Grade 12 Chemistry - Development of the Atom PDF

9Grade 12 Chem Development of the Atom Democritus - 460 BCE - Speculated (intuition) that matter is composed of particles called atoms - Thousands of years later it was proven by scientists, now they use STM- scanning tunnel...

9Grade 12 Chem Development of the Atom Democritus - 460 BCE - Speculated (intuition) that matter is composed of particles called atoms - Thousands of years later it was proven by scientists, now they use STM- scanning tunnel microscopes John Dalton - 1805 - Made atomic theory - Made a modern theory saying that matter is composed of indestructible, invisible atoms which are identical for one element but different from other elements - All elements are different → proton number - He called it the Billiard ball model - His theory explains 3 laws Law of conservation of mass → Atoms are never created or destroyed Law of multiple proportions → when atoms combine to become stable they can do it in more than one way (H2O or H2O2) Law of definite composition → always going to get the same composition of mass (in water there will always be the same composition of water and oxygen) - Theory All matter is composed of extremely small particles called atoms Atoms can not be created or destroyed Atoms can not be subdivided or changed into another All atoms of the same element are the same and different from other elements Atoms combine in small whole-number ratios William Crookes New - Cathode-ray tube - Helped in the discovery of the electron, but did not know it and j.j thomson took credit it for the discovery of the lectron because he understood what made the gas to glow J.J Thomson - 1897  - Believed in the billiard ball model of the atom - Used a cathode-ray tube which was developed by crooks to study rays - In his studies he was credited with discovering the electrons (negative particles) which is embedded in a positive material - The kind of element is determined by the number of electrons in the atom - Raisin bun model - E. Goldstein NEW - Discover the existence of positively charged rays. He discovered the positively charged particles, but rutherford, took credit and called them protons Hantaro Nahaoka NEW - Proposed a model similar to rutherdords - Large positive sphere surrounded by a ring a negative electrons Robert Milikan - Discovered a charge and mass of electron - Charged oil drops - When the x-rays touched the oil droplets turning them negative some of them would float (-1) or some of the would rise (-2) - He was able to calculate the mass of the electron Ernest Rutherford - 1911 - Believed in the raisin bun model - During radioactive decay, alpha particles are positively charged and discovered the nucleus - Tested by using alpha rays to go through a piece of gold foil and thought that the rays would go directly through it - When it hit the rays it shot in all different directions giving him the idea that atoms are made up of empty space with a dense center called the nucleus - If particles get too close together they are deflected (same charges repel, different charges attract) -  - James Chadwick - 1932 - Credited with the discovery of the neutron - Performed experiments similar to the gold foil experiment (using beryllium) to determine the mass of the nucleus - When the masses of the nuclei were compared to the masses of the protons for an element they noticed that the mass of the protons were about half of the mass of the nucleus which led them to discover that neutrons existed Neils Bohr - 1913 - Applied electricity and thermal energy to hydrogen gas - He tried to explain why the electron does not belong in the nucleus of the atom - Observed that hydrogen atoms emitted light when they were excited by the additional energy - He directed the light through a glass prism with a screen behind it and observed lines of only certain colours of light - Observation led to the thought that electrons orbit the nucleus in definite energy levels - Planetary model of the atom (electrons orbit around the nucleus) - Max number of electrons in each orbit and electrons jump to a higher level when energy is absorbed or dropped to a lower level when energy is emitted - The further the electrons are from the nucleus the more energy it has Francis Aston - 1919 - Discovery of isotopes - Using a mass spectrometer (a modified cathode tube) noticed that the masses of the atoms of the same elements can have different masses Atomic model today (BR diagrams) - Atom has a nucleus which contains protons and neutrons - Electrons orbit around the nucleus in circular panthers of specific energy values - Determine protons and neutrons by looking at periodic table - Showing the arrangement of electrons around the nucleus - Make sure to pair up electrons in clockwise circular motion - In the first 4 orbits you put 2,8,8,18 -  - Origins of Quantum Theoy and Bhor model 1. The origins of Quantum Theory Rutherford asked 4 questions: 1. What is the details structure of the atom 2. WHat keeps the nucleus together 3. How exactly are the electrons arranged around the nucleus? 4. What are the negatively charged electrons not attracted and therefore spiral into the nucleus which is positively charged? Spectroscopy: Developed by Robert Bunsen & Gustav Specific elements produce a flame of a specific colour. When there was a new colour determined a new element was discovered Electromagnetic Spectrum: Developed by James Maxwell He proposed that light is an electromagnetic wave that is made of electric and magnetic fields. Classical theory of light Photoelectric Effect Developed by Hienrich Hertz He observed that when light strikes the surface of a metal, electrons are emitted. The colour of the light is what determines the energy of the emitted electrons  Quantum Theory Discovred by Max Planck - He saw that red hot is colder than white hot That when light is absorbed or emitted by atoms in a soiled, the energy is absorbed or emitted in bundles. He called these bundles Quanta Photon Discovered by Albert Einstein - Used plank’ideas of quantum energy to explain the photoelectric effect - He said that light consists of a stream of energy packets which he called photons - Each photon and a curtain frequency/light/energy - E=HF - h=Plank’s constant f=frequency - ONe photon of energy emits one electron, and energy needs tob e strong enough for the electron to be ejected - 2. Bohr Model of the atom Explain why the electrons dont fall into the nucleus and stay around the proton and nucleus - Bohr came up with 8 postulates (need to know this_ 1. Electrons can be found confined to st energy levels. He called these energy levels orbits or shells. 2. Each energy level consists of certain amounts of energy 3. The orbit closest to the nucleus has the least amount of energy 4. Energy increases as orbits get farther away from the nucleus 5. In order for an electron to occupy an orbit it must have identical energy as the orbit 6. Electrons cannot occupy space between orbit 7. Electrons do not radiate energy as they orbit the nucleus 8. If an electron makes a jump form higher orbit to a lower orbit energy is released. If an electron makes a jump from a lower orbit to a higher orbit, energy is absorbed An electron is in its ground state when it is in its lowest possible energy level based on the energy it posses (ex: H only has one electron level so that is in its ground state but to go to level to the electron has to absorb energy to move to level 2 and it is called an excited state although it is not stable in this state so it will want to come back down- all og levels are grounded ) - When it is in its excited state the electron just moves orbits. Making it excited - Emit energy =Grounded  - Absorb energy +excted When a electron is ubstale it is in an excites state, and it is unstable and sontaneoulsy jumpers back to itts grounds tate When making the jump from high energy to ground state, electron gives off energy in the form of a photon of light which corresponds to the colured lines seen in emission spectrum. Excited state would be in a dark line, but when it comes into ground sate, the emission spectrum, show the bright line. Hydrogen Emmision Spectrum (dont really need to know) Bohr concreted the H atom. He calculated the energy that an electron would have to possess in order to occupy each of the energy levels. From these calculations, Bohr was abe to show that a jump of the H electron from n=5 to n=2 emitted energy corresponding to the blue light Bhor calculations only worked for hydrogen and could not account for the spectra of multi electron atoms. Lesson 3 Quantum Mechanics - The quantum theory could not answer questions regarding the behabiour of electrons in the tom - Wave mechanics or quantum mechanics was developed by louis de brogile, schrodinger and heisenberg. - By 1923 the dual nature of light was well understood. Electrons as waves They looked at the electron as a wave, and a lot of thing starsted clicking - Louis de brogile used planck's and einstein’s work E=hf and e=mc2 to support his hypothesis QUantum Mechaincs - Scrodinger discovered that LOuis de Brogile’s electron wave could be understood to understand the behaviour of the electron inside the atom - *He Propsed that electrons are stable circular waves around the nucleus and therefore, they could have certain energies - Heisenberg, showed mathematically that there are definmite limits to our ability to know both where a particle and its speed at a givem time - Heisnberg uncertainly principle  - Impossible to now both the momentum and po stion of an electron at the same time - This is were orbital came from - Heinsberg said that it is a impossible to find an electron Schrodinger wave equation: Has a finite # of solutions → quantized energy levels Describes the electron as a 3-dimensional wave in the electric field of a positively charged nucleus Defines probability of finding an electron (orbitals) Schrodinger found it - Electron cna only have certain energies - Schrodingers wave mechanics included all four quantum numbers nad produced the erngies of all electron oribitals Orbitals - Higher changeo f finding an electron closer to the nucleus - 4 different types - are regions of space where is 90% probability of finding an e Atomic Orbitals - Principal quantum number (n) - energy level of the electros The four different types of orbitlas 1. S orbitals (know the shapes) - There is one 2 oribial for every energy level. - Each s orbital can hold 2 electrons - Called the 1s, 2s, 3s, etc.. orbitals o 2. P Orbitals (know the shapes) - Start at the second energy level - 2p, 3p, 4p - 3 different orientations (shapes) - Each can hold 2 electrons all of equal energy - They come in a set of three 3. D Orbitals - Start at the third energy level - 3d, 4, 5d, etdc - 5 different orientations (shapes) - Each can 2 electrons - 4. F orbitals - Start at the fourth energy level - 4f, 5f, etc - Have seven different orientations (shapes) - 2 electrons per shape Summary Table # shapes Max electrons Stars at energy level? s 1 2 1 p 3 6 2 d 5 10 3 f 7 14 4 Wave Mechanical Model (1920s) - Particle-wave duality: Electrons exhibit both particle and wave properties. - Nucleus surrounded by pulsating electron waves. - Wave mechanics help determine the size, shape, and direction of electron regions. - High probability electron regions are called orbitals. - Orbitals are described by quantum numbers, which indicate electron location. - Quantum numbers are derived by solving the wave equation for each atom. - how they fall in place Quantum Numbers 1. Principal Quantum Numbers - Energy levels - Size of the orbital - N2 = # of orbitals in the energy level 2. Secondary Quantum Number (0>l>3) - It he;ps in describing the sptital distribution and shape of electron cloud around the nucleus - l = 0: s orbital - l = 1: p orbital - l = 2: d orbital  - l = 3: f orbital - - 3. Magnetic Quantum number ( - Was created because of the when magnetic was placed they found foun single line spilt into new lines The magnetic quantum number (denoted as ml) specifies the orientation of an electron's orbital in space relative to an external magnetic field. The value of ml depends on the secondary quantum number (l) and can range from -l to +l, including zero. For each value of l, there are multiple possible orientations of the orbital: ○ If l = 0 (s orbital), ml can only be 0 (since there's only one orientation). ○ If l = 1 (p orbital), ml can be -1, 0, or +1 (three possible orientations). ○ If l = 2 (d orbital), ml can be -2, -1, 0, +1, or +2 (five orientations). These values define the number of possible orientations an orbital can have in three-dimensional space, and this becomes important in the presence of a magnetic field, where different orientations have slightly different energies. Speicfie sthe exact energy level 4. Spin Quantum Number (ms =+½ or -½ ) - Has only two different possible numbers - Added to account for different kinds of magnetism - Ferromagnetism Was added for strong substances (Fe, Co, and Ni) - Paramagnetism is weak - Each electron in a orbital would have option spins - Pauli Exclusion Principle, One electron spins clock wise nad the other spins counter clockwise - All electrons have all four wuatums numbers Pauli Exlusion Prinple  - No two electrons in an atom can have the same 4 quantum numbefrs - Each e- has a unique “address” 1. Principal # → energy level 2. Anular momentant → # sublevel (s, p, d,f) 3. Magnetic # → orbital 4. Spin # → electron Electrons in Atoms As you move away from the first energy level the Energy level diagrams - Show the energies of electrons in various orbitals - Electrons inan atom have different engries Pauli exclusion Principle - Each orbital can hold Two electrons with opposite spins Aufbau Prinple - Electrons will fall into lowest energy sublevels first and fill from there - Electrons fill the lowest energy orbitals first - “Lazy tenant rule” - In order to fill the 2p orbitals 1s has to be filled - Only 7 energy levels Hund’s Rule - Within sublevel palace one e- per orbital before pairing them - This is kinda of like when filling out dots in an electron E.g O The subsprict is the number of electrons within the level This xist because you cant do bhor rutherford diagram until 20, but now you can! Long hand Short hand Inner shells valence shells valence shells Repersentive elements (Ramon A) S and p levels Transition metals The valence shell would be d f orbtials  When filling, think about when it's going - S and p block smae as period number - D block 1 less - F block 2 less You have to replace something with a electron that it fully stable (Noble Gases) E.g Germainum Ar 4s2 3d104p2 The argon replace the other three energy levels Stability - Full Energy elvels - Cases where not all sub levels are energy are filled up, only s p d, are filled up - Half-full sublevel (most important in d, f) - When they are one away from being half filled they are very unstable elements - Expections for Aufbau and Hund's rule E.x Cooper: Expect: [ar] 4s23d9 This happens because cooper is very unstable What acc happens: [ar] 4s13d10 in this state Chromium Expect: [Ar] D4 d9 f6 f13 these are the ones that have exceptions It gives it more stability Ion Formation When a metal atom lose or gains electrons to become stable It becomes isoelectronic (same) with noble gases Loess → goes behind Gains → goes in front (like bhor rutherford you know this) Only applies for roman numerals A not B Saying that when something losses Anions (negative) Electron Configuration Add the gained electrons to the total # of electrons found in the neutral atom and draw the electronic configuration Write the electron configuration for the closest noble gas O2-=Ne [He]2s22p6 or [Ne] Cations Electronic Configuration Show the electron configuration of the neutral atom and then remove electrons from the highest quantum number orbital zn = [Ar]4s23d10 Zn2+: [Ar] 3d10 (stable) Zn2+:[Ar]4s23d8 (not stable) YOu have to remove electrons from 4s (see how it goes backwards) Ga: [Ar]4s23d104p1 (you remove from p first because its highest ) Ga3+: [Ar]3d10 All Roman Number B and metals are gaining electrons. NANd positively Magnetism Associated with electron spin and the presence of unpaired electrons 1. Ferromagnetism: - Based on the properties of a collection of atoms - Each atom acts like a little magnet - They influence each other to form groups (domains in which all the atoms are oriented in the same direction) - They all go in the same direction 2. Paragamenstim - Due to unpaired electrons within substances where domains do not form - Based on the magnetism of individuals atoms - Weaker Lesson 5 - Application of Quantum Mechanics Quatum technologies are use in larsers, microchips, MRI machines, Ct sances and other things Laser Technology -1917 (Einstein) Amplification of photos) - Einstein established the theoretical foundations for lasers. - L.a.s.e.r : Light Ampilfication by Stimulated Emission of Radiation - Laser- Adevice that produces an intense beam of light  1. The light is completely monochromatic. That means it is all the same wavelength (and therefore the same colour). 2. The light is made of coherent waves (waves vibrate in the smae direction at he same time). 3. Waves are parallel to each other. MRI (Magnetic Resonance Imaging) - Provides a much more detail picture of the body (compared to x-ray) - Can see soft tissues because it has a higher water continent - Can tel lthe difference between soft tissues and cancer - Powerful magnent lines up all the protons of hydrogen atoms - A pulse of radio waves is then sent out. This causes the portions to scatter - As the protons realign back in place, they emit a signal - Doesnt use ionizing radtion → safer for the body X-ray - They can damage tissues - X-rays are extremely high energy photons (have to wear protivce wear) - Radiation can cause cancers - Ct scan uses X-ray and can create Ct of CAT scan - CT- Computer tomography - Produces images through series of thin X-ray sections through the body - Patient is moved slowly through the CT machine as x-ray machines rotates around the,. - -