carbon 12.011 Atomic Number # of protons (and also # of electrons) Name of Element Chemical symbol Atomic Mass The weight Of carbon atom or average weight of all isotopes What is the difference between atoms?

13. Special Form: ISOTOPE Atoms of the same element containing different numbers of neutrons in the nucleus • Some give off radiation – used to: • Trace atoms through a reaction or an organism • Treat cancer • Date very old, once living organisms

14. 2 examples of Isotopes

15. Recap • What are the three subatomic particles and their charges? • What is the only actual difference between gold and mercury? • What is the atomic mass of lead? • What is an isotope?

16. Organization of Matter • Atoms usually do not occur alone, but exist with other atoms as: • Elements (all of the same atoms) • Molecules or Compounds • Same or different atoms bonded

17. MATTER: anything that has mass & takes up space Protons Neutrons Electrons ATOMS Elements: Shown in Periodic Table pure bonded Molecules or Compounds Organic Inorganic Does not contain C-C bonds C-C, C-H bond LIFE!

18. Elements • A substance which cannot be split into simpler substances by a chemical rxn. • A grouping of the same type of atoms • ORDER MATTERS! • More than 100 elements exist (shown in the periodic table) • Carbon Elements:

19. Elements of Life 92 naturally occurring elements Elements Found in Living Organisms N CHOPS (macronutrients) C HOPKINS Ca Fe Mg B Mn Cu Cl Mo Zn

20. MATTER: has mass & takes up space Protons Neutrons Electrons ATOMS Elements: Shown in Periodic Table pure bonded Molecules or Compounds Organic Inorganic Does not contain C-C bonds C-C, C-H bond LIFE!

21. Why do atoms bond? • An atom wants to have a complete outer shell of electrons. To do this, it can… • Share electrons with another atom • Give away its extra electrons • Steal extra electrons • Bonds store ENERGY

22. Remember: an atom is  when its outer orbital is filled For the “smaller” elements, the outer shell holds 8. The electron clouds Increase in size as you go across the periodic table. HOWEVER, the trend stays the same for the number needed to complete the outermost shell.

23. BONDING: depends on number of electrons in the outermost orbital shell • Covalent Bond- atoms share a pair of electrons sometimes share 2 (double bond) or 3 (triple bond) pairs • Ionic Bond- One atom (very unstable) gives 1, 2 or 3 electrons away to another atom. The atom that loses electrons becomes positively charged. The atom that gains the electrons becomes negatively charged. • The opposite charges cause the atoms to “bond” together (opposites attract).

24. Molecules  bonded atoms “Molecule” is often used to refer to an individual grouping. “singular”

25. aspirin sucrose Compounds: also, 2 or moreatomsbonded together.Often referred to as a larger conglomerate of bonded molecules • Order matters! • These 2 cmpnds are made of the same 3 atoms, but in a different arrangement. • The arrangement is responsible for their affects.

26. Bonding • IONIC • COVALENT • Polar • Nonpolar • VAN DER WAALS • HYDROGEN

27. The periodic table shows the pattern of electrons each element has in its outermost shell

28. Ions • Atoms that have lost or gained electrons cation – positively charged ion – lost electrons anion – negatively charged ion – gained electrons • Ionic bonds are weaker than covalent bonds • Hold less energy in the bond

29. Cations versus Anions • Atom that loses electrons • Positively charged ion • Elements in Groups 1, 2, and 3 tend to lose electrons • Metallic elements tend to form positive ions Example: Ca Lose 2 electrons +2 charge • Atom that gains electrons • Negatively charged ion • Elements in groups 5, 6, and 7 tend to gain electrons • Nonmetallic elements tend to form anions Example: Cl Gain 1 electron -1 charge

30. IONIC BONDING Na (sodium) is very unstable because it only has one e- in its outer orbital. Cl’s (chlorine) outer orbital is almost filled. Na gives its lonely e- to Cl. Na become Na+ Cl becomes Cl- Their opposite charges cause them to be attracted to one another- This is an ionic bond.

31. Ionic Compounds • Metals react with nonmetals forming ionic compounds • Salts • Held together by electrostatic forces Example: + attracted to – • Most are crystalline solids at room temperature • When dissolved in water they conduct electricity Ex. Sodium Chloride or Table Salt Na+ + Cl- • Dissociate(break apart) in water, producing free ions

32. Electrolyte • A solution that conducts electricity. • Term for salts, specifically ions. • The term electrolyte means that this ion is electrically-charged and moves to either a negative (cathode) or positive (anode) electrode • Ions that move to the cathode (cations) are positively charged • Ions that move to the anode (anions) are negatively charged

33. Covalent Bonds • Bonds formed when atoms share electrons • Atoms with 4 or 5 e-’s tend to share • Each e- spends part of its time around one nucleus and then around the other nucleus. • Sharing e-’s completes the valence shell

34. Example of Covalent Bonding- Water Example of a Covalent Bond

35. Covalent Bonds • Bonds between non-metals • Poor conductors of electricity • Do not dissociate easily in water • Two Types: Polar: Unequal Sharing Non-Polar: Equal Sharing

36. Van der Waals forces • Attraction between oppositely charged areas of adjacent molecules • Remember, e- are constantly swarming around the nucleus. At times, there may be a moment of asymmetry • All of the electrons might be on one side • This sets up a “dipole” weaker than covalent bonds and ionic bonds

38. RECAP • What is an ion? • Why is it important that atoms bond? • What causes atoms to bond? • Explain the difference between an ionic bond and a covalent bond • What are Van der Waals forces?

39. IONIC Transfers/takes electrons Positive/Negative charges Weak bond Ex: NaCl (salt) COVALENT Shares electrons Neutral Strong bond Ex: H2O, CO2, NH3 BONDS

43. WATER – Why so Great?

44. Water! • What is the chemical formula for water? • How much water covers the Earth? • How much of your body is water? • Is there water in food? • How long could you live without water? • H2O • 75% • 60-70% • Yes! • 3 days

45. Properties of Water • Phases: Solid, Liquid, Gas • Polarity • Hydrogen bonds • Adhesion • Cohesion • Making Mixtures • Solutions • Suspensions • Making Acids and Bases

46. Water Density • Ice is less dense than liquid water • When water freezes air is trapped within the frozen ice making the cube larger and less dense • Benefits: • Fish and plant life can survive in liquid layers of water under ice

47. PHASE CHANGES: the closeness and speed of the compounds

48. Polarity Hydrogen ends become slightly positive • Water is polar • Although the compound is neutral overall there is a shift of charge within the compound The much larger atom, Oxygen, pulls more on the shared e- This end of the compound becomes slightly more negative.

49. Hydrogen Bonding • Due to polarity, water compounds attract to one another • Slightly negative oxygen attracts slightly positive hydrogen from another compound • This attraction among water is COHESION. • Water is also attracted to other materials. This is ADHESION.

50. COHESION ADHESION Water compounds attract To glass molecules And form a meniscus Water compounds attract To one another- causes water to “bead”