Chem 113 notes - General chemistry, Organic Chemistry and Physical chemistry PDF

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General chemistry, Organic Chemistry and Physical chemistry ...

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Lecture 1: Isotopes – atoms having the same atomic number but different mass number. - Same group of elements have the same outer shell configuration - Outermost is known as valence electron and inner shell is “core electrons” - Metalloids are bond between metal to non-metal. - Group 17 and 18 are halogens and non- metal - Fission – means large nucleus broken into 2 small nuclei - Fusion – is reaction when 2 small nuclei combine to make a bigger nucleus.

Lecture 2: ATOM MODELS 1. Plum pudding model debunked by Rutherford 2. Rutherford – nucleus in the middle with electrons around it – and repulsion is discovered. 3. Bohr model – nucleus in the middle has orbits distance - Energy of electrons are quantized - Electrons are attracted to nucleus by electrostatic - Energy level with higher energy are further from the nucleus 4. Quantum mechanic – electrons have bother particle and wave properties - Probability of finding the -e in a particular region of space can be calculated - Energy level is divided into subshells which has orbitals S orbital P orbital – py,px,pz D orbital – dxy, dyz, dz2 Describing -e - Most electrons in atom happens when -e occupy the lowest energy orbitals, - Each orbital can accommodate 2 -e 1s2, 2s2, 2p6, 3s2,3p6, 4s2, 3d10, 4p6 Paramagnetic – is when a species has unpaired -e Diamagnetic – is when a species has all -e pairs Ground state- lowest possible energy level Excited state – when atom absorbs energy and -e is promoted to a higher energy level After elements 20 – valence -e are those in the partly filled level and the most recently filled s- level.

Lecture 3: Periodic trends - Atomic properties depends on how strongly valence -e attract to the nucleus. - Electrostatic – force of attraction between +ve and -ve charge - The strength of the attraction depends on : size of +ve charge and no. of protons in nucleus, distance between the valence -e and the nucleus, no. of core -e (shielding)

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When comparing elements in the same row of the periodic table, valence -e are in the same shell so the size of the nuclear charge is the key factor in determining the trend.

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When elements are in different periods in the periodic table the valence -e are in different shells so the distance of the shell form the nucleus and the amount of shielding are both key factors to explain the trend.

Accounting for the trends: - Atomic radius – is the distance between nucleus of atom and its outer -e. - Across period – valence -e in the same shell so distance to nucleus is the same. - Valence -e in the same shell so no. of core -e is the same (shielding effect of the core -e is the same) - Across a period the valence -e become more strongly held so the radius decreases. – increased +ve charge increases attraction between nucleus and the -e so -e are drawn tighter around the nucleus = harder to remove -e - Down a group – the valence -e become less strongly held so the radius increase – decreased attraction of outer -e to the nucleus and the increased no. of shell. = results easier to remove -e. -

Isoelectronic – have the same no. of -e and -e arrangement

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Cation - +ve charged – always smaller than a neutral atoms

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Anion - -ve charged – larger than neutral atom.

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Electronegativity – power of an atom to attract -e to itself when its part of a molecule . Across (left to right) – increase. Down (top to bottom) – decrease

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Step to consider when accounting for trends: Write electron configuration for atom/ ions Identify similarities or difference If valence shell is the same then compare nuclear charge If the valence shell is different then compare then compare to distance to nucleus and shielding by core -e

Lecture 4: Counting atoms - Avagardo’s number = 6x1025 Equations n(mol) = m (g)/ M (gmol-1) Step 1: Mass to moles Step 2: Moles to Moles Step 3: Moles to mass

m (g)= n(mol) x M(gmol-1 )

c(molL-1) = n (mol)/ V (L) Step 1: change mass to moles or change concentration to moles ( depends on the question) Step 2: relate moles (calculated) to moles (unknown) using equation Step 3: Change moles to mass or change moles to concentration (depends on the question)

Dilution = V(initial) / V(final)

c(final) = c x d ( dilution)

Lecture 5: Bonding Chemical bond – is a lasting attraction between atoms that enables the formation of chemical compounds, Force a = Charge A x charge B / distance -

A bond is formed – when attractive force and the repulsive forces are balanced

There are 3 types of bonding 1) Ionic – electrostatic attraction between +ve and -ve ion. -It could be metal or non-metal 2) Covalent – bond formed when -e are shared between the nuclei of 2 atoms - occurs when there is little differences in tendency of the bonded atoms to gain or lose -e eg non-metals

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3) Metallic- occurs between metal atoms Valence -e on metals are loosely held and tend to be delocalized and shared among large number of metal atoms. Metal atoms without valence -e are cations – so bond can be described as “attraction between metal cations and surrounding valence -e”

Polarity – is when the spread of -e are uneven Dipole – occurs when one end becomes more partially negative than the other. Lewis structure show the distribution of all valence -e in the molecule. How to draw Lewis structure. 1) Count the valence -e 2) Assemble the bonding framework using single bonds 3) Place three non-bonding pairs of -e on each outer atom except H 4) Assign the remaining valence -e to inner atoms 5) If central atom doesn’t have at least 8 -e convert a lone pair to a double bond (octet) Octet – is a set of four -e associated with an atom Atoms from period 3 are able to accommodate more than 8 -e then its no octet rule.

Shape of molecules - There is no direct correlation between the formula of a molecule and its shape - Molecular shape – is determined by the relative position of the atom. Different types of shapes - Linear - Trigonal planar - Tetrahedral - Trigonal pyramid - Octahedral

To determine the shape: - Use the lewis structure of the molecule to determine the no. of -e regions around the central atom - Then decide the geometry (shape) that will minimize repulsion between the -e clouds - Count the no. of bonds and describe the shape they form (when there are lone pairs) - Always consider the number of -e clouds on the central atom before discussing the lone pairs of -e To workout the shape: 1) Draw lewis structure of the molecule 2) Count the number of -e regions on the central atom ( bonding +non-bonding ) 3) Determine the arrangement of -e pairs around the central atom 4) Count no. of bonded regions of -e 5) Determine the shape of the molecule 6) Justify the answer by 1. Stating no. of e clouds 2. Stating no. of bonds 3. Describing the shape that will be observed. Bond angles – determined by the geometry around the central atom Linear 180o. Trigonal planar 120o Tetrahedral 109o. Octahedral 90o Trigonal bipyramid 90o and 120o.

Lecture 6 – polarity molecules - If 2 different atom has different electronegativity this will result polarity and uneven spread of -e Polar bonds – depends on the pulling power of the atoms on the -e electronegativity Polar molecules – depends on the arrangements of the polar bonds – if they are evenly arranged the central atom the polarity will cancel. This leads to bond dipole – one end is -ve and the other is +ve Polar covalent – is explained by looking at the nature of the covalent bond Non-polar – opposite ends of the molecule do not have different charges

Polar molecules - Bond dipoles/ dipole moment – is charge imbalance To predict polarity: - Determine the polarity at the bond – which end is partially +ve and which end is partilly -ve. - Determine the shape of the molecule - Decide whether the shape of molecules cancels the bond polarity. Lecture 7 and 8 – organic molecules Constitutional isomers – are compounds with the same molecular formula but different connectivity Structural isomers – they have the same molecular formula but different order of attachment of atoms Number of isomers drawn? Method 1) Draw the simplest long chain 2) Draw a new chain with 1 less C atom and add extra C atom as a side chain 3) Consider the other unique position for the side -chain 4) Repeat 3, then 2 and 3 until no more unique isomers can be drawn

To name: 1) Length of the longest continuous carbon chain 2) Name the position of side-chains 3) Names and position of functional groups. Side-chain names: Methyl CH3 Ethyl CH2CH3 Propyl CH2CH2CH3 Butyl CH2CH2CH2CH3 To name: 1) Parent name 2) Identify the side chain and give the lowest possible number. Cycloalkanes – ring structures Naming cycloalkanes – when the ring is lager than its side chains, number the ring to give the lowest possible number Alkene – double bond CnHn.

Alkyne- triple bond CnH2n-2

The multiple bond must have the lowest position number in the carbon ring

Method naming: -

Name the longest carbon chain containing double bond on the molecule shown Identify the side chains and their position

Stereoisomerism Optical isomers (enantiomer) – have non- superimposable mirror images - These occur when there is the same order of attachment of atoms but different spatial orientation - Alkanes don’t have geometric isomers because there is free rotation around single bonds Cycloalkane – can have geometric isomers because the ring means freedom of rotation is lost – which means that if there is more than one group attached to the ring there is a potential for 2 different isomers to form Wedge = out of the page Dash = into the page Geometric isomers (ciz – tranz) – have different orientations of groups around a bond with no free rotation Geometric isomers – Cis or tranz - These occur when the arrangement of atoms in space can’t be change by rotation around the C-C bonds, such as in cycloalkanes and alkene - Some alkenes have geometric isomers since there is no free rotation about the double bond Method: - Orientation of atoms in the parent chain determines if the alkene is cis or trans - When the carbon atoms of the parent chain atoms are on the same side of the double bond called cis, on opposite sides its trans Enantiomers - Have atoms with the same points of attachment But they have non-superimposable mirror images - The molecule is chiral (4 different substituents is chiral center ) - They are different molecules A chiral center in molecules will have 4 different atoms or groups of atoms attached Functional groups and names: Method: 1) Number C atoms as denoted in longest C chain and side chain (prefix) 2) Nature of the C-C bond ie single, double – ene, triple – yne (infex) 3) Class of compound – functional group ending (suffix)

Functional group must take the lowest number possible in the chain Polarity rules for functional groups: 1) Where there is more than 1 functional group, priorities apple 2) Highest priority have name ending 3) Where there are multiple functional groups -keep the wholename of the parent chain and add di-,tri-, tetra- , penta, etc in front of the functional group suffix 4) Note extra requirements for esters, amides, and amines

Priority rule Carboxylic acid > acid chloride > ester> amide> aldehyde> Ketone> alcohol > amine > alkene> alkyne> alkane> haloalkane

Pre-Lecture 9 There are 3 types of bonding 1) Ionic bonding 2) Covalent bonding 3) Metallic bonding

Four types of solids Structure Particles Ionic Ions Metallic Cations and delocalized Ions Network Atoms Molecular Molecules

Interactions Ionic bond Metallic bond

Examples NaCl, CaF2 Cu, Fe

Covalent bond Intermolecular forces

C ( graphite) CO2, H2O

Melting and boiling points - Change of state (solid to liquid or liquid to solid) requires energy to separate particles - Stronger bonds require more energy

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Conductance Needs mobile charged particles to carry the charge

Hardness – depends on how easily the particles can be displaced from one another by physical pressure Solubility – Dissolving one substance in another depends on separating particle in the solute and in the solvent

Intermolecular – is forces between molecules The stronger the intermolecular force the higher the melting and boiling points Intermolecular forces – are electrostatic attractions between +ve and -ve charges - Dipole- dipole force – polar molecules align themselves so that +ve end of one molecules faces towards -ve end of an adjacent molecule Intermolecular forces are electrostatic attractions between dipoles and are much weaker than covalent bonds 1) Temporary dipole – dipole attraction - Created by the random movement of -e in a molecule causing uneven distribution of -ve charge which induces a dipole in another molecule creating attractive forces between them 2) Permanent dipole – attraction - Is when polar molecules have permanent dipoles - Attraction is between partially +ve end and partially -ve end 3) Hydrogen bonding - Dipole -dipole interaction between hydrogen atom which is bonded to a N, O, or F in one molecule. Lecture 9 What influences the size of temporary dipole attractions? - The size of electron and the shape - Forces increase in strength with the size of the molecule Temporary dipole depends on how easy it is to distort the -e cloud of a molecule

Permanent dipole – when polar molecules are attracted to each other by permanent dipoledipole forces . Comparing forces to account for boiling points – to account for boiling points differences or predict boiling point differences 1) Compare the size of temporary dipole-dipole attractions 2) Check whether the molecules are polar or non-polar – if they are polar then the dipole – dipole attractions are added to the temporary dipole- dipole attraction 3) Check for hydrogen bonding – remember that it need to be bonded to O,N, or F

Molecular solids – weakest - Has intermolecular forces between molecules

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Has low melting points Soft – weak forces Don’t conduct because particles has no charge

Ionic compounds - Held together by strong attraction between +ve and -ve charges - Can’t be separated - Loss of -e is an endothermic process Properties of Ionic compound 1 - High melting points - Brittle - Conducts electricity in liquid solution but not in solid - Has high melting points -

Ionic solids are hard but brittle

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They are brittle when pressure is applied

Dissolving ionic solids - Known as salts

Properties of metal 1 - High melting points - Hard but malleable and ductile - Conducts in solid and liquid - Arrangement +ve ions and delocalized valence -e Properties of metal 2 - Melting points are high - Metallic bond is strong – strength depends on size of ions and charge of ions Properties of metal 3 - Metals are good conductors of heat and electricity – in solid and liquid state - Because of delocalized valence -e are able to move when metal is placed between the terminals

Strategy for answers - Identify the type of solids - Describe the structure - Related structural/ bonding feature to relevant property

Network solids (covalent) - High melting points

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Hard Don’t conducts. Eg. Diamond and graphite

Lecture 10: Energy and Chemical processes. Spontaneously – once they have started they continue without involvement of outside factors Non- spontaneously – can be made to occur with different energy conditions Predicting spontaneity 1st law of thermodynamics – total energy of universe is constant and can’t be changed. Can’t be created or destroyed it can only be transferred. 2nd law of thermodynamics – in a spontaneous process, the entropy of the universe is increasing – driving power of natural charge. Thermodynamics definitions Thermodynamics – understanding and predicting chemical and physical change Enthalpy – total heat content of a system – given symbol H Enthalpy change – of a system us the heat transferred when a reaction is carried out at constant pressure. Endothermic – energy is absorbed Exothermic – energy is released Temperature – measures how hot or cold something compared to other things is a measure of the average kinetic energy of the particle Entropy – measure of randomness or disorder Gibbs energy – combines enthalpy and entropy of a system – change in gibbs energy to predict spontaneity Kelvin- temperature scale for thermodynamics Internal energy (U) – sum of all the forms of energy in a system Exothermic (release heat) H(final)- H(initial) = 0 so enthalpy is +ve Calculating enthalpy change 1) From bond energy data – the net energy (enthalpy) of reaction comes from making and breaking chemical bonds 2) Using hess’s law

Lecture 11: Entropy and spontaneity

1. Standard enthalpy formation - formation of a 1 mol compound 2. Standard enthalpy of combustion – burning 1 mol of a compound in oxygen under standard conditions 3. Standard enthalpies of vaporization, fussion Using Heats of formation to calculate enthalpy of reactions - Since enthalpy is a state functions, it’s determined only by the initial state of the reaction and final state of the reaction, not the process that occurs in between TiCl4(l) + 2H2O TiO2 (s) + 4HCl Step 1 TiCl4(l)  Ti(s) + 2Cl2(g) H2O  2H2 + O2 Step 2 Ti(s) + O2 (g)  TiO2 (s) 2H2 + 2Cl2(g)  4HCl Lecture 12: Oxidation and reductions - If these reaction electrons transferred from one species to another - Oxidation – occurs when electrons are removed/ lost from species - Reductions – occurs when electrons are added. Oxidant is reduced during reactions Reductant – gives -e

Lecture 13: Acid and base -

Know acid and conjugate base Describe weak and strong acid – know how to relate it to pH Relate Ka,kb, pka and pkb

Kb=kw/ka where kw =1x10^-14 To convert pka to ka: Pka=-logka and ka=10^-pka To convert ka to pka Pka = -logka Salt dilution - Salt dissolves one or both of the ions may react with water and affect the pH of the solution Neutral solution - Cation and anion are from strong acid and base

Acidic solution - Anion is from strong acid and cation is from weak base Basic solution - Anions from weak acid and cation from a strong base To determine the pH of the salt solution, begin by writing the two equation- one for dissolving the other for reaction

Buffer - Solution which minimizes change in pH when some strong acid or base is added - Contains a large amount of both weak acid and its conjugate base - Optimum weak base Note: buffer solution ,the acid, and the conjugate base are both additions to the solution so the conjugate base component does not just come from the acid ionization Titration curves: - Shows how pH solution change at the volume of titrant added during a titration - Account the differences in the pH at the beginning of each titration curve in terms of the species present - Account the differences in equivalent points by writing the equation for each titration reaction and considering the species present at each equivalent point. - Similarities in the pH after the equivalence point of each titration - Explain why there is or isn’t buffer region...