Chapter 1 1.1 Atoms, Molecules and Chemical ... amazonaws com
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 X1023 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 C1V1= C2V2
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 (
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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=p1+p2+p3+p4 (each p is an individual pressure) Mole Fraction (X): Parts/million (ppm) and Parts/billion (ppb) PA= XApTotal 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 Liquid or solid Aqueous Solution Gas
Equation n= m/M n=cV n=pV/nT
2.5 Molecular View of Gases
Ekinetic =1/2mv2 Eavekinetic= 3RT/2NA
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 NO2 ( 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) O3 , SO2 , SO3 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
SI unit: kgm2/ss
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
qsurr=-qsys
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
Wsurr=-Wsys
First Law of Thermodynamics Energy(E) is exchanged in two ways : heat or work ∆E=q+w
∆Esys=-∆Esurr 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 Standard Enthalpy of formation (∆Hof): formation reaction Formation reaction= 1 mol of a substance produced from elements 1/2N2(g) +1/2 O2(g) N(g) ∆Hof reaction =Σcoeff products∆Hofproducts -Σcoeff 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-1 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= 2.18x10-18J/ n2 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(ml) (ms) -magnetic quantum number ml: indexes restrictions l=1 ml: -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 ms: indexes the behavior -ms are +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 xray radiation coming from the sun. N2+hvN2+ +eO2+hv O+O Reactions are unstable and will recombine releasing energy in the form of heat; N2++e-N2 +heat O+OO2 +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 180240nm 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): HH++eIEH: 2.18x10^-18 J He+ He2+ +eIEHe+: 8.72 x10^ -18 J 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 ml 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 1s2s2p3s3p4s3d4p5s4d5p6s 4f5d6p7s5f6d7p 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 an atom’s electrons are distributed in its orbitals 3 ways: -complete specification of quantum numbers -shorthand notation from which the quantum numbers can be inferred -diagrammatic representation of orbital energy levels and their occupancy Electron-Electron Repulsion Hund’ Rule: The most stable electron configuration involving orbitals of equal energies is the one with the maximum numbers of electrons with the same spin orientation Orbitals with Nearly Equal Energies Orbitals Atomic Numbers Affected 4s, 3d 24, 29 5s, 4d 41,42,44,45,46,47 6s,5d, 4f 57,58,64,78,79 6d, 5f 89,91,92,93,96,110,111 Electron Configuration of Ions Atoms and Ions that have the same number of electrons are Isoelectric(n-1) Magnetic Properties of Atoms Diamagnetic: An atom or ion with all electrons paired is not attracted by string magnets Paramagnetic: An atom or ion with unpaired electrons is attractd to strong magnets Excited States -Ground state configuration is most stable arrangement of electrons 5.4 Periodicity of Atomic Properties As the principal quantum number n increases, atomic orbitals become larger and less stable. As the atomic number Z increases, any given atomic orbital becomes smaller and more stable. ATOMIC RADII Atomic size decreases from left to right and increases top to bottom of the periodic table.
-electron in the outermost occupided orbital, ns, feels an effective nuclear charge that changes very little across these blocks. Atomic size remains nearlu constant across each row of d and f blocks IONIZATION ENERGY First ionization energy increases from left to right across each row and decreases from top to bottom of each column of the periodic table -Atom absorbs a photon, the gain in energy promotes an electron to a less stable orbital. As electrons move into less stable orbitals, they have less electrical attraction for the nucleus. ELECTRON AFFINITY Electron affinity tends to become more negative from left to right across a row of the periodic table. -neutral atom can add an electron to form an anion. The energy change when an electron is added to an atom Sizes of Ions -Atomic cation is always smaller than the corresponding neutral atom -Atomic anion is always larger than the neutral atom 5.5 Energetic Of Ion Compounds 1. Sublimation 2. Bond energy 3. Ionization energy 4. Electron affinity 5. Condensation = 6. Formation E=kq1q2/r The lattic energy is the energy required to separate an ionic compound into its ions. The born-haber cycle is an example of the use of Hess’s law. 5.6 Ions and Chemical Periodicity Alkali metals (ns1) Alkaline Earth Metals (ns2) p-block elements
Chapter 6 6.1 Overview of Bonding -Electrical potential energy as a function if charges (q1, q2) and their separation distance - electrical energy between changed species is proportional to the magnitudes of the charges and inversely proportional to the distance between them. -Charges of opposite sign attract one another , but like charges repel 3 interactions -electrons and nuclei attract one another. Attractive interactions release energy, so an electron attracted to a nucleus is at a lower energy, therefore more stable, than free electron -electrons repel each other raising the energy and reducing the stability of a molecule -nuclei repel each other so these interactions also reduce the stability of a molecule The electrons and nuclei in a molecule balance these three interactions in a way that gives the molecule its greatest possible solubility. -Balance when electrons are concentrated between the nuclei called Covalent Bond Covalent Bond: mutual attraction of the bonding electrons to the two nuclei -Bond length : separation distance where the molecule is most stable -Bond length and strength are important properties of bonds that describe the characteristics of chemical bonds Electronegativity - A greater difference in electronegativity leads to a more polar bond -increase from lower left to upper right. -metals have lower electronegativity , nonmetals have high electronegativity’s 6.2 Lewis Structure -Formal charge= (valence electrons from free atom )-(valence electrons to that the atom in the lewis structure) - Most atoms other than hydrogen are most stable when associated with an octect of electrons -Only the valence electrons appear in lewis structures -Most likely lewis structures has the lowest formal charges 6.3 Molecular Shapes: Tetrahedral Systems SN 4= tetrahedral electron group geometry -Tetrahedron -trigonal pyramid -bent shape
- The steric number of an inner atom is the sum of the numbers of ligands plus the number of lone pairs -An electron group can be two electrons in a single bonds, four electrons in a double bond, six electrons in a triple bond, a pair of non-bonding electrons, or a single electron -molecular shape describes how the ligands, not the electron groups are arranged in space -Molecular shapes can be dervived from the steric number and number of ligands bonded to a central atom 6.4 Other Molecular Shapes SN ELECTRON GROUP GEOMETRY 2 Linear 3 Trigonal planar) 5 trigonal bipyramidal 6 Octahedral -linear -linear -bent shaped -Trigonal planar -sqaure planar -t shaped -trigonal bipyramidal -seesaw -square pyramidal -octahedral 6.5 Properties of Covelent Bonds Bond Angles Electron electron repulsion generated by non-bonding pairs is always greater than that generated by bonding pairs -lone pairs decrease bond angles Dipole Moments -affected by electronegativity differences and molecular geometry -diple moment: molecule with asymmetrical distrution of electron density (μ) Bond Length -affected by atomic radii, bond multiplicity, effective nuclear changes, and electronegativity differences
Bond Energy - bond strength increase as more electrons are shared between the atoms -bond strength increases as the electronegativity difference (∆x) between bonded atoms increase -bond strength decreases as bonds become longer Chapter 7 7.1 localized bonds 2 types of bonds: localized bond - between 2 atoms or localized in a single atom. -orbital of two types: 1. Bonding orbitals that has high electron density between two atoms 2.any atom is restricted to the region around a single atom, either in a bond to another atom or non bonding orbital Delocalized bonds Orbital overlap Bonding orbitals are constructed by combining atomic orbitals from adjacent atoms -Waves with the same size , generate a new wave= same sign waves add -aplitudes with opposite signs wave subtract= new wave amplitude is smaller, node Orbital overlap: when two orbitals are superimposed , one result is a new orbital that is composite if originals. -only valence orbitals are needed to describe bonding 7.2 Hybridization of Atomic Orbitals Charge Cloud Hybridization 2 sp 3 sp2 4 sp3 5 sp3d 6 sp3d2 7.3 Multiple Bonds σ bonds and π bonds - a sigma bond has a high electron density distributed symmetrically along the bond axis - pi bond has a high electron density concentrated above an below the bond axis 7.4 Molecular Orbittal Theory: Diatomic Molecules -the total number of molecular orbitals produced by a set of interacting atomic orbitals is equal to the number of interacting atomic orbitals -a positive MO results from positive overlap of atomic orbital wave functions and stabilizes the bond -antibonding MO results from negative overlap of atomic orbital wave functions and destabilizes the bond
Bond order= ½(number of electrons in bonding MO-number of electrons in antibonding MOs) -Para-magnetism arises from unpaired electrons 7.5 Three center π Orbitals Delocalized π orbitals Π-orbitals Non-bonding MO(πnb) Π* orbitals -Delocalized π orbitals result from side-by-side overlap of atomic π orbitals on more than two adjacent. -hybrid orbitals can be used to construct the molecular framework. Molecular orbitals can then be used to construct the π bonding system 7.6 Extended π Systems Ex. Vitamin A - A delocalized π system is present whenever p orbitals on more than two adjacent atoms are in position for side-by-side overlap -Molecules that have alternating single and double bonds are said to have conjugated π systems. -Delocalized π systems result increased stability and often-in coloured compounds. 7.7 Band Theory of Solids electrical conductivity: metallic bonding is the energy bands that arise because of the close spacing between orbital energies. Semiconductor & metalloids: band gap full in homo empty in lumo. n-type: band gap homo full and a bit in lumo (doped semiconductor) p-type: band gap lumo empty and homo is not completely full (doped semiconductor) -Metal conducts heat and electricity because their band gaps are so small -semiconductors have intermediate band gaps -a doped semiconductor has almost the same band structure as the pure material, but different electron populations in its bands Chapter 8 8.1 Effects of Intermolecular Forces Intermolecular attractive forces tend to hold molecules together in a condensed phase, but molecules that are moving fast enough can overcome these forces and move freely in the gas phase. -Average kinetic energy of motion is large enough, molecules remain separated from one another and the substance is a gas -intermolecular attractive forces are large enough, molecules remain close to one enough: liquid or solid
Melting and Boiling Points Melting points can be used as indicators of the strengths of intermolecular forces. -Average kinetic energy of molecular motion increases with absolute temperature Boiling point of a substance is the temperature at which -average kinetic energy of molecular motion balances the attractive energy of intermolecular attractions -pressure is 1 bar, that the temp is the normal boiling point Vaporization: Conversion of liquid into gas ( liquid vapourizes when milecules leave the liquid phase faster than they are captured from the gas. Condensation is gas to liquid -molecules in liquid able to move freely even though they do not escape from the liquid. -liquid cooled, molecular energies decrease -Temperature below melting point, the molecules become locked in place and liquid solidifies -pressure 1 bar, temperature is normal melting point -liquid freezes when liquid molecules have too little energy of motion to slide past one another -Solid melts when its molecules have enough energy of motion to move freely past one another -Intermolecular forces determine the normal boiling point and melting point -larger the intermolecular forces, the higher the melting point Boiling points and melting points depends on -strengths of intermolecular forces because the rate of escape and capture depend on the balance between molecular kinetic energies and intermolecular force attraction -large intermolecular forces must raid to a higher temperature before its molecules have sufficient kinetic energies to overcome those forces -small intermolecular forces must be cooled to a low temperature before its molecules have small enough kinetic energies to coalesce into a condensed phase 8.2 Types of Inter molecular Forces Magnitudes of intermolecular forces depend on whether each interacting particle is : CHARGED , POLAR, NONPOLAR -if both charged than force is ionic Ion-Dipole Forces Strongest intermolecular force is between ions and permanent dipole. Dispersion Forces Weakest Forces between two molecules but forces are collectively strong; -exist because the electron cloud of any molecule distorts easily -net attractive forces among molecules generated by all these induced charge inbalances -magnitude depends how easy it is to distort the electron cloud of a molecule. Polarizability is ease of distortion because distortion of an electron cloud generates a temporary polarity within the molecule (examining boiling points)
-Boiling point increase with the total number of electrons; larger and more polarizable -Molecular size increases with chain length and individual size of atoms Induced Dipole -ion approached a non-polar molecule, electron cloud in the molecule is distorted by electrostatic attraction or repulsion to the ion. -results; dipole in the non-polar molecule ion has induced a dipole in the molecule. Ion and induced dipole create a force of attraction Dipolar forces Dispersion forces in all molecules, some substances remain liquid at much higher temperatures than accounted for by dispersion force alone. Hydrogen Bonding Forces -highly electronegative atom with a lone pair of electrons shares its non-bonding electrons with a positively polarized hydrogen atom on another molecule. -exists among species having a polar bond between H and O, N, or F Inter molecular forces NAME OF ION Ion-dipole
Hydrogen bond
Dipole-dipole Ion-induced dipole
Dipole-induced dipole
Dispersion (London)
APPROX. POTENTIAL ATTRACTION ENERGY RANGE (KJ/MOL) 40-600 Charged ion is attracted to the oppositly charged end of the dipole Ex. Cl- -- H20 5-50 + charged H atom is attracted to lone pair on an electronegative atom Ex. (CH3)2CO—H2O 5-25 Mutual attraction of two dipole Ex. HCl--HCl 3-15 The ion induces a dipole in the other species, and is then attracted to it Ex.Ca 2+ --O2 2-10 A dipole induces a dipole in the other species, and is then attracted to it Ex. HCl—Cl2 0.1-5 Momentary shifts in electron clouds produce momentary dipoles, which attract each other Ex. He—HE
8.3 Liquids Capillary action , Meniscus, Viscosity, Surface Tension Vapour Pressure: Equillibrium: p=pvap -Substances with higher intermolecular forces have higher surface tensions, higher viscosities, and lower vapour pressures 8.4 Forces in Solids Solid type Molecular Attractive Dispersion Force Example Ar Melting 84 Point (k)
Molecular Dispersion+ dipolar HCl 158
Molecular Dispersion+dipolar+ H-Bonding H2O 273
Metallic Delocalized bonding Cu 1357
Network Ionic Covalent Electric al SiO2 NaCl 1983 1074
-Forcs in solids may include dispersion, dipolar, H-bonding, metallic, covalent, and ionic 8.5 Order in Solids Solid materials may be : -Amorphous Solid: Glass -Crystalline Solid Simple Cubic Body Centred Cubic Face Centred Cubic Unit cell is the building block of crystalline solids Close-Packed structure: cubic and Hexagonal Crystalline defects -One way to discuss ionic structures is to identify a crystal lattice for one set of ions and then describe how the other ion pack within the lattice of the first set -crystalline defects can profoundly alter the properties of a solid such as layers of iron atoms from sliding past one another and hardens iron into steel 8.6 Phase Changes Any phase change that results in increased molecular mobility requires that intermolecular forces be overcome. energy input. -Molar enthalpy of vaporization ∆Hvap: heat needed to vapourize 1 mole of a substance at its normal boiling point -Molar enthalpy of fusion, ∆Hfus: heat needed to melt 1 mole of substance at its normal melting point -Exothermic because reverse of endothermic phase change.
-Phase diagram shows which phases are the most stable at any conditions of temperature and pressure -Two phases are in equilibrium on a line on a phase diagram, and three phases coexist at the triple point Diagram of Phase Diagram:
Chapter 9 9.1 The Nature of Solutions -Solution is a homogenous mixture of two or more substances Components of Solutions Solvent determines the phase of solution, but solutes may be substantances that would normally exist in different phases. Solution Concentration -Any concentration value is a ratio of amounts Solvent type Gas Liquid Gas Diving gas (He,O2) Humid air (N2,O2,H2O) Liquid Carbonated Gas Vodka (H2O, CO2) (H2O,C2H5OH) Solid H2 Storage alloy Plastic(PVC (La, Ni, H2) dioctyphthalate)
Solid Air above I2 (N2, O2,I2) Saline solution (H2O,NaCl) Steel (Fe,C,Mn)
-Concentrations are expressed as amount of solute divided by the amount of solvent, or the amount of solution
9.2 Determinants of Solubility -miscible to insoluble -substance dissolves, but there is a limit to the amount of solute that will dissolve in a given amount of solvent -when limit has been reached saturated -concentration of a saturated solution is solubility of the substance in that particular substance in that particular solvent at a specified temperature -LIKE dissolves LIKE -substances that dissolve in each other usually have similar types of inter molecular interactions -metals do not dissolve in water -a few metals react with water and several react with aqueous solids but no metal will simply dissolve in water - any solution of another metal dissolves in mercury is called amalgam( metals close to mercury such as tin, gold are soluble in mercury) Solubility of Salts Ionic solids (salts) conatin cation and anion held in 3D lattices by strong electrical attractions 9.3 Characteristics of Aqueous Solutions -Molar enthalpy of solution (∆Hsol) is the net energy flow resulting from 3 factors; 1.seperating the ions, for which energy equal to the lattice energy must be supplied(quite large) 2.addiitional energy is needed to move solvent molecules apart to make room for the dissolving ions.Hydrogen bonds break (less energy required) 3. solvation of the ions by solvent molecules. Water molecules cluster around each ion,to give attractive ion-dipole interactions. Releases energy -enthalpies of solution and of dilution depend on the lattice energy of the salts as well as the hydration energies of its ions. - when a concentrated solution is diluted by adding water, energy is released and the temperature of the solution increases. The molar change accompanying process is enthalpy of dilution Gas-solution Equilibria - the solubilty of a gas in liquid depends in the partial pressure of the gas. Henry’s Law : [gas(aq)]eq= KH(pgas)eq 9.4 Colligative Properties - for a broad range of solutes, colligative properties depend on the amount of solute but not on the nature of the solute Presence of solute molecules caused changes in four common properties of solutions: vapour pressure, freezing point, boiling point, and osmotic properties -known as colligative properties
Vapour Pressure Reduction -vapour pressure drops when a solute is added to a liquid. Solute decreases the concentration of solvent molecules in the gas phase by reducing the rates of both evaporation and condensation. Raoult’s Law: Pvap,solution= XApvap,A Total vapour pressure above the solution is a sum of contributions from both solvent and solute ; solvent (A) Solute (B) Pvap,solution= XApvap,A + XBpvap,B Distillation -Differences in vapour pressure can be used to separate liquid mixtures by fractional distillation -azeotropes: liquid mixture is far from ideal and can not separate the components by distillation Boiling and Freezing Points -Boiling Point elevation and freezing point depression result from a decrease in solution vapour pressure compared with the pure solvent Van’t Hoff Factor ,i: -depression of the freezing point and elevation of boiling point
Kf : freezing point depression constant Kb : boiling point elevation constant Osmosis Movement of solvent molecules through a semipermeable membrane. -thin, pliable sheet of material, perforated with molecular scale holes -pressure increased needs to equalize the transfer rates = osmotic pressure (Π) pressure difference -pressure is exterted in both sides of semipermeable membrane but Π is the extra pressure that is exterted on the solution to maintain dynamic equilibrium Reverse Osmosis: used to purify water because liquid passes through semipermeable membrane is pure solvent. Determination of Molar Mass
Chapter 10 10.1 Hydrocarbons Alkanes – ane Alkenes –ene Alkynes –yne 1Methane 2Ethane 3Propane 4Butane 5Pentane 6Hexane 7Heptane 8Octane 9Nonane 10Decane IUPAC Names Methyl Ethyl Propyl 1-Methylethyl
Structure -CH3 -CH2CH3 -CH2CH2CH3 CH3CHCH3 -CH2CH2CH2CH3 CH3 - CH2CHCH3 CH3CHCH2CH3 CH3 CH3CCH3
butyl 1-Methylpropyl 1,1-Dimethylethyl
Cyclopropane (C3H6)
Cyclobutuane (C4H8)
Cyclopentane (C5H10)
Cyclohexane (C6H12)
-Cl Chloro -Br Bromo -F Fluoro -I Iodo -NH2 Amino -NO2 Nitro -OH Hydroxy 10.2 Aromatic Compounds
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 C1V1= C2V2
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=p1+p2+p3+p4 (each p is an individual pressure) Mole Fraction (X): Parts/million (ppm) and Parts/billion (ppb) PA= XApTotal 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 Liquid or solid Aqueous Solution Gas
Equation n= m/M n=cV n=pV/nT
2.5 Molecular View of Gases
Ekinetic =1/2mv2 Eavekinetic= 3RT/2NA
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 NO2 ( 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) O3 , SO2 , SO3 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
SI unit: kgm2/ss
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
qsurr=-qsys
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
Wsurr=-Wsys
First Law of Thermodynamics Energy(E) is exchanged in two ways : heat or work ∆E=q+w
∆Esys=-∆Esurr 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 Standard Enthalpy of formation (∆Hof): formation reaction Formation reaction= 1 mol of a substance produced from elements 1/2N2(g) +1/2 O2(g) N(g) ∆Hof reaction =Σcoeff products∆Hofproducts -Σcoeff 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-1 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= 2.18x10-18J/ n2 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(ml) (ms) -magnetic quantum number ml: indexes restrictions l=1 ml: -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 ms: indexes the behavior -ms are +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 xray radiation coming from the sun. N2+hvN2+ +eO2+hv O+O Reactions are unstable and will recombine releasing energy in the form of heat; N2++e-N2 +heat O+OO2 +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 180240nm 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): HH++eIEH: 2.18x10^-18 J He+ He2+ +eIEHe+: 8.72 x10^ -18 J 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 ml 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 1s2s2p3s3p4s3d4p5s4d5p6s 4f5d6p7s5f6d7p 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 an atom’s electrons are distributed in its orbitals 3 ways: -complete specification of quantum numbers -shorthand notation from which the quantum numbers can be inferred -diagrammatic representation of orbital energy levels and their occupancy Electron-Electron Repulsion Hund’ Rule: The most stable electron configuration involving orbitals of equal energies is the one with the maximum numbers of electrons with the same spin orientation Orbitals with Nearly Equal Energies Orbitals Atomic Numbers Affected 4s, 3d 24, 29 5s, 4d 41,42,44,45,46,47 6s,5d, 4f 57,58,64,78,79 6d, 5f 89,91,92,93,96,110,111 Electron Configuration of Ions Atoms and Ions that have the same number of electrons are Isoelectric(n-1) Magnetic Properties of Atoms Diamagnetic: An atom or ion with all electrons paired is not attracted by string magnets Paramagnetic: An atom or ion with unpaired electrons is attractd to strong magnets Excited States -Ground state configuration is most stable arrangement of electrons 5.4 Periodicity of Atomic Properties As the principal quantum number n increases, atomic orbitals become larger and less stable. As the atomic number Z increases, any given atomic orbital becomes smaller and more stable. ATOMIC RADII Atomic size decreases from left to right and increases top to bottom of the periodic table.
-electron in the outermost occupided orbital, ns, feels an effective nuclear charge that changes very little across these blocks. Atomic size remains nearlu constant across each row of d and f blocks IONIZATION ENERGY First ionization energy increases from left to right across each row and decreases from top to bottom of each column of the periodic table -Atom absorbs a photon, the gain in energy promotes an electron to a less stable orbital. As electrons move into less stable orbitals, they have less electrical attraction for the nucleus. ELECTRON AFFINITY Electron affinity tends to become more negative from left to right across a row of the periodic table. -neutral atom can add an electron to form an anion. The energy change when an electron is added to an atom Sizes of Ions -Atomic cation is always smaller than the corresponding neutral atom -Atomic anion is always larger than the neutral atom 5.5 Energetic Of Ion Compounds 1. Sublimation 2. Bond energy 3. Ionization energy 4. Electron affinity 5. Condensation = 6. Formation E=kq1q2/r The lattic energy is the energy required to separate an ionic compound into its ions. The born-haber cycle is an example of the use of Hess’s law. 5.6 Ions and Chemical Periodicity Alkali metals (ns1) Alkaline Earth Metals (ns2) p-block elements
Chapter 6 6.1 Overview of Bonding -Electrical potential energy as a function if charges (q1, q2) and their separation distance - electrical energy between changed species is proportional to the magnitudes of the charges and inversely proportional to the distance between them. -Charges of opposite sign attract one another , but like charges repel 3 interactions -electrons and nuclei attract one another. Attractive interactions release energy, so an electron attracted to a nucleus is at a lower energy, therefore more stable, than free electron -electrons repel each other raising the energy and reducing the stability of a molecule -nuclei repel each other so these interactions also reduce the stability of a molecule The electrons and nuclei in a molecule balance these three interactions in a way that gives the molecule its greatest possible solubility. -Balance when electrons are concentrated between the nuclei called Covalent Bond Covalent Bond: mutual attraction of the bonding electrons to the two nuclei -Bond length : separation distance where the molecule is most stable -Bond length and strength are important properties of bonds that describe the characteristics of chemical bonds Electronegativity - A greater difference in electronegativity leads to a more polar bond -increase from lower left to upper right. -metals have lower electronegativity , nonmetals have high electronegativity’s 6.2 Lewis Structure -Formal charge= (valence electrons from free atom )-(valence electrons to that the atom in the lewis structure) - Most atoms other than hydrogen are most stable when associated with an octect of electrons -Only the valence electrons appear in lewis structures -Most likely lewis structures has the lowest formal charges 6.3 Molecular Shapes: Tetrahedral Systems SN 4= tetrahedral electron group geometry -Tetrahedron -trigonal pyramid -bent shape
- The steric number of an inner atom is the sum of the numbers of ligands plus the number of lone pairs -An electron group can be two electrons in a single bonds, four electrons in a double bond, six electrons in a triple bond, a pair of non-bonding electrons, or a single electron -molecular shape describes how the ligands, not the electron groups are arranged in space -Molecular shapes can be dervived from the steric number and number of ligands bonded to a central atom 6.4 Other Molecular Shapes SN ELECTRON GROUP GEOMETRY 2 Linear 3 Trigonal planar) 5 trigonal bipyramidal 6 Octahedral -linear -linear -bent shaped -Trigonal planar -sqaure planar -t shaped -trigonal bipyramidal -seesaw -square pyramidal -octahedral 6.5 Properties of Covelent Bonds Bond Angles Electron electron repulsion generated by non-bonding pairs is always greater than that generated by bonding pairs -lone pairs decrease bond angles Dipole Moments -affected by electronegativity differences and molecular geometry -diple moment: molecule with asymmetrical distrution of electron density (μ) Bond Length -affected by atomic radii, bond multiplicity, effective nuclear changes, and electronegativity differences
Bond Energy - bond strength increase as more electrons are shared between the atoms -bond strength increases as the electronegativity difference (∆x) between bonded atoms increase -bond strength decreases as bonds become longer Chapter 7 7.1 localized bonds 2 types of bonds: localized bond - between 2 atoms or localized in a single atom. -orbital of two types: 1. Bonding orbitals that has high electron density between two atoms 2.any atom is restricted to the region around a single atom, either in a bond to another atom or non bonding orbital Delocalized bonds Orbital overlap Bonding orbitals are constructed by combining atomic orbitals from adjacent atoms -Waves with the same size , generate a new wave= same sign waves add -aplitudes with opposite signs wave subtract= new wave amplitude is smaller, node Orbital overlap: when two orbitals are superimposed , one result is a new orbital that is composite if originals. -only valence orbitals are needed to describe bonding 7.2 Hybridization of Atomic Orbitals Charge Cloud Hybridization 2 sp 3 sp2 4 sp3 5 sp3d 6 sp3d2 7.3 Multiple Bonds σ bonds and π bonds - a sigma bond has a high electron density distributed symmetrically along the bond axis - pi bond has a high electron density concentrated above an below the bond axis 7.4 Molecular Orbittal Theory: Diatomic Molecules -the total number of molecular orbitals produced by a set of interacting atomic orbitals is equal to the number of interacting atomic orbitals -a positive MO results from positive overlap of atomic orbital wave functions and stabilizes the bond -antibonding MO results from negative overlap of atomic orbital wave functions and destabilizes the bond
Bond order= ½(number of electrons in bonding MO-number of electrons in antibonding MOs) -Para-magnetism arises from unpaired electrons 7.5 Three center π Orbitals Delocalized π orbitals Π-orbitals Non-bonding MO(πnb) Π* orbitals -Delocalized π orbitals result from side-by-side overlap of atomic π orbitals on more than two adjacent. -hybrid orbitals can be used to construct the molecular framework. Molecular orbitals can then be used to construct the π bonding system 7.6 Extended π Systems Ex. Vitamin A - A delocalized π system is present whenever p orbitals on more than two adjacent atoms are in position for side-by-side overlap -Molecules that have alternating single and double bonds are said to have conjugated π systems. -Delocalized π systems result increased stability and often-in coloured compounds. 7.7 Band Theory of Solids electrical conductivity: metallic bonding is the energy bands that arise because of the close spacing between orbital energies. Semiconductor & metalloids: band gap full in homo empty in lumo. n-type: band gap homo full and a bit in lumo (doped semiconductor) p-type: band gap lumo empty and homo is not completely full (doped semiconductor) -Metal conducts heat and electricity because their band gaps are so small -semiconductors have intermediate band gaps -a doped semiconductor has almost the same band structure as the pure material, but different electron populations in its bands Chapter 8 8.1 Effects of Intermolecular Forces Intermolecular attractive forces tend to hold molecules together in a condensed phase, but molecules that are moving fast enough can overcome these forces and move freely in the gas phase. -Average kinetic energy of motion is large enough, molecules remain separated from one another and the substance is a gas -intermolecular attractive forces are large enough, molecules remain close to one enough: liquid or solid
Melting and Boiling Points Melting points can be used as indicators of the strengths of intermolecular forces. -Average kinetic energy of molecular motion increases with absolute temperature Boiling point of a substance is the temperature at which -average kinetic energy of molecular motion balances the attractive energy of intermolecular attractions -pressure is 1 bar, that the temp is the normal boiling point Vaporization: Conversion of liquid into gas ( liquid vapourizes when milecules leave the liquid phase faster than they are captured from the gas. Condensation is gas to liquid -molecules in liquid able to move freely even though they do not escape from the liquid. -liquid cooled, molecular energies decrease -Temperature below melting point, the molecules become locked in place and liquid solidifies -pressure 1 bar, temperature is normal melting point -liquid freezes when liquid molecules have too little energy of motion to slide past one another -Solid melts when its molecules have enough energy of motion to move freely past one another -Intermolecular forces determine the normal boiling point and melting point -larger the intermolecular forces, the higher the melting point Boiling points and melting points depends on -strengths of intermolecular forces because the rate of escape and capture depend on the balance between molecular kinetic energies and intermolecular force attraction -large intermolecular forces must raid to a higher temperature before its molecules have sufficient kinetic energies to overcome those forces -small intermolecular forces must be cooled to a low temperature before its molecules have small enough kinetic energies to coalesce into a condensed phase 8.2 Types of Inter molecular Forces Magnitudes of intermolecular forces depend on whether each interacting particle is : CHARGED , POLAR, NONPOLAR -if both charged than force is ionic Ion-Dipole Forces Strongest intermolecular force is between ions and permanent dipole. Dispersion Forces Weakest Forces between two molecules but forces are collectively strong; -exist because the electron cloud of any molecule distorts easily -net attractive forces among molecules generated by all these induced charge inbalances -magnitude depends how easy it is to distort the electron cloud of a molecule. Polarizability is ease of distortion because distortion of an electron cloud generates a temporary polarity within the molecule (examining boiling points)
-Boiling point increase with the total number of electrons; larger and more polarizable -Molecular size increases with chain length and individual size of atoms Induced Dipole -ion approached a non-polar molecule, electron cloud in the molecule is distorted by electrostatic attraction or repulsion to the ion. -results; dipole in the non-polar molecule ion has induced a dipole in the molecule. Ion and induced dipole create a force of attraction Dipolar forces Dispersion forces in all molecules, some substances remain liquid at much higher temperatures than accounted for by dispersion force alone. Hydrogen Bonding Forces -highly electronegative atom with a lone pair of electrons shares its non-bonding electrons with a positively polarized hydrogen atom on another molecule. -exists among species having a polar bond between H and O, N, or F Inter molecular forces NAME OF ION Ion-dipole
Hydrogen bond
Dipole-dipole Ion-induced dipole
Dipole-induced dipole
Dispersion (London)
APPROX. POTENTIAL ATTRACTION ENERGY RANGE (KJ/MOL) 40-600 Charged ion is attracted to the oppositly charged end of the dipole Ex. Cl- -- H20 5-50 + charged H atom is attracted to lone pair on an electronegative atom Ex. (CH3)2CO—H2O 5-25 Mutual attraction of two dipole Ex. HCl--HCl 3-15 The ion induces a dipole in the other species, and is then attracted to it Ex.Ca 2+ --O2 2-10 A dipole induces a dipole in the other species, and is then attracted to it Ex. HCl—Cl2 0.1-5 Momentary shifts in electron clouds produce momentary dipoles, which attract each other Ex. He—HE
8.3 Liquids Capillary action , Meniscus, Viscosity, Surface Tension Vapour Pressure: Equillibrium: p=pvap -Substances with higher intermolecular forces have higher surface tensions, higher viscosities, and lower vapour pressures 8.4 Forces in Solids Solid type Molecular Attractive Dispersion Force Example Ar Melting 84 Point (k)
Molecular Dispersion+ dipolar HCl 158
Molecular Dispersion+dipolar+ H-Bonding H2O 273
Metallic Delocalized bonding Cu 1357
Network Ionic Covalent Electric al SiO2 NaCl 1983 1074
-Forcs in solids may include dispersion, dipolar, H-bonding, metallic, covalent, and ionic 8.5 Order in Solids Solid materials may be : -Amorphous Solid: Glass -Crystalline Solid Simple Cubic Body Centred Cubic Face Centred Cubic Unit cell is the building block of crystalline solids Close-Packed structure: cubic and Hexagonal Crystalline defects -One way to discuss ionic structures is to identify a crystal lattice for one set of ions and then describe how the other ion pack within the lattice of the first set -crystalline defects can profoundly alter the properties of a solid such as layers of iron atoms from sliding past one another and hardens iron into steel 8.6 Phase Changes Any phase change that results in increased molecular mobility requires that intermolecular forces be overcome. energy input. -Molar enthalpy of vaporization ∆Hvap: heat needed to vapourize 1 mole of a substance at its normal boiling point -Molar enthalpy of fusion, ∆Hfus: heat needed to melt 1 mole of substance at its normal melting point -Exothermic because reverse of endothermic phase change.
-Phase diagram shows which phases are the most stable at any conditions of temperature and pressure -Two phases are in equilibrium on a line on a phase diagram, and three phases coexist at the triple point Diagram of Phase Diagram:
Chapter 9 9.1 The Nature of Solutions -Solution is a homogenous mixture of two or more substances Components of Solutions Solvent determines the phase of solution, but solutes may be substantances that would normally exist in different phases. Solution Concentration -Any concentration value is a ratio of amounts Solvent type Gas Liquid Gas Diving gas (He,O2) Humid air (N2,O2,H2O) Liquid Carbonated Gas Vodka (H2O, CO2) (H2O,C2H5OH) Solid H2 Storage alloy Plastic(PVC (La, Ni, H2) dioctyphthalate)
Solid Air above I2 (N2, O2,I2) Saline solution (H2O,NaCl) Steel (Fe,C,Mn)
-Concentrations are expressed as amount of solute divided by the amount of solvent, or the amount of solution
9.2 Determinants of Solubility -miscible to insoluble -substance dissolves, but there is a limit to the amount of solute that will dissolve in a given amount of solvent -when limit has been reached saturated -concentration of a saturated solution is solubility of the substance in that particular substance in that particular solvent at a specified temperature -LIKE dissolves LIKE -substances that dissolve in each other usually have similar types of inter molecular interactions -metals do not dissolve in water -a few metals react with water and several react with aqueous solids but no metal will simply dissolve in water - any solution of another metal dissolves in mercury is called amalgam( metals close to mercury such as tin, gold are soluble in mercury) Solubility of Salts Ionic solids (salts) conatin cation and anion held in 3D lattices by strong electrical attractions 9.3 Characteristics of Aqueous Solutions -Molar enthalpy of solution (∆Hsol) is the net energy flow resulting from 3 factors; 1.seperating the ions, for which energy equal to the lattice energy must be supplied(quite large) 2.addiitional energy is needed to move solvent molecules apart to make room for the dissolving ions.Hydrogen bonds break (less energy required) 3. solvation of the ions by solvent molecules. Water molecules cluster around each ion,to give attractive ion-dipole interactions. Releases energy -enthalpies of solution and of dilution depend on the lattice energy of the salts as well as the hydration energies of its ions. - when a concentrated solution is diluted by adding water, energy is released and the temperature of the solution increases. The molar change accompanying process is enthalpy of dilution Gas-solution Equilibria - the solubilty of a gas in liquid depends in the partial pressure of the gas. Henry’s Law : [gas(aq)]eq= KH(pgas)eq 9.4 Colligative Properties - for a broad range of solutes, colligative properties depend on the amount of solute but not on the nature of the solute Presence of solute molecules caused changes in four common properties of solutions: vapour pressure, freezing point, boiling point, and osmotic properties -known as colligative properties
Vapour Pressure Reduction -vapour pressure drops when a solute is added to a liquid. Solute decreases the concentration of solvent molecules in the gas phase by reducing the rates of both evaporation and condensation. Raoult’s Law: Pvap,solution= XApvap,A Total vapour pressure above the solution is a sum of contributions from both solvent and solute ; solvent (A) Solute (B) Pvap,solution= XApvap,A + XBpvap,B Distillation -Differences in vapour pressure can be used to separate liquid mixtures by fractional distillation -azeotropes: liquid mixture is far from ideal and can not separate the components by distillation Boiling and Freezing Points -Boiling Point elevation and freezing point depression result from a decrease in solution vapour pressure compared with the pure solvent Van’t Hoff Factor ,i: -depression of the freezing point and elevation of boiling point
Kf : freezing point depression constant Kb : boiling point elevation constant Osmosis Movement of solvent molecules through a semipermeable membrane. -thin, pliable sheet of material, perforated with molecular scale holes -pressure increased needs to equalize the transfer rates = osmotic pressure (Π) pressure difference -pressure is exterted in both sides of semipermeable membrane but Π is the extra pressure that is exterted on the solution to maintain dynamic equilibrium Reverse Osmosis: used to purify water because liquid passes through semipermeable membrane is pure solvent. Determination of Molar Mass
Chapter 10 10.1 Hydrocarbons Alkanes – ane Alkenes –ene Alkynes –yne 1Methane 2Ethane 3Propane 4Butane 5Pentane 6Hexane 7Heptane 8Octane 9Nonane 10Decane IUPAC Names Methyl Ethyl Propyl 1-Methylethyl
Structure -CH3 -CH2CH3 -CH2CH2CH3 CH3CHCH3 -CH2CH2CH2CH3 CH3 - CH2CHCH3 CH3CHCH2CH3 CH3 CH3CCH3
butyl 1-Methylpropyl 1,1-Dimethylethyl
Cyclopropane (C3H6)
Cyclobutuane (C4H8)
Cyclopentane (C5H10)
Cyclohexane (C6H12)
-Cl Chloro -Br Bromo -F Fluoro -I Iodo -NH2 Amino -NO2 Nitro -OH Hydroxy 10.2 Aromatic Compounds
