OCR Physical Chemistry Past Paper PDF 2024-2025

Part I Physical Chemistry  Lecture 1 Matter and Energy I. Introduction Everything around you is composed of matter. Besides matter, energy is the other major component of our universe. II. Matt...

Part I Physical Chemistry  Lecture 1 Matter and Energy I. Introduction Everything around you is composed of matter. Besides matter, energy is the other major component of our universe. II. Matter Matter is anything that occupies space and has mass. Some matter is easy to see (water, wood), others are difficult (air, dust). The most basic building block of matter is the atom. What is Matter? Atoms, ions and molecules Anything that has mass and takes up space. Two forms: Element – distinctive building blocks of matter that make up every material substance Compound – two or more different elements held together by chemical bonds What is Matter? Organic compounds Compounds containing carbon atoms combined with each other and with atoms of one or more other elements such as hydrogen, oxygen, nitrogen, sulfur, phosphorus, chlorine, and fluorine. Inorganic compounds All compounds not classified as organic compounds. II. Atoms and Molecules atoms: submicroscopic particles that are the fundamental building blocks of all matter. Sometimes, atoms are bonded together to form molecules. molecules: two or more atoms joined to one another in specific geometric arrangements. II. Atomic and Molecular Matter II. Actual Images of Atoms and Molecules II. States of Matter Matter can be classified by its state. solid: closely-packed particles with fixed locations liquid: closely-packed particles, but free to move around gas: great distances between particles with free movement II. States of Matter II. The Solid State II. Properties of Different States II. Pure Substances and Mixtures Matter can be classified by its composition. pure substance: matter composed of only one type of atom or molecule mixture: matter composed of two or more different types of atoms or molecules which may vary in proportion II. Elements element: a pure substance that cannot be broken down into simpler substances  II. Compounds compound: a pure substance composed of two or more elements in fixed definite proportions. II. Mixtures Most matter exists in this form. heterogeneous: varied composition from one region to another homogeneous: uniform composition throughout II. Classification by Composition II. Sample Problem Classify the following as a pure substance or mixture. Further classify them as an element, compound, homogeneous, or heterogeneous. a) blood b) sugar c) mercury in a thermometer d) chicken noodle soup III. Distinguishing Matter We use physical and chemical properties to tell the difference between samples of matter. physical property: a property a substance displays without changing its composition chemical property: a property a substance displays only by changing its composition III. Boiling Point of Water At the boiling point, water is converted to steam, but steam is just a different form of water. III. An Iron Nail Rusts When iron rusts, it must react and incorporate oxygen to become a new compound. III. Physical/Chemical Changes Physical/chemical changes are closely related to definitions of physical/chemical properties. physical change: matter changes its appearance, but not its composition chemical change: matter changes its composition Chemical changes occur through chemical reactions in which reactants become products. III. Physical/Chemical Changes IV. There is No New Matter In ordinary chemical reactions, matter is neither created nor destroyed. Known as Conservation of Mass. The Law of Conservation of Matter Matter is not destroyed It only changes form There is no “away” – atoms are not destroyed, just rearranged. What are some examples of matter changing form?  V. Energy Physical and chemical changes are accompanied by energy changes. energy: the capacity to do work work: results from a force acting on a distance V. Two Types of Energy potential energy (PE): energy due to the position or composition of the object kinetic energy (KE): energy due to motion of the object An object’s total energy is the sum of its PE and KE Energy Kinetic Potential Wind Water behind a dam Electicity Gasoline in your car Flowing water Unlit match V. Energy Conversions The Law of Conservation of Energy states that energy is neither created nor destroyed. Energy can change from one form to another or transferred from one object to another. V. Specific Types of Energy Electrical energy is the energy associated with the flow of electrical charge. Thermal energy is the energy associated with motions of particles of matter. Chemical energy is a form of PE associated with positions of particles in a chemical system. V. Energy Unit Conversions There are three common units for energy. V. Sample Problem The complete combustion of a wooden match produces about 512 cal of heat. How many kilojoules are produced? V. System and Surroundings When describing energy changes, we need reference points. system: object of study surroundings: everything else Systems with high PE tend to change such that their PE is lowered.  V. Energy Diagrams Chemical reactions can either be exothermic or endothermic. exothermic: release energy to surroundings endothermic: absorb energy from surroundings V. Sample Problem Identify the following changes as exothermic or endothermic. a) Water freezing into ice. b) Propane burning. c) Isopropyl alcohol evaporating from skin. VI. Thermal Energy Atoms and molecules of matter are in constant, random motion, which is the source of thermal energy. More motion = more thermal energy. Is there a way to easily measure this motion? VI. Temperature and Heat Temperature is the measure of the thermal energy of a substance. The hotter an object, the greater the motion of its particles, and the greater the thermal energy. Heat is the transfer or exchange of thermal energy caused by a temperature difference. VI. Temperature Scales VI. Temperature Conversions The formulas below allow conversion between different temperature units. VI. Sample Problem Convert 67 °F to kelvin and degrees Celsius. VII. Heating a Substance When you heat a substance, its temperature changes. The amount of change depends on the substance. heat capacity: quantity of heat needed to raise the temp of substance by 1 °C specific heat capacity: quantity of heat needed to raise temp of 1 g of substance by 1 °C VII. Specific Heat Capacities VII. Energy and Heat Capacity Heat absorbed and temperature change are directly related as shown in the equation below. VII. Sample Problem Calculate the heat necessary to warm a 3.10 g sample of copper from -5.0 °C to 37.0 °C if the specific heat capacity of copper is 0.385 J/g °C. VII. Sample Problem A sample of lead (C = 0.128 J/g °C) absorbs 11.3 J of heat, rising in temperature from 26 °C to 38 °C. Find the mass of the sample in grams. VII. Sample Problem A 328-g sample of water absorbs 5.78 kJ of heat. If the water sample has an initial temperature of 35.3 °C, what will be its final temperature? Note that C = 4.18 J/g °C for water.  Second Law of Thermodynamics In every transformation, some energy is converted to heat You cannot break even in terms of energy quality Waste energy is low quality and cannot be reused Second Law of Thermodynamics What are some other examples of the Second Law of Thermodynamics? Water is heated due to energy loss from the flowing water a 20-25% of the chemical energy in gasoline is converted to mechanical energy. The rest is lost into the environment as low quality heat energy. 5% of electricity is changed into useful light. 95% is lost as low-quality heat.  Photosynthesis is the process of converting solar energy into chemical energy stored in food CO2 + H20 ---> C6H12O6 + O2  Respiration is the process of releasing chemical energy stored in food to be used by living things. C6H12O6 + O2 ---> CO2 + H20 THE SI system and units of measurement Units of Measurement Measurement Measurement is used to measure quantities. Quantity is something that has magnitude, size, or amount (volume). In the late 18th century, scientists used the metric system. The metric system is a precursor to the SI System. Scientists all over the world use a single measurement system called Le Systeme International d’Units, abbreviated SI. SI system of measurement It is based on the metre-kilogram-second system and replaces both the foot-pound-second system and the centimetre-gram-second system. There are seven SI base units, i.e. metre, kilogram, second, mole, ampere, Kelvin and candela. SI Base units The most common base units that we will study include: Quantity Quantity Symbol Unit name Unit abbreviation Length l meter m Mass m kilogram kg Time t second s Temperature T Kelvin K SI Base units To enable the measurement of quantities larger or smaller than the base units or derived units, the SI Units system also includes a set of prefixes. The use of a prefix makes a unit larger or smaller. The ranges of SI unit prefixes are listed in the tables 1 and 2 Table 1 Prefix Symbol Function Divided by deci d 10-1 10 centi c 10-2 100 milli m 10-3 1000 micro µ 10-6 1000000 nano n 10-9 1000000000 pico p 10-12 1000000000000 femto f 10-15 1000000000000000  Table 2 Prefix Symbol Function Multiply by deca da 101 10 hector h 102 100 kilo k 103 1000 mega M 106 1000000 Giga G 109 1000000000 tera T 1012 1000000000000 peta P 1015 1000000000000000 Length The SI standard unit for length is the meter. A distance of 1m is about the width of an average doorway. To express longer distances, the kilometer (km) is used. One kilometer is equal to 1000 m. To express shorter distances, the centimeter (cm) is used. One centimeter is equal to 1/100 of a meter. Length can be measured using a meter stick or rulers. Mass Mass is a measure of the quantity of matter. The standard unit for mass is the kilogram (kg). The gram (g), which is 1/1000 of a kg is used for measuring masses of small objects. Mass is typically measured using a balance. Mass is different from weight. Weight is a measure of the gravitational pull on matter (Newton). The weight of an object increases as gravity acts on it. Time The standard unit of measurement for time is the second (s). Time can be measured using stop watches, clocks, count down timers, and other time pieces. Larger amounts of time are measured in minutes and hours. There are 60 seconds in one minute. There are 60 minutes in one hour. Given that there are 24 hours in one day, how many seconds are there in one day? Temperature The standard unit of measurement for temperature is degrees Kelvin (K). Temperature can also be measured in degrees Celsius (°C) and degrees Fahrenheit (°F). To convert degrees Celsius (°C) to degrees Fahrenheit (°F) multiply by 1.8 and then add 32. To convert degrees Fahrenheit to degrees Celsius, subtract 32 and then divide by 1.8 ° K= °C + 273 Temperature is measured using a thermometer. (measures the degree of heat or coolness) Derived units Derived units are combinations of base units. They are produced by multiplying or dividing standard units. The derived units we will study include: Quantity Unit Quantity Unit symbol abbreviation Area A square meter m2 Volume V cubic meter m3 kilograms per Density D kg/m3 cubic meter  Area Area is length times the width. It is expressed as square meters. Area can also be expressed as cubic centimeters. What is the area of a rectangle that has an a length of 6 cm and a width of 16 cm? Volume Volume is the amount of space occupied by an object. The derived SI unit for volume is cubic meters. The cubic meter is rather large, so a more common unit of cubic centimeters is more commonly used. Non-SI units are also used to measure volume such as the milliliter (mL) and the liter (L), which is 1000 cm3. There are 1000 mL in 1 L. Beakers, flasks, and graduated cylinders are often used to determine the volume of liquids. density Density is the ratio of mass to volume, or mass divided by volume. It can be written: density=mass/volume or D=m/V Density is a characteristic physical property of a substance that does not depend on the size of the sample. As the mass of an object increases, its volume increases.  Lecture 2 Atoms, Elements, Molecules, Ions, and Compounds 70  Atomic Theory of Matter Protons, neutrons, electrons Postulates of Dalton’s Atomic Theory Later found indivisible 1.)All matter is composed of indivisible atoms. An atom is an extremely small to be untrue. particle of matter that retains its identity during chemical reactions. 2.)An element is a type of matter composed of only one kind of atom, each atom of a given element having the same properties. Mass is one such property. Thus the atoms of a given element have a characteristic mass. Later found all atoms of the same element does not have to have the same mass. Atoms have different isotopes that have the same # protons but different # neutrons and hence different 71 mass. Note: #protons gives identity of atom. Atomic Theory of Matter Postulates of Dalton’s Atomic Theory 3.) A compound is a type of matter composed of atoms of two or more elements chemically combined in fixed proportions. – The relative numbers of any two kinds of atoms in a compound occur in simple ratios. – Water, for example, consists of hydrogen and oxygen in a 2 to 1 ratio (2H: 1O) for all molecules of water. 72 Atomic Theory of Matter i.e., solid sodium mixed with chlorine Postulates of Dalton’s Atomic Theory gas forms a new substance, salt, with 4.) A chemical reaction consists of the totally different rearrangements of the atoms present in the properties from the starting materials. reacting substances to give new chemical combinations present in the substances formed by the reaction (new chemical with different properties). 2 Na (s) + Cl2 (g)  2 NaCl (s) Once again, later 5.) Atoms are not created, destroyed, or broken found indivisible into smaller particles by any chemical reaction. to be untrue. Protons, neutrons, electrons 73  Law of Constant Composition Joseph Proust (1754–1826) Also known as the law of definite proportions. The elemental composition of a pure substance never varies. The relative amounts of each element in a compound doesn’t vary. H N NH3 ammonia ammonia always has 3 H and 1 N. Law of Conservation of Mass The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place. 3H2 + N2 2NH3 ammonia The atoms on the right all appear on the left Atomic Theory of Matter The Structure of the Atom – Although Dalton postulated that atoms were indivisible, experiments at the beginning of the 1900’s showed that atoms themselves consist of particles. – Experiments by Ernest Rutherford in 1910 showed that the atom was mostly “empty space.” 76 Atomic Theory of Matter – These experiments showed that the atom consists of two kinds of particles: a nucleus, the atom’s central core, which is positively charged and contains most of the atom’s mass, and one or more electrons. – Electrons are very light, negatively charged particles that exist in the region around the atom’s positively charged nucleus. Nucleus (+) - e 77  Atomic Theory of Matter The nuclear model of the atom. – Rutherford’s famous gold leaf experiment gave credibility to the theory that the majority of the mass of the atom was concentrated in a very small nucleus. – Positively charged alpha particles were directed at a metal foil. Only 1/8000 were deflected indicating that the nucleus was extremely small and positively charged. Only those alpha particles that directly hit the nucleus were deflected; the rest passed through. 78 Radioactivity Three types of radiation were discovered by Ernest Rutherford:  particles, attracted to negative electrode, so they have a positive charge, much more mass than negative stuff (turn out to be He nuclei)  particles, attracted to positive electrode, so they have a negative charge, 1000s of times less massive (turn out to be electrons coming from nucleus).  rays, no charge, no mass, Discovery of the Nucleus Ernest Rutherford shot  particles at a thin sheet of gold foil and observed the pattern of scatter of the particles. The Nuclear Atom Virtually all the particles went straight through Most of the atom essentially empty A few particles deflected, some straight back. A very small part of the atom is very dense, impenetrable. The mass must be concentrated. Atomic Theory of Matter – Most of the mass of an atom is in the nucleus; however, the nucleus occupies only a very small portion of the space in the atom. – The nucleus of an atom is composed of two different kinds of particles: protons (+) and neutrons (neutral). – An important property of the nucleus is its positive electric charge. 82 The Nuclear Atom Rutherford postulated a very small, dense nucleus with the negative electrons around the outside of the atom. Most of the volume of the atom is empty space.  Subatomic Particles Protons and electrons are the only particles that have a charge. Protons and neutrons have essentially the same mass. The mass of an electron is so small we ignore it.  Atomic Theory of Matter – A proton is the nuclear particle having a positive charge equal to that of the electron’s (a “unit” charge) and a mass more than 1800 times that of the electron. It is for this reason that we refer to H as a pure proton. – The number of protons in the nucleus of an atom is referred to as its atomic number (Z) and gives the identity of an element. All species that have same #p have the same properties. H Z=1 1p, 1e- neutral species: #p = #e- Na Z=11 11p, 11e- Cl Z=17 17p, 17e- - Na+ Z=11 11p, 10e- Cl Z=17 17p, 18e- #p remains constant in species; + charge, more p than e- #e- can vary and dictates the charge of species - charge, more e- than p Valence Electrons & Bohr Diagrams Atomic Structure Atoms have a nucleus that contains Protons and Neutrons Electrons are contained in shells that surround the nucleus An atom is made of mostly empty space Protons have a positive charge Electrons have a negative charge Neutrons are Neutral Valence Electrons Each electron shell can hold a Electron Number of certain number of electrons Shell Electrons Electron shells are filled from the inside out 1 2 Noble Gases have full outer 2 8 electron shells All other elements have 3 8 partially filled outer electron 4 18 shells 5 18 6 32 7 32 Valence Electrons The electrons in the outer most electron shell are called valence electrons The shell containing electrons that is furthest from the nucleus is called the valence shell The number of electron shells with electrons is the same as the period number Noble Gas Stability Noble gases are usually unreactive This is because they have full valence shells An element with a full valence shell is a happy element  For two atoms to join together atoms must gain, lose or share electrons Elements with full valence shells do not easily gain or lose electrons Noble Gas Stability Atoms want to gain stability Atoms will try to gain or lose electrons to have a full valence shell Metals try to lose electrons Non-Metals try to gain electrons Becoming An Ion Electrons are negatively charged Protons are positively charged Neutral atoms do not have a charge because the number of protons is the same as the number of electrons When atoms gain or lose electrons they become positively or negatively charged An atom with a charge is called an Ion Bohr Models Niels Bohr created a visual model of the atom to make them easy to understand A Bohr Model contains a central nucleus surrounded by electron shells For each model you state the number of protons and neutrons in the nucleus and draw a dot on the electron shells for each electron Atomic Theory of Matter – An element is a substance whose atoms all have the same atomic number (Z). The #protons defines the identity of an atom and can be found on the periodic table (large number in top of element box). – The neutron is a nuclear particle having a mass almost identical to that of a proton, but no electric charge. The charge of the nucleus comes from the #protons. The atoms may have different masses because of different #neutrons (isotopes). – Summary of masses and charges of the three fundamental particles: 94  The mass number (A) is the total number of protons and neutrons in a nucleus. A = #p + #n = Z + #n How many neutrons does sodium 23 have? A = 23, Z = 11 (number on periodic table) A = Z + #n 23 = 11 + #n #n = 23 - 11 = 12 A nuclide is an atom characterized by a definite atomic number and mass number. The shorthand notation for a nuclide consists of its symbol with the atomic number, Z, as a subscript on the left and its mass number, A, as a superscript on the left. A Z E sodium  23 23 23 + 11 Na 11 Na 11p, 12 n, 11e- 11p, 12 n, 10e- 95 Symbols of Elements Elements are symbolized by one or two letters. Atomic Number All atoms of the same element have the same number of protons: The atomic number (Z) Atomic Mass The mass of an atom in atomic mass units (amu) is approximately the total number of protons and neutrons in the atom. What is the nuclide symbol for a nucleus that consists of 17 protons, 18 neutrons, and 17 electrons? What’s the element? 17 p  atomic number on periodic table for chorine A = #p + #n = 17 + 18 = 35 35 17Cl Note: #e- = #p; therefore, neutral species How many protons, neutrons, and electrons are in the following nucleus 80 35 Br A = #p + #n 80 = 35 + #n #n = 80 – 35 = 45 n 99  Atomic Theory of Matter – Isotopes are atoms whose nuclei have the same atomic number (Z) but different mass numbers (A); that is, the nuclei have the same number of protons but different numbers of neutrons thereby causing them to have different masses. – Chlorine, for example, exists as two isotopes: chlorine-35 and chlorine-37. 37 Frac abund = 24.229% 35 17 Cl Frac abund = 75.771% Mass = 34.97 amu 17 Cl Mass = 36.97 amu 17p, 18 n, 17e- 17p, 20 n, 17e- – The fractional abundance is the fraction of a sample of atoms that is composed of a particular isotope. (0.75771)(34.97 amu) + (0.24229) (36.97 amu) = 35.45 amu Note: The mixture of isotope masses make up the actual mass of the element given on periodic 100 table. Cl has a mass of 35.45 amu which is based on the two isotopes of Cl-35 and Cl-37. Atomic Weights Calculate the atomic weight of boron, B, from the following data: ISOTOPE ISOTOPIC MASS (amu) FRACTIONAL ABUNDANCE B-10 10.013 0.1978 (19.78%) B-11 11.009 0.8022 (80.22%) Note: fractional abundances must add to 1 (100%) B-10: 10.013 amu x 0.1978 = 1.980 B-11: 11.009 amu x 0.8022 = 8.831 10.811 = 10.811 amu ( = atomic wt.) Note: mass on periodic table matches 10.811 amu (weighted average of isotopes) 101 Shapes of Atomic Orbitals for Electrons Four different kinds of orbitals for electrons based on those derived for a hydrogen atom Denoted s, p, d, and f s and p orbitals most important in organic and biological chemistry s orbitals: spherical, nucleus at center p orbitals: dumbbell-shaped, nucleus at middle d orbitals: elongated dumbbell-shaped, nucleus at center Orbitals and Shells part 1 Orbitals are grouped in shells of increasing size and energy Different shells contain different numbers and kinds of orbitals Each orbital can be occupied by two electrons Orbitals and Shells part 2 First shell contains one s orbital, denoted 1s, holds only two electrons Second shell contains one s orbital (2s) and three p orbitals (2p), eight electrons Third shell contains an s orbital (3s), three p orbitals (3p), and five d orbitals (3d), 18 electrons Atomic Structure: Electron Configurations Ground-state electron configuration (lowest energy arrangement) of an atom lists orbitals occupied by its electrons. Rules: 1. Lowest-energy orbitals fill first: 1s 2s 2p 3s 3p 4s 3d (Aufbau (“build-up”) principle) 2. Electrons act as if they were spinning around an axis. Electron spin can have only two orientations, up  and down . Only two electrons can occupy an orbital, and they must be of opposite spin (Pauli exclusion principle) to have unique wave equations 3. If two or more empty orbitals of equal energy are available, electrons occupy each with spins parallel until all orbitals have one electron (Hund's rule).  Writing Atomic Electron Configurations Two ways of writing configs. One is called the spdf spdf notation notation. for H, atomic number = 1 1 no. of 1s electrons value of l value of n  Writing Atomic Electron Configurations Two ways of writing configs. ORBITAL BOX NOTATION Other is called for He, atomic number = 2 the orbital box Arrows 2 1s notation. depict electron 1s spin One electron has n = 1, l = 0, ml = 0, ms = + 1/2 Other electron has n = 1, l = 0, ml = 0, ms = - 1/2 The Periodic Table In 1869, Dmitri Mendeleev discovered that if the known elements were arranged in order of atomic mass (A), they could be placed in horizontal rows such that the elements in the vertical columns had similar properties. – periodic table - tabular arrangement of elements in rows and columns, highlighting the regular repetition of properties of the elements. – periodic law – states that certain sets of physical and chemical properties recur at regular intervals (periodically) when the elements are arranged according to increasing atomic number (Z). – Note: eventually changed from atomic mass to atomic number because of a couple of anomalies. Figure: A modern form of the periodic table. anomalies The Periodic Table Periods and Groups – A period consists of the elements in any one horizontal row of the periodic table. – A group consists of the elements in any one column of the periodic table (similar properties/structure). – The groups are usually numbered (North American uses roman numbers and A/B; IUPAC 1-18). – The eight “A” groups are called main group (or representative) elements. The Periodic Table Periods and Groups – The “B” groups are called transition elements. – The two rows of elements at the bottom of the table are called inner transition elements. – Elements in any one group have similar properties because their outer shells have the same number of valence electron (discuss in later sections).  The Periodic Table Periods and Groups – The elements in group IA (except H) - alkali metals – The elements in group IIA - alkaline earth metals, – The group VIIA elements - halogens – The group VIIIA elements – noble gases (monoatomic) – Diatomic elements – H2, N2, O2, F2, Cl2, Br2, I2 – Most species are solids at room temperature; H2, N2, O2, F2, Cl2, and noble gases are gases; Br2 and Hg are liquids.  Metals, Nonmetals, and Metalloids – generally, left of staircase are metals, touching staircase are metalloids, right of staircase are nonmetals. This is important for determining bond type, using proper terminology, and making decisions. Metallic character Metallic character nonmetals metals  Elements of life Elements required for living organisms (pretty much all organisms). Red, most abundant blue, next most abundant Green, trace amounts.  Lecture 3 Chemical Formulas Chemical Reaction Reaction Rates and Equilibrium Chemical Formulas Molecular and Ionic Substances The chemical formula of a substance is a notation using atomic symbols with subscripts to convey the relative proportions of atoms of the different elements in a substance. – aluminum oxide, Al2O3 2Al:3O ratio – sodium chloride, NaCl 1Na:1Cl ratio – calcium nitrate, Ca(NO3)2 1Ca:2NO3- ratio or 1Ca:2N:6O ratio 118  Chemical Formulas Molecular and Ionic Substances involves ionic bond – transfer electrons between C:C atoms – attraction between charged particles – involves covalent bond – share electrons typically metal/nonmetal or polyatomic ions + - Na Cl between atoms – typically nonmetal/nonmetal Molecular substances – A molecule is a definite group of atoms that are chemically bonded together through sharing of electrons (covalent bonding, generally nonmetal-nonmetal including H). – A molecular substance is a substance that is composed of molecules, all of which are alike. – A molecular formula gives the exact number of atoms of elements in a molecule (i.e. C2H6O). – Structural formulas show how the atoms are bonded to one another in a molecule. i.e. ethanol (C2H6O) has a structural formula of CH3CH2OH 119 Ionic substances – Although many substances are molecular, others are composed of ions (charged particles) that have transferred electrons and have ionic bonding; occurs generally with metal-nonmetal interactions. – An ion is an electrically charged particle obtained from an atom or chemically bonded group of atoms by adding or removing electrons. Na +  1e-  Cl - – Sodium chloride is a substance made up of ions. Na+ Cl- 120 Ionic substances – When an atom gains extra electrons, it becomes a negatively charged ion, called an anion (more electrons than protons). i.e, Cl- – An atom that loses electrons becomes a positively charged ion, called a cation (more protons than electrons). i.e., Na+ – An ionic compound is a compound composed of cations and anions. NaCl CaBr2 Na2SO4 CO2 121 Ions in Aqueous Solution Many (not all) ionic compounds (ionic bond/m-nm) dissociate into independent ions when dissolved in water NaCl (s)  Na+(aq) + Cl-(aq) Soluble salt charges particles Soluble ionic compounds dissociate 100% - referred to as strong electrolytes – breaks into charged particles until reaches saturation point. 122 Molecules in Aqueous Solution Most molecular (covalent bond) compounds dissolve but do not dissociate into ions, exception acids. C6H12O6 (s)  C6H12O6 (aq) no charges particles; remains whole These compounds are referred to as nonelectrolytes; no charged particles; soluble to saturation point but no ions formed. How would sodium sulfate dissolve based on bonding? Na2SO4 (s)  2Na+(aq) + SO42-(aq) ionic bond dissociates while covalent bonds in sulfate remain intact 123 Chemical Substances; Formulas and Names Ionic compounds – Most ionic compounds contain metal and nonmetal atoms (as well as polyatomic ions); for example, NaCl. – You name an ionic compound by giving the name of the cation followed by the name of the anion with -ide. Sodium chloride, NaCl Calcium Iodide, CaI2 Potassium Bromide, KBr – We give the monatomic ion name for the cations and anions when naming compounds. A monatomic ion is an ion formed from a single atom. 124 How do we get the charge for ions? – Most of the main group metals form cations with the charge equal to their roman group number. – The charge on a monatomic anion for a nonmetal equals the roman group number minus 8. – Most transition elements form more than one ion, each with a different charge (exceptions Cd2+, Zn2+, Ag+). – Other important elements with variable charge Pb4+, Pb2+ Sn4+, Sn2+ As5+, As3+ Sb5+, Sb3+ 1+ 4- 0 2+ 3+ 4+ 3- 2- 1- varies 125  Rules for naming monatomic ions – Monatomic cations are named after the element. For example, Al3+ is called the aluminum ion. – If there is more than one cation of an element (charge), a Roman numeral in parentheses denoting the charge on the ion is used. This often occurs with transition elements. Na+ sodium ion Ca2+ calcium ion Fe2+ iron (II) ion Fe3+ iron (III) ion Older name: higher ox state (charge) – ic, / lower, -ous Fe3+ ferric ion Fe2+ ferrous ion Cu2+ cupric ion Cu+ cuprous ion Hg2+ mercuric ion Hg22+ mercurous ion For the names of the monatomic anions, use the stem name of the element followed by the suffix – ide. For - example bromine, the anion is called bromide ion, Br. 126 The formula of an ionic compound is written by giving the smallest possible whole-number ratio of different ions in the substance. ions and charges formula Sodium chloride Na+ Cl- NaCl + 1Na = 1+ charge - 1Cl = 1- charge Based on the balanced charge of the ions Iron (III) sulfate 3+ Fe SO4 2- Fe2(SO4)3 and balancing the Roman number tells charge of transition 3+ 2Fe = 6+ charge 4 2- 3SO = 6- charge overall charge on metal balanced the compound by Chromium (III) oxide Cr O 3+ 2- Cr2O3 adjusting the number of ions, a formula is written. Calcium nitrate Ca2+ NO3- Ca(NO3)2 Note the sum of all the charges Generally, you Sodium phosphate Na+ PO43- must equal zero, can crisscross Na3PO4 and you do of the charge not display one ion the as the Strontium oxide 2+ Sr O 2- charges subscriptinonthethe SrO final formula. second ion, reducing when possible. 127 Naming Ionic Binary Compounds NaF - sodium fluoride LiCl - lithium chloride MgO - magnesium oxide manganese (II) bromide MnBr2 - CuCl2 - copper (II) chloride or cupric chloride 128 Chemical Substances; Formulas and Names Polyatomic ions – A polyatomic ion is an ion consisting of two or more atoms chemically bonded together and carrying a net electric charge. We name the compounds the same way we just discussed except each polyatomic ion has a particular name. – Books typically have a table that lists common polyatomic ions. Most are oxo anions – consists of oxygen with another element (central element). NO3- nitrate SO42- sulfate NO2- nitrite SO32- sulfite Most groups have –ate, -ite endings and differ by #O. Mn, Br, Cl, I have per- -ate, -ate, -ite, hypo- -ite. 129 Ions You Should Know O22- - Peroxide Polyatomic ions PO43- - Phosphate NH4+ - Ammonium OH- - Hydroxide PO33- - Phosphite CN- - Cyanide CO32- - Carbonate SO42- - Sulfate HCO3- - Bicarbonate or Hydrogen Carbonate SO32- - Sulfite ClO4- - perchlorate N3- - azide ClO3- - chlorate NO3- - nitrate ClO2- - chlorite NO2- - nitrite - ClO- - hypochlorite C2H3O2- or CH3COO - acetate Hg22+ - mercury (I) or mecurous Cr2O72- - dichromate S2O32- - thiosulfate CrO42- - chromate SCN- - thiocyanate C2O42- - oxalate CNO- - cyanate HSO4- - bisulfate or hydrogen MnO4- - permanganate sulfate H2PO4- - dihydrogen phosphate 130 Ca(ClO)2 calcium hypochlorite Ba(OH)2 barium hydroxide Cr2(SO4)3 chromium (III) sulfate potassium perchlorate KClO4 Fe3(PO4)2 iron (II) phosphate or ferrous phosphate 131  Chemical Substances; Formulas and Names Molecular compounds – Binary compounds composed of two nonmetals are usually molecular and are named using a prefix system (name same as ionic except must indicate how many atoms are present using mono, di, tri, etc.). No charges (share electrons) involved with molecular compounds, but we typically put more metallic compound first. Which way is the correct way to write the following formula based on putting the more metallic compound first? NF3 F3N 132  Chemical Substances; Formulas and Names Binary molecular compounds – The name of the compound has the elements in the order given in the formula. – You name the first element using the exact element name. – Name the second element by writing the stem name of the element with the suffix “–ide.” – If there is more than one atom of any given element, you add a prefix (di, tri, tetra, penta, hexa, hepta, octa, etc.) 133  Binary molecular compounds – N2O3 dinitrogen trioxide – SF4 sulfur tetrafluoride – ClO2 chlorine dioxide – SF sulfur hexafluoride 6 – Cl2O7 dichlorine heptoxide – HCl (g) hydrogen chloride Since this is a gas, we name using molecular rules; however, if acid we have other rules. Name this compound but think about bonding: MgCl2 magnesium chloride; ionic comp, no prefix Older names: water - H2O, ammonia – NH3, hydrogen sulfide – H2S, nitric oxide – NO, hydrazine – N2H4 134  Chemical Substances; Formulas and Names Acids – Acids are traditionally defined as compounds with a potential H+ as the cation. – Binary acids consist of a hydrogen ion and any single anion in aqueous solution. For example, HCl (aq) is hydrochloric acid. Binary acid: hydrostemic acid – An oxoacid is an acid containing hydrogen, oxygen, and another element. An example is HNO3, nitric acid. The oxoacids are a derivation of the oxoanions we discussed earlier. 135  oxoacids Anion prefix/suffix acid prefix/suffic per- -ate ion per- -ic acid If you learn the oxoanions, you can -ate ion -ic acid easily adapt to naming the oxoacids: -ite ion -ous acid -ate  -ic and –ite  -ous hypo- -ite ion hypo- -ous acid NO3- nitrate ion HNO3 nitric acid NO2- nitrite ion HNO2 nitrous acid ClO4- perchlorate ion HClO4 perchloric acid For some species there is a change in spelling in the name. SO42- sulfate ion H2SO4 sulfuric acid PO43- phosphate ion H3PO4 phosphoric acid 136 Chemical Reactions: Equations Writing chemical equations – A chemical equation is the symbolic representation of a chemical reaction in terms of chemical formulas. – For example, the burning of sodium and chlorine to produce sodium chloride is written 2Na  Cl 2 2NaCl Reactants (consumed) Products  (produced) – The reactants (consumed; left side of reaction) are starting substances in a chemical reaction. The arrow means “yields.” The formulas on the right side of the arrow represent the products (produced). 137 Chemical Reactions: Equations Writing chemical equations – In many cases, it is useful to indicate the states of the substances in the equation (s, g, l, aq). – When you use these labels, the previous equation becomes 2Na(s )  Cl 2 ( g ) 2NaCl(s ) We write above the arrow any conditions for the reaction such as pressure, catalyst, heat, etc. A reaction gives a recipe for the amount of reactants needed to produce the amount of products. Species with no coefficient have an understood coefficient of 1. 138  Chemical Reactions: Equations Writing chemical equations – The law of conservation of mass dictates that the total number of atoms of each element on both sides of a chemical equation must match. The equation is then said to be balanced. CH 4  2O 2 CO 2  2H 2O We must have the same number of atoms on both sides for a reaction to be considered balanced and obeying the law of conservation of mass. To balance a reaction: 1. First, balance the atoms for elements that occur in only one substance on each side of the reaction. In this problem, O is involved with two substances on the product side; therefore, I will wait on balancing O until later. C & H are only in one species on both sides so I will balance them first. C needs no changes because there are one on each side, but H needs a 2 in front of H2O to balance the 4H on the reactants side. 2. Now that we have changed the coefficient of one of the O on the product side, it is easier to balance the O. We determine that we need a 2 coefficient on the O2 to balance the O on both sides at 4. Now the equation is balanced with 1C, 4O, and 4H on both sides. 139 Chemical Reactions: Equations Caution: For formulas that have subscripts, you must account for all atoms especially when dealing with parentheses for polyatomic species. For example, Fe2(SO4)3: has 2-Fe, 3x1 = 3-S, 3x4 = 12-O Caution: Remember that you can’t change the subscripts in formulas to balance equations; you may only change coefficients. If you change the subscripts, you are changing the substance. 140 Chemical Reactions: Equations O 2  2 PCl 3 2 POCl 3 P4  6N 2O P4O6  6N 2 Technique to handle odd numbers: determine number needed and divide by subscript of species. Next, you multiple the entire equation by the subscript to obtain whole numbers. [ 2 As2S 3  O2 As 2O 3  3 SO 2 ] 2 As 2S3  9 O2 2 As 2O3  6 SO2 141 The Mole The mole is a number, 6.02 x 1023. When you have a mole of something you have 6.02 x 1023 of them. The weight of a mole of one type of atoms is the atomic weight from the periodic table in grams. The mole allows you to count atoms by weighing. For example, how many carbon atoms are there in 24 grams of carbon? From the periodic table each mole has 12 grams. So there are 2 moles or 2 x 6.02x1023 atoms of carbon or 12.04x1023 carbon atoms. The Mole Formula Units To find out how much a mole of molecules weigh you just add up the atomic weights of the atoms in the molecule. How much does one mole of NaOH weigh? Add the atomic weights of Na, O, and H you get 23 + 16 + 1 = 40 grams. How much does a mole of H2O weigh? Add 2 times 1 for H and 16 for O to get 18 grams. How much does a mole of C3H8 weigh? (3 x 12) + (8 x 1 ) = 36 + 8 = 44 grams.  Exothermic and Endothermic Reactions Chemical changes that give off energy are called exothermic reactions. Chemical changes that absorb energy are called endothermic reactions. Exothermic Reactions In an exothermic reaction, heat is released. the energy of the products is less than the energy of the reactants. heat is a product. C(s) + 2 H2(g) CH4(g) + 18 kcal Copyright © 2005 by Pearson Education, Inc. Publishing as Benjamin Cummings 148 Endothermic Reactions In an endothermic reaction Heat is absorbed. The energy of the products is greater than the energy of the reactants. Heat is a reactant (added). N2(g) + O2 (g) + 43.3 kcal 2NO(g) Copyright © 2005 by Pearson Education, Inc. Publishing as Benjamin Cummings 149 Summary Reaction Energy Heat Type Change in Reaction Endothermic Heat absorbed Reactant Exothermic Heat released Product 150 Rate of Reaction Reaction rate is the speed at which reactant is used up. is the speed at which product forms. increases when temperature rises because reacting molecules move faster providing more colliding molecules with energy of activation. 151 Reaction Rates Kinetics: the study of how fast reactions take place Some reactions are fast (photosynthesis) Some reactions are slow (conversion of diamond to graphite) Rate of Reaction Expressed as either: Rate of disappearance of reactants (decrease or negative) OR Rate of appearance of products (increase or positive) Average Reaction Rate Copyright McGraw-Hill 2009 Average Reaction Rate Equation A B rate =   [A] or  [B] t t Why the negative on [A]? Average Reaction Rate Br2(aq) + HCOOH(aq)2Br(aq) + 2H+(aq) + CO2(g) Note: Br2 disappears over time Copyright McGraw-Hill 2009 Average Reaction Rate Br2(aq) + HCOOH(aq) 2Br(aq) + 2H+(aq) + CO2(g) Copyright McGraw-Hill 2009 Calculate Average Rate Avg. rate = [Br2 ]  [Br2 ] final initial t  t final initial Avg. rate = 0.0101 M  0.0120 M = 3.80 M /s 50 s  0 s Average Rate Average rate depends on time interval Plot of [Br2] vs time = curve Plot of Rate vs [Br2] = straight line Copyright McGraw-Hill 2009  Rate Constant Using data from Table 14.1 - what can you conclude? Time (s) [Br2] Rate (M/s) 50 0.0101 3.52 x 105 250 0.00596 1.75 x 105 Copyright McGraw-Hill 2009 Rate Constant Answer: When the [Br2] is halved; the rate is halved Rate is directly proportional to [Br2] rate = k [Br2] k = proportionality constant and is constant as long as temp remains constant Rate Constant Calculate the value of the rate constant for any set of data and get basically the same answer! k = rate / [Br2] 3.52 x 105 M/s k   3.5x 103 /s 0.0101 M Dependence of Reaction Rate on Reactant Concentration Rate law expression For the general equation: aA + bB cC + dD rate law = k[A]x[B]y k = proportionality constant x and y = the order of the reaction with respect to each reactant Order Exponents represent order Only determined via experimental data 1st order - rate directly proportional to concentration 2nd order - exponential relationship 0 order - no relationship Sum of exponents (order) indicates overall reaction order Determining Rate Law Exp. [A] (M) [B] (M) Initial In experiment 1 and 2; [B] is Rate (M/s) constant; [A] doubles and rate doubles - the reaction is 1st order 1 0.10 0.015 2.1 x 104 with respect to [A] In experiment 1 and 3; [A] is 2 0.20 0.015 4.2 x 104 constant; [B] doubles but the rate quadruples! This means that the reaction is 2nd order with respect to 3 0.10 0.030 8.4 x 102 [B] Calculate the Rate Constant Rate = k[A] [B]2 The rxn is 1st order w/ respect to [A] The rxn is 2nd order w/ respect to [B] The rxn is 3rd order overall (1 + 2) 2.1  104 M /s  9.3 M 2  s1 [rate] k  [A] [B]2 (0.10 M) (0.015 M ) 2 Dependence of Reaction Rate on Temperature Most reactions occur faster at a higher temperature. How does temperature alter rate? Collision Theory Particles must collide in order to react The greater frequency of collisions, the higher the reaction rate Only two particles may react at one time Many factors must be met: Orientation Energy needed to break bonds (activation energy) Collision Theory Copyright McGraw-Hill 2009 Collision Theory Though it seems simple, not all collisions are effective collisions Effective collisions: a collision that does result in a reaction An activated complex (transition state) forms in an effective collision Activation Energy Copyright McGraw-Hill 2009 The Arrhenius Equation The dependence of the rate constant of a reaction on temperature can be expressed k  Ae Ea / RT Ea = activation energy R = universal gas constant A = frequency factor T = Kelvin temp Catalysis Catalyst - a substance that increases the rate of a chemical reaction without being used up itself Provides a set of elementary steps with more favorable kinetics than those that exist in its absence Many times a catalyst lowers the activation energy Reaction Pathway with Catalyst Copyright McGraw-Hill 2009 Biological Catalysts Enzymes: large protein molecule that contains one or more active sites where interactions with substrates occur Enzymes are highly specific (lock and key) Enzyme-Substrate Complex Copyright McGraw-Hill 2009 Summary of Factors That Increase Reaction Rate 177 Chemical Equilibrium Consider the following reactions: CaCO3(s) + CO2(aq) + H2O(l) Ca2+(aq) + 2HCO3-(aq)..(1) and Ca2+(aq) + 2HCO3-(aq) CaCO3(s) + CO2(aq) + H2O(l)..(2) Reaction (2) is the reverse of reaction (1). At equilibrium the two opposing reactions occur at the same rate. Concentrations of chemical species do not change once equilibrium is established. Expression for Equilibrium Constant Consider the following equilibrium system: wA + xB ⇄ yC + zD [C]y [D] z Kc = [A] w [B] x The numerical value of Kc is calculated using the concentrations of reactants and products that exist at equilibrium. Expressions for Equilibrium Constants Examples: [NH 3 ] 2 N2(g) + 3H2(g) ⇄ 2NH3(g); Kc = [N 2 ][H 2 ]3 PCl5(g) ⇄ PCl3(g) + Cl2(g); Kc = [PCl3 ][Cl 2 ] [PCl5 ] CH4(g) + H2O(g) ⇄ CO(g) + 3H2(g); Kc = [CO][H 2 ]3 [CH 4 ][H 2 O] Calculating Equilibrium Constant Example-1: 1.000 mole of H2 gas and 1.000 mole of I2 vapor are introduced into a 5.00- liter sealed flask. The mixture is heated to a certain temperature and the following reaction occurs until equilibrium is established. H2(g) + I2(g) ⇄ 2HI(g) At equilibrium, the mixture is found to contain 1.580 mole of HI. (a) What are the concentrations of H2, I2 and HI at equilibrium? (b) Calculate the equilibrium constant Kc.  Ci=n/v= 1/5= 0.2 Calculating Equilibrium Constant for reaction: H2(g) + I2(g) ⇄ 2HI(g) X= 1.58 HI moles C(HI) equi = 1.58/5= 0 ———————————————————————————— 2x = 0.316 H2(g) + I2(g) ⇄ 2 HI(g) ———————————————————————————— X= 0.316/2=0.158 Initial n: 1 1 0.000 -x -x 2x Change in n: -0.79 -0.79 + 1.58 Equilibrium, n 0.21 0.21 1.58 ———————————————————————————— C equi = 0.21/ = 0.042 [HI] 2 2 Kc = = (0.316) = 57 [H 2 ][I 2 ] (0.042) 2 Calculating Equilibrium Constant Example-2: 0.500 mole of HI is introduced into a 1.00 liter sealed flask and heated to a certain temperature. Under this condition HI decomposes to produce H2 and I2 until an equilibrium is established. An analysis of the equilibrium mixture shows that 0.105 mole of HI has decomposed. Calculate the equilibrium concentrations of H2, I2 and HI, and the equilibrium constant Kc for the following reaction: H2(g) + I2(g) ⇄ 2HI(g), Calculating Equilibrium Constant 2x= 0.10 The reaction: H2(g) + I2(g) ⇄ 2HI(g), proceeds from X=0.105 right to left. ———————————————————————————— H2(g) + I2(g) ⇄ 2HI(g) ———————————————————————————— Initial [ ], M: 0.000 0.000 0.500 x x -2x Change in [ ], M: +0.0525 +0.0525 -0.105 Equil’m [ ], M 0.0525 0.0525 0.395 ———————————————————————————— [HI] 2 (0.395) 2 Kc = = = 56.6 [H 2 ][I 2 ] (0.0525) 2  Lecture 4 Physical States of Matter Intermolecular forces Solutions and Acids  The three states of matter. Silberberg, Principles of Chemistry Gases Because gases have so much space between the particles they have properties that are dependent on one another. Solids and Liquids Because the particles are so much closer in liquids and solids, there are chances for particles to attract (or repel). This and the mass of the particles are main factors in determining the properties of solids and liquids. Some properties are boiling and melting points, surface tension, vapor pressure, and crystalline structure. Solids Solids may have a definite structure and are called crystalline. Solids that have no regular shape are called amorphous. The hexagonal structure of ice. The striking beauty of crystalline solids.  The crystal lattice and the unit cell. lattice uni point t cel l uni t cel l portion of a 3-D portion of a 2-D lattice lattice Phase Changes Energy released sublimination melting vaporizing solid liquid gas freezing condensing Energy absorbed Water (1) Why Study Water: - Life exists in an aqueous environment. - Intra- and extra-cellular fluids very similar to sea water. (2) Characteristics of Water  The water molecule is bent.  O (3.5) is more electronegative than H (2.2). H2O is highly polar with a net dipole. has 33% ionic character.   O O H H H H  104.5oC  H2O interacts through hydrogen bonds. b.p. (oC) m.p. Water 100 0 MeOH 65 ‒98 BuOH 117 Butane –0.5 1 Å (Angstrom) = 10‒8 cm = 0.1 nm Hydrogen bond: A type of dipole force. An electrostatic attractive force between molecules. Found when a H atom is bonded to N, O or F. Weaker than covalent bonds (23 kJ/mol in water). C—C: 350 kJ/mol, O—H: 470 kJ/mol. Can a H atom bonded to C form a H-bond? The answer is No (for the moment). C (2.5) is only slightly more electronegative than H.  Hydrogen bonding in ice. Each H2O forms 4 H-bonds. In liquid water, each H2O forms, on average, 3.4 H-bonds. This is why less dense ice floats in liquid water, and why water expands when frozen.  Common H-bonds in biological systems:  Good solvent for polar and ionic material such as salts, small polar molecules (amino acids, sugars, nucleic acids …) and the exterior of proteins.  Solubility enhanced for molecules that can form H-bonds i.e. hydroxy, keto, carboxy, amino groups, and these are termed hydrophilic (water loving) groups.  Solubility reduced for molecules that cannot form H-bonds (i.e. hydrocarbons such as alkanes, alkenes), and these are termed hydrophobic (water fearing) compounds.  Nonpolar Polar Amphipathic phospholipids (3) Aqueous (Water-based) Solutions Water is a polar solvent.  Solvation of crystalline salt: Water breaks up the salt crystal lattice by H-bonding with the individual ions. hydration Entropy-driven: G < 0  H ≥ 0, S ≫ 0 (of ions) DeltaG= deltaH –Tdelta S  Nonpolar gases (O2, CO2) and nonpolar compounds (benzene, hexane) are not soluble in water. H > 0 breaking up the water H-bonds. S < 0 cage-like, highly ordered H2O molecules around the hydrophobic compounds. Overall, G > 0. This is why we need a water soluble O2 carrier, hemoglobin.  Amphipathic (Amphiphilic) compounds contain both polar (hydrophilic) & non-polar (hydrophobic) regions (ex. fatty acids) In aqueous solution: water hydrates the hydrophilic portion but excludes the hydrophobic region to give micelles and lipid bilayers in a process that reduces entropy (S) or increases the order of the lipid molecules, but increases the entropy of surrounding water (basis of cell membranes) Formation of micelles Not favorable G > 0 Favorable G < 0 Hydrophobic Interactions (4) Hydrophobic Interactions The tendency of hydrophobic (lipophilic) groups to form intermolecular aggregates in an aqueous medium, and analogous intramolecular interactions. The name may be misleading, not repulsion between water and hydrophobic compounds, per se, but effects of the hydrophobic groups on the water-water interactions are the driving force for the hydrophobic interactions. Hydrophobic interaction is driven by S ≫ 0.  Hydrophobic Interactions G > 0 Hexane S < 0 Hexane O O O O O O O O O O O O O O O O O O O O O O G < 0 O O O O O O O O H2O S > 0 O H2O Spontaneous The transfer of nonpolar solutes from non-aq. solvent to water. Note that water molecules surrounding the non-polar solutes are highly ordered (red).  O O O O Add oil drops O O O O OO O O O O O O O O H 2O O O O O Step 1 Step 2 Less ordered bulky water solvent Supose there are 11 water molecules, The rests (2 molecules) are Highly ordered water molecules all disordered. still disodered (9 molecules) surround oil drops. When the oil drops come closer, they will spontaneouly collapse onto themselves to become a larger drop. We know this from experience. Step 3 O O O Now, 3 previously orderd water molecules O become disordered, increasing "disorder" O O in the system (S > 0), and leading to O O O a state of lower G (G < 0). O O Hydrophobic Interaction and Protein Folding: Hydrophobic interactions are the most important noncovalent force that will cause the linear polypeptide to fold into a compact structure. However, it is not the interactions between side chains of hydrophobic amino acids per se (mainly van der Waals) that induce the strong interaction, but the increase in entropy gained by the removal of hydrophobic surface area from ordered solvating water. The aggregation of the hydrophobic surfaces gives the tightly packed core of a protein. O folding O Phe O Leu Phe O O O S > 0 O O Leu O O O (5) van der Waals Interactions ① Dispersion forces (London forces): Intermolecular, attractive force arising from temporary dipoles. b.p. of noble gases: He ‒269°C Ne ‒246°C Ar ‒186°C Original temporary Induced dipole dipole (5) van der Waals Interactions ② Dipole-dipole interactions: Intermolecular, attractive force arising from permanent dipoles. ethane b.p. 184.5 K + – fluoroethane b.p. 194.7 K Permanent dipole ③ All molecules experience dispersion forces. ④ Dipole-dipole interactions are not an alternative to dispersion forces. - They occur in addition to them.  (6) Noncovalent interactions are: -individually weak, but their cumulative effects are critical in biological functions. (8) Solutes Affect the Colligative Properties of Aqueous Solutions: [H2O]eff is lower in solution than in pure water. Colligative: depending on the number of particles (as molecules) and not on the nature of the particles Other colligative properties of solution: b.p.  m.p.  (8) Solutes Affect the Colligative Properties of Aqueous Solutions: [H2O]eff is lower in solution than in pure water.  Osmosis: water moves from a region of high to low water conc.  Osmolarity: The osmotic concentration of a solution expressed as osmoles of solute per liter of solution. The number of osmotically Osmolarity is dependent on the active particles in solution number of particles in solution but independent of the nature 1 mM NaCl of the particles. Osmolarity of a simple solution = (molarity) X (the number of particles per molecule) 1 M glucose solution – 1 OsM 1 M NaCl – 2 OsM 1 M Na2SO4 – 3 OsM Total osmolarity = 2 mOsM isotonic hypertonic hypotonic Water moves fairly freely across the cell membrane. Not all molecules can cross the membrane (semi-permeable). Examples of Osmosis 1. Fresh water fish in salt water 2. A cell in distilled water 3. Blood 4. Bacteria & salt 5. Wilted lettuce Q: Why do the cells store energy as polysaccharides, not as individual glucose molecules?  Ionization of Water, Weak Acids, and Weak Bases (1) Ionization of Water: Slight tendency of H2O to undergo reversible ionization to hydrogen ion (proton) and hydroxide ion: H2O ⇄ H+ + OH‒ BUT, free protons do not exist in solution: protons are hydrated to hydronium ions: H2O + H+ ⇄ H3O+ (very fast due to proton hopping) The ionization is expressed by an equilibrium constant (Keq) H2O ⇄ H+ + OH‒ [H+][OH‒] Keq = [H2O] At 25 °C, [H2O] = 55.5 M. ∴ (55.5 M)(Keq) = [H+][OH‒] = Kw (ion product of water) Keq is 1.8 x 10-16 M as determined by electro-conductivity exp. Kw = (55.5 M)(1.8 x 10‒16 M) = [H+][OH‒] = 1.0 x 10‒14 M2 At neutral pH, [H+] =[OH‒] and [H+][OH-] = [H+]2  [H+] = 1.0 x 10‒7 M in pure water at 25°C. Since Kw is constant, when [H+] > 10-7 M ⇒[OH‒] < 10-7 M, and when [H+] < 10-7 M ⇒[OH‒] > 10-7 M. (2) The pH Scale 1 pH = log [H+] = ‒ log [H+] pH 7 is neutral, [ H+] = [OH‒] pH < 7 is acidic, [ H+] > [OH‒] pH > 7 is basic, [OH‒] > [ H+] pH + pOH = 14  The pH must be controlled in an organism; where the breakdown in pH regulation can lead to serious metabolic disturbances: The pH of blood is normally kept within 7.35~7.45. Outside the narrow range, the organism can not function. The pH of the cytosol of most cells is ~ 7.4, however, in the lysosomal organelles the pH is ~ 5.0. This is the pH at which the degradative enzymes (proteases) of the lysosome function best, and they are actually inactive at cytosolic pH! What is the pH of 0.1 M NaOH solution ? (p.62) NaOH is a strong base and completely ionized in dilute aq. solution. [OH‒] = 0.1 M. From Kw = [H+][OH‒] = 1.0  10‒14 M2, [H+] = 10-14 M2/0.1 M = 10‒13 M.  pH = ‒log10‒13 = 13. Or, from [OH‒] = 0.1 M, pOH = ‒log 10‒1 = 1. Since pH + pOH = 14, pH of the solution is 13. What is [OH‒] in a solution with [H+] of 1.3 x 10‒4 M? [OH‒] = Kw/[H+] = 10‒14 M2/1.3  10-4 M = 0.769....  10‒10 M = 7.7  10‒11 M (3) Weak Acids and Weak Bases  Acids – proton (H+) donors; Bases – proton acceptors HA ⇄ H+ + A‒ (A‒ : conjugate base) [H+] [A‒]  Keq = Ka (dissociation constant) = [HA]  Stronger acids larger Ka lower pKa (‒ logKa) ⇒ phosphoric acid (H3PO4): pKa = 2.34; monohydrogen phosphate (HPO4‒2): pKa = 12.4 Conjugate base of a strong acid is a weak base.  (4) pKa and Titration Curve - How do pH values of an acetic acid solution vary with added [OH‒]? a titration curve - Constructed by: a) experiment b) H-H equation midpoint pH of titration = pKa of corresponding acid: b/c pH = pKa when [HA] = [A‒] * Slope lower near midpoint When [HA] = [A‒], pH is relatively insensitive to addition of strong acid or base i.e. buffered solution Buffering capacity is maximal when pH = pKa. The useful range of a buffer is within one pH unit of its pKa. Above this, the pH will change rapidly. * Slope lower near midpoint When [HA] = [A‒], pH is relatively insensitive to addition of strong acid or base i.e. buffered solution. Buffering capacity is maximal when pH = pKa. The useful range of a buffer is within one pH unit of its pKa. Above this, the pH will change rapidly. (5) Buffers Buffer: A system whose pH changes only slightly when small amounts of acid or base is added. A buffer ordinarily consists of a weak acid and its conjugate base, present in roughly equal amounts (at pH = pKa of the acid) Used to control the pH within a system How buffer works? (Fig 2-19) Buffer works because the added H+ or OH‒ ions are consumed and do not directly affect the pH. HAc + OH‒ ⇄ H2O + Ac‒ Ac‒ + H+ ⇄ HAc (6) The Henderson-Hasselbalch (H-H) Equation The pH of a solution, and the concentration of an acid and its conjugate base are related by the H-H equation: [A‒] pH = pKa + log [HA] When the molar concentration of an acid (HA) and its conjugate base (A‒) are equal ([A‒] = [HA]), [A‒]/[HA] = 1; and log[A‒]/[HA] = log1 = 0 So the pH of the solution simply equals the pKa of the acid. When [A‒] > [HA], pH > pKa. When [A‒] < [HA], pH < pKa. (7) Polyprotic Acids: Substances that have more than one acid/base group. H3PO4 ⇄ H2PO4‒ + H+ ⇄ HPO42‒ + H+ ⇄ PO43‒ + H+ pKa1 = 2.14 pKa2 = 6.86 pKa3 = 12.4 Example: 1.00 mole of phosphoric acid (H3PO4) and 1.75 moles of NaOH are added to 1 L of water. Calculate the pH. Step 1: 1 mol of H3PO4 + 1.75 mol OH- 1 mol H2PO4- + 1 mol H2O + 0.75 mol OH- Step 2: 1 mol of H2PO4- + 0.75 mol OH- 0.25 mol H2PO4- + 0.75 mol HPO42- + 0.75 mol H2O Step 3: In the end, we have 0.25 moles of H2PO4- and 0.75 moles of HPO42-, we can calculate the pH using the H-H equation: Step 4: Look up the pKa of the reaction: pKa for H2PO4- ⇄ HPO42- + H+, is 6.86 Step 5: Calculate [HA] = [H2PO4-]: 0.25 mol/1 L = 0.25 M Step 6: calculate [A-] = [HPO42-] : 0.75 mol/1 L = 0.75 M Therefore: pH = pKa+ log[A-]/[HA] = 6.86 + log((0.75 M)/(0.25 M)) = 7.34 Example: The addition of a 0.01 mL drop of 1 M HCl to 1 L of water will change the pH from 7 to 5. A small concentration of buffer can alter this so that there is virtually no change in the pH, even with much larger amounts of acid added! Q13, p.74. A buffer contains 0.010 mol of lactic acid (pKa=3.86) and 0.050 mol of sodium lactate per liter. (a) Calculate the pH of the buffer. (b) Calculate the change in pH when 5 mL of 0.5 M HCl is added to 1 L of the buffer. (c) What pH change would you expect if you added the same quantity of HCl to 1 L of pure water (pH=7). (8) Biological buffers: HCO3‒ : H2CO3 pKa = 6.35 HPO42‒ : H2PO4‒ pKa = 6.86 pKa of amino acid, histidine = 6.0 Common buffers used in lab: Buffer pKa (25oC) Effective pH Range succinate (pK1) 4.21 3.2-5.2 acetate 4.76 3.6-5.6 citrate (pK2) 4.76 3.0-6.2 malate (pK2) 5.13 4.0-6.0 succinate (pK2) 5.64 5.5-6.5 MES 6.10 5.5-6.7 carbonate (pK1) 6.35 6.0-8.0 citrate (pK3) 6.40 5.5-7.2 imidazole 6.95 6.2-7.8 MOPS 7.14 6.5-7.9 phosphate (pK2) 7.20 5.8-8.0 HEPES 7.48 6.8-8.2 Trizma (Tris) 8.06 7.5-9.0 glycine (pK2) 9.78 8.8-10.6 carbonate (pK2) 10.33 9.5-11.1