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CHEM 1001 (14)


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Carleton University
CHEM 1001
Robert Burk

Chapter 1 1.1 Atoms, Molecules and Chemical Formulas 1.2 Measurements in Chemistry Physical Properties: size, colour and mass Quantitative Measurements of properties Every Measurement gives a numerical result of 3 aspects: Magnitude Units Precision: exactness of a measurement Accuracy: close a measurement is to the true value Chemical Properties: chemical transformations 1.3 Chemical Problem Solving 1.4 Counting Atoms: The Mole Mole and Avogadro’s Number Avogadro’s Number : 6.022 X10 23 items/mol Molar mass: (Fractional Abundance) (isotopic molar mass) Mass- Mole-Atom Conversions: 1.5 Amounts of Compounds Molar mass 1.6 Aqueous Solutions Solvent: substances used to dissolve solutes Solute: pure substance dissolved in solution Polyatomic ions remain intact Concentration (c ) amount of solute in a solution Any solutions is formed by dissolving an ionic solid in water conducts electricity Dilutions: amount of solute remains the same but the volume of the solution increases RESULTS IN A : solution of lower molarity The number of moles of solute does not change during the dilution C1 1= C2 2 Concentrated acid is always added to the water 1.7 Writing Chemical Equations The number of atoms of each element is conserved in any chemical reaction 1.8 The stoichiometry of Chemical Reactions Stoichiometry is the study of the amounts of materials consumed and produced in chemical reactions ( ) 1.9 Yields of Chemical Reactions Under practical conditions, chemicals almost always produce smaller amounts of products than the amount predicted by stoichiometric analysis. 3 Reasons: 1. Many reactions stop before reaching completion ( do not go to completion because they reach dynamic equilibrium) 2. Competing reactions often consume some of the starting materials 3. When product of a reaction is purified and isolated, some of it is inevitably lost during the collection process (impossible to get all the product out of the container) 1.10 The Limiting Reactant The reactant that runs our is called limiting reactant ( imits the amount of product that can be made). Other starting materials are the excess reactant Limiting reactant: smallest value of moles divided by coefficient Chapter 2 2.1 Pressure The pressure of the atmosphere can be measured with the instrument called the barometer Units of pressure : standard atmosphere ( atm) pressure that will support a column of mercury 760mm in height. Manometer: gases exert pressure on both liquid surfaces. The difference in pressure exerted by two gases. 2.2 Describing Gases Boyle found: Volumr of the trapped gas is inversely proportional to the total pressure applied bu mercury plus the atmosphere. V 1/Pgas ( fixed temperature and amount) Volume of a gas is directly proportional to its temperature: Vgas  Tgas ( fixed pressure and amount) Gas volume is proportional to the amount of gas Vgas  ngas (fixed pressure and temperature) Ideal Gas Equation Universal Gas constan: R= 8.314 LkPa/mol K or 0.08314 L bar/mol K Ideal Gas law: pV=nRT As pressure increases= Volume increases Solving quantitative problems about gases at moderate temperatures and pressures requires only one equation. The ideal gas equation. 2.3 Gas Mixture Dalton’s law of partial pressure Ptotal 1p 2p +3 4 (each p is an individual pressure) Mole Fraction (X): Parts/million (ppm) and Parts/billion (ppb) PA= XA Total In a mixture of gases, each gas contributes to the total pressure the pressure that it would exert if the gas were present in the container by itself. 2.4 Gas Stiochiometry Substance Equation Liquid or solid n= m/M Aqueous Solution n=cV Gas n=pV/nT 2.5 Molecular View of Gases 2 E kinetic=1/2mv E = 3RT/2N avekinetic A At a given Temperature, all gases have the same kinetic energy distribution Ideal gas: negligible molecular sizes Negligible intermolecular forces The volume occupied by the molecules of an ideal gas is negligible compared with the volume of its container. The energies generated by forces among ideal gas molecules are negligible compared with molecular kinetic energies. 2.6 Additional Gas Properties Root Mean square : Vave(3RT/M) 1/2 Diffusion: movement of one type of molecule through molecules of another type Effusion: movement of molecules escaping from a container into a vacuum 2.7 Non-ideal (real) Gases Van der Waals Equation: 2.8 Chemistry of the Earths Atmosphere Troposphere : NO NO 2 red-brown gas that can be seen in the atmosphere over large cities, absorbes energy from sunlight and decomposes into NO molecules and oxygen molecules) O 3 SO 2 SO 3 Photochemical Smog: mixture of all these pollutants Dynamic Equilibrium:rate of evaporation equals the rate of condensation Vapour Pressure of any substance increases rapidly with temperature because the kinetic energies of the molecules increase as the temperature rises. Humidity= (100%) (pH2o/pvapour,H2O) Chapter 3 3.1 Types of Energy Energy: ability to do work Work: displacement of an object against an opposing force Kinetic Energy : moving objects Potential Energy: Stored energy Thermal Energy: Content of a hot object Radiant energy: content of electromagnetic radiation (light/ infrared radiation) Kinetic and Potential Energies 2 s SI unit: kgm /s Potential energy= gravitational energy Electrical Energy Electrical forces lead to electrical potential energy. Energy is released when oppositely charged objects move close together whereas energy must be supplied to pull oppositely charged objects apart. Thermal Energy Total energy of these random movements Average molecular kinetic energy increases as temperature increases. Liquid phase atoms and molecules still have energies of translation, rotation and vibration. Thermal energy is the cause and temperature is the effect. Radiant Energy Energy transfers and transformations: When atoms/molecules collide with one another, energy transfers cause some molecules to speed up and others to slow down. Energy transformations accompany chemical reactions: Reaction releases energy, chemical energy is converted into other forms of energy. 3.2 Thermodynamics - Quantitative study of energy transfers and transformations System: describe and study by it self Surroundings: everything else System is separated from its surroundings by a boundary across from which matter/molecules can move Conservation of Energy Law: energy is neither created nor destroyed Heat A system can exchange thermal energy with its surroundings. Amount of energy that is transferred is heat (q). Measured in Joules. ∆T depends on the q the amount of heat transferred, direction of the heat flow, ∆T absorbs heat :+ ∆T releases: - ∆T depends on the identity of the material. Expressed by Molar Heat Capacity q surr-q sys Work Energy used to move an object against an opposing force (w) Amount of work depends on the magnitude of the force that must be overcome and the amount of movement or displacement. w=Fd W surr-W sys First Law of Thermodynamics Energy(E) is exchanged in two ways : heat or work ∆E=q+w ∆E sysE surr Law: ∆Etotal=∆Esys +∆Esurr=0 State and Path Functions Properties: only on the conditions that describe the system before and after the transformation State Function: How the changes occur Path Function: how a change takes place Key concept: the value of a state function does not depend on the path taken or on the rate of the change. Thermodynamic Path Function Energy is a state function Heat and Work are Path Functions 3.3 Energy Changes in Chemical Reactions Energy is released during the reaction Reversing the direction of the reaction changes the sign of the energy changes. Features of Reaction Energies When a chemical reaction releases energy. ∆E has a negatives sign. Reaction absorbs energy. ∆E has a positive sign. ∆Ereaction= ΣBEbonds broken – ΣBEbondsf formed 3.4 Measuring Energy Changes: Calorimetry Calorimeters : measure heat flow that accompany chemical processes. Chemical reaction takes place within a calorimeter, resulting in a heat flow between the chemicals and the calorimeter. The temperature of the calorimeter rises or falls in response to this heat flow. qchemicals = -qcalorimeter release heat, qchemicals is negative  EXOTHERMIC gains heat, qcalorimeter is positive , temperature rises Absorbs heat, qchemicals is positive  ENDOTHERMIC Loses heat(qcalorimeter is negative) , temperature falls qcalorimeter =Ccal∆T Calculating Energy Change ∆Emolar=∆E/n 3.5 Enthalpy F=pA Wsys=-p∆Vsys∆H =qp ∧ ∧ H=E+pV Enthalpy (H)= Heat flow at constant pressure Hess Law: Overall ∆H= Sum of ∆H for individual steps Reactants (net reaction) Products (decomposition reactions) Elements (formation reactions) Products o Standard Enthalpy of formation (∆H )f formation reaction Formation reaction= 1 mol of a substance produced from elements 1/2N 2g) +1/2 O 2g) N(g) ∆H of reactioneff products∆H fproductscoeff reactants ∆Hof reactants ∆Hreaction= ∆E reaction + RT∆ngases Key Concepts: A formation reaction produces 1 mole of a chemical substance from the elements in their standard states. The enthalpy change for any overall process is equal to the sum if the enthalpy changes for any set of steps that leads from the starting materials to be products, Enthalpy is a thermodynamic state function that describes heat flow at constant pressure. Chapter 4 4.1 Characteristics of Atoms Fundamental characteristics of atoms: Atoms possess mass: Atoms contain positive nuclei Atoms contain electrons Atoms occupy volume The volume of an atom is determined by the size of its electron cloud. Atoms have various properties Atoms attract one another Atoms can combine with one another: form chemical bonds with one another to construct molecules. 4.2 Characteristics of light Studying the structure of atoms by electromagnetic radiation. Light is a form of this radiation. Light has wave aspects Light is a wave like property. Frequency(v): number of wave crests passing a point on space in one second. Unit s or Hz Wavelength (λ): distance between successive wave crests. Units m or nm Amplitude: height of a wave Amplitude of a light wave measures the intensity of the light. Λv=c Photoelectric Effect Total energy of a beam of light depends on its intensity. Photoelectric effect : energy of light depends on its frequency and intensity. Light comes on bundles of photons. Eeach photon has an energy that is directly proportional to the frequency. Ephoton=hvphoton Ephoton: energy of light Vphoton: frequency E=hv , v=c/λ , E= hc/λ Electron kinetic energy= Photon energy –Binding energy. Higher intensity means more photons but not more energy per photon Light has particle Aspects Light has some properties of waves and some properties of particles. 4.3 Absorption and Emission Spectra when light interacts with free atoms , interactions reveal infomormation about electrons bound to individual atoms. Light and energy Absorption of photons by free atoms has two different results, depending on the energy of the photons. When an atom absorbs a photons suffieciently high energy, an electron is ejected,a process called photoionization. 2 types of results: Transferred to a higher energy state  Excited State -give up their excess energy to return to loswer energy states lowest energy state of an atom, which is the most stable state  groung state Attractive electrical forces hold a bound of electrons within an atom ad energy must be supplied to remove a bound electron from an atom. Lower the energy state, more energy must be supplied to remove the electron.-> energy changes that are measured relative to the energy of free electrons. Energy of a free stationary electron is zero. Exchange of energy between atoms and light is that energy is conserved. Therefore, change in energy of the atom exactly equals the energy of the photon ∆Eatom=±hvphoton When an atom absorbs a photon, the atom gains the photons energy so ∆Eatom is + Atom emits a photon, atom loses the photons energy so ∆Eatom is – As an atom returns to its ground state it must lose exactly the amount of energy that is originally gained,However the excited atoms lose excess energy involving small energy changes do the frequency of emitted photons are often lower than absorbed photons. Atomic Spectra Absorption spectrum: unique for each gas because different atoms absorb different energies of photons. Emission spectrum: intensity of light emitted as a function of frequency. Each Frequency absorbed or emitted by an atom corresponds to a particular energy change for the atom. Quantization of Energy When an atom absorbs light of frequency v, the light beam loses energy hv, and the atom gains the amount of energy. Therefore, a photon with high enough energy can cause an atom to lose one of electrons.Absorption of a photon recults in an energy gain for an electron in the atom energy change for the atom equals the energy change for an atomic electron. En= 2.18x10 J/ n 2 4.4 Properties of Electrons Electrons have particle like and wave like properties. -Each electron has the same mass and charge. -Electrons behave like magnets: occurs due to property called spin. -Electrons have wave properties. The momentum of a particle is the porfuct of its mass and speed; p=mv Λparticle= h/mν 4.5 Quantization and Quantum Numbers Each atomic energy level is associated with a specific three-dimensional atomic orbital. Principal Quantum Number (n): - indexes energy. Each electron in an atom can be assigned a value n that is a positive integer that correlates with the energy of the electron. Most stable energy for an atomic electron corresponds to n=1 and each one higher value of n is less stable energy state. -Size of an atomic orbital: energy of an electron correlated with its distribution in space. -Higher the principal quantum number the more energy the electron has and the greater its average distance from the nucleus. n: Increases, energy of the electron increases, orbital gets bigger, electron is less tightly bound to the atom. Azimuthal Quantum Number (l) -Shape of atomic orbitals -l correlates with the number of preferred axis in a particular orbital and thereby identifies the orbital shape. Value of l 0 1 2 3 Orbital letter s p d f Magnetic and Spin Orientation Quantum Numbers(m) (m ) l s -magnetic quantum number m: indeles restrictions l=1 m:l-1, 0, +1 -electron has magnetism associated with a property called spin. -Magnetism is directional, so the spin of an electron is directional -Spin orientation is quantisized: electron spin in 2 ways ; UP or DOWN -Spin orientation quantum number m : indsxes the behavior -m sre +1/2 and -1/2 4.6 Shapes of Atomic Orbitals -An atom that contains many electrons can be described by superimposing(adding together) the orbitals for all of its electrons to obtain the overall size and shape of the atom. -Chemical properties of an atom are determined by behavior of their electrons because atomic electrons are described by orbitals, the interactions of electrons can be described in terms of orbital interactions -Two characteristics of orbitals determine how electrons interact : SHAPES & ENERGIES -Quantum number n&l determine the size and shape of an orbital. -n increases the size of the orbital increases -l increases, shape of the orbital becomes more elaborate Orbital Despictions 3 types of plots : Electron density plot : represents the electron distribution in an orbital as a two dimentional plot. Does not show it three dimentional Orbital Density picture: Take a lot of time and care Electron contour drawing: all details of electron density inside the surface is lost -Value for r where electron density drops to zero called Node -Electron has wave like properties- this wave has zero amplitude at the node, but non-zero amplitude on either side of the node. Thus, electrons can move across the node without actually existing at the node. -Number of nodes increases as s and n increases. However 1s orbitals have no nodes, a 2s orbital has one node. Orbital Size In any particular atoms; -orbitals get larger as the value of n increases -all orbitals with the same principal quantum number are similar in size -Each orbital becomes smaller as nuclear charge increases; As the positive charge of the nucleus increases, the electrical force exerted by the nucleus on the negatively charged electrons increase as well, and electrons become more tightly bound. Thus reducing the radius of the orbital, orbitals shrink in size as atomic number increases. 4.7 Sunlight and the Earth Thermosphere (heat released) -85km molecules of nitrogen and oxygen absorb x- ray radiation coming from the sun. N2+hvN2 +e-+ O2+hv O+O Reactions are unstable and will recombine releasing energy in the form of heat; + N2 +e-N2 +heat O+OO2 +heat -Aurora : atoms and molecules that reach excited states by gaining energy from some external source return to their ground states by emitting photons.Absorption by atoms and molecules removes high energy photons. -Result: intensity of high energy light decreases as sunlight moves down through the atmosphere towards the Earth’s surface Reactions in the Ozone Layer Stratosphere, solar radiation generates an abundance of ozone (O3) -Ozone layer forms in two steps; First photon with a wavelength between 180- 240nm break on O2 molecules into two atoms of oxygen. O2+hv O+O -Second, oxygen molecules capture one of the oxygen atoms to form an ozone molecule O2+O O3 +heat -Second step occurs twice for each O2, giving the overall balanced process for ozone formation 3O2+hv 2O3 +Heat Why id the ozone layer confined to one region of the atmosphere? -production of ozone requires both a source of oxygen atoms and frequent collsions between the atoms and the molecules that make up the atmosphere. Lower than 20 km O2 dissociation does not occur Higher than 35km plenty of light to dissociate O2, but molecular density and rate od molecular collisions are too low. O3+hv O+O2 Interactions of molecules and light in the ozone layer result in a delicate balance that holds the ozone concentration at a relatively constant value; -photons with wavelengths 180-240nm break apart O2molecules -photons with wavelengths of 200-340nm break apart O3 molecules -oxygen atoms combine with O2 molecules can produce O3 molecules and heat Greenhouse Effect Gases in troposphere: carbon dioxide, water vapor and methane Known as Green house gases. -Glass panes of greenhouses heep the greenhouse warm, greenhouse gases moderate temperature changes from night to day by absorbing some of the infrared photons emitted by the earth. Following the absorption the gases reemit still longer wavelengths photons some where plants reabsorb them. -Incoming sunlight -Out coming: infrared radiation emitted by the Earth Chapter 5 5.1 Orbitals Energies A hydrogen atom can absorb a photon and change from its most stable state (ground state) to a less stable state (exited state). The Effect of Nuclear Charge The stability of an orbital can be determined by measuring the amount of energy required to remove an electron completely: Ionization energy (IE): HH +e- IEH: 2.18x10^-18 J He  He +e- IE : 8.72 x10^ -18 J He+ LARGER THE IE HAS STRONGER GROUND STATE AND STABLE. Effects of Other Electrons Multi electron atom, each electron affects the properties of all the other electrons. A given orbital is less stable in a multielectron atom than it is in the single electron ion with the same nuclear charge. Screening This electron- electron repulsion cancels a portion of attraction between the nucleus and the incoming electron. Partial cancellation is called screening. -incomplete screening can be seen in the ionization energies of hydrogen atoms, helium atoms and helium ions. -Without screening , the IE of a helium atom would be the same as that of a helium ion -With complete screening, one helium electron would compensate for one of the protons in the nucleus -Electrons in compact orbitals pack around the nucleus more tightly than do electrons in large, diffuse orbitals. -RESULT: the effectiveness in screening the nuclear charge decreases as the orbital size increase. -Size of an orbital increases with n, and an electron ability to screen decrease as n increases -Higher the value of the quantum number l, the more that orbital is screened by electrons in smaller, more stable orbitals - Electrons with the same l value but different values of l do not screen one another effectively . -ex. Electrons occupy different p orbitals that have the same n value = same energies known as degenerate, the amount of mutual screening is slight -Quantitative information about energies of atomic orbitals is obtained using photoelectron spectroscopy  principles of the photoelectric effect to gaseous atoms. 5.2 Structure of the Periodic Table Pauli Exclusion Principle: EACH ELECTRON IN AN ATOM HAS A UNIQUE SET OF QUANTUM NUMBERS AUFBAU Principle: ELECTRONS ARE PLACED INTO ATOMIC ORBITALS BEGINNING WITH THE LOWEST ENERGY ELECTRONS FOLLOWED BY SUCCESSIVELY HIGHER ENERGY ELECTRONS. -stable: electron occupy the lowest energy orbitals available: ground state configuration of an atom by placing electron in the orbitals starting with the most stable in energy and moving progressively upward. -most stable: quantum number are not already assigned to another electron Order of orbital filling 1s2s2p3s3p4s3d4p5s4d5p6s 4f5d6p7s5f6d7p Valence Electrons -An electron is spatially when it occupies one of the largest orbitals of the atom. -Energetically when it occupies one of the least stable occupied orbitals of the atom -electrons in less stable (higher energy) orbitals are thus more chemically active than electrons in more stable orbitals. - Accessible electrons: Valence electrons participate in chemical reactions -Inaccessible electrons: Core electrons Orbital size increases and orbital stability decreases as the principal quantum number n gets longer. -Valence electrons are all those of highest principal quantum number plus those in partially filled d and f orbitals. 5.3 Electron Configuration Electron Configuartion: a complete specification of how
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