Master Note PDF - Chemistry Notes

Quick note: ONENOTE has better access to certain details and the lab covers the rest! 1.1.1 - A - Elements, Compounds and Mixtures Matter can be classified into two broad categories: Pure substances - constant composition ○ Set of properties such as melting point, color, boiling...

Quick note: ONENOTE has better access to certain details and the lab covers the rest! 1.1.1 - A - Elements, Compounds and Mixtures Matter can be classified into two broad categories: Pure substances - constant composition ○ Set of properties such as melting point, color, boiling point, etc. Mixtures Classification of Matter Examples Matter Classification Filtered Tea homogenous Selenium element Aluminum oxide compound Freshly squeezed orange juice heterogenous Stainless steel homogenous Table salt compound 1.1.1 - B - Separating Techniques Filtration - separate the components of a mixture containing an undissolved solid in a liquid Using gravity or applying vacuum Evaporation - separate components of a mixture with a dissolved solid in a liquid Distillation - separate components of a liquid mixture (miscible liquid) by a process of heating and cooling, which exploits the differences in the volatility of each of the components Chromatography - most versatile separation technique especially for small amount and anything in gas stage Two components ○ Stationary phase (e.g. silica, paper, etc.) ○ Mobile phase (the mixture) More soluble component would move faster Examples: Mixture Separation Technique Vinegar (5% acetic acid in water) Distillation Loose tea leaves in tea Filtration Copper sulfate (blue crystal) dissolved in water Evaporation 1.1.2 - Kinetic Molecular Theory Kinetic Molecular Theory (KMT) is a theory that explains the state of matter and is based on the idea that matter is composed of tiny particles that are always in motion. The theory applies mostly to gas. It specifically models a gas called an ideal gas. States of Matter Gas Particles have enough energy to move freely Molecules come into contact with one another only when they randomly collide Attraction forces between the particles are not strong enough to hold them together Liqui d Particles are constantly in contact but have enough energy to keep changing position relative to one another Attraction forces between particles are strong enough to keep the molecules relatively close together but not strong enough to prevent them from moving past one another. Solid Particles do not have enough energy to move Constantly in contact and in fixed positions relative to one another Attraction forces between atoms or molecules are strong enough to keep the molecules together and to prevent them from moving past one another Plas ma Ionized gas that is mainly nuclei Kinetic Molecular Theory (Ideal Gas) There are 5 assumptions that is made in order for a gas to be ideal. The volume of an individual gas molecule is negligible compared to the volume of its container There are neither attractive nor repulsive forces between gas molecules Gas molecules move randomly in all directions, in straight lines  Gas molecules do not lose kinetic energy when they collide with each other or other boundaries The kinetic energy of gas molecules is directly proportional to temperature Particle Models of Different States 1.1.3 - Kinetic Energy and Temperature Temperature is a measure of the average kinetic energy of particles. As the substances absorb energy, particles of a solid vibrate more, particles of a liquid vibrate more and move faster, while in gas they move faster. Heating Curve of Water Temperature Scale Kelvin (K) is the base unit of temperature measurement. Proportional to the average kinetic energy of particles and is considered an absolute scale. Absolute zero - implies that at this temperature, the particles cannot transfer any kinetic energy on collisions. Fahrenheit is another one measured by body temperature, 100 Fahrenheit means a fever! 1.2.1 - Atomic Model Models of the Atom Song: Atomic Theory Song Chemistry Democritus (500 B.C.) Greek philosopher and a "scientist" He proposed that everything is composed of "atoms" which are physically indivisible The atoms are indestructible and have always been in motion There are infinite number of atoms, and kinds of atoms, which differ in shape, and size. John Dalton (1803) - The Billiard Ball Model Dalton proposed the Atomic Theory in 1805 which stated that 1. All matter is made of atoms. Atoms are indivisible and indestructible. 2. All atoms of a given element are identical in mass and properties. 3. Compounds are formed by a combination of two or more different kinds of atoms. 4. A chemical reaction is a rearrangement of atoms. Dalton's theory identified chemical elements as a specific type of atom. Dalton’s view of the atom is that of a solid sphere, similar to a billiard ball. J. J. Thomson (1897) - The Plum Pudding Model In 1887, Thomson used a cathode ray tube and proposed that atoms contain negatively charged particles that have mass. He called them ‘corpuscles’, later called electrons. Thomson received the Nobel Prize in 1906 for his discovery. He theorized that the electrons were much like raisins in pudding, where the pudding is the positive particle. See Link for Cathode Ray Experiment: http://highered.mheducation.com/sites/0072512644/student_view0/chapter2/animations_center.html# Ernest Rutherford (1911) He theorized that using Thomson's model, alpha particles should pass through atoms unaffected.  See following link: http://phet.colorado.edu/simulations/rutherford-scattering/rutherford-scattering.jnlp Rutherford was shocked that the a particles bounced backwards. He concluded that the atom contained a tiny, dense core that was positively charged. See following links: ○ Rutherford's Experiment: Nuclear Atom ○ https://www.youtube.com/watch?v=XBqHkraf8iE  Niels Bohr (1912) - The Energy Level Model/ The Solar System Model Bohr studied the light produced when atoms were excited by hear or electricity Bohr describe electrons location as "orbits" & jump to a higher orbit when excited These orbits are specific distance from the nucleus Erwin Schrodinger (1926) Developed the "Schrodinger Equation" to model the atom which he won a Nobel Prize in Physics for "Electron Cloud Model" James Chadwick (1932) Discovery of neutron Chadwick observed that when beryllium-9 was exposed to alpha particles, particles with the same mass as protons but no charge were given off 1.2.2 - Isotopes Isotopes Isotopes are atoms of the same element with different numbers of neutrons. Examples: Cl - 35 and Cl - 37 Examples: uranium Isotopes # of protons # of neutrons # of electrons Abundance U - 232 92 140 92 trace U - 233 92 141 92 trace U - 234 92 142 92 0.005% U - 235 92 143 92 0.720% U - 236 92 144 92 trace U - 238 92 146 29 99.3% Radioisotopes Unstable form of isotope that emit radiation to transform into a more stable form Example: U - 235 to Th-231 Example of Usage of Radioisotopes Radioisotope Industrial Applications Co-60 Radiation therapy to prevent cancer I-131 Locate brain tumour, monitor cardiac, liver and thyroid activity Na-24 Study blood circulation Ir-192 Test integrity of boilers and aircraft parts U-235 Nuclear power plant 1.2.3 - Mass Spectra Mass Spectrometry (MS) Used to determine the relative atomic mass of an element from its isotopic composition The result of the mass spectrometer is shown in the form of a so-called "Mass Spectra" Basics: Imagine that there is a giant magnet that can ionized molecules to a molecular ion Ions produced when just a single electron is removed from the molecule is called molecular ion, M+  There will be fragmentation pattern because the molecule can be broken apart into small fragments when bombarding by high energy electrons. Mass Spectrometer The deflection or path of an ion in the mass spectrometer depends on: The absolute mass of the ion The charge of the ion The strength of the magnetic field The velocity (or speed) of the ion which is controlled by the strength of the electric field Common Fragmentation Patterns You would realized that the number is just the "molecular mass" of the fragmented ions How does a mass spectrum look like? Example: Pentane The tallest line in the stick diagram (in this case at m/z = 43) is called the base peak. This is usually given an arbitrary height of 100, and all the other heights will be relative to the base peak. Example: Pentan-3-one Example: Propane Example: Propanoic Acid Isotopes are present in some halogen-compounds e.g. chlorine. It is possible that we can use the MS spectrum to determine the relative abundance of each of the isotopes from the element. Mass Spectrometry Infographic: http://www.compoundchem.com/wp-content/uploads/2015/05/Mass-Spe ctrometry-Common-Mass-Spectra-Fragments.png Using Mass Spectrometry Data to find Average Atomic Mass Example 1: Use mass spectrometry data below to calculate the atomic mass of neon. Isotope Mass of atom in amu % abundance in nature (relative to C-12 = 12.0000) Neon - 20 19.9924 90.48 Neon - 21 20.9938 0.27 Neon - 22 21.9914 9.25 Example 2: Use mass spectrometry data below to calculate the atomic mass of gallium. Isotope Mass of atom in amu % abundance in nature (relative to C-12 = 12.0000) Gallium - 69 68.9256 60.11 Gallium - 71 70.9247 39.89 Example 3: Example 1: A result from the recording of the mass spectrometer is given below. Calculate iron's relative atomic mass with 2 decimal places. (6 digits cause of the original concept) 1.3.1 - Emission Spectra Electromagnetic radiation comes in different forms of different energy. All EM waves travel at the same speed (c) but can be distinguished by their wavelengths (𝜆). The wavelengtyh of the light controls the colour of the light. The relationship between the frequency (f), speed and wavelength is according to the following formula. = 𝜆 Electromagnetic Spectrum Atomic Emission Spectra In nature, the electrons in an atom tend to be arranged in such a way that the energy of the atom is as low as possible. Ground State: the lowest energy state of the atom Excited State: a state where its potential energy is higher than the ground state When an atom at ground state are given energy, the electrons absorb the energy and move to a higher energy level. These energy levels of the electrons in atoms are quantized, meaning again that the electron must move from one energy level to another in discrete steps rather than continuously. An atom in the excited state is not stable. When it returns back to the ground state, it releases the energy that had previously gained in the form of EM radiation. The energy of light can be calculated by =ℎ 𝜆 where h is the Planck's constant. (h = 6.63x 10-34Js) An atomic emission spectrum is the patterns of lines formed when light passes through a prism to separate it into different frequencies of light it contains. As different elements have different line spectra, they can be like barcodes to identify unknown elements. Here are images for examplation: This is the emission spectrum where what happens is when you make an electron excited in these elements the color that shows up is this! For example this is why neon is pink! 1.3.2 & 1.3.3 - Hydrogen Emission Spectrum & Bohr's Atomic Model When an electric current is passed through a gas tube that contains hydrogen gas at low pressure the tube gives off blue light. When this light is passed through a prism, four narrow bands of bright light are observed. When the hydrogen spectrum is expanded further than the visible spectrum, patterns of lines in both ultraviolet and infrared regions are detected as well. They are named as different series. The lines in the hydrogen emission spectrum form regular patterns and can be represented by a simple equation. Each line can be calculated from combination of simple whole numbers. When unexcited, hydrogen's electron is in the first energy level - the closest to the nucleus. When energy is supplied to the atom, the electron is excited into a higher energy level, or even removed from the atom altogether. The electron would tend to lose energy again by falling back down to a lower level. It can be done in many different ways. The amount of the energy emitted is referred as quantum of light or photon. Depending on the transition, this could result in many different series of lights. According to the Bohr model, the electrons encircle the nucleus of the atoms in specific allowable paths called orbits. Then the electron is in one of these orbits, its energy is fixed. The electron is not allowed to occupy any of the spaces in between the orbits. Analogy: a ladder Success of the Bohr Model: Explained the stability of the atom Explained the atomic line spectrum of the hydrogen atom Introduced the concept of stationary energy levels. Introduced a model of the atom that could be easily visualized Failures of Bohr Model: Did not explain the following: Line spectra for many electron atoms Electron configurations of many electron atoms The difference in energies of electrons occupying the same energy level. The shapes and characteristics of molecules. Difference between Atomic Absorption and Emission Line Spectra Other things to note: Electrons further from the nucleus require less energy to jump, and as electrons get excited they emit light! 1.3.4 - A - The Quantum Mechanical Model of the Atom The Bohr Model was an attempt to explain the energy states of electrons in atoms. It was based on quantization: the idea that electrons existed in discrete levels. According to Bohr, the emission spectra of hydrogen consisted of narrow lines because the wavelength of these lines corresponded to the differences in allowable energy levels. However, this model was limited by several problems and incorrect assumptions: 1. The model could not predict the emission spectra of elements containing more than one electron. It was only successful with the hydrogen atom. 2. It assumed the electron was a subatomic particle in fixed orbit about the nucleus. 3. It could not account for the effect of electric and magnetic fields on the spectral lines of atoms and ions. 4. It could not explain molecular bonding and geometry 5. Heisenberg's Uncertainty Principle states that it is impossible to precisely know the location and momentum of an electron simultaneously. Bohr's model stated that electron exhibited fixed momentum in specific circular orbits. Max Planck (1900) Proposed that energy is in discrete quantized amount of packets, rather than in a continuous unbroken wave  Each quantum is equal to the frequency of the radiation multiplied by a universal constant Photoelectric Effect (1905) Discovered in 1887 by the German scientist Hertz  Einstein proposed that a beam of light is not a wave but rather a collection of discrete wave packets (photon) Bohr Model (1913) De Broglie Wave Model (1929) (Watch this: Particles and waves: The central mystery of quantum mechanics - Chad Orzel) Wave Particle Duality ○ Electrons have wave like properties ○ Electrons do not follow a specific orbit, they are now in a wave pattern  ○ "Orbital" circumference has to fit whole waves (1, 2, 3) and not part waves (1.5) ○ Therefore, there can only be orbit of certain sizes (a.k.a. Quantum) Heisenberg Uncertainty Principle (1932) Due to wave nature of matter, it is impossible to predict both the position and momentum of an electron with certainty Schrodinger's Atomic Model - Orbitals (1933) Derived an equation that tells the PROBABILITY that an electron is at particular point  Hence Bohr's orbits were replaced by orbital: a region of probable location of electrons Examples of orbitals Orbits VS Orbitals Orbits Orbitals By Bohr Through a series of quantum mechanics theory 2 dimensional ring 3 dimensional space Electron is a fixed distance Electrons are a variable distance from nucleus from nucleus 2, 8, 18, ….. Per orbit 2 electrons per orbital 1.3.4 - B - Atomic Orbitals Atomic orbitals are regions of space where the probability of detecting an electron is high. They are categorized by their shapes. Each shape is labelled with a letter (s, p, d, f, g, h, etc.) When a magnetic field is applied to the sources, the single lines of the emission spectra split. The external magnetic field seems to be interacting with the electrons and changing the energy of the system slightly. Magnetic fields interact with objects what have a magnetic moment depending on the angle and direction. As a result, it can be derived that orbitals also have their angular moment. Quantum Numbers ○ Each electron surrounding a nucleus is described by a set of quantum numbers ○ Principle Quantum Numbers ○ Second Quantum Number / Angular Momentum Quantum Number ○ Magnetic Quantum Number ○ Spin Number ○ These quantum numbers also describe the behaviour of the electron as a wave ○ Example: 2, 1, -1, -1/2 Principle Quantum Numbers (n) ○ Indicates the energy level and distance from the nucleus of an electron (same as period number)  ○ Possible values: n = 1, 2, 3, 4,…. ○ The greatest number of electrons possible in each energy level is 2n2 Secondary Quantum Number (l) ○ Also known as Angular Momentum Quantum Number ○ Scientists hypothesized that there were sublevels within the main energy levels. ○ Each of these sublevels (orbitals) has a different shape or region with most likelihood of finding an electron ○ Possible values: l = 0 to n-1  Values of 0 1 2 3 4 l Letter s p d f g used Magnetic Quantum Number (ml) ○ Describes an orbital's orientation ○ Each sublevel orbital s, p, d, f has its own shape, but can have a different orientation about the nucleus. ○ Possible values: ml = -l to +l Spin Number (ms) ○ Also known as Magnetic Spin Number ○ Specifies the direction in which an electron is spinning (either up or down)  ○ Possible values: ms = -1/2 or +1/2 Pauli Exclusion Principle ○ Because of the spin, no two electrons in an atom can have the same four quantum numbers ○ Therefore: ○ Maximum 2 electrons in an orbital ○ Electrons within the same atomic orbital must have opposite spin One s orbital 1 x 2e- = 2 e- Combined three p 3 x 2e- = 6 e- orbitals Combined five d orbitals 5 x 2e- = 10 e- Combined seven f 7 x 2e- = 14 orbitals e- More explanation into it! One electron moves up one moves down, two per bubble! 1.3.5 - A - Electron Configurations According to Bohr, electrons are arranged in shells Example: O 2, 6 S 2, 8, 6 Ca 2, 8, 8, 2 Sc 2, 8, 9, 2 However, he couldn't explain why the third shell capacity is 18 electrons while in scandium after having 9 electrons, it goes to the 4th shell. Electron Configuration would be able to do that Electron Configurations Electron Configuration is a designation of how electrons are distributed among various orbitals. Single Electron System For hydrogen, each sublevel has the same energy in each energy level. For example, Multi Electron System Inner electrons act as shielding electron to the outer electrons from the positive charge of the nucleus. Example: Lithium Example: nitrogen Example: Sodium (Bohr: 2, 8, 1) Example: potassium Example: calcium Example: scandium Rules for Assigning Electrons to Orbitals The Pauli Exclusion Principle ○ Only two electrons may occupy the same orbital and these electrons must have opposing spins ○ Crudely, this could mean that two electrons cannot occupy the same place at the same time with the same orientation The Aufbau Principle ○ Aufbau is German for "build up" ○ Electrons enter the orbital with the lowest energy first until it is full. This minimizes the energy of the atom ○ Start filling orbitals from the lowest to the highest ○ If two or more orbitals exist at the same energy level, do not pair the electrons until you need to. The Hund's Rule ○ Electrons in the same sublevel will not pair up in an orbital until the sublevel is half filled ○ All unpaired electrons in the orbitals must have the same spin (up or down) ○ For example: nitrogen ○ For example: carbon Blocks of Orbital in Periodic Table In basic terms we have sets within the shells. Labeled as S, P,D,F, and further out we go more do those that are unlocked. The ending of one full shell, that subset, has or should be equal to next subset (first) in the next shell. So when you are filling it out usually just fill out the next shell, before you unlock the full subset of the current shell. When you do unlock the full one, you have to make sure that if there is space electrons will want to spread out so don’t draw up and down. Each arrow is one electron. To write it quicker then you have to just write the shell number, the subset letter IN ORDER, and in subscript the amount of electrons. If you want to be quicker you can write noble gasses to sum up a portion of it as you can see above. REMEMBER COPPER AND CHROMIUM ARE BOTH IRREGULAR. THOSE COLUMNS ARE DIFFERENT! 1.3.5 - B - Electron Configurations in Special Cases and Ions Electron Configuration of Neutron Atoms Example 1: phosphorus Example 2: titanium Example 3: argon Example 4: Chromium Special case 1: when there are 4 electrons on the d subshell, one electron will be borrowed from the s subshell so that the d orbital can be half full Example 5: copper Special case 2: when there are 9 electrons on the d subshell, one electron will be borrowed from the s subshell so that d subshell can be full Example 6: silver Electron Configurations of Ions Example 1: magnesium ion Example 2: chloride ion Example 3: lithium ion Dealing with Ions of Transition Metals Although d orbital is filled before the s orbital, electrons are taken from the s orbital when an ion is formed. This is because a full s subshell is on a slightly higher energy level than a partially full d subshell Example 1: Fe+2 Example 2: Fe+3 Example 3: Zn+2 With transition metals they lose electrons and gasses gain electrons. So you either add on or subtract. Now a cool concept is that if d has space then s will give it. This is because the further the subset is away from the first subset of the shell the more space it has. This allows it to have a bigger capacity. Now with transition metals the d doesn't lose electrons, rather it is the s. Because it is further away! If it is a noble gas in itself then you can’t shorten it to itself, if it’s another element that got ionized then you can still use noble gasses. 1.3.6 - Ionization Energies Ionization energy (IE) is the minimum energy (in kJ/mol) required to remove an electron from a gaseous atom in its ground state. The amount of energy in kJ needed to strip 1 mole of electrons from 1 mole of gaseous atom. Always an endothermic process First Ionization Energy Second and Third Ionization Energy In general, each ionization should require more energy than the previous electron because of the nuclear charge of the atom. Convergence of Ionization Energy As the energy increases, the energy levels in the hydrogen atom become more closely spaced until they converge at a point, which corresponding to the electron leaving the atom. The electron is free to move around and no longer under the electrostatic influence of the nucleus once it reaches the convergence limit. That is ionization. The energy difference between this energy level and the ground state (n=1) is called ionization energy Example 1: For the atomic emission spectrum of hydrogen, the convergence limit occurs at a frequency of 3.27 x 1015 Hz. Calculate the ionization energy for a single hydrogen atom and for a mole of hydrogen atoms. Example 2: The emission spectrum of helium has its lowest wavelength at around 390nm in the visible spectrum and lowest wavelength at 50nm in the EM radiation spectrum. What is the ionization energy for a mole of helium atoms? 1.3.7 - Successive Ionization Energy Successive ionization energies of an element are the amounts of energy required to remove all the electrons from one mole of an element in the gaseous state, one mole of electrons at a time. It depends on the energy level the electron is on and the number of electrons in the valence shell As electrons on the same energy level are removed one at a time, the ionization energy gradually increases, because electrons are taken from a more positively charged ion  Successive ionization energy provides evidence about electron configuration of an atom Example: Nitrogen Nitrogen has five electrons on its valence shell The first five successive ionization energy increase gradually because electrons being removed are from the same energy level There is a dramatic increase between 5th and 6th ionization energy, because the 5th and 6th electron are on different energy levels Example 2: Magnesium