AP Biology. Chapter 2

Transcription

1 AP Biology Chapter 2

2 Matter is anything that has weight and takes up space 1. Mass is a measure of how much matter is present in a body 2. Weight is a measure of the gravitational force exerted on an object

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4 Elements are the essential building blocks of all matter An element is either a single atoms or a collection of atoms that all have the same atomic number 1 1 Au atom =elemental gold or a pound of Au atoms= elemental gold A compound is made up of more than one type of atom bonded in a particular proportion NaCl=1 sodium atom to every 1 chlorine atom H 2 SO 4 is made up of Hydrogen, Sulfur, and oxygen atoms in a 2:1:4 ratio

5 Elements cannot be broken down into simpler substances by chemical reaction while compounds can

6 ELEMENTS All matter is made from elements There are 92 elements that occur naturally All elements are named with one or two letter symbols

7 There are about 25 elements essential for life C (carbon), O (oxygen), H (hydrogen), and N (nitrogen) make up 96% of living matter. Ca (calcium), P (phosphorous), K (potassium), S (sulfur), Na (sodium), Cl (chlorine), and Mg (magnesium) make up the vast majority of the 4% of remaining living matter There are then trace elements that are necessary in extremely small quantities. B, Cr, Co, Cu, F, I, Fe, Mn,, Mo, Se, Si, Sn,, V, and Zn are all trace elements.

8 Combining elements creates compounds Elements must bond to create compounds Pure compounds have specific properties that may be very different from any of the elements that they are made of (emergent properties)

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10 Atoms Atoms are the smallest piece of an element that still has the chemical and physical properties of that element All atoms are made up of subatomic particles All atoms of a given element have similar chemical properties and exactly the same number of protons

11 Subatomic Particles Subatomic particles of interest in biology include 1. Neutrons- found in the nucleus of the atom and they have no charge 2. Protons- found in the nucleus of the atom and they have a positive charge 3. Electrons- found orbiting the nucleus, these have a negative charge and are responsible for bonding

12 The Atom

13 Neutrons and protons have almost all of the matter (mass) of an atom their weight is and dalton respectively. Electrons have a negligible mass of about.002 dalton but a negative charge that is equal in amount to the positive charge on a proton

14 Neutral Atoms Have equal numbers of electrons and protons.

15 Numbers Atomic number refers to the number of protons in a nucleus Atomic mass number refers to the mass of an atom (it is always very close to the combined number of neutrons and protons in the nucleus) To find the number of neutrons an atom has, you only need to subtract the atomic number from the atomic mass number. The weight that is left over is almost entirely neutrons The number of electrons is equal to the number of protons in a neutral atom but in charged atoms the number of protons and electrons are not equal. Ca 2+ for instance has two fewer electrons than protons causing the 2+ charge. O 2- on the other hand has two more electrons than protons.

16 The number of protons in an atom determine its identity The numbers of neutrons in a type of atom may fluctuate changing the mass number of the atom. The number of electrons in an atom may vary changing the charge of an atom

17 Ions and Isotopes Ions are atoms that have different numbers of protons and electrons leading to a charge Isotopes are atoms that have different numbers of neutrons than the normal atom of that type. The atomic weight listed on the periodic table is an average mass of all of the isotopes of a given element.

18 Formal writing of chemical information 23 Na The 23 at the top is the The +1 above notes that number of neutrons there is one more proton plus protons. The 11 than electrons in this says there are 11 atom meaning that 11- protons so there must 1=10 electrons justifying be 12 neutrons. the +1 charge.

19 Isotopes Isotopes of one element have the same atomic number but different numbers of neutrons In nature collections of one element will consist of atoms of multiple isotopes Different isotopes react chemically in the same way. Some isotopes are unstable and radioactive (radioisotopes)

20 Three isotopes of carbon

21 Radioisotopes Radioisotope proportions can be used to a. date the age of rocks, fossils, formerly living material, ice layer, etc. b. determine what type of environment existed in prehistory Radioisotopes have half lives which means the amount of time required for one half of any amount of the radioactive element to decay into a different element. Therefore the presence of radioisotopes can be measured and used to determine the age of a material

22 Radioactive tracers Radioactive isotopes can be used to get a look at specific tissues (radioactive iodine can be ingested at low levels, it concentrates es in the thyroid emitting radiation that allows doctors to get an image of the thyroid) Radioactive tracers can be used to make enzymes or other chemicals then placed in the body to see where they go and how they behave. Radioactive isotopes can be incorporated into biological chemicals to determine that chemicals shape and behavior. (that is how the structure of DNA was determined) Radiation released by radioisotopes is hazardous meaning that it damages important chemical bonds in cells The damaging effects of radiation allow it to be used to intentionally cause damage to tissue like cancer tumors.

23 gamma camera image

24 Electrons These are light little negatively charged particles which orbit the nucleus of atoms These are the subatomic particles that are responsible for bonding. These have potential energy based on their position in relation to the nucleus.

25 Electron energy Electrons exist in fixed potential energy states known as energy levels or energy shells. The differences between the potential energy between the shells that an electron can occupy are known as quanta and determine what amount of energy that atom can absorb or emit when the electron changes levels. Electron shells nearer the nucleus are the lowest energy shells, shells that are farther from the nucleus are higher in potential energy.

26 Energy absorption and emission Electrons may move between energy levels or states. There may be several states that an electron can occupy, but that atom can only absorb or emit energy in amounts that are equal to the difference between these energy levels. The difference in energy between level 1 and 2 is what is absorbed when an electron moves from level one to level two. That is also the amount of energy released when the electron falls back to level 1. The electron could also do the same thing between levels 2 and 3, or levels 1 and three, but greater amounts of energy are involved in moving an electron between levels 1 and 3.

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28 Electrons determine the chemical behavior of an atom

29 To understand atomic behavior, you must understand electron configuration The electron configuration is the hierarchy of electron organization. Electrons are organized into patterns known as shells which are made up of one or more orbitals. Electrons will fill the lowest (innermost) shells first. The first shell contains two electrons, the next shell contains eight then the third shell contains eight. Only the outermost shell participates in bonding

30 Every electron shell is made up of one or more electron orbitals An electron orbital is the three dimensional space that an electron is most likely to occupy Shell one is made up of the 1s orbital Shell two is made up of the 2s orbital and three different 2p orbitals Shell three is made up of the 3s, and three 3p orbitals

31 Use this graphic Shell 1= 1s Shell 2= 2s Shell 3= 2p 3s 3p Each blank represents an orbital which can hold two electrons. You can see from the chart that the first shell only holds one orbital and therefore, one electron pair. The second and third shells hold four orbitals or 8 electrons each.

32 Only the outermost shells react Only the outermost shells are involved in bonding, they are known as valence shells. Because only the valence shells are involved in bonding, there is no sense in diagramming lower full shells for the purposes of determining reactivity Atoms bond to fill their valence (outer) electron shell, to do this they may share, steal, or kick off electrons. Stealing or kicking off electrons leads to the formation of an ion, while sharing leads to a strong interaction between two atoms known as a covalent bonds Lewis Dot structures are diagrams that show the number of available valence electrons and thereby likely reactions of an atom

33 The lewis dot structure Please note that oxygen has eight electrons all total but only six are shown here. The other two are in the full 1s orbital and are not available for bonding so they are not shown here. The other six have filled the orbitals as shown on the next slide.

34 Notice that O has eight total electrons but only the six valence electrons are shown on the lewis dot structure because they are the ones that determine chemical behavior Shell 1= 1s ll Shell 2= 2s ll 2p ll l l Shell 3= 3s 3p

35 Electron influence on bonding patterns Atoms with a full valence shell are satisfied so to speak and are e not reactive. Helium, argon, neon, krypton all have full valence shells and therefore, no need to lose, gain, or share electrons. Other atoms are trying to gain full shells. It is very common to see atoms with one valence electron lose that electron to have a full shell while atoms with seven electrons often steal an electron to t have eight. Atoms with two valence electrons often lose them creating a 2+ charge while atoms with six gain two more creating a 2- charge. Atoms with 3, 4, or 5 electrons tend to form more covalent bonds by sharing electrons they do have with another atom. Atoms with equal numbers of valence electrons tend to behave in similar ways in regard to reacting resulting in the groups found on the periodic table.

36 Valence shells that are full are known as octets, therefore the tendency of an atom to want eight valence electrons is called the octet rule Valence formation is the driving force determining bonding patterns and types.

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38 Bond Types Covalent bonds are bonds in which two or more atoms share electrons to complete their octets. These are the strongest bonds Ionic bonds are bonds in which atoms that have a positive charge and atoms that have a negative charge are electromagnetically attracted to one another to balance out the charge. Hydrogen bonds are weak electromagnetic attractions between a hydrogen atom involved in a covalent bond and a slightly negatively charged atom from another molecule.

39 Chemical bonds are simply attractions that hold atoms together All of the atoms that are covalently bonded together make up one covalent molecule The lowest whole number proportion of atoms gives you the formula for one ionic molecule.

40 Covalent bonds Covalent bonds are bonds that are created when at least one pair of electrons is shared between two atoms. H:H Two pairs of electrons may be shared (double covalent bond) ::O::O:: Three pairs of electrons may be shared between two atoms to achieve an octet. :N:::N: Single bonds allow atoms to rotate in relation to one another while double and triple bonds result in a rigid relationship.

41 Covalent bonds hold energy Covalent bonds either require an energy input to break them or give off energy when broken (actually they do both)

42 Covalent bond numbers are usually easy to predict The number of unpaired electrons that an atom has tends to be the number of covalent bonds the atom will make (there are exceptions)

43 Covalent bonds Polar Covalent Bonds are bonds in which the two atoms involved in the bond do not share electrons equally. Nonpolar Covalent Bonds are bonds in which there is a relatively equal sharing of electrons used in bonding by the atoms involved. This is based on electronegativity.

44 Electronegativity The ability of an atom to attract and then hold electrons. Oxygen has a high electronegativity and hydrogen has a low negativity. The result is the electrons in a water molecule tend to stay with the oxygen more than the hydrogen giving the oxygen a slightly negative charge and the hydrogen a slightly positive charge.

45 Ionic bonds Ionic bonds are really electrostatic attractions between two oppositely charged ions. Negatively charged ions (anions) are attracted to positively charged ions cations creating ionic bonds. The lowest proportion of cations to anions is the formula for an ionic compound. Ionic bonds are strong in the crystal form but very weak in polar liquids

46 Electronegativity review Equal attraction to electrons between two atoms equals a nonpolar covalent bond Electronegativity differences between two atoms that share electrons result in polar covalent bonds. Electronegativity differences between two atoms that are very large results in the electron being taken from the less electronegative atom by the more electronegative atom creating an anion (negative ion) and a cation (positive ion). These can then be held together by the electrostatic forces between them creating an ionic bond.

47 The importance of weak chemical bonds Hydrogen bonds, ionic bonds in solution, Van der Waals and hydrophobic interactions are all included They are temporary associations allowing for a quick response that is self terminating Can form between two chemicals or between separate parts of the same chemical Help stabilize the three dimensional shapes of some molecules

48 Hydrogen bonds are weak attractions between the slightly positive hydrogen on one polar molecule and the slightly negative atom of another polar molecule Covalent bonds are about 20 times stronger than Hydrogen bonds Hydrogen bonds are electrostatic attractions between oppositely charged portions of two polar molecules Often seen between a H and an O or N as in H 2 O or NH 3 (between two molecules not within the molecule) These are the bonds responsible for the pleasure of a belly flop!

49 Van der Waals interactions These are incredibly brief attractions created between adjacent atoms due to slight electrostatic charges produced as a result of electron distribution assymetry

50 Each molecule has a characteristic shape and size For many biological molecules, their function is determined by their shape. If there are only two atoms in a molecule, the shape of that molecule is linear. Each additional atom increases the potential complexity in shape of the resultant molecule Covalent bonds and complete valence pairs arrange themselves as far away from one another as possible (tetrahedron)

51 Chemical reactions The process of making and breaking chemical bonds Matter is not gained or lost, just rearranged Reactions can proceed until all reactants are used up Most are reversible!

52 Equilibrium The concentration difference at which the forward and reverse reaction rates are equal This does NOT mean that the amount of products and reactants are equal, the rate of reaction is equal.

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