Biochem Exam 1.pdf

 Biochem Lec1-BuildingBlocks of Molecules Pl Side Notes Matter- physical material of the universe nothing that has · Main...

 Biochem Lec1-BuildingBlocks of Molecules Pl Side Notes Matter- physical material of the universe nothing that has · Main components thingsof life > - proteins , Nodez Acids Carbohydrates lipids , , mass and occupies space volume Mass-amount of mtter present energetics > -> Metabolism/free energy Pure substance- uniform &C H N 0 * chemical composition , , , compounds Mixture-variable comp can be physically separated Molecule-smallest unit of pure substance that retains · HO zhydrogen loxy , : wath properties made up of atoms · NaCl I sortium Ichloum table salt LotticOy hyd boxy Scarbon 12 glucose · Rep of Molecules. Medeclar. Structural formula : Bull/sticke Model · Space filling H20 · ⑭ O I It Summary :  Biochem Leck: P2 Main Eide Notes things nucleus protons as proton/neutron electrons outside - - , are about the same Size but elections Atomic #- > A of protons to make neutral It has same # O are much smaller of electrons · Mass - protons + neutrons in a element · Isotopes same element but diff neutrons onections is protons - mass = neutrons Atomic Weight-> Aug mass of of · an element bottom · Molecular Weight- 40 2. 146 180 = neutrons/mass # differ istipes if neutral Summary  electron config. sholds 2....... ↑ pholds - ↑ & O. d holds L shoks 14 & Hof elements * sblock > - Phlock S Di I Abou is 252 zpb 3523pb ns24p4d14414 552 spl do f 65 6pt d 75 +pl f - 4/S  Biochemilec1 P3. Valence shell- Horizonta 1 Row > - - periods (15) Outer most election-vertical ROW (7) periods shell of an element highest energy main 1-0 Erp) Valence electron- groups An election In the Neutral element # protins - : A elections Valence shell Of more unstable atom shells-electrons further from nucleus an Energy el · are higher energy 4 things define elections lot 1 12l Shell 2, 3 dif "determine crags. , · subshell shape s p, 3 Orbital , 4. Spin - the more 4 Orbitals hold you divide subshell 2 clea Higher energy it is Peric (n) = Shall group Subshell/specific cif) : block (s p , , are the same : Valence electrons but Spide -be diff-12, ,3 4 , 5 6 7 8 , dire , , I J  Biochem : Lec 1 PS a &S' great colorless · · Molecules More stable than atoms · How the octet many elections to fill Lewis dot symbol dots to of Valence elections · uses show on around then again Bunding Am can grin/duse transfor elections Sonic bonding covalent > - sharing : Fi "  BC 212 Lecture 1: Building Blocks of Molecules Part 1/6: Introduction and Definitions Allison Lamanna Lecture 1 Part 1/6 Suggested Reading and Problems (8th ed.) Chapter 1, sections 1-5 Chapter 1: 1.26, 1.50, 1.52 Learning Goals What are atoms, elements, molecules, and compounds? What are the most common elements found in biological systems? Biochemistry: The Chemistry of Life Understanding the molecules and reactions of biological systems To understand a Components of Life: system: Proteins Deconstruct to Nucleic Acids component parts Carbohydrates Study isolated Lipids components Reconstruct from Energetics of Life: isolated components Metabolism Free Energy  Elements found in cells Structure à Function What makes up the structure of biological molecules? Nitrogen Carbon Oxygen Hydrogch , , , Definitions Matter: Physical material of the universe; anything that has mass and occupies space (volume) Mass: Amount of matter present Pure substance: Uniform chemical composition Mixture: Variable composition; can be physically separated Molecule: Smallest unit of pure substance that retains properties; made up of atoms Atom: Smallest particle of matter (element) Element: Homogeneous, pure substance, only one kind of atom Compound: Homogeneous, pure substance, multiple kinds of atoms Elements O2 Cu Compounds H2O (2 hydrogens, 1 oxygen) water NaCl (1 sodium, 1 chlorine) table salt C6H12O6 (6 carbon, 12 hydrogen, 6 oxygen) glucose  Which of the following is a compound? Na O2 C NH3 Lecture Question 1-1-1  There are 118 known elements. https://iupac.org/what-we-do/periodic-table-of-elements/  Which element below is not present in significant amounts in the human body? Carbon (C) Nitrogen (N) Oxygen (O) Fluorine (F) Lecture Question 1-1-2 Representations of molecules  How many nitrogen (N) atoms are in 2 molecules of caffeine (C8H10N4O2)? 2 4 8 10 Lecture Question 1-1-3 Lecture 1 Part 1/6 Summary Elements are made of atoms that can combine to form molecules or compounds. The most prevalent elements in the human body are carbon (C), nitrogen (N), oxygen (O), and hydrogen (H).  BC 212 Lecture 1: Building Blocks of Molecules Part 2/6: Subatomic Particles and Isotopes Allison Lamanna Lecture 1 Part 2/6 Suggested Reading and Problems (8th ed.) Chapter 2, sections 1-2, 3 (no calculations) Chapter 2: 2.42, 2.44, 2.46 Learning Goals What are subatomic particles, and how do they define the properties of atoms? What are isotopes?  Subatomic particles Why do specific atoms come together to form specific molecules? Why do those molecules have specific properties? Atomic Structure Electron cloud Nucleus: protons and neutrons Atomic size: 10-10 m (0.0001 µm)  If the atom is the size of Michigan stadium… If a proton is the size of the Epcot ball… …an electron is the size of your head. …its nucleus is the size of a fly on the “M”. Definitions Neutral atom: Element with equal numbers of protons and electrons, uncharged Atomic number: Number of protons in an element, top number on the periodic table (Z) Isotopes: Atoms with the same number of protons (atomic number) but different numbers of neutrons Mass number: Number of protons + neutrons in an element (A) Atomic weight: Average mass of an element, based on isotope abundance  Periodic Table of the Elements Atomic Number: 1 8 20 30 26 82 H O Ca Zn Fe Pb 1.0079 15.999 40.078 65.39 55.847 207.2 Definitions Neutral atom: Element with equal numbers of protons and electrons, uncharged Atomic number: Number of protons in an element, top number on the periodic table (Z) Isotopes: Atoms with the same number of protons (atomic number) but different numbers of neutrons Mass number: Number of protons + neutrons in an element (A) Atomic weight: Average mass of an element, based on isotope abundance Mass number (A) = Protons + neutrons A Z E Element symbol Atomic number (Z) = A Protons  Spri Hydrogen Isotopes 1 H 1.0079 # Neutrons Mass number Symbol Name 0 1+0=1 1 1H Hydrogen-1 (protium) 1 1+1=2 2 1H Hydrogen-2 (deuterium) 2 1+2=3 3 1H Hydrogen-3 (tritium) A major isotope used in radiology studies is technetium-99 (99Tc). How many protons and neutrons are in an atom of Tc-99? How many electrons are associated with the neutral atom? O43 Tc O (99) 43 p, 43 n, 43 e find Neutrono 43 p,O56 n, 43 e to Pro- emer Lecture Question 1-2-1 has 99 p, 99 n, 99 e samel 43 p, 99 n, 99 e Dande  Periodic Table of the Elements Atomic Number: 1 8 20 30 79 82 H O Ca Zn Au Pb 1.0079 15.999 40.078 65.39 196.97 207.2 Atomic Weight Atomic weight is the average mass of an element, based on the relative abundance of its isotopes. (You do not need to calculate this!) Molecular Weight 1 6 8 H C O 1.0079 12.01 15.999 Molecular weight is the sum of the atomic weights for all of the atoms in a molecule. What is the molecular weight of water (H2O) ? 18 u 2*1 + 16 = 18 u What is the molecular weight of glucose (C6H12O6) ? 180 u 6*12 + 12*1 + 6*16 = 180 u  -and carbon-14. 2 isotopes of carbon are carbon-12 O Which of the following numbers are identical for these neutral isotopes: Atomic #, Mass #, # protons, # neutrons, and # electrons? ① 6 C & Got 12.011 Atomic #, mass # Atomic #, # protons, # electrons O Mass #, # neutrons Lecture Question 1-2-2 Mass #, # protons, # neutrons Lecture 1 Part 2/6 Summary An atom is made of subatomic particles (protons, electrons, and neutrons), and its properties are defined primarily by its atomic number. A neutral atom has an equal number of protons and electrons. Isotopes are atoms with the same number of protons (atomic number) but different numbers of neutrons. The atomic weight of an element depends on the relative abundance of its isotopes.  BC 212 Lecture 1: Building Blocks of Molecules Part 3/6: Electronic Structure and the Periodic Table Allison Lamanna Lecture 1 Part 3/6 Suggested Reading and Problems (8th ed.) Chapter 2, sections 4, 6, 8 Chapter 2: 2.28, 2.58, 2.64, 2.68, 2.76, 2.94 Learning Goals How is the periodic table organized? How do valence electrons define the properties of atoms?  Periodic Table Elements are defined by their atomic number (number of protons). Elements are listed in order by atomic number. Elements with similar chemical properties appear at regular (periodic) intervals. https://iupac.org/what-we-do/periodic-table-of-elements/ Periodic Table Structure  Periodic Table Structure Horizontal rows are called “periods”. Vertical columns are called ”groups” or “families”. 18 groups/families 7 periods  I Groups/Families VIII II III IV V VI VII Periods Electronic structure of the atom Elements have a specific atomic number (number of protons). In neutral atoms, the number of protons = number of electrons. Each neutral element has a specific number of electrons. The electrons define the chemical properties of elements. Where are the electrons located in the element?  How many electrons does one neutral atom of oxygen (O) have? 6 8 10 16 Lecture Question 1-3-1  Energy shells An electron’s location is related to its energy. Electrons further from nucleus are higher energy (more unstable) Electrons closer to nucleus are lower energy (more stable) Defining an electron’s location 1. Shell: (principle quantum number) n=1, 2, 3, etc. 2. Subshell: (shape) s, p, d, f 3. Orbital: division of subshell, holds 2 electrons 4. Spin: which position in the orbital (up, down) Each electron must have a different location (as defined by these four values). Energy shells n=3 n=2 n=1 Higher shell (principle quantum) number (n) means higher energy (more unstable). Only so many electrons can fit in each shell.  sphere Subshells and orbitals s “dumb-bell” p “clover-leaf” d Higher subshell (more separated lobes) means higher energy (more unstable). Electron location in atoms Shell/subshell defines the energy of the electron Bigger shell numbers = higher energy Within a shell, larger subshells = higher energy Higher energy = further from nucleus (Orbitals within a subshell have the same energy = degenerate)  What is the maximum number of electrons that can fit in the 2nd shell (n=2)? 2 6 8 10 Lecture Question 1-3-2  Periodic table organization s begins with n=1 p begins with n=2 d begins with n=3 f begins with n=4 Location in the periodic table tells you shell/subshell of highest energy electron for each element Period = shell (n) Group = subshell/specific “block” (s, p, d, f) s begins with n=1 p begins with n=2 d begins with n=3 f begins with n=4 2 electrons 6 electrons 10 electrons 14 electrons Main group elements = s + p blocks (main group (main group elements) elements) s2 p6 Definitions Valence shell: Outermost electron shell of an element (highest energy) Valence electron: An electron in the valence shell of an atom The period of an element is its valence shell. The group specifies the subshell and how many valence electrons an element has. For main-group elements, the valence shell can have up to 8 (s2 + p6) total electrons. Valence electrons (main groups) Group IA (1) Subshell Valence Element electrons H (1) s1 1 Li (3) s1 1 Na (11) s1 1 K (19) s1 1 Group VIIIA Subshell Valence (18) Element electrons He (2) s2 2 Ne (10) s2p6 8 Ar (18) s2p6 8 Kr (36) s2p6 8 Elements with similarly filled valence shells have similar physical and chemical properties. Strontium (Sr) from nuclear fallout can replace calcium (Ca) in bones, because they both have the same number of valence electrons. How many valence electrons do they both have? 20 38 Ca Sr 40.078 87.62 one (s1) two (s2) three (s2p1) four (s2p2) Lecture Question 1-3-3  Valence electrons (main groups) Period 2 Li Be B C N O F Ne elements Valence 1 2 3 4 5 6 7 8 electrons  Do you expect Ca to be more similar to K or Mg? 20 19 12 Ca K Mg 40.078 39.098 24.305 K Mg neither Lecture Question 1-3-4 Lecture 1 Part 3/6 Summary The periodic table is organized by the location of electrons in shells (periods; n=1, 2, 3, etc.) and subshells (groups; s: 1A/2A and p: 3A-8A). The periodic table is organized by the number of valence electrons. The location of the outermost (valence) electrons determines their energy and the kind of reactions and bonding an element will undergo. For main-group elements, the group tells you the number of valence electrons.  BC 212 Lecture 1: Building Blocks of Molecules Part 4/6: This section provided as a Periodic Properties video with IVQs Allison Lamanna to complete S Lecture 1 Part 4/6 Suggested Reading and Problems (8th ed.) Chapter 2, section 4; Chapter 3, section 4 Chapter 2: 2.52; Chapter 3: 3.50 Learning Goals What are the major periodic properties of elements in the periodic table? What are ionization energy and electronegativity?  Periodic Table Structure s1 s2 s2p1 s2p2 s2p3 s2p4 s2p5 s2p6 Period 2 Li Be B C N O F Ne elements Valence 1 2 3 4 5 6 7 8 electrons The periodic table is organized by number of valence electrons. Valence electrons determine physical and chemical properties. 1-8 left to Right Valence Electrons numbered  Periodic property: Metallic Character nthe ↳ middl Brittle ↳ Powde as Metallic characteristics: transmit heat and electricity, form into wires and sheets, have metallic luster  Metallic character trends The most metallic elements are the bottommost and leftmost on the periodic table. TheMos - Fromo top The Elements of Group VA (5A, 15) As arsenic Sb antimony N nitrogen P phosphorous Bi bismuth Periodic property: Atomic size Period As period increases, atoms have more (larger) shells filled with 1 electrons. 2 Group 3 1A 2A 3A 4A 5A 6A 7A 4 Within the same period (shell), the electrons are the same distance from the 5 nucleus. But, as atomic number increases, the positive charge (# of protons) increases, and the increased positive 6 charge pulls the electrons closer in.  Atomic size trends The largest elements are the bottommost and leftmost on the periodic table. THE Periodic property: Atomic size, density As atomic size increases, density increases. (Density is the amount of matter in a given volume.) Atomic Size and Densities: Density 0.18 g/L Increase Density of air 0.90 g/L 1.29 g/L 1.78 g/L When gases are less dense than air, they rise. 3.75 g/L When gases are more dense 5.89 g/L than air, they sink. 9.73 g/L https://www.youtube.com/watch?v=QLrofyj6a2s&feature=player_embedded Which of the following elements is the largest? Mg L P ↓ S Cl o  Periodic property: Ionization Energy Ionization Energy: how much energy it takes to remove an electron El à El+ + e- The elements with the highest ionization energy are the topmost and rightmost on the periodic table. opporat  Periodic property: Ionization Energy and Reactivity As ionization energy increases, reactivity decreases. As ionization &v (Reactivity represents how easily an electron is lost Reactivity its opp/ so atomic during a reaction.) As same density sizei The most reactive elements are the bottommost and leftmost Reactivity on the periodic table. moh Group 1 metal “M”: 2M + 2H2O à 2MOH + H2 & https://www.youtube.com/watch?v=QSZ-3wScePM&feature=player_embedded Periodic property: Ionization Energy electronegativy is sure the e and ab Electronegativity ion As ionization energy increases, electronegativity increases. (Electronegativity represents how tightly elements hold electrons.) The elements with the highest electronegativity are the topmost - and rightmost on the periodic table. Electronegativity Which of the following elements has the greatest ionization energy? C N O F  BC 212 Lecture 2: Molecular Structure and Interactions with Water Part 1/6: Covalent Bonding and Lewis Structures Allison Lamanna Lecture 2 Part 1/6 Suggested Reading and Problems (8th ed.) Chapter 4, sections 1-3, 6-7 Chapter 4: 4.47, 4.49, 4.56 Learning Goals How many bonds and lone pairs do you expect a certain element to have in a molecule? How is covalent bonding assessed in a polyatomic ion?  Covalent bonding - 2 atoms that share election to make them an stable Covalent bond based on attractive force between nuclei (+) and shared electrons (-) Lewis structures: Has tochar A “lone pair” of electrons F + F F F or F F F2 Example: Use the Lewis symbol to determine how many electrons carbon (C) needs for a stable octet. Carbon is in group IVA Na C C It has 4 valence electrons (2 in 2s, 2 in 2p) It needs 4 more electrons for a stable octet C has 4 unpaired electrons available to bond. C can make 4 bonds to get 4 additional shared electrons and have a stable octet.  How many bonds to fill the octet? IVA VA VIA VIIA VIIIA Group (14) (15) (16) (17) (18) Valence electrons 4 5 6 7 8 Expected bonds 4 3 2 1 0 (octet) Howmoreens ? Lone pairs 0 1 2 3 4 2 I pair = Example CH4 NH3 H2O HF Ne compound Lewis structures: H H O H F H C H N H H H H H Sulfur (S) makes 2 bonds to fill its octet. How many lone pairs does S have when A it makes 2 bonds? 2 pairs j - · : · A 16 S 39.098 1 2 3 4 Lecture Question 2-1-1 Lewis structure of water (H2O) O H Oxygen is in group VIA Hydrogen is in group IA It has 6 valence electrons It has 1 valence electron (2 in 2s, 4 in 2p) (1 in 1s) It needs 2 more electrons It needs 1 more electron for a stable octet for a stable duet It will form 2 bonds It will form 1 bond O + H + H = O H O H H H 1 O + 2 H = H2O (water) Lewis structure of molecular oxygen (O2) O Oxygen is in group VIA O + O = O O It has 6 valence electrons (2 in 2s, 4 in 2p) It needs 2 more electrons O O for a stable octet A “double bond” It will form 2 bonds 2 pair of elections  How many bonds for “X”2? VA VIA VIIA Group (15) (16) (17) Valence electrons 5 6 7 Expected bonds 3 2 1 (octet) Lone pairs 1 2 3 (per atom) Example N2 O2 F2 compound Lewis structures: N N O O Neverwhat F F More A “triple bond” A “double bond” A “single bond” Rules for Lewis structures Each atom has a full octet (duet for H) The number of electrons in the compound must equal the sum of valence electrons of all the atoms If overall charge is negative, that is number of electrons added to total valence electrons. If overall charge is positive, that is number of electrons subtracted from total valence electrons. Atoms should have the expected number of bonds If there is a charge, then atoms are not making the expected number of bonds. For each negative charge, an atom is making one less bond (has one more unshared electron). E For each positive charge, an atom is making one more bond (has one less unshared electron).- Note: there are some cases where atoms can exceed the octet (extra bonds)  Lewis Structure of CO2 O C O does itobey Total electrons = 4 + 2x6 = rulese. 4 + 12 = 16 electrons Carbon = group IVA 4Cllone pairsAr = 8 electrons Na C 4 valence electrons 4 bonds (2 double bonds) 4 bonds = 8 electrons Oxygen = group VIA 8 + 8 = 16 electrons 6 valence electrons O Carbon makes 4 bonds, 2 bonds no lone pairs counterections bands · Each oxygen makes 2 do they make bonds, 2 lone pairs  2- Lewis Structure of CO32- O C Carbon = group IVA O O 4 valence electrons 4 bonds Total electrons = 4 + 3x6 + 2 = 4 + 18 +&2 = 24 electrons Oxygen = group VIA 8 lone pairs = 16 electrons 6 valence electrons 4 bonds (1 double bond) = 8 2 bonds electrons Polyatomic ion (2-) = 16 + 8 = 24 electrons 2 extra electrons Carbon makes 4 bonds, no lone pairs 2 atoms with less One oxygen makes 2 bonds, 2 lone sharing (one less pairs bond, one more Two oxygens make 1 bond, 3 lone lone pair) pairs = 2 atoms with less sharing  H H CH4 H C H NH4+ H N H H H Total electrons = 4 + 1x4 = 8 Total electrons = 5 + 1x4 -1 = 8 No lone pairs, 4 bonds = 8 No lone pairs, 4 bonds = 8 Carbon makes 4 bonds, no Nitrogen makes 4 bonds, no lone pairs lone pairs = one atom with Each hydrogen makes 1 more sharing bond Each hydrogen makes 1 bond  Which of the following Lewis structures is incorrect? - & Cl ~ H O Cl C H H o C88 N N o H H H d CH2Cl2 HCN NH3 ⑰ A B 145 C 7 Lecture Question 2-1-2  Which of the following Lewis structures is incorrect? X F o ~ F 00 O - & 8 F O O C & O H C : N% 0 N O so.. F F H ↓ CH2F2 ⑳A V 4+ CN- B 2" + ⑭ NF3 C Y - - S Lecture Question 2-1-3 Exceptions to the octet rule Period 3 elements (have d orbitals) can sometimes have more than an octet Group VA can have 5 bonds Group VIA can have 6 bonds 3- 2- O O PO43- O P O SO42- O S O O O 54 00  O 3- 2- O O PO43- O P O SO42- O S O O O Total electrons = 5 + 4x6 +3 = 32 Total electrons = 6 + 4x6 +2 = 32 11 lone pairs, 5 bonds = 32 10 lone pairs, 6 bonds = 32 P makes 5 bonds, no lone pairs S makes 6 bonds, no lone pairs One oxygen makes 2 bonds, 2 2 oxygens make 2 bonds, 2 lone lone pairs pairs 3 oxygens make 1 bond, 3 lone pairs = 3 atoms with less sharing 2 oxygens make 1 bond, 3 lone pairs = 2 atoms with less sharing An incorrect SO4 Lewis structure 2- The case for expanding the octet of S 2- 2- O O O S O O S O O O S makes 4 bonds, no lone pairs, 2 more than expected 4 oxygens make one bond, 3 lone pairs = 4 atoms sharing less than expected Overall charge is -2, so only expect 2 atoms sharing less than expected Lecture 2 Part 1/6 Summary Lewis structures are used to represent shared electrons (line bonds) and unshared electrons (lone pairs) in molecules. The number of covalent bonds and lone pairs an element has depends on its valence electrons (group), such that it has a full octet. Polyatomic ions have more electrons (negative charge, atoms share fewer) or less electrons (positive charge, atoms share more) than neutral molecules. Some period 3 elements (P, S) have more than 8 valence electrons in compounds (phosphate, sulfate). The group is the maximum bond number.  BC 212 Lecture 2: Molecular Structure and Interactions with Water Part 2/6: 3D Shapes of Molecules Allison Lamanna Lecture 2 Part 2/6 Suggested Reading and Problems (8th ed.) Chapter 4, section 8 Chapter 4: 4.66, 4.67 Learning Goals What general 3D shapes can molecules have? Molecular shape is based on electron repulsion VSEPR: Valence Shell Electron Pair Repulsion Electrons are all negatively charged. Electron pairs (bonded or lone) will spread out around an atom to minimize repulsion of negative charge from other electron pairs. Shape description is relative to a central atom. Electrons in double or triple bonds are treated as a single electron group. CENTRAL ATOM GEOMETRY 2 bond groups 3 bond groups 4 bond groups 180o 120o 109.5o LINEAR TRIGONAL PLANAR TETRAHEDRAL-going back Cominga CO2 SO3 CH4 I H A T > - O O H O C O S O & H C H  Build it! CH4: TETRAHEDRAL H2CO: TRIGONAL PLANAR HCN: LINEAR 4 bond groups 3 bond groups 2 bond groups 109.5o 120o 180o H O H C H C N C H H H H  What is the central atom geometry of ammonia (NH3)? Cl Linear - H C H Trigonal O C NA H Planar N H Tetrahedral X H H T A bas -S B C lam Dair Lecture Question 2-2-1 Lecture 2 Part 2/6 Summary The three dimensional orientation of the electron clouds around a bonded atom can be described as linear (180o bond angles), trigonal planar (120o bond angles), or tetrahedral (109.5o bond angles).  BC 212 Lecture 2: Molecular Structure and Interactions with Water Part 3/6: Polarity Allison Lamanna Lecture 2 Part 3/6 Suggested Reading and Problems (8th ed.) Chapter 4, sections 9-10 Chapter 4: 4.30, 4.34, 4.74, 4.76 Learning Goals How do you describe bond polarity? How do you determine the overall polarity of a molecule? Definitions Polar: Has an uneven charge distribution Nonpolar: Electrons are evenly distributed Dipole: Difference in charge across a bond or molecule (i.e. positive end and negative end) -  Bond polarity (unequal sharing) Determine electronegativity difference between atoms 0.5 1.9 Nonpolar Polar Ionic covalent covalent Nonpolar covalent bond is < 0.5 Polar covalent bond from 1.9 to 0.5 Ionic bond > 1.9 DEN Nonpolar covalent: Electronegativities: CH4 0.4 Polar covalent: NH3 0.9 Ionic: NaCl 2.1  Dipole Uneven charge distribution across a bond (polar covalent) Atom with higher electronegativity has greater electron density and negative charge os Partin -  Which compound has ionic metal + metal bonds? non tract CO e Sub MgO 3 5-1 2 1 371 9 ⑧ -.. =.. H2O O None of the above. O O Lecture Question 2-3-1  Which compound has nonpolar covalent bonds? less than 5 0. count the H2O Don't * HF Subscript & justfle & NCl3 Singers None of the above. Lecture Question 2-3-2 Molecular Polarity ↳ Single O C O Pa =" s S make O O asymential O O S O F F S H H C H H Nonpolar Polar (Symmetric) always (Asymmetric) have to have polar pones 34 Molecular Polarity => & HO side more on ton o side more Imagine converting CH4 to NH3 Convert CH4 into NH3 (include the lone pair) Electronegativity difference: 3.0 (N) – 2.1 (H) = 0.9 3 bonds are polar; no polarity for lone pair position Which face of the molecule is positive and which is negative? Which way would the overall dipole point? NH3 Central atom Name of the geometry is shape (just tetrahedral the atoms) = PYRAMID Imagine converting NH3 to H2O Convert NH3 into H2O (include two lone pairs) Electronegativity difference: 3.5 (O) – 2.1 (H) = 1.4 2 bonds are polar; no polarity for lone pairs Which face of the molecule is positive and which is negative? Which way would the overall dipole point? H2 O Central atom Name of the geometry is shape (just tetrahedral the atoms) = BENT  Cl Cl O C Cl Cl H H Polar Nonpolar (Asymmetric) (Symmetric) - Molecular polarity pattern table IVA VA VIA VIIA VIIIA Group (14) (15) (16) (17) (18) Valence electrons 4 5 6 7 8 Expected bonds 4 3 2 1 0 (octet) Lone pairs 0 1 2 3 4 Example CH4 NH3 H2O HF Ne compound nonpolar H one e or H Molecular All the O H C H NO Polarity Same H H H F H H H Central Atom Geometry tetrahedral tetrahedral tetrahedral linear Molecular Symmetry symmetric - asymmetric asymmetric asymmetric ↑ When the central atom is not bonded to only one kind of atom… vs. CH2Cl2 H2CO C2H2 Polar Nonpolar (Asymmetric) & (Symmetric)  Which molecule is polar? I F F S F C F N C F F F H H bolar CF4 A H2CS B No Ponds- > notpolar O NF C 3 S 0. Lecture Question 2-3-3 Lecture 2 Part 3/6 Summary The polarity of a covalent bond is determined by the electronegativity difference between the atoms. The atom that holds the electrons more tightly (more electronegative) has a partial negative charge while the other atom has a partial positive charge, creating a dipole. The three dimensional orientation of bond dipoles determines the overall polarity of the molecule. Symmetric molecules tend to be nonpolar because bond dipoles cancel out.  BC 212 Lecture 2: Molecular Structure and Interactions with Water Part 4/6: Intermolecular Forces Allison Lamanna Lecture 2 Part 4/6 Suggested Reading and Problems (8th ed.) Chapter 8, sections 1-2 Chapter 8: 8.34ab, 8.35abce, 8.37 Learning Goals What are the properties of the states of matter, and which changes of state are exothermic vs. endothermic? What are the major kinds of molecular forces and their relative strengths? How do molecular shape and polarity determine intermolecular forces? How do intermolecular forces determine physical properties? Intermolecular forces from https://www.slideshare.net/angelatoh1/kinetic-particle-theoryppt Hold molecules of pure substance together in states of matter Molecules are most tightly held together in solid, then liquid, and have least interactions in gas The stronger the IMFs, the more energy (higher temperature) required to break the interactions and change states STATES OF MATTER Solid Liquid Gas - - - - Definitions OEndothermic: Thermal energy (heat) is absorbed or required for the process EExothermic: Thermal energy (heat) is released by - -the process Xchange Enthalpy (DH): O amount of thermal energy added to (+) or released (-) by the process Changing states of matter - ENDOTHERMIC: requires heat (thermal energy) to be added EXOTHERMIC: releases heat (thermal energy) ENTHALPY (DH): amount of thermal energy added (+) or released (-)  The Most is If yolrogen Bonding Intermolecular c M,. k War. Dipole-dipole interactions (polar molecules) 5 - St positive end is of one molecule is attracted to the negative end of another molecule  London dispersion forces (induced dipole-induced dipole) (nonpolar molecules) charged electron cloud becomes unevenly distributed influences surrounding molecules to form temporary - dipoles - more interacting surface area = stronger force avoid Makes attraction They'll neg-neg Repulsion > move to - ↳ but of Merahbor nts Increasing VIII Atomic 2 He V VI VII Size 4.0026 7 8 9 10 Increasing N O F Ne Dispersion 1 14.007 15.999 18.998 20.180 Force 15 16 17 18 P S Cl Ar 6 30.974 32.066 35.453 39.948 33 34 35 36 As Se Br Kr 74.922 78.96 79.904 83.80 51 52 53 54 Sb Te I Xe 1 121.75 127.60 126.90 131.29 to port More thermal energy 83 84 85 86 Bi Po At Rn Im forces an stronge 208.98 (209) (210) apart (222)  Cl2 Br2 l2 0Gas at OLiquid 0 at Solid at room temp. room temp. room temp. weakest < - - - - - - IMFs - - - - - - - > strongest Intermolecular  (NaCl) (H2O) (CO) (O2)  Which compound has the highest melting point? < Nonpolar I ~ Br2 O 2 8. O KBr 2 Tonic ↑ HBr Oof 0 2. O Lecture Question 2-4-1 Lecture 2 Part 4/6 Summary Exothermic changes of state release energy. Endothermic changes of state require energy input. Dipole-dipole interactions are stronger than London dispersion forces, but hydrogen bonds are stronger than dipole-dipole interactions. Molecular polarity determines the kind of intermolecular forces molecules have, which determines their properties. The amount of energy required to melt solids and boil liquids is determined by the strength of the intermolecular forces.  BC 212 Lecture 2: Molecular Structure and Interactions with Water Part 5/6: Hydrogen Bonding Allison Lamanna Lecture 2 Part 5/6 Suggested Reading and Problems (8th ed.) Chapter 8, sections 2 Chapter 8: 8.34c, 8.35df, 8.36 Learning Goals What is a hydrogen bond, and what makes it unique? What are hydrogen bond donors and acceptors? Hydrogen bonding H-bond donor - vollone Dipok par Internation oin ( ↓ H-bond acceptor Small, highly Generic H-bond: electronegative atoms —X : H—X— where “X” can be N, O, or F |||||||| - Hydrogen bonding Ethanol (CH3CH2OH) CH3 H 2C O H d- d+  Which molecule has hydrogen bonding as its intermolecular force? has * Any Molecule that HCl Hydrogen attached directly to * NH3 Oxygen/nition 00 - CH4 Has donors / com pairs CO2 Lecture Question 2-5-1 Formaldehyde (H2CO) and water H-bond acceptor No H-bond donor  Which molecule will hydrogen bond with water? HBr NH3 : C2H2 H2S Lecture Question 2-5-2  Which arrow is pointing to the hydrogen bond donor? O B H H C O A H H O C Nitro OX , attened*a Hydrogcor I always U cakedit off thats to Lecture Question 2-5-3 to. Lecture 2 Part 5/6 Summary Hydrogen bonds are the strongest of the intermolecular forces, because there is some orbital overlap. I - Hydrogen bonding requires an H bonded to a small, highly electronegative atom (O, N, F). Hydrogen bonds require a donor and an acceptor and are directional. The hydrogen bond donor has the H and the hydrogen bond acceptor has the lone pair.  BC 212 Lecture 2: This section Molecular Structure and provided as a Interactions with Water video with IVQs to complete Part 6/6: Solutions and their Properties Allison Lamanna Lecture 2 Part 6/6 Suggested Reading and Problems (8th ed.) Chapter 9, sections 1-3, 8 (no calculations), 10 (no osmotic pressure), 11 Chapter 9: 9.34, 9.68, 9.79, 9.80, 9.81a, 9.92a Learning Goals What is solubility, and how is impacted by intermolecular forces? How do you define the properties of a solution? What are electrolytes and osmosis? What are osmolarity and iso-, hyper-, and hypotonic solutions? How is dialysis similar to and different from osmosis? Definitions Mixture: Variable composition, either uniform (homogeneous) or not uniform (heterogeneous) Solution: Homogeneous mixture of particles the size of typical ions or small molecules Colloid: Homogeneous mixture of particles larger than a typical ion or small molecule Solute: Substance dissolved in a solvent. The solvent is typically present in larger amounts and the solution retains its properties/state. Solubility: Maximum amount of a substance that will dissolve at a given temperature and pressure Intermolecular forces determine solubility and vs. “like dissolves like”: intermolecular forces between solvent and solute must be similar enough in strength to overcome intra-substance interactions  Is a molecule polar? Does it have polar bonds? (electronegativity) What is the 3D shape? Are the polar bonds oriented in space such that they cancel each other out? (symmetry of shape) Polar molecules interact with other polar molecules Nonpolar molecules interact with other nonpolar molecules Hydration of ionic/polar solutes in water Ionic Polar Whole molecules hydrated Individual ions hydrated Water is polar, so interacts well with other polar substances Which molecule will NOT dissolve in water? - HBr - NH3 O CH4 NaBr - ↓ Definitions Soluble: Substance will dissolve significantly in solvent Insoluble: Substance will not dissolve in solvent Concentration: Amount of solute dissolved per unit solution Saturated: Contains the maximum amount of a solute that will dissolve at a given temperature and pressure Supersaturated: Contains more than the maximum amount of a solute that will dissolve at a given temperature and pressure Unsaturated: Contains less than the maximum amount of a solute that will dissolve at a given temperature and pressure Precipitation: Insoluble solid is removed from solution  A supersaturated solution https://youtu.be/FcxZ9DyOaUk  Electrolytes Substances that conduct electricity when dissolved in water Conductance depends on the concentration of ions in solution Substances that ionize or dissociate in water: (s) = solid NaCl (s) à Na+ (aq) + Cl- (aq) (aq) = aqueous or dissolved in water) https://www.youtube.com/watch?v=APUdc_dusI4 Electrolytes Strong electrolytes: Ionic compounds table salt: NaCl (s) à Na+ (aq) + Cl- (aq) Weak electrolytes partially ionize: ex. acids CH3COOH (s) ⇌ CH3COO- (aq) + H+ (aq) Nonelectrolytes do not ionize: polar covalent molecules H2CO (s) ⇌ H2CO (aq) sucrose: C12H22O11 (s) ⇌ C12H22O11 (aq)  Electrolytes and osmolarity Note: the more ions an ionic compound (salt) dissolved into, the greater the osmolarity Osmolarity: concentration of all dissolved particles in solution (can be multiple solutes, ions, etc.) NaCl (s) à Na+ (aq) + Cl- (aq) 2 ions per unit Al2(SO4)3 (s) à 2 Al3+ (aq) + 3 SO42- 5 ions per unit Assume solutions of the following binary compounds have the same concentration. Which has the greatest osmolarity? CaCO3 K2O MgSO4 AlPO4 Osmosis and Dialysis Both involve molecules passing through a semi- permeable membrane Osmosis: only solvent can pass through the membrane Dialysis: Solvent and some small molecules can pass through the membrane (depends on the size of the pore in the membrane); usually water and ions can pass through but larger biomolecules (i.e. proteins) cannot  Osmosis occurs across the membrane of red blood cells Normal plasma has an osmolarity of 0.3 (seawater ~ 1.0) Hypotonic solution Isotonic solution Hypertonic solution Lower osmolarity Similar osmolarity Higher osmolarity Solvent flows into cells Even exchange of solvent Solvent flows out of cells Q8 Cells are swollen Cells are normal Cells are shriveled How does seawater compare to normal plasma, and what would happen to red blood cells if they were suspended in seawater? Osmolarities: Plasma 0.3 and Seawater 1.0 Seawater is hypertonic, and RBCs will swell. Seawater is isotonic, and RBCs will be normal. Seawater is hypertonic, and RBCs will shrivel. Seawater is hypotonic, and RBCs will swell. Lecture 2 Part 6/6 Summary Solutions are homogenous mixtures. Aqueous solutions require that the solute have similar intermolecular forces to water (ions, polar molecules). Concentration defines the amount of solute in solution. Electrolytes are solutes that dissociate into ions in solution. Osmolarity is the concentration of all particles in solution. Semipermeable membranes can allow solvent (osmosis) or some solute (dialysis) particles to equilibrate between two solutions. Hypertonic solutions have greater osmolarity; hypotonic solutions have smaller osmolarity; isotonic solutions are the same.  ↑ Lecture 1 Part 4/6 Summary The periodic table is organized by the number ofmust metal valence electrons, giving rise to periodic trends: Ioniu T must iX metallic character, atomic size, and ionization Muster · - ↓ size BiggestAtomic energy (related to reactivity and electronegativity). - electricity : form into wire and sheets ↓ reactive Metals trunser heat and must ~ L Ionization energy and electronegativity define how hard it is for atoms to lose electrons. left to Right top bottom > Right toleft - to top bottom & lonization electronegativity · ↓ · & Metal size : Atomic · Reactivity  S BC 212 Lecture 3: Reactions of Molecules Part 1/6: Chemical Equations and Stoichiometry Allison Lamanna Lecture 3 Part 1/6 Suggested Reading and Problems (8th ed.) Chapter 5, sections 1-2 Chapter 5: 5.26, 5.28, 5.30, 5.70 Learning Goals How does the law of conservation of matter apply to chemical reactions? What is a balanced chemical reaction and how is it used? What is a mole (mol)?  Chemical Reactions Reactants Products Hydrogen + Oxygen à Water 12 2 H2 (g) + OO2 (g) à s 8 2 H2O (l) 4 gas gas liquid 4 Law of conservation of matter: Atoms are not created or destroyed in chemical reactions: balance on each side of “à” using coefficients (“=”) Note: Don’t balance by changing subscripts! H2O2 (hydrogen peroxide) is not the same substance as H2O (water) Ehanges the of compound Subscript changed Examples: Pb (s) solid; NaCl (aq) aqueous=dissolved in water  Balance the following reaction: i _Na 2 (s) + 2 _H2O (l) à 2 _NaOH (aq) + _H I 2 (g) is 2 - 2 4 2 H2O, 2 H2 2 Na, 2 NaOH 2 Na, 2 H2O, 2 NaOH The reaction is balanced as written. Lecture Question 3-1-1 How much material? 6.022 x 1023 atoms or molecules = 1 mole (mol) Also called Avogadro’s number (NA)  Same number of atoms or molecules: 1 mole (mol) 29 82 #ofg in Cu Pb Imole mo 63.55 O 207.2 - NaCl (table salt) = 58.4 g 63.55 g 207.2 g Aspirin (C9H8O4) = 180.2 g Note: Mass (g) is not the same, but the number of atoms or molecules is! Stoichiometry Use reaction coefficients to determine how much of a substance reacts or is produced, usually in moles Coefficients can be used to make ratios 00 2 H2 (g) + O2 (g) à 2 H2O (l) 2 "#$ %! 1 "#$O&! 1 "#$ &! 2 "#$ %! & - 2 "#$ %! & 2 "#$ %!  Stoichiometry (just for reference) 2 H2 (g) + O2 (g) à 2 H2O (l) 1 8 2 "#$ %! 2 "#$ %! & 1 "#$ &! 2 "#$ %! & 1 "#$ &! 2 "#$ %! H O 1.008 16.00 Note: can convert to g using molecular weight H2 = 2 g/mol O2 = 32 g/mol H2O = 18 g/mol Conservation of mass: both sides are 36 g 4 g H2 (g) + 32 g O2 (g) à 36 g H2O (l) You will not be asked to convert from g to mol or calculate molecular weights.  How many mol of O2 are required 1716 to fully react with 10 mol of H2? 2 ↓ 5 1 8 2 H2 (g) + O2 (g) à 2 H2O (l) H O 1.008 16.00 1 mol 5 mol motor = 5 mol on e 10 mol 2 you 20 mol wantch one 02 Jop Lecture Question 3-1-2 Lecture 3 Part 1/6 Summary The coefficients of balanced chemical reactions can be used to create stoichiometric ratios and calculate amounts of reactants and products by the law of conservation of matter. A mole (mol) is a specific number of atoms or molecules (6.022 x 1023) that is used for reactions on the gram (rather than atomic) scale. Definitions: Reactants: Substances that come together to react Products: Substances produced by the reaction  BC 212 Lecture 3: Reactions of Molecules Part 2/6: Redox Reactions Allison Lamanna Lecture 3 Part 2/6 Suggested Reading and Problems (8th ed.) Chapter 5, sections 5-6 Chapter 5: 5.51, 5.54, 5.56, 5.58, 5.60, 5.72 Learning Goals How do you define oxidation and reduction? What is a redox reaction? How do you determine the oxidation number of an atom in a compound? How do you determine if an atom is oxidized or reduced in a reaction?  Combustion Reactions & Redox - Methane + Oxygen à Carbon Dioxide + Water CH4 (g) + O2 (g) à 00 CO2 (g) + H2O (l) ↑ Balance using coefficients: 40 CH4 (g) + 2 O2 (g) à CO2 (g) + 2 H2O (l) 40 4H Combustion: “burn” in O2 For biological compounds, =produce CO2 and H2O Equivalent to metabolic reactions Fully “oxidize” carbon-containing compounds  Redox reactions Oxidation = loss of electrons (increase oxidation #) able to Form bond to highly electronegative O always Oxygen More electrons shared pull away ; always going has Lose bond to H (low electronegativity) of Super electropasitic o win Oxyge control Opposite Oxygen more ad opp Reduction = gain of electrons (decrease oxidation #) Lose bond to highly electronegative O Less electrons shared Gain bond to H (low electronegativity) if Someone guins Comes in pairs lost 17 If something is oxidized, something else must be Redox- bustcan someone Selections) so reduced! (Electrons must be accounted for)  Redox LEO says GER OIL RIG Lose Electrons Oxidation Oxidation Is Loss Gain Electrons Reduction Reduction Is Gain Will see lion , f As Rox http://animal-jam-clans.wikia.com/wiki/File:Lion-roar-front- https://s.hdnux.com/photos/20/04/75/4214654/9/ view-wallpaper-4.jpg 940x940.jpg 16  atomi c ona For poly Oxidation Numbers X six You anddow Assigned to determine elements’ propensity to lose or gain electrons in a chemical reaction Rules: Not Sharing Ome kind of Atom Ox # of uncombined atom or homoatomic = 0 Ox # of monatomic ion = its charge 1st column 2ndstumn Ox # of Group OIA atoms = +1, Group IIA = +2 Ox # of H = +1, O = -2 (except for peroxides) Sum of ox # = overall charge of compound If it's neutral > - its If rinvatimic > - adds up to charge  Oxidation numbers for carbon 08 ** S CO2 + 4- highe (+4) + 2× (-2) = 0 (+4) (-2) Has to be - 4 bas CH4 A IS +4 and 4 so +4 -y = 0 (-4) + 4× (+1) = 0 (-4) (+1)2 has tob--2 C6H12O6 ↓ Oxy (0) + 12× (+1) + 6× (-2) = 0 (0) (+1) (-2) bas H+ = 0 0 So 1 Note: Ox #s do not have to be integers. = Looking ahead: Metabolic reactions release energy by oxidizing carbon to CO2. The more reduced carbon is to start, the more oxidation can occur (and more energy produced).  F What is the oxidation state of carbon in steric acid, C18H36O2? (4) -1 7 - 18 + 36 1 78 2 + 33 0 - ⑫ -. -1 E18 0 181132 = - 32 182 Rules: = -2 - - 18 18 Ox # of uncombined atom or A non-integer ⑧ homoatomic = 0 between -1 and -2 Ox # of monatomic ion = its charge (-1 < ? < -2) Ox # of Group IA atoms = +1, Group IIA = +2 Ox # of H = +1, O = -2 (except for peroxides) Sum of ox # = overall charge of Lecture Question 3-2-1 compound. Redox Reactions CH4 (g) + 2 O2 (g) à CO2 (g) + 2 H2O (l) (-4)(+1) (0) (+4)(-2) (+1)(-2) Carbon is oxidized (-4 à +4) CH4 is reducing agent/reductant Oxygen is reduced (0 à -2) O2 is oxidizing agent/oxidant Classic combustion reaction (burn in oxygen) Core reaction of metabolism is oxidation of glucose C6H12O6 (s) + 6 O2 (g) à 6 CO2 (g) + 6 H2O (l) (0)(+1)(-2) (0) (+4)(-2) (+1)(-2)  elect oxi lose Red grin What is oxidized in the metabolism of lactic acid (C3H6O3)? What is the steo oxidant? pedi get Reduced = ⑤ C3H6O3 (aq) + (NAD)+ (aq) à C3H4O3 (aq) + (NAD)H (aq) + H+ (aq) = - O “NAD” is the O “NADH” has one more H O H abbreviation than NAD ⑭ H 3C C for a larger H 3C C C OH molecule C OH Oxidized (cofactor) OH O just lactic acid pyruvic acid I lactic acid, NAD+ NAD+, lactic acid lactic acid, lactic acid NAD+, NAD+ Lecture Question 3-2-2 Lecture 3 Part 2/6 Summary Reduction is the gain of electrons while oxidation is the loss of electrons. Redox reactions represent gain (red.) or loss (ox.) of electrons from an element in a reaction; these can be accounted for using oxidation numbers. Atoms that gain electrons in a reaction decrease in oxidation number while atoms that lose electrons increase in oxidation number. Definitions: Oxidizing agent (oxidant): Reactant responsible for oxidation, is reduced Reducing agent (reductant): Reactant responsible for reduction, is oxidized  BC 212 Lecture 3: Reactions of Molecules Part 3/6: Enthalpy of Reactions Allison Lamanna Lecture 3 Part 3/6 Suggested Reading and Problems (8th ed.) Chapter 7, section 3, p.186-7 (definitions) Chapter 7: 7.23, 7.26a, 7.27a Learning Goals What is enthalpy? How do you define exothermic vs. endothermic reactions? Enthalpy (DH) and reactions Enthalpy (DH) of reaction is thermal energy released or absorbed per mol (of coefficients) Combustion of glucose releases energy; energy is a product; exothermic reaction; negative DH C6H12O6 (s) + 6 O2 (g) à 6 CO2 (g) + 6 H2O (l) + 686 kcal/mol C6H12O6 (s) + 6 O2 (g) à 6 CO2 (g) + 6 H2O (l) DH = -686 kcal/mol energy Released being Enthalpy (DH) and reactions enter Exit - - ↑ niDH - nas DH Enthalpy (DH) is the difference between the chemical energy of products and reactants: Hproducts – Hreactants = DH  When NaCl dissolves in water, the solution decreases in temperature. Which set of terms below best heat in describes this process? sucking T enco Eexo-- Exothermic, negative DH Exothermic, positive DH Endothermic, negative DH Endothermic, positive DH O Lecture Question 3-3-1  Energy in food Cal = kcal; combustion used to determine Calories CxHyOz + O2 à CO2 + H2O Average Energy: Snickers: Carbohydrates ~ 4 kcal/g Carbohydrates = 35 g x 4 kcal/g = 140 kcal Fats ~ 9 kcal/g Fats = 14 g x 9 kcal/g = 126 kcal Proteins ~ 4 kcal/g Proteins = 4 g x 4 kcal/g = 16 kcal Total = 280 kcal Lecture 3 Part 3/6 Summary Enthalpy is the thermal energy released or absorbed in a reaction. An exothermic reaction releases energy and has a negative enthalpy. An endothermic reaction requires energy input and has a positive enthalpy.  BC 212 Lecture 3: Reactions of Molecules Part 4/6: Spontaneity and Thermodynamics Allison Lamanna Lecture 3 Part 4/6 Suggested Reading and Problems (8th ed.) Chapter 7, section 4-5 Chapter 7: 7.18, 7.20, 7.22, 7.32, 7.36, 7.38, 7.40 Learning Goals What is entropy and which processes increase or decrease entropy? When is a process spontaneous? What is free energy and how do you define exergonic vs. endergonic processes? Spontaneous reaction Once started, proceeds without external stimuli falling dropping * free cruser a Free Energy (DG) determines if a process is spontaneous; represents energy available (for “work”) Exergonic reaction - T ⑤ releases free energy (negative DG) - # ⑤ = Endergonic reaction requires free energy (positive DG) - g free energy Key biochemical question: Will a reaction release free energy (that can be used by the cell)? Is it spontaneous? Or, does a reaction require free energy (that must come from somewhere else in the cell)? Is it non- spontaneous? the direction free energy determines Note: If a reaction is non-spontaneous in the forward direction, then it is spontaneous in the reverse direction! (This is very important for reversible biochemical reactions!) DGforward = -(DGreverse) What is free energy (DG)? Energy available for a process Difference between free energy of products and reactants: Gproducts – Greactants = DG Depends on 3 things: Enthalpy (DH) Temperature (T) Entropy (DS)  What is entropy (DS)? Measure of disorder in a system Temperature dependent (increase temp, increase disorder) Low entropy High entropy mixed up organized All Increase Entropy = Positive DS Decrease Entropy -Negative DS organized Mixed Super disorganized Solid à Liquid à Gas Increasing disorder, positive DS Examples of increasing entropy (DS) & Change of state (s) à (l) à (g) More gas molecules in the products more entropy Reactants break apart into more, simpler particles O+ 6 O2 (g) à 6 CO2 (g) + 6 H2OO C6H12O6 (s) (l) More molecules Simpler molecules ②O (s) à (l) (gas molecules stay the same) entropy DS increase in  Select the process below that represents an increase in entropy: mag Gaseous water (steam) condenses into liquid water gas 8 Reaction of solid sodium with water: Solve 2 Na (s) O + 2 H2O (l) à 2 NaOH (aq) 00 + H2 (g) Synthesis of (ATP): (ADP) (aq) + H3PO4 (aq) à (ATP) (aq) + H2O (l) Lecture Question 3-4-1  Measuring free energy (DG) DG = DH – TDS Negative DG is spontaneous Negative DH favors spontaneity Positive DS (when –TDS is negative) favors spontaneity high startsthen Exergonic Endergonic Starts down low then goe Negative DG nigh Positive DG Whether or not a reaction will go forward (is spontaneous) depends on Enthalpy and Entropy (and Temperature) Entropy (Disorder) Free Energy Enthalpy (Heat) (Temperature-dependent) (Spontaneity) Spontaneous Heat released Increased disorder (Free Energy released) Non-spontaneous Heat absorbed Less disorder (Free Energy required) Depends on Temperature Heat released Less disorder (spontaneous at low T) Depends on Temperature Heat absorbed Increased disorder (spontaneous at high T) (Increasing Temperature (T) increases the importance (the impact) of Entropy)  Which of the following processes are spontaneous? Combustion: always heat released, always breaks - down into CO2 and H2O (more disorder) à- >always spontaneous not dependent ontemp Water melts: endothermic (heat absorbed), changes state from (s) à (l) so more disorder à sometimes spontaneous, only when temperature is high enough (above melting temperature) tomp dependent Reverse process is spontaneous when the forward process is not: at temperatures where melting is not spontaneous, freezing is spontaneous! -  NH3 is produced by the following reaction under 3 separate conditions (A, B, C). N2 (g) + 3 H2 (g) à 2 NH3 (g) Is this reaction exergonic or endergonic? down hill so ex goes Exergonic ⑧ A Endergonic B Neither C Energy Lecture Question 3-4-2 Reaction Progress  A general rule about the free energy of metabolic pathways Catabolic pathways are a Anabolic pathways are a series of reactions that series of reactions that break down molecules. build up molecules. They They usually have a net usually have a net RELEASE of free energy. REQUIREMENT of free energy. O'Connor, C. M. & Adams, J. U. Essentials of Cell Biology. Cambridge, MA: NPG Education, 2010. Coupled reactions An endergonic reaction is coupled to a more exergonic reaction, so that the sum of the reactions is exergonic Overall exergonia euohi , crall don hi e Aà C is overall exergonic  Coupled metabolic reactions Catabolic reactions break down biological molecules to release free energy Anabolic reactions build up biological molecules using up free energy Biochemical energy is captured in the molecule “ATP” Lecture 3 Part 4/6 Summary Entropy is the measure of disorder in a system or process. Processes that increase disorder have increased entropy. Free energy determines if a process is spontaneous and depends on enthalpy, entropy, and temperature. Exergonic processes release free energy, and endergonic processes require free energy input. Definition: Free energy (DG): Difference in available energy between the products and reactants; + if energy absorbed and – if energy released  BC 212 Lecture 3: Reactions of Molecules Part 5/6: Reaction Kinetics Allison Lamanna Lecture 3 Part 5/6 Suggested Reading and Problems (8th ed.) Chapter 7, section 5-6 Chapter 7: 7.21, 7.42, 7.46, 7.75 Learning Goals What is activation energy? How do you define reaction rate? How do temperature, concentration, and catalysts affect reaction rates? Activation energy Additional energy necessary for a reaction to occur (get started) bigger speed bump fast Slower How more > - energy (E ) Activation energy act O Reaction Rate How fast will a reaction occur? (Kinetics) - Measure change in concentration of reactants or products over change in time For a reaction to occur, particles must collide in the correct orientation and with enough energy - - Activation energy Determines the rate of the reaction Larger activation energy = slower rate Two sets of conditions for the reaction (blue and red): The red path has the smaller activation energy and the - faster rate - - bump Slower big Speed small spood bump faster 3 factors that affect reaction rate 1. Temperature – increasing temperature increases reaction rate (more collisions with more energy) - - - 2. Concentration – increasing concentration increases reaction rate (more collisions) - - 3. Catalyst : a substance that increases the rate of reaction but is not used up in the reaction. It lowers the activation energy for the reaction. - Inhibitor: slows down reaction rate - Effect of a catalyst ↳ Slow one Easter ↑ Catalyst is the tonnel Note: A catalyst lowers the activation energy for and increases the rate of both the forward and reverse reactions.  NH3 is produced by the following reaction under 3 separate conditions (A, B, C). N2 (g) + 3 H2 (g) à 2 NH3 (g) Which condition results in the fastest rate? A A B B C C on Energy The rate is the same for all. Lecture Question 3-5-1 Reaction Progress Lecture 3 Part 5/6 Summary The size of the activation barrier (activation energy) is related to the rate of a reaction. Reaction rate, or kinetics, is how fast a reaction occurs (change over time). Reaction rates can be increased by increasing temperature, concentration, or adding a catalyst. Definitions: Activation energy: Energy required to start any reaction; determines reaction rate Catalyst: Speeds up the reaction rate but is unchanged over the course of the reaction  BC 212 Lecture 3: Reactions of Molecules Part 6/6: Equilibrium Allison Lamanna Lecture 3 Part 6/6 Suggested Reading and Problems (8th ed.) Chapter 7, sections 7, 9, p.203 (summary graphic) Chapter 7: 7.60, 7.61, 7.63, 7.64, 7.65, 7.70 Learning Goals What is equilibrium and what is a reversible reaction? What kinds of conditions can change the direction of a reaction? Equilibrium When the rate of the forward reaction = the rate of the reverse => reaction Note: Rate is dependent on concentration and changes as the concentrations of reactants and products change Reversible reaction: can proceed in either the forward or reverse direction, depending on the conditions A+B!C+D Equilibrium concentrations Rate of forward reaction = Rate of reverse reaction the Reactant concentration DOES NOT EQUAL product only is Rate - concentration! equal Comparing reactant and product concentration tells you which direction is favored (has gone to a greater extent) Equilibrium constant Keq The ratio of product concentrations over reactant concentrations at a given temperature/pressure [%&'()*